Advances in catalyst deactivation

Bartholomew , Calvin H.

Argyle, Morris D.

2016

MDPI - Multidisciplinary Digital Publishing Institute

Catalyst deactivation, the loss over time of catalytic activity and/or selectivity, is a problem of great and continuing concern in the practice of industrial catalytic processes. Costs to industry for catalyst replacement and process shutdown total tens of billions of dollars per year. While catalyst deactivation is inevitable for most processes, some of its immediate, drastic consequences may be avoided, postponed, or even reversed. Accordingly, there is considerable motivation to better understand catalyst decay and regeneration. Indeed, the science of catalyst deactivation and regeneration has been developing rapidly as evidenced by the considerable literature addressing this topic, including 21,000 journal articles, presentations, reports, reviews, and books; and more than 29,000 patents for the period of 1980 to 2012. This developing science provides the foundation for continuing, substantial improvements in the efficiency and economics of catalytic processes through development of catalyst deactivation models, more stable catalysts, and regeneration processes. This special issue focuses on recent advances in catalyst deactivation and regeneration, including advances in (1) scientific understanding of mechanisms; (2) development of improved methods and tools for investigation; and (3) more robust models of deactivation and regeneration.

Chemical Engineering

English

9783038421870 (print)

9783038421887 (online)

Advances in Catalyst
Deactivation
Edited by

Calvin H. Bartholomew and Morris D. Argyle

Printed Edition of the Special Issue Published in Catalysts

www.mdpi.com/journal/catalysts

Calvin H. Bartholomew and Morris D. Argyle (Eds.)

Advances in Catalyst Deactivation

This book is a reprint of the Special Issue that appeared in the online, open access journal,
Catalysts (ISSN 2073-4344) from 2013–2015 (available at:
http://www.mdpi.com/journal/catalysts/special_issues/catalyst-deactivation).
Guest Editors
Calvin H. Bartholomew
Brigham Young University
USA
Morris D. Argyle
Brigham Young University
USA
Editorial Office
MDPI AG
Klybeckstrasse 64
Basel, Switzerland
Publisher
Shu-Kun Lin
Managing Editor
Mary Fan

1. Edition 2016
MDPI • Basel • Beijing • Wuhan • Barcelona
ISSN 978-3-03842-187-0 (Hbk)
ISSN 978-3-03842-188-7 (PDF)

© 2016 by the authors; licensee MDPI, Basel, Switzerland. All articles in this volume are
Open Access distributed under the Creative Commons Attribution license (CC BY), which
allows users to download, copy and build upon published articles even for commercial
purposes, as long as the author and publisher are properly credited, which ensures maximum
dissemination and a wider impact of our publications. However, the dissemination and
distribution of physical copies of this book as a whole is restricted to MDPI, Basel,
Switzerland.

III

Table of Contents
List of Contributors ............................................................................................................... V
About the Guest Editors...................................................................................................... VII
Preface .................................................................................................................................IX
Morris D. Argyle and Calvin H. Bartholomew
Heterogeneous Catalyst Deactivation and Regeneration: A Review
Reprinted from: Catalysts 2015, 5(1), 145-269
http://www.mdpi.com/2073-4344/5/1/145 .............................................................................. 1
Erling Rytter and Anders Holmen
Deactivation and Regeneration of Commercial Type Fischer-Tropsch
Co-Catalysts—A Mini-Review
Reprinted from: Catalysts 2015, 5(2), 478-499
http://www.mdpi.com/2073-4344/5/2/478 .......................................................................... 129
Gary Jacobs, Wenping Ma and Burtron H. Davis
Influence of Reduction Promoters on Stability of Cobalt/g-Alumina Fischer-Tropsch
Synthesis Catalysts
Reprinted from: Catalysts 2014, 4(1), 49-76
http://www.mdpi.com/2073-4344/4/1/49 ............................................................................ 152
Rahman Gholami, Mina Alyani and Kevin J. Smith
Deactivation of Pd Catalysts by Water during Low Temperature Methane Oxidation Relevant
to Natural Gas Vehicle Converters
Reprinted from: Catalysts 2015, 5(2), 561-594
http://www.mdpi.com/2073-4344/5/2/561 .......................................................................... 180
Jose Antonio Calles, Alicia Carrero, Arturo Javier Vizcaíno and Montaña Lindo
Effect of Ce and Zr Addition to Ni/SiO2 Catalysts for Hydrogen Production through Ethanol
Steam Reforming
Reprinted from: Catalysts 2015, 5(1), 58-76
http://www.mdpi.com/2073-4344/5/1/58 ............................................................................ 214

IV
Giuseppe Trunfio and Francesco Arena
Deactivation Pattern of a “Model” Ni/MgO Catalyst in the Pre-Reforming of n-Hexane
Reprinted from: Catalysts 2014, 4(2), 196-214
http://www.mdpi.com/2073-4344/4/2/196 .......................................................................... 233
Emmanuel Skupien, Rob J. Berger, Vera P. Santos, Jorge Gascon, Michiel Makkee,
Michiel T. Kreutzer, Patricia J. Kooyman, Jacob A. Moulijn and Freek Kapteijn
Inhibition of a Gold-Based Catalyst in Benzyl Alcohol Oxidation: Understanding
and Remediation
Reprinted from: Catalysts 2014, 4(2), 89-115
http://www.mdpi.com/2073-4344/4/2/89 ............................................................................ 252
Vincenzo Vaiano, Diana Sannino, Ana Rita Almeida, Guido Mul and Paolo Ciambelli
Investigation of the Deactivation Phenomena Occurring in the Cyclohexane Photocatalytic
Oxidative Dehydrogenation on MoOx/TiO2 through Gas Phase and in situ DRIFTS Analyses
Reprinted from: Catalysts 2013, 3(4), 978-997
http://www.mdpi.com/2073-4344/3/4/978 .......................................................................... 279

V

List of Contributors
Ana Rita Almeida: Faculty of Science & Technology, University of Twente, PO Box 217,
Meander 225, 7500 AE, Enschede, The Netherlands.
Mina Alyani: Department of Chemical & Biological Engineering, University of British
Columbia, 2360 East Mall, Vancouver, BC V6T 1Z3, Canada.
Francesco Arena: Dipartimento di Ingegneria Elettronica, Chimica e Ingegneria Industriale,
Università degli Studi di Messina, Viale F. Stagno D’Alcontres 31, I-98166 Messina, Italy;
Istituto CNR-ITAE “Nicola Giordano”, Salita S. Lucia 5, I-98126 Messina, Italy.
Morris D. Argyle: Chemical Engineering Department, Brigham Young University, Provo,
UT 84602, USA.
Calvin H. Bartholomew: Chemical Engineering Department, Brigham Young University,
Provo, UT 84602, USA.
Rob J. Berger: Anaproc c/o Delft University of Technology, Julianalaan 136, 2628 BL Delft,
The Netherlands.
Jose Antonio Calles: Department of Chemical and Energy Technology, Rey Juan Carlos
Universtity, c/Tulipán, s/n, Móstoles 28933, Spain.
Alicia Carrero: Department of Chemical and Energy Technology, Rey Juan Carlos
Universtity, c/Tulipán, s/n, Móstoles 28933, Spain.
Paolo Ciambelli: Department of Industrial Engineering, University of Salerno, Via Giovanni
Paolo II, 132; 84084, Fisciano, SA, Italy.
Burtron H. Davis: Center for Applied Energy Research, 2540 Research Park Drive,
Lexington, KY 40511, USA.
Jorge Gascon: Catalysis Engineering, Chemical Engineering Department, Delft University of
Technology, Julianalaan 136, 2628 BL Delft, The Netherlands.
Rahman Gholami: Department of Chemical & Biological Engineering, University of British
Columbia, 2360 East Mall, Vancouver, BC V6T 1Z3, Canada.
Anders Holmen: Department of Chemical Engineering, Norwegian University of Science
and Technology (NTNU), N-7491 Trondheim, Norway.
Gary Jacobs: Center for Applied Energy Research, 2540 Research Park Drive, Lexington,
KY 40511, USA.
Freek Kapteijn: Catalysis Engineering, Chemical Engineering Department, Delft University
of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands.
Patricia J. Kooyman: Catalysis Engineering, Chemical Engineering Department, Delft
University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands.
Michiel T. Kreutzer: Product & Process Engineering, Chemical Engineering Department,
Delft University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands.
Montaña Lindo: Department of Chemical and Energy Technology, Rey Juan Carlos
Universtity, c/Tulipán, s/n, Móstoles 28933, Spain.
Wenping Ma: Center for Applied Energy Research, 2540 Research Park Drive, Lexington,
KY 40511, USA.

VI
Michiel Makkee: Catalysis Engineering, Chemical Engineering Department, Delft University
of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands.
Jacob A. Moulijn: Catalysis Engineering, Chemical Engineering Department, Delft
University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands.
Guido Mul: Faculty of Science & Technology, University of Twente, PO Box 217, Meander
225, 7500 AE, Enschede, The Netherlands.
Erling Rytter: Department of Chemical Engineering, Norwegian University of Science and
Technology (NTNU), N-7491 Trondheim, Norway˗ SINTEF Materials and Chemistry,
N-7465 Trondheim, Norway.
Diana Sannino: Department of Industrial Engineering, University of Salerno, Via Giovanni
Paolo II, 132; 84084, Fisciano, SA, Italy.
Vera P. Santos: Catalysis Engineering, Chemical Engineering Department, Delft University
of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands.
Emmanuel Skupien: Catalysis Engineering, Chemical Engineering Department, Delft
University of Technology, Julianalaan 136, 2628 BL Delft, The Netherlands.
Kevin J. Smith: Department of Chemical & Biological Engineering, University of British
Columbia, 2360 East Mall, Vancouver, BC V6T 1Z3, Canada.
Giuseppe Trunfio: Dipartimento di Ingegneria Elettronica, Chimica e Ingegneria Industriale,
Università degli Studi di Messina, Viale F. Stagno D'Alcontres 31, I-98166 Messina, Italy
Vincenzo Vaiano: Department of Industrial Engineering, University of Salerno, Via
Giovanni Paolo II, 132; 84084, Fisciano, SA, Italy.
Arturo Javier Vizcaíno: Department of Chemical and Energy Technology, Rey Juan Carlos
Universtity, c/Tulipán, s/n, Móstoles 28933, Spain.

VII

About the Guest Editors
Calvin H. Bartholomew, Professor Emeritus of Chemical
Engineering at Brigham Young University (BYU), has taught and
mentored students at BYU in catalysis, materials, and catalyst
deactivation for 42 years. He is an active researcher in
heterogeneous catalysis and a recognized authority on FischerTropsch synthesis and catalyst deactivation; he has co-authored over
140 journal articles, 20 chapters, four books, and three patents. He is
co-author with Dr. Robert Farrauto of Fundamentals of
Industrial Catalytic Processes, a leading handbook and textbook.
Together with Professors Bill Hecker and Morris Argyle of BYU, he has taught short courses
on “Heterogeneous Catalysis,” “Fischer Tropsch Synthesis,” and “Catalyst Deactivation” to
more than 700 professionals from industry and academe. He has worked at four companies
and consulted with more than 70 company clients on catalyst and support design, FischerTropsch synthesis, selective catalytic reduction of nitrogen oxides, FT reactor design,
BTL/GTL process design, and litigation relating to catalyst failure.
Morris D. Argyle is an Associate Professor of Chemical
Engineering at Brigham Young University (BYU). After earning his
bachelor’s degree, he became interested in catalysis while working
as the process engineer for one of the fluid catalytic cracking units
at the Exxon Baytown Texas Refinery. After completing graduate
school at the University of California, Berkeley, he joined the
Department of Chemical and Petroleum Engineering at the
University of Wyoming, where he became an Associate Professor
and served as Department Head before joining the faculty at BYU in
2009. His research interests include metal oxide catalysts for
oxidative dehydrogenation of light alkanes, high temperature water gas shift catalysts for
hydrogen production, Fischer-Tropsch catalysis, plasma reactions, and carbon capture
techniques. He also shares Professor Calvin Bartholomew's interest in catalyst deactivation.
He has co-authored 38 journal articles, one book chapter, and five patents.

IX

Preface
Catalyst deactivation, the loss over time of catalytic activity and/or selectivity, is a
problem of immense and ongoing concern in the practice of industrial catalytic processes.
Costs to industry for catalyst replacement and process shutdown total tens of billions of
dollars per year. While catalyst deactivation is inevitable for most processes, some of its
immediate, drastic consequences may be avoided, postponed, or even reversed. Accordingly,
there is considerable motivation to better understand catalyst decay and regeneration. Indeed,
the science of catalyst deactivation and regeneration and its practice have been expanding
rapidly, as evidenced by the extensive growth of literature addressing these topics. This
developing science provides the foundation for continuing, substantial improvements in the
efficiency and economics of catalytic processes through development of catalyst deactivation
models, more stable catalysts, and regeneration processes.
This special issue focuses on recent advances in catalyst deactivation and regeneration,
including advances in: (1) scientific understanding of mechanisms; (2) development of
improved methods and tools for investigation; and (3) more robust models of deactivation and
regeneration. It consists mainly of topical reviews.
The editors thank Keith Hohn, Editor-in-Chief, for the opportunity to organize this
special issue and Mary Fan, Senior Assistant Editor, and the staff of the Catalysts Editorial
Office for their significant support, encouragement, and patience. We would also like to thank
the reviewers of the submitted manuscripts for their invaluable recommendations, and the
contributing authors for their hard work in revising their manuscripts several times in order to
meet the high standards of this special issue. The quality of the published work appears to
have rewarded these efforts.
Calvin H. Bartholomew and Morris D. Argyle
Guest Editors

1

Heterogeneous Catalyst Deactivation and Regeneration:
A Review
Morris D. Argyle and Calvin H. Bartholomew
Abstract: Deactivation of heterogeneous catalysts is a ubiquitous problem that causes loss of
catalytic rate with time. This review on deactivation and regeneration of heterogeneous catalysts
classifies deactivation by type (chemical, thermal, and mechanical) and by mechanism (poisoning,
fouling, thermal degradation, vapor formation, vapor-solid and solid-solid reactions, and
attrition/crushing). The key features and considerations for each of these deactivation types is
reviewed in detail with reference to the latest literature reports in these areas. Two case studies on
the deactivation mechanisms of catalysts used for cobalt Fischer-Tropsch and selective catalytic
reduction are considered to provide additional depth in the topics of sintering, coking, poisoning,
and fouling. Regeneration considerations and options are also briefly discussed for each
deactivation mechanism.
Reprinted from Catalysts. Cite as: Argyle, M.D.; Bartholomew, C.H. Heterogeneous Catalyst
Deactivation and Regeneration: A Review. Catalysts 2015, 5, 145-269.
1. Introduction
Catalyst deactivation, the loss over time of catalytic activity and/or selectivity, is a problem of
great and continuing concern in the practice of industrial catalytic processes. Costs to industry for
catalyst replacement and process shutdown total billions of dollars per year. Time scales for
catalyst deactivation vary considerably; for example, in the case of cracking catalysts, catalyst
mortality may be on the order of seconds, while in ammonia synthesis the iron catalyst may last for
5–10 years. However, it is inevitable that all catalysts will decay.
Typically, the loss of activity in a well-controlled process occurs slowly. However, process
upsets or poorly designed hardware can bring about catastrophic failure. For example, in steam
reforming of methane or naphtha, great care must be taken to avoid reactor operation at excessively
high temperatures or at steam-to-hydrocarbon ratios below a critical value. Indeed, these conditions
can cause formation of large quantities of carbon filaments that plug catalyst pores and voids,
pulverize catalyst pellets, and bring about process shutdown, all within a few hours.
While catalyst deactivation is inevitable for most processes, some of its immediate, drastic
consequences may be avoided, postponed, or even reversed. Thus, deactivation issues (i.e., extent,
rate, and reactivation) greatly impact research, development, design, and operation of commercial
processes. Accordingly, there is considerable motivation to understand and treat catalyst decay.
Over the past three decades, the science of catalyst deactivation has been steadily developing, while
literature addressing this topic has expanded considerably to include books [1–4], comprehensive
reviews [5–8], proceedings of international symposia [9–14], topical journal issues (e.g., [15]), and
more than 20,000 U.S. patents for the period of 1976–2013. (In a U.S. patent search conducted in
November 2013 for the keywords catalyst and deactivation, catalyst and life, and catalyst and

2
regeneration, 14,712, 62,945, and 22,520 patents were found respectively.) This area of research
provides a critical understanding that is the foundation for modeling deactivation processes,
designing stable catalysts, and optimizing processes to prevent or slow catalyst deactivation.
The purpose of this article is to provide the reader with a comprehensive overview of the
scientific and practical aspects of catalyst deactivation with a focus on mechanisms of catalyst
decay, prevention of deactivation, and regeneration of catalysts. Case studies of deactivation and
regeneration of Co Fischer-Tropsch catalysts and of commercial catalysts for selective catalytic
reduction of nitrogen oxides in stationary sources have been included.
2. Mechanisms of Deactivation
There are many paths for heterogeneous catalyst decay. For example, a catalyst solid may be
poisoned by any one of a dozen contaminants present in the feed. Its surface, pores, and voids may
be fouled by carbon or coke produced by cracking/condensation reactions of hydrocarbon
reactants, intermediates, and/or products. In the treatment of a power plant flue gas, the catalyst can
be dusted or eroded by and/or plugged with fly ash. Catalytic converters used to reduce emissions
from gasoline or diesel engines may be poisoned or fouled by fuel or lubricant additives and/or
engine corrosion products. If the catalytic reaction is conducted at high temperatures, thermal
degradation may occur in the form of active phase crystallite growth, collapse of the carrier
(support) pore structure, and/or solid-state reactions of the active phase with the carrier or
promoters. In addition, the presence of oxygen or chlorine in the feed gas can lead to formation of
volatile oxides or chlorides of the active phase, followed by gas-phase transport from the reactor.
Similarly, changes in the oxidation state of the active catalytic phase can be induced by the
presence of reactive gases in the feed.
Thus, the mechanisms of solid catalyst deactivation are many; nevertheless, they can be grouped
into six intrinsic mechanisms of catalyst decay: (1) poisoning, (2) fouling, (3) thermal degradation,
(4) vapor compound formation and/or leaching accompanied by transport from the catalyst surface
or particle, (5) vapor–solid and/or solid–solid reactions, and (6) attrition/crushing. As mechanisms
1, 4, and 5 are chemical in nature while 2 and 6 are mechanical, the causes of deactivation are
basically threefold: chemical, mechanical, and thermal. Each of the six basic mechanisms is
defined briefly in Table 1 and treated in some detail in the subsections that follow, with an
emphasis on the first three. Mechanisms 4 and 5 are treated together, since 4 is a subset of 5.
2.1. Poisoning
Poisoning [3,16–22] is the strong chemisorption of reactants, products, or impurities on sites
otherwise available for catalysis. Thus, poisoning has operational meaning; that is, whether a
species acts as a poison depends upon its adsorption strength relative to the other species
competing for catalytic sites. For example, oxygen can be a reactant in partial oxidation of ethylene
to ethylene oxide on a silver catalyst and a poison in hydrogenation of ethylene on nickel. In
addition to physically blocking of adsorption sites, adsorbed poisons may induce changes in the
electronic or geometric structure of the surface [17,21]. Finally, poisoning may be reversible or

3
irreversible. An example of reversible poisoning is the deactivation of acid sites in fluid catalytic
cracking catalysts by nitrogen compounds in the feed. Although the effects can be severe, they are
temporary and are generally eliminated within a few hours to days after the nitrogen source is
removed from the feed. Similar effects have been observed for nitrogen compound (e.g., ammonia
and cyanide) addition to the syngas of cobalt Fischer-Tropsch catalysts, although these surface
species require weeks to months before the lost activity is regained [23]. However, most poisons
are irreversibly chemisorbed to the catalytic surface sites, as is the case for sulfur on most metals,
as discussed in detail below. Regardless of whether the poisoning is reversible or irreversible, the
deactivation effects while the poison is adsorbed on the surface are the same.
Table 1. Mechanisms of catalyst deactivation.
Mechanism

Type

Poisoning

Chemical

Fouling

Mechanical

Thermal degradation
and sintering

Thermal
Thermal/chemical

Vapor formation

Chemical

Vapor–solid and
solid–solid reactions

Chemical

Attrition/crushing

Mechanical

Brief definition/description

Strong chemisorption of species on catalytic sites which
block sites for catalytic reaction
Physical deposition of species from fluid phase onto the
catalytic surface and in catalyst pores
Thermally induced loss of catalytic surface area, support
area, and active phase-support reactions
Reaction of gas with catalyst phase to produce volatile
compound
Reaction of vapor, support, or promoter with catalytic
phase to produce inactive phase
Loss of catalytic material due to abrasion; loss of internal
surface area due to mechanical-induced crushing of the
catalyst particle

Many poisons occur naturally in feed streams that are treated in catalytic processes. For
example, crude oil contains sulfur and metals, such as vanadium and nickel, that act as catalyst
poisons for many petroleum refinery processes, especially those that use precious metal catalysts,
like catalytic reforming, and those that treat heavier hydrocarbon fractions in which the sulfur
concentrates and metals are almost exclusively found, such as fluid catalytic cracking and residuum
hydroprocessing. Coal contains numerous potential poisons, again including sulfur and others like
arsenic, phosphorous, and selenium, often concentrated in the ash, that can poison selective
catalytic reduction catalysts as discussed later in Section 4.3.3.1. As a final example, some poisons
may be added purposefully, either to moderate the activity and/or to alter the selectivity of fresh
catalysts, as discussed as the end of this section, or to improve the performance of a product that is
later reprocessed catalytically. An example of this latter case is lubricating oils that contain
additives like zinc and phosphorous to improve their lubricating properties and stability, which
become poisons when the lubricants are reprocessed in a hydrotreater or a fluid catalytic
cracking unit.
Mechanisms by which a poison may affect catalytic activity are multifold, as illustrated by a
conceptual two-dimensional model of sulfur poisoning of ethylene hydrogenation on a metal
surface shown in Figure 1. To begin with, a strongly adsorbed atom of sulfur physically blocks at

4
least one three- or fourfold adsorption/reaction site (projecting into three dimensions) and three or
four topside sites on the metal surface. Second, by virtue of its strong chemical bond, it
electronically modifies its nearest neighbor metal atoms and possibly its next-nearest neighbor
atoms, thereby modifying their abilities to adsorb and/or dissociate reactant molecules (in this case
H2 and ethylene molecules), although these effects do not extend beyond about 5 atomic units [21].
A third effect may be the restructuring of the surface by the strongly adsorbed poison, possibly
causing dramatic changes in catalytic properties, especially for reactions sensitive to surface
structure. In addition, the adsorbed poison blocks access of adsorbed reactants to each other
(a fourth effect) and finally prevents or slows the surface diffusion of adsorbed reactants (effect
number five).

Figure 1. Conceptual model of poisoning by sulfur atoms of a metal surface during
ethylene hydrogenation. Reproduced from [8]. Copyright 2006, Wiley-Interscience.
Catalyst poisons can be classified according to their chemical makeup, selectivity for active
sites, and the types of reactions poisoned. Table 2 lists four groups of catalyst poisons classified
according to chemical origin and their type of interaction with metals. It should be emphasized that
interactions of Group VA–VIIA elements with catalytic metal phases depend on the oxidation state
of the former, e.g., how many electron pairs are available for bonding and the degree of shielding
of the sulfur ion by ligands [16]. Thus, the order of decreasing toxicity for poisoning of a given
metal by different sulfur species is H2S, SO2, SO42í, i.e., in the order of increased shielding by
oxygen. Toxicity also increases with increasing atomic or molecular size and electronegativity, but
decreases if the poison can be gasified by O2, H2O, or H2 present in the reactant stream [21]; for
example, adsorbed carbon can be gasified by O2 to CO or CO2 or by H2 to CH4.
Table 2. Common poisons classified according to chemical structure.
Chemical type

Examples

Groups VA and VIA

N, P, As, Sb, O, S, Se, Te

Group VIIA

F, Cl, Br, I

Toxic heavy metals and ions
Molecules that adsorb with
multiple bonds

As, Pb, Hg, Bi, Sn, Cd,
Cu, Fe
CO, NO, HCN, benzene, acetylene,
other unsaturated hydrocarbons

Type of interaction with metals

Through s and p orbitals; shielded
structures are less toxic
Through s and p orbitals; formation
of volatile halides
Occupy d orbitals; may form alloys
Chemisorption through multiple
bonds and back bonding

5
Table 3 lists a number of common poisons for selected catalysts in important representative
reactions. It is apparent that organic bases (e.g., amines) and ammonia are common poisons for
acidic solids, such as silica–aluminas and zeolites in cracking and hydrocracking reactions, while
sulfur- and arsenic-containing compounds are typical poisons for metals in hydrogenation,
dehydrogenation, and steam reforming reactions. Metal compounds (e.g., of Ni, Pb, V, and Zn) are
poisons in automotive emissions control, catalytic cracking, and hydrotreating. Acetylene is a
poison for ethylene oxidation, while asphaltenes are poisons in hydrotreating of petroleum residuum.
Table 3. Poisons for selected catalysts in important representative reactions.
Catalyst

Reaction

Silica–alumina, zeolites

Cracking

Nickel, platinum, palladium

Hydrogenation/dehydrogenation

Iron, ruthenium

Steam reforming of
methane, naphtha
Ammonia synthesis

Cobalt, iron

Fischer–Tropsch synthesis

Noble metals on zeolites

Hydrocracking
Ethylene oxidation to
ethylene oxide
Oxidation/selective
catalytic reduction
Oxidation of CO and hydrocarbons

Nickel

Silver
Vanadium oxide
Platinum, palladium
Cobalt and
molybdenum sulfides

Hydrotreating of residuum

Poisons

Organic bases, hydrocarbons,
heavy metals
Compounds of S, P, As, Zn, Hg,
halides, Pb, NH3, C2H2
H2S, As
O2, H2O, CO, S, C2H2, H2O
H2S, COS, As, NH3, metal
carbonyls
NH3, S, Se, Te, P
C2H2
As/Fe, K, Na from fly ash
Pb, P, Zn, SO2, Fe
Asphaltenes, N compounds, Ni, V

Poisoning selectivity is illustrated in Figure 2, a plot of activity (the reaction rate normalized to
initial rate) versus normalized poison concentration. “Selective” poisoning involves preferential
adsorption of the poison on the most active sites at low concentrations. If sites of lesser activity are
blocked initially, the poisoning is “antiselective”. If the activity loss is proportional to the
concentration of adsorbed poison, the poisoning is “nonselective.” An example of selective
poisoning is the deactivation of platinum by CO for the para-H2 conversion (Figure 3a) [24] while
Pb poisoning of CO oxidation on platinum is apparently antiselective (Figure 3b) [25], and arsenic
poisoning of cyclopropane hydrogenation on Pt is nonselective (Figure 3c) [26]. For nonselective
poisoning, the linear decrease in activity with poison concentration or susceptibility (ı) is defined
by the slope of the activity versus poison concentration curve. Several other important terms
associated with poisoning are defined in Table 4. Poison tolerance, the activity at saturation
coverage of the poison, and resistance (the inverse of deactivation rate) are important concepts that
are often encountered in discussions of poisoning including those below.

6

Figure 2. Three kinds of poisoning behavior in terms of normalized activity
versus normalized poison concentration. Reproduced from [8]. Copyright 2006,
Wiley-Interscience.
Table 4. Important Poisoning Parameters.
Parameter

Definition

Activity (a)

Reaction rate at time t relative to that at t = 0
Negative slope of the activity versus poison concentration curve [ı = (a í 1)/C (t)].
Measure of a catalyst’s sensitivity to a given poison
Susceptibility of a given catalyst for a poison relative to that for another poison
Inverse of the deactivation rate. Property that determines how rapidly a catalyst
deactivates
Activity of the catalyst at saturation coverage (some catalysts may have negligible
activity at saturation coverage)

Susceptibility (ı)
Toxicity
Resistance
Tolerance (a(Csat))

The activity versus poison concentration patterns illustrated in Figure 2 are based on the
assumption of uniform poisoning of the catalyst surface and surface reaction rate controlling, i.e.,
negligible pore-diffusional resistance. These assumptions, however, are rarely met in typical
industrial processes because the severe reaction conditions of high temperature and high pressure
bring about a high pore-diffusional resistance for either the main or poisoning reaction or both. In
physical terms, this means that the reaction may occur preferentially in the outer shell of the
catalysts particle, or that poison is preferentially adsorbed in the outer shell of the catalyst particle,
or both. The nonuniformly distributed reaction and/or poison leads to nonlinear activity versus
poison concentration curves that mimic the patterns in Figure 2 but do not represent truly selective
or antiselective poisoning. For example, if the main reaction is limited to an outer shell in a pellet

7
where poison is concentrated, the drop in activity with concentration will be precipitous. Pore
diffusional effects in poisoning (nonuniform poison) are treated later in this review.

Figure 3. (a) CO poisoning of para-H2 conversion over a Pt foil, reproduced from [24],
copyright 1974, Wiley-VHC; (b) effect of lead coverage on the rate of CO oxidation of
Pt film, reproduced from [25], copyright 1978, Elsevier; (c) rate constants of
cyclopropane hydrogenolysis over a Pt film as a function of the amount of AsH3
adsorbed, reproduced from [26], copyright 1970, Elsevier.
As sulfur poisoning is a difficult problem in many important catalytic processes (e.g.,
hydrogenation, methanation, Fischer–Tropsch synthesis, steam reforming, and fuel cell power
production), it merits separate discussion as an example of catalyst poisoning phenomena. Studies of
sulfur poisoning in hydrogenation and CO hydrogenation reactions have been thoroughly
reviewed [8,21,27–31]. Much of the previous work focused on poisoning of nickel metal catalysts
by H2S, the primary sulfur poison in many important catalytic processes, and thus provides some
useful case studies of poisoning.
Previous adsorption studies [28–30] indicate that H2S adsorbs strongly and dissociatively on
nickel metal surfaces. The high stability and low reversibility of adsorbed sulfur is illustrated by the
data in Figure 4 [28], in which most of the previous equilibrium data for nickel are represented on a

8
single plot of log (PH2S/PH2) versus reciprocal temperature. The solid line corresponds to the
equilibrium data for formation of bulk Ni3S2. Based on the equation ǻG = RT ln(PH2S/PH2) = ǻH í
TǻS, the slope of this line is ǻH/R, where ǻH = í75 kJ/mol and the intercept is íǻS/R. Most of the
adsorption data lie between the dashed lines corresponding to ǻH = í125 and í165 kJ/mol for
coverages ranging from 0.5 to 0.9, indicating that adsorbed sulfur is more stable than the bulk
sulfide. Indeed, extrapolation of high temperature data to zero coverage using a Tempkin
isotherm [29] yields an enthalpy of adsorption of í250 kJ/mol; in other words, at low sulfur
coverages, surface nickel–sulfur bonds are a factor of 3 more stable than bulk nickel–sulfur bonds.
It is apparent from Figure 4 that the absolute heat of adsorption increases with decreasing coverage
and that the equilibrium partial pressure of H2S increases with increasing temperature and
increasing coverage. For instance, at 725 K (450 °C) and ș = 0.5, the values of PH2S/PH2 range from
about 10í8 to 10í9. In other words, half coverage occurs at 1–10 ppb H2S, a concentration range at
the lower limit of our present analytical capability. At the same temperature (450 °C), almost
complete coverage (ș > 0.9) occurs at values of PH2S/PH2 of 10í7–10í6 (0.1–1 ppm) or at H2S
concentrations encountered in many catalytic processes after the gas has been processed to remove
sulfur compounds. These data are typical of sulfur adsorption on most catalytic metals. Thus, we
can expect that H2S (and other sulfur impurities) will adsorb essentially irreversibly to high
coverage in most catalytic processes involving metal catalysts.
Two important keys to reaching a deeper understanding of poisoning phenomena include
(1) determining surface structures of poisons adsorbed on metal surfaces and (2) understanding
how surface structure and hence adsorption stoichiometry change with increasing coverage of the
poison. Studies of structures of adsorbed sulfur on single crystal metals (especially Ni) [3,28,32–38]
provide such information. They reveal, for example, that sulfur adsorbs on Ni(100) in an ordered
p(2 × 2) overlayer, bonded to four Ni atoms at S/Nis < 0.25 and in a c(2 × 2) overlayer to two Ni
atoms for S/Nis = 0.25–0.50 (see Figure 5; Nis denotes a surface atom of Ni); saturation coverage of
sulfur on Ni(100) occurs at S/Nis = 0.5. Adsorption of sulfur on Ni(110), Ni(111), and higher index
planes of Ni is more complicated; while the same p(2 × 2) structure is observed at low coverage,
complex overlayers appear at higher coverages—for example, at S/Nis > 0.3 on Ni(111) a
(5 3 × 2)S overlayer is formed [32–34]. In more open surface structures, such as Ni(110) and
Ni(210), saturation coverage occurs at S/Nis = 0.74 and 1.09 respectively; indeed, there is a trend of
increasing S/Nis with decreasing planar density and increasing surface roughness for Ni, while the
saturation sulfur concentration remains constant at 44 ng/cm2 Ni (see Table 5).
Reported saturation stoichiometries for sulfur adsorption on polycrystalline and supported
Ni catalysts (S/Nis) vary from 0.25 to 1.3 [28]. The values of saturation coverage greater than
S/Nis = 0.5 may be explained by (1) a higher fractional coverage of sites of lower coordination
number, i.e., atoms located on edges or corners of rough, high-index planes (Table 5); (2) enhanced
adsorption capacity at higher gas phase concentrations of H2S in line with the observed trend of
increasing saturation coverage with increasing H2S concentration in Figure 4; and/or (3)
reconstruction of planar surfaces to rougher planes by adsorbed sulfur at moderately high
coverages and adsorption temperatures.

9

Figure 4. Equilibrium partial pressure of H2S versus reciprocal temperature (values of
ǻHf based on 1 mole of H2S); open symbols: ș = 0.5–0.6; closed symbols: ș = 0.8–0.9.
Reproduced from [28]. Copyright 1982, Academic Press.

Figure 5. Schematic view of sulfur adsorbed on a Ni(100) surface at a (a) S/Nis = 0.25
in a p(2 × 2) structure and (b) S/Nis = 0.50 in a c(2 × 2) structure. Reproduced
from [39]. Copyright 2001, Elsevier.

10
Table 5. Sulfur Adsorption Densities on Various Crystal Faces of Nickel a.
Crystal face

Sulfur conc. at
saturation, ng·S/cm2

Number of S
atoms/cm2 (×1015)

Number of Ni
atoms/cm2
(×1015)

S atoms per surface
Ni atom

(111)
(100)
(110)
(210)
Polycrystalline

47 ± 1
43 ± 1
44.5 ± 1
42 ± 1
44.5 ± 1

0.86
0.80
0.82
0.78
0.82

1.8
1.6
1.1
0.72
—

0.48
0.50
0.74
1.09
—

a

Data from [35].

The first effect would be favored, and in fact is observed, for supported metals of higher
dispersion [28]. The second effect may explain the typically lower observed values of S/Nis for
single crystal Ni, which are measured at extremely low pressures (high vacuum) relative to the
higher values of S/Nis for polycrystalline and supported Ni, typically measured at orders of
magnitude higher pressure; thus, in the case of the single crystal studies, the surface is not in
equilibrium with gas phase H2S/H2.
The third effect, reconstruction of nickel surfaces by adsorbed sulfur, has been reported by a
number of workers [28,32,33,36–38]; for example, McCarroll and co-workers [37,38] found that
sulfur adsorbed at near saturation coverage on a Ni(111) face was initially in a hexagonal pattern,
but upon heating above 700 K reoriented to a distorted c(2 × 2) (100) overlayer. Oudar [36]
reported that sulfur adsorbed on a Ni(810) surface caused decomposition to (100) and (410) facets.
During adsorption of H2S at RT, Ruan et al. [33] observed surface restructuring of Ni(111) from a
p(2 × 2) at low coverage to a missing-row (5 3 × 2)S terrace structure (0.4 monolayer) sparsely
covered with small, irregular islands composed of sulfur adsorbed on disordered nickel; upon
annealing to 460 K for 5 min, the islands ordered to the (5 3 × 2)S phase and their size increased,
suggesting further diffusion of Ni atoms from the terraces. The reconstruction of Ni (111)
involving ejection and migration of Ni atoms was attributed to compressive surface stresses
induced by sulfur adsorption; the role of compressive surface stress due to sulfur coverages
exceeding 0.3 was confirmed by Grossmann et al. [32]. From these and similar studies, it is
concluded that at moderately high temperatures (300–600 K) and coverages greater than 0.3,
restructuring by sulfur of different facets of Ni to rougher, more open, stable structures is probably
a general phenomenon. Thus, reconstruction probably accounts at least in part for observed
increases in saturation S coverage with decreasing Ni site density.
The nature of reconstruction of a surface by a poison may depend on its pretreatment. For
example, in a scanning tunneling microscopy (STM) study of room temperature H2S adsorption on
Ni(110), Ruan and co-workers [40] found that the S/Ni structure at saturation varied with the initial
state of the surface, i.e., whether clean or oxygen covered. Beginning with a clean Ni(110) surface,
oxygen adsorbs dissociatively to form a (2 × 1)O overlayer at 1/2 monolayer coverage (Figure 6a);
this is accompanied by a homogeneous nucleation of low-coordinated -Ni-O- rows along the [001]
direction. As the oxygen-covered surface is exposed stepwise to 3 and then 8 Langmuirs (L) of
H2S, oxygen atoms are removed by reaction with hydrogen to water; the surface is first roughened,

11
after which white islands and black troughs having a c(2 × 2) structure are formed as sulfur atoms
replace oxygen atoms (Figure 6b). Upon exposure to 25 L of H2S, the c(2 × 2) islands dissolve,
while low-coordinated rows (periodicity of 1) form in the [001] direction, developing into ordered
regions with a periodicity of 4 in the [ͳ Ͳ] direction (Figure 6c). After exposure to 50 L of H2S
(Figure 6d), a stable, well-ordered (4 × 1)S structure appears, a surface clearly reconstructed
relative to the original Ni(110). Moreover, the reconstructed surface in Figure 6d is very different
from that observed upon direct exposure of the Ni(110) to H2S at room temperature, i.e., a c(2 ×
2)S overlying the original Ni(110) (similar to Figure 5b); in other words, it appears that no
reconstruction occurs by direct exposure to H2S at room temperature, rather only in the presence of
O2 (or air). This emphasizes the complexities inherent in predicting the structure and stability of a
given poison adsorbed on a given catalyst during a specified reaction as a function of different
pretreatments or process disruptions, e.g., exposure to air.
In the previous discussion of Figure 4, íǻHads was observed to decrease with increasing sulfur
coverage; data in Figure 7 [41] show that –ǻHads decreases with increasing gas-phase H2S
concentration and coverage. However, in contrast to the data in Figure 4, those in Figure 7 [41]
show that at very high H2S concentrations and high adsorption temperatures, íǻHads falls well
below the íǻHformation of bulk Ni3S2; at the same time, the S/Nis ratio approaches that of Ni2S3. This
is a unique result, since all of the data obtained at lower temperatures and H2S concentrations [28]
show íǻHads to be greater than íǻHformation of Ni3S2.
From the above discussion, the structure and stoichiometry of sulfur adsorbed on nickel
evidently are complex functions of temperature, H2S concentration, sulfur coverage, and
pretreatment, phenomena that account at least in part for the complex nature of nickel poisoning by
sulfur. Could one expect similar complexities in the poisoning of other metals? Probably, since
poisoning of nickel is prototypical, i.e., similar principles operate and similar poisoning behaviors
are observed in other poison/metal systems, although none have been studied to the same depth
as sulfur/nickel.
Since one of the necessary steps in a catalytic reaction is the adsorption of one or more
reactants, investigation of the effects of adsorbed sulfur on the adsorption of other molecules, can
provide useful insights into the poisoning process [21,28]. Previous investigations [28,42–48]
indicate that both H2 and CO adsorptions on nickel are poisoned by adsorbed sulfur. For example,
thermal desorption studies of CO from presulfided Ni(100) [44] reveal a weakening of the CO
adsorption bond and a rapid, nonlinear decline in the most strongly bound ȕ2 state (bridged CO)
with increasing sulfur coverage, corresponding to a poisoning of about 8–10 Ni atoms for bridged
CO adsorption per adsorbed sulfur atom at low sulfur coverage (see Figure 8); moreover, the ȕ2 CO
species is completely poisoned at about 0.2–0.4 mL of sulfur relative to a saturation coverage of
0.5 mL. Hydrogen adsorption is poisoned in a similar nonlinear fashion. On the other hand, the
coverage of the ȕ1 state (linear CO) is constant with increasing sulfur coverage. The sharp
nonlinear drop in CO and hydrogen adsorptions at low sulfur coverages has been interpreted in
terms of a combination of short-range electronic and steric effects operating over a range of less
than 5 atomic units [13]. The different effects of sulfur on ȕ1 and ȕ2 states of CO have important

12
implications for sulfur poisoning in reactions involving CO; that is, sulfur poisoning can affect
reaction selectivity as well as activity [28].

Figure 6. A series of in situ scanning tunneling microscope (STM) images recorded
after exposure of Ni(110) to oxygen and then progressively higher exposures of H2S:
(a) (2 × 1)O overlayer; (b) white islands and black troughs with a c(2 × 2)S structure
after exposure to 3 and 8 L of H2S; (c) 25 L, islands transform to low-coordinated
rows in the [001] direction; and (d) 50 L, stable, well-ordered (4 × 1)S. Reproduced
from [40]. Copyright 1992, American Physical Society.

Figure 7. Sulfur chemisorption isosteres on a Ni/Į-Al2O3 catalyst at high temperatures
and high H2S concentrations. Reproduced from [41]. Copyright 1999,Elsevier.

13
Because sulfur adsorbs so strongly on metals and prevents or modifies the further adsorption of
reactant molecules, its presence on a catalyst surface usually effects substantial or complete loss
of activity in many important reactions. This is illustrated by the data in Figure 9 showing the
steady-state methanation activities of Ni, Co, Fe, and Ru relative to the fresh, unpoisoned surface
activity as a function of gas phase H2S concentration. These data indicate that Ni, Co, Fe, and Ru
all suffer 3–4 orders of magnitude loss in activity at 15–100 ppb of H2S, i.e., their sulfur tolerances
are extremely low. Moreover, the sharp drop in activity with increasing H2S concentration suggests
highly selective poisoning. Nevertheless, the rate of sulfur poisoning and hence sulfur resistance
varies from catalyst to catalyst and is apparently a function of catalyst composition [28] and
reaction conditions [49]. Indeed, it is possible to significantly improve sulfur resistance of Ni, Co,
and Fe with catalyst additives such as Mo and B that selectively adsorb sulfur. Because the
adsorption of sulfur compounds is generally rapid and irreversible, surface sulfur concentrations in
catalyst particles and beds are nonuniform, e.g., H2S adsorbs selectively at the entrance to a packed
bed and on the outer surface of catalyst particles, making the experimental study and modeling of
sulfur poisoning extremely difficult.
There are other complications in the study of sulfur poisoning. For example, the adsorption
stoichiometry of sulfur in CO hydrogenation on Ni is apparently a function of the temperature,
H2/CO ratio, and water partial pressure [49]. Moreover, at high CO partial pressures sulfur may be
removed from the surface as COS, which is not as strongly adsorbed as H2S. At low temperature
conditions, e.g., those representative of Fischer–Tropsch synthesis or liquid phase hydrogenations,
the gas phase concentration of H2S in poisoning studies must be kept very low, i.e., below 0.1–5
ppm, to avoid formation of bulk metal sulfides—a phenomenon that seriously compromises the
validity of the results. Thus, the importance of studying poisoning phenomena in situ under realistic
reaction conditions, at low process-relevant poison concentrations, and over a process-representative
range of temperature and concentration conditions is emphasized.
As mentioned earlier, there are a number of industrial processes in which one intentionally
poisons the catalyst in order to improve its selectivity. For example, Pt-containing naphtha reforming
catalysts are often pre-sulfided to minimize unwanted cracking reactions. On basic Pt/KL zeolite
catalysts, these short term, low concentration exposures are beneficial to produce Pt ensemble sizes
that promote aromatization, while longer term or higher concentration exposures poison the catalyst
both by forming Pt-S bonds and producing large crystallites that block pores, as shown by
transmission electron microscopy (TEM) and X-ray absorption fine structure spectroscopy (EXAFS),
and favor only dehydrogenation [50–53]. Other examples are sulfur added to Fischer-Tropsch
catalysts that have been reported to have either beneficial or negligibly harmful effects, which are
important considerations in setting the minimum gas clean-up requirements [27,30,54–56]. S and P
are added to Ni catalysts to improve isomerization selectivity in the fats and oils hydrogenation
industry, while S and Cu are added to Ni catalysts in steam reforming to minimize coking. In
catalytic reforming, sulfided Re or Sn is added to Pt to enhance the dehydrogenation of paraffins to
olefins while poisoning hydrogenolysis/coking reactions. V2O5 is added to Pt to suppress SO2
oxidation to SO3 in diesel emissions control catalysts.

14

Figure 8. Area under thermal programmed desorption spectra for H2 and the Į, ȕ1, ȕ2,
and total CO adsorption curves, as a function of sulfur precoverage. Reproduced
from [44]. Copyright 1981, Elsevier.

Figure 9. Relative steady-state methanation activity profiles for Ni (Ɣ), Co (ǻ), Fe (ż),
and Ru (Ƒ) as a function of gas-phase H2S concentration. Reaction conditions: 100 kPa,
400 °C, 1% CO/99% H2 for Co, Fe, and Ru, 4% CO/96% H2 for Ni. Reproduced
from [28]. Copyright 1982, Academic Press.

15
2.2. Fouling, Coking, and Carbon Deposition
2.2.1. Fouling
Fouling is the physical (mechanical) deposition of species from the fluid phase onto the catalyst
surface, which results in activity loss due to blockage of sites and/or pores. In its advanced stages, it
may result in disintegration of catalyst particles and plugging of the reactor voids. Important examples
include mechanical deposits of carbon and coke in porous catalysts, although carbon- and coke-forming
processes also involve chemisorption of different kinds of carbons or condensed hydrocarbons that
may act as catalyst poisons. The definitions of carbon and coke are somewhat arbitrary and by
convention related to their origin. Carbon is typically a product of CO disproportionation while
coke is produced by decomposition or condensation of hydrocarbons on catalyst surfaces and typically
consists of polymerized heavy hydrocarbons. Nevertheless, coke forms may vary from high molecular
weight hydrocarbons to primarily carbons such as graphite, depending upon the conditions under
which the coke was formed and aged. A number of books and reviews treat the formation of
carbons and coke on catalysts and the attendant deactivation of the catalysts [1,4,57–62].
The chemical structures of cokes or carbons formed in catalytic processes vary with reaction
type, catalyst type, and reaction conditions. Menon [62] suggested that catalytic reactions
accompanied by carbon or coke formation can be broadly classified as either coke-sensitive or
coke-insensitive, analogous to Boudart’s more general classification of structure-sensitive and
structure-insensitive catalytic reactions. In coke-sensitive reactions, unreactive coke is deposited on
active sites, leading to activity decline, while in coke-insensitive reactions, relatively reactive coke
precursors formed on active sites are readily removed by hydrogen (or other gasifying agents).
Examples of coke-sensitive reactions include catalytic cracking and hydrogenolysis; on the other
hand, Fischer–Tropsch synthesis, catalytic reforming, and methanol synthesis are examples of
coke-insensitive reactions. On the basis of this classification, Menon [62] reasoned that the
structure and location of a coke are more important than its quantity in affecting catalytic activity.
Consistent with Menon’s classification, it is also generally observed that not only structure and
location of coke vary but also its mechanism of formation varies with catalyst type, e.g., whether it
is a metal or metal oxide (or sulfide, sulfides being similar to oxides). Because of these significant
differences in mechanism, formation of carbon and coke is discussed below separately for
supported metals and for metal oxides and sulfides.
2.2.2. Carbon and Coke Formation on Supported Metal Catalysts
Possible effects of fouling by carbon (or coke) on the functioning of a supported metal catalyst
are illustrated in Figure 10. Carbon may (1) chemisorb strongly as a monolayer or physically
adsorb in multilayers and in either case block access of reactants to metal surface sites, (2) totally
encapsulate a metal particle and thereby completely deactivate that particle, and (3) plug microand mesopores such that access of reactants is denied to many crystallites inside these pores.
Finally, in extreme cases, strong carbon filaments may build up in pores to the extent that they
stress and fracture the support material, ultimately causing the disintegration of catalyst pellets and

16
plugging of reactor voids. For example, in steam methane reforming (SMR) catalysts, which are
typically nickel supported on alumina with alkaline earth oxides, the carbon can diffuse through and
begin to grow filaments from the back side of the nickel particles (structural type 3 in Table 6)
especially at high reaction temperatures and low steam to methane ratios, which push the nickel
particles off the support surface. Thermal or mechanical shock can then cause the carbon filaments to
fall off the support, thus permanently deactivating the catalyst [8,60]. However, the behavior is
complex because for other reaction conditions and other metals, the filaments may grow from the
top surface of the metal particles or the carbon may diffuse into the metal and form bulk
carbides [8].
An example of recent interest for biomass reactions that points to the complex interaction
between the active metal and the support during carbon deposition is the steam reforming of light
alcohols and other oxygenates, in which deactivation occurs primarily through coking. For
traditional SMR catalysts (e.g., Ni/MgAl2O4) the coke is believed to originate primarily from
alkene formation [63,64]. However, for the case of Ni/La2O3 catalysts, carbon appears to form at
the interface between the active metal and the support to block the active phase [65].
Mechanisms of carbon deposition and coke formation on metal catalysts from carbon monoxide
and hydrocarbons, including methane during SMR for hydrogen production [4,57–61], are
illustrated in Figures 11 and 12. Different kinds of carbon and coke that vary in morphology and
reactivity are formed in these reactions (see Tables 6 and 7). For example, CO dissociates on
metals to form CĮ, an adsorbed atomic carbon; CĮ can react to Cȕ, a polymeric carbon film. The
more reactive, amorphous forms of carbon formed at low temperatures (e.g., CĮ and Cȕ) are
converted at high temperatures over a period of time to less reactive, graphitic forms [60]

Figure 10. Conceptual model of fouling, crystallite encapsulation, and pore plugging of
a supported metal catalyst owing to carbon deposition. Reproduced from [8]. Copyright
2006, Wiley-Interscience.

17

Figure 11. Formation, transformation, and gasification of carbon on nickel (a, g, s refer
to adsorbed, gaseous, and solid states respectively). Reproduced from [60]. Copyright
1983, Elsevier.

Figure 12. Formation and transformation of coke on metal surfaces (a, g, s refer to
adsorbed, gaseous, and solid states respectively); gas phase reactions are not considered.
Reproduced from [60]. Copyright 1983, Elsevier.
Table 6. Forms and Reactivities of Carbon Species Formed by Decomposition of CO
on Nickel a.

1.
2.
3.
4.
5.

Structural type

Designation

Adsorbed, atomic (surface carbide)
Polymeric, amorphous films or filaments
Vermicular filaments, fibers, and/or whiskers
Nickel carbide (bulk)
Graphitic (crystalline) platelets or films

CĮ
Cȕ
Cv
CȖ
Cc

a

Ref. [60].

Temp.
formed, °C
200–400
250–500
300–1000
150–250
500–550

Peak temp. for
reaction with H2, °C
200
400
400–600
275
550–850

18
Table 7. Carbon Species Formed in Steam Reforming of Hydrocarbons on Nickel Catalysts a.
Attribute

Encapsulating film

Whisker-like

Pyrolytic carbon

Formation

Slow polymerization of
CnHm radicals on Ni
surface, into
encapsulating film

Diffusion of C through Ni
crystal, nucleation and
whisker growth with Ni
crystal at top

Thermal cracking of
hydrocarbon; deposition
of C precursors
on catalyst

Effects

Progressive deactivation

No deactivation of Ni
surface. Breakdown of
catalyst and increasing ǻP

Encapsulation of catalyst
particle; deactivation and
increasing ǻP

Temp. range, °C

<500

>450

>600

Critical parameters

Low temperature, low
H2O/CnHm, low
H2/CnHm, aromatic feed

High temperature, low
H2O/CnHm, no enhanced
H2O adsorption, low
activity, aromatic feed

High temperature, high
void fraction, low
H2O/CnHm, high
pressure, acidic catalyst

a

Ref. [60].

It should also be emphasized that some forms of carbon result in loss of catalytic activity and
some do not. For example, at low reaction temperatures (<300–375 °C) condensed polymer or
ȕ-carbon films and at high temperatures (>650 °C) graphitic carbon films encapsulate the metal
surfaces of methanation and steam reforming catalysts [60]. Deactivation of steam reforming
catalysts at high reaction temperatures (500–900 °C) may be caused by precipitation of atomic
(carbidic) carbon dissolved in the Ni surface layers to a depth of more than 50–70 nm [62,66]. If it
accumulates on the metal surface (at high or low temperatures), adsorbed atomic carbon can
deactivate metal sites for adsorption and/or reaction. For example, Durer and co-workers [67]
demonstrated that carbon atoms residing in the fourfold hollow sites of Rh(100) block the
adsorption of hydrogen (and hence could block sites for hydrogenation). In the intermediate
temperature range of 375–650 °C, carbon filaments (Figure 13) are formed by precipitation of
dissolved carbon at the rear side of metal crystallites, causing the metal particles to grow away
from the support [57]. Filament growth ceases when sufficient carbon accumulates on the free
surface to cause encapsulation by a carbon layer; however, encapsulation of the metal particles
does not occur if H2/CO or H2O/hydrocarbon ratios are sufficiently high. Thus, carbon filaments
sometimes formed in CO hydrogenation or steam reforming of hydrocarbons would not necessarily
cause a loss of intrinsic catalyst activity unless they are formed in sufficient quantities to cause
plugging of the pores [60] or loss of metal occurs as the carbon fibers are removed during
regeneration [68,69]. However, in practice, regions of carbon forming potential in steam reforming
must be carefully avoided, since once initiated, the rates of filamentous carbon formation are
sufficiently high to cause catastrophic pore plugging and catalyst failure within a few hours to days.

19

Figure 13. Electron micrograph of 14% Ni/Al2O3 having undergone extensive carbon
deposition during CO disproportionation at 673 K, PCO = 4.55 kPa (magnification of
200,000). Courtesy of the BYU Catalysis Laboratory.
The rate at which deactivation occurs for a given catalyst and reaction depends greatly on reaction
conditions—especially temperature and reactant composition. A fundamental principle for
coke-insensitive reactions on metals (e.g., methanation, Fischer–Tropsch synthesis, steam
reforming, catalytic reforming, and methanol synthesis) is that deactivation rate depends greatly on
the difference in rates of formation and gasification of carbon/coke precursors, i.e., rd = rf – rg. If
the rate of gasification, rg, is equal to or greater than that of formation, rf, carbon/coke is not deposited.
Rates of carbon/coke precursor formation and gasification both increase exponentially with
temperature, although the difference between them varies a great deal with temperature because of
differences in preexponential factors and activation energies. Thus, carbon/coke formation is
avoided in regions of temperature in which precursor gasification rate exceeds deposition rate. This
is illustrated in Figure 14, an Arrhenius plot for rates of formation and hydrogenation of alpha and
beta carbons on nickel during CO methanation. Since at temperatures below 600 K (1/T > 1.66 ×
10í3 Kí1) the rate of CĮ gasification exceeds that of CĮ formation, no carbon is deposited. However
above 600 K, CĮ accumulates on the surface since the rate of CĮ formation exceeds that of CĮ
gasification. As CĮ accumulates (at 600–700 K), it is converted to a Cȕ polymeric chain or film that
deactivates the nickel catalyst; however, above 700 K (1/T < 1.43 × 10í3 Kí1), the rate of Cȕ
hydrogenation exceeds that of formation and no deactivation occurs. Thus, the “safe” regions of
methanation for avoiding deactivation by carbon are below 600 K and above 700 K; of course, these
regions will vary somewhat with reactant concentrations and catalyst activity. A similar principle
operates in steam reforming, i.e., at a sufficiently low reaction temperature, the rate of hydrocarbon
adsorption exceeds the rate of hydrocracking and a deactivating polymer film is formed [70];
accordingly, it is necessary to operate above this temperature to avoid deactivation.

20

Figure 14. Rates of formation (log scale) and hydrogenation of CĮ and Cȕ versus
reciprocal temperature. Reproduced from [60]. Copyright 1983, Elsevier.
In steam reforming, filamentous carbon formation rate is a strong function of reactant
hydrocarbon structure; for example, it decreases in the order acetylenes, olefins, paraffins, i.e., in
order of decreasing reactivity, although activation energies for nickel are in the same range
(125–139 kJ) independent of hydrocarbon structure and about the same as those observed for
formation of filamentous carbon from decomposition of CO [60]. This latter observation suggests
that the reactions of CO and different hydrocarbons to filamentous carbon proceed by a common
mechanism and rate-determining step—probably the diffusion of carbon through the metal
crystallites [60].
The rate at which a carbon or coke is accumulated in a given reaction under given conditions can
vary significantly with catalyst structure, including metal type, metal crystallite size, promoter, and
catalyst support. For example, supported Co, Fe, and Ni are active above 350–400 °C for filamentous
carbon formation from CO and hydrocarbons; the order of decreasing activity is reportedly Fe > Co >
Ni [60]. Pt, Ru, and Rh catalysts, on the other hand, while equally or more active than Ni, Co, or Fe
in steam reforming, produce little or no coke or carbon. This is attributed to reduced mobility and/or
solubility of carbon in the noble metals, thus retarding the nucleation process. Thus, it is not surprising
that addition of noble metals to base metals retards carbon formation; for example, addition of Pt in Ni
lowers carbon deposition rate during methanation, while addition of Cu or Au to Ni substantially
lowers carbon formation in steam reforming [60,71]. In contrast to the moderating effects of noble
metal additives, addition of 0.5% Sn to cobalt substantially increases the rate of carbon filament
formation from ethylene [72], an effect desirable in the commercial production of carbon
filament fibers.
Since carbon formation and gasification rates are influenced differently by modifications in
metal crystallite surface chemistry, which are in turn a function of catalyst structure, oxide
additives or oxide supports may be used to moderate the rate of undesirable carbon or coke
accumulation. For example, Bartholomew and Strasburg [73] found the specific rate (turnover
frequency) of filamentous carbon deposition on nickel during methanation at 350 °C to decrease in
the order Ni/TiO2 > NiAl2O3 > Ni/SiO2, while Vance and Bartholomew [74] observed CĮ

21
hydrogenation rates at 170 °C to decrease in this same order (the same as for methanation at
225 °C). This behavior was explained in terms of promotional or inhibiting effects due to
decoration of metal crystallites by the support, for example silica, inhibiting both CO dissociation
and carbon hydrogenation. This hypothesis is consistent with observations [75,76] that silica
evaporated on metal surfaces and supported metals inhibits formation of filamentous carbon.
Similarly Bitter and co-workers [77] observed rates of carbon formation in CO2/CH4 reforming to
decrease in the order Pt/Ȗ-Al2O3ĺPt/TiO2 > Pt/ZrO2; while 90% of the carbon deposited on the
support, the authors linked deactivation to carbon accumulated on the metal owing to an imbalance
between carbon formed by methane dissociation and oxidation by chemisorbed CO2. The rate of
formation of coke in steam reforming is delayed and occurs at lower rates in nickel catalysts
promoted with alkali or supported on basic MgO [78].
Since formation of coke, graphite, or filamentous carbon involves the formation of C-C bonds on
multiple atoms sites, one might expect that coke or carbon formation on metals is structure-sensitive,
i.e., sensitive to surface structure and metal crystallite size. Indeed, Bitter and co-workers [77] found
that catalysts containing larger Pt crystallites deactivate more rapidly than those containing
small crystallites. Moreover, a crystallite size effect, observed in steam reforming of methane on
nickel [60,78], appears to operate in the same direction, i.e., formation of filamentous carbon occurs
at lower rates in catalysts containing smaller metal crystallites.
In summary, deactivation of supported metals by carbon or coke may occur chemically, owing
to chemisorption or carbide formation, or physically and mechanically, owing to blocking of
surface sites, metal crystallite encapsulation, plugging of pores, and destruction of catalyst pellets by
carbon filaments. Blocking of catalytic sites by chemisorbed hydrocarbons, surface carbides, or
relatively reactive films is generally reversible in hydrogen, steam, CO2, or oxygen. Further details of
the thermodynamics, kinetics, and mechanisms of carbon and coke formation in methanation and steam
reforming reactions are available in reviews by Bartholomew [60] and Rostrup-Nielsen [70,78]. In
recent reviews addressing deactivation of Co catalysts by carbon during Fischer-Tropsch
synthesis [79,80], the same or similar carbon species, e.g., Į, ȕ, polymeric, and graphitic carbons,
are observed on Co surfaces as on Ni; moreover, poisoning or fouling of the Co surfaces with ȕ,
polymeric, and graphitic carbon layers are found to be major causes of deactivation.
2.2.3. Coke Formation on Metal Oxide and Sulfide Catalysts
In reactions involving hydrocarbons, coke may be formed in the gas phase and on both noncatalytic
and catalytic surfaces. Nevertheless, formation of coke on oxides and sulfides is principally a result of
cracking reactions involving coke precursors (typically olefins or aromatics) catalyzed by acid
sites [81,82]. Dehydrogenation and cyclization reactions of carbocation intermediates formed on
acid sites lead to aromatics, which react further to higher molecular weight polynuclear aromatics
that condense as coke (see Figure 15). Reactions 1–3 in Figure 15 illustrate the polymerization of
olefins, reactions 4–8 illustrate cyclization from olefins, and reactions 9–14 illustrate chain reaction
formation of polynuclear aromatics that condense as coke on the catalyst surface. Because of the
high stability of the polynuclear carbocations (formed in reactions 10–13), they can continue to

22
grow on the surface for a relatively long time before a termination reaction occurs through the back
donation of a proton.
From this mechanistic scheme (Figure 15), it is clear that olefins, benzene and benzene
derivatives, and polynuclear aromatics are precursors to coke formation. However, the order of
reactivity for coke formation is clearly structure dependent, i.e., decreases in the order polynuclear
aromatics > aromatics > olefins > branched alkanes > normal alkanes. For example, the weight
percent coke formed on silica–alumina at 500 °C is 0.06, 3.8, 12.5, and 23% for benzene,
naphthalene, fluoranthene, and anthracene respectively [83].
Coking reactions in processes involving heavy hydrocarbons are very complex; different kinds
of coke may be formed and they may range in composition from CH to C and have a wide range of
reactivities with oxygen and hydrogen, depending upon the time on stream and temperature to
which they are exposed. For example, coke deposits occurring in hydrodesulfurization of residuum
have been classified into three types [84]:
(1) Type I deposits are reversibly adsorbed normal aromatics deposited during the first part of
the cycle at low temperature.
(2) Type II deposits are reversibly adsorbed asphaltenes deposited early in the coking process.
(3) Type III deposits result from condensation of aromatic concentrates into clusters and then
crystals that constitute a “mesophase.” This crystalline phase is formed after long reaction
times at high temperature. This hardened coke causes severe deactivation of the catalyst [84].

Figure 15. Cont.

23

Figure 15. Coke-forming reactions of alkenes and aromatics on oxide and sulfide catalysts:
(a) polymerization of alkenes, (b) cyclization from alkenes, and (c) formation of polynuclear
aromatics from benzene. Reproduced from [8], Copyright 2006, Wiley-Interscience.

24
In addition to hydrocarbon structure and reaction conditions, extent and rate of coke formation
are also a function of the acidity and pore structure of the catalyst. Generally, the rate and extent of
coke formation increase with increasing acid strength and concentration. Coke yield decreases with
decreasing pore size (for a fixed acid strength and concentration); this is especially true in zeolites
where shape selectivity plays an important role in coke formation. For example, coke yield in fluid
catalytic cracking is only 0.4% for ZSM-5 (pore diameters of 0.54 × 0.56 nm) compared to 2.2%
for Y-faujasite (aperture diameter of 0.72 nm) [82]. However, in pores of molecular diameter, a
relatively small quantity of coke can cause substantial loss of activity. It should be emphasized that
coke yield can vary considerably into the interior pores of a catalyst particle or along a catalyst bed,
depending upon the extent to which the main and deactivation reactions are affected by film mass
transport and pore diffusional resistance.
The mechanisms by which coke deactivates oxide and sulfide catalysts are, as in the case of
supported metals, both chemical and physical. However, some aspects of the chemistry are quite
different. The principal chemical loss of activity in oxides and sulfides is due to the strong
adsorption of coke molecules on acidic sites. However, as discussed earlier, strong acid sites also
play an important role in the formation of coke precursors, which subsequently undergo
condensation reactions to produce large polynuclear aromatic molecules that physically coat
catalytic surfaces. Physical loss of activity also occurs as coke accumulates, ultimately partially or
completely blocking catalyst pores as in supported metal catalysts. For example, in isomerization
of cis-butene on SiO2/Al2O3 [85] catalyst deactivation occurs by rapid, selective poisoning of
strong acid sites; coke evolved early in the reaction is soluble in dichloromethane and pyridine and
is slightly aromatic. Apparently, the blocking of active sites does not significantly affect porosity or
catalyst surface area, as SiO2/Al2O3 contains relatively large mesopores.
In the case of supported bifunctional metal/metal oxide catalysts, different kinds of coke are
formed on the metal and the acidic oxide support, e.g., soft coke (high H/C ratio) on Pt or Pt–Re
metals and hard coke (low H/C ratio) on the alumina support in catalytic reforming [86]. In this
case, coke precursors may be formed on the metal via hydrogenolysis, following which they
migrate to the support and undergo polymerization and cyclization reactions, after which the larger
molecules are dehydrogenated on the metal and finally accumulate on the support, causing loss of
isomerization activity. Mild sulfiding of these catalysts (especially Pt–Re/alumina) substantially
reduces the rate of hydrogenolysis and the overall formation of coke on both metal and support; it
especially reduces the hard coke, which is mainly responsible for deactivation.
Several recent studies [82,87–97] have focused on coke formation during hydrocarbon reactions
in zeolites including (1) the detailed chemistry of coke precursors and coke molecules formed in
zeolite pores and pore intersections (or supercages) and (2) the relative importance of adsorption on
acid sites versus pore blockage. The principal conclusions from these studies can be summarized as
follows: (1) the formation of coke and the manner in which it deactivates a zeolite catalyst are
shape-selective processes, (2) deactivation is mainly due to the formation and retention of heavy
aromatic clusters in pores and pore intersections, and (3) while both acid-site poisoning and pore
blockage participate in the deactivation, the former dominates at low coking rates, low coke
coverages (e.g., in Y-zeolite below 2 wt%), and high temperatures, while the latter process

25
dominates at high reaction rates, high coke coverages, and low temperatures. Thus, pore size and pore
structure are probably more important than acid strength and density under typical commercial process
conditions. Indeed, deactivation is typically more rapid in zeolites having small pores or apertures
and/or a monodimensional structure [95]. Figure 16 illustrates four possible modes of deactivation of
HZSM-5 by carbonaceous deposits with increasing severity of coking [95].

Mode

Mode

Mode

Mode

Figure 16. Schematic of the four possible modes of deactivation by carbonaceous deposits
in HZSM-5: (1) reversible adsorption on acid sites, (2) irreversible adsorption on sites with
partial blocking of pore intersections, (3) partial steric blocking of pores, and (4) extensive
steric blocking of pores by exterior deposits. Adapted from [95].
These conclusions (in the previous paragraph) are borne out, for example, in the study by
Cerqueira and co-workers [97] of USHY zeolite deactivation during methylcyclohexane transformation
at 450 °C, showing the following:
(1) Coke is probably mainly formed by rapid transformation of toluenic C7 carbenium ions
with lesser contributions from reactions of cyclopentadiene, C3–C6 olefins, and aromatics.
(2) Soluble coke consists of polynuclear aromatic clusters containing three to seven five- and
six-membered rings having a typical compositions of C30H40 to C40H44 and having
dimensions of 0.9 × 1.1 nm to 1.1 × 1.5 nm, i.e., sizes that would cause them to be trapped
in the supercages of Y-zeolite.
(3) At short contact times, coking is relatively slow and deactivation is mainly due to acid-site
poisoning, while at long contact times, coking is much faster because of the high
concentrations of coke precursors; under these latter conditions coke is preferentially
deposited at the outer pore openings of zeolite crystallites and deactivation is dominated by
pore-mouth blockage.
That coke formed at large contact times not only blocks pores and/or pore intersections inside
the zeolite, but also migrates to the outside of zeolite crystallites, where it blocks pore entrances,
has been observed in several studies [91,93,94,97]. However, the amount, structure, and location of
coke in ZSM-5 depends strongly on the coke precursor, e.g., coke formed from mesitylene is

26
deposited on the external zeolite surface, whereas coking with isobutene leads to largely paraffinic
deposits inside pores; coke from toluene, on the other hand, is polyaromatic and is deposited both
on external and internal zeolite surfaces [91].
2.3. Thermal Degradation and Sintering
2.3.1. Background
Thermally induced deactivation of catalysts results from (1) loss of catalytic surface area due to
crystallite growth of the catalytic phase, (2) loss of support area due to support collapse and of catalytic
surface area due to pore collapse on crystallites of the active phase, and/or (3) chemical transformations
of catalytic phases to noncatalytic phases. The first two processes are typically referred to as
“sintering”. The third is discussed in the next section under solid–solid reactions. Sintering processes
generally take place at high reaction temperatures (e.g., > 500 °C) and are generally accelerated by the
presence of water vapor.
Most of the previous sintering and redispersion work has focused on supported metals.
Experimental and theoretical studies of sintering and redispersion of supported metals published
before 1997 have been reviewed fairly extensively [8,98–107]. Three principal mechanisms of
metal crystallite growth have been advanced: (1) crystallite migration, (2) atomic migration, and
(3) (at very high temperatures) vapor transport. The processes of crystallite and atomic migration
are illustrated in Figure 17. Crystallite migration involves the migration of entire crystallites over
the support surface, followed by collision and coalescence. Atomic migration involves detachment
of metal atoms or molecular metal clusters from crystallites, migration of these atoms over the
support surface, and ultimately, capture by larger crystallites. Redispersion, the reverse of
crystallite growth in the presence of O2 and/or Cl2, may involve (1) formation of volatile metal
oxide or metal chloride complexes that attach to the support and are subsequently decomposed to
small crystallites upon reduction and/or (2) formation of oxide particles or films that break into
small crystallites during subsequent reduction.

Figure 17. Two conceptual models for crystallite growth due to sintering by
(A) atomic migration or (B) crystallite migration. Reproduced from [8], Copyright 2006,
Wiley-Interscience.
There is controversy in the literature regarding which mechanism of sintering (or redispersion)
operates at a given set of conditions. Logically, atomic migration would be favored at lower

27
temperatures than crystallite migration, since the higher diffusivities of atoms or small cluster
would facilitate their migration, whereas the thermal energy necessary to induce motion of larger
crystallites would only be available at higher temperatures. Moreover, migration of small
crystallites might be favorable early in the sintering process but unfavorable as crystallites become
larger. However, focusing on only one of the three sintering mechanisms (and two dispersion
mechanisms) is a simplification that ignores the possibility that all mechanisms may occur
simultaneously and may be coupled with each other through complex physicochemical processes,
including the following: (1) dissociation and emission of metal atoms or metal-containing
molecules from metal crystallites, (2) adsorption and trapping of metal atoms or metal-containing
molecules on the support surface, (3) diffusion of metal atoms, metal-containing molecules and/or
metal crystallites across support surfaces, (4) metal or metal oxide particle spreading, (5) support
surface wetting by metal particles, (6) metal particle nucleation, (7) coalescence of, or bridging
between, two metal particles, (8) capture of atoms or molecules by metal particles, (9) liquid
formation, (10) metal volatilization through volatile compound formation, (11) splitting of
crystallites in O2 atmosphere owing to formation of oxides of a different specific volume, and (12)
metal atom vaporization. Depending upon reaction or redispersion conditions, a few or all of these
processes may be important; thus, the complexity of sintering/redispersion processes is emphasized.
In general, thermal sintering processes are kinetically slow (at moderate reaction temperatures)
and irreversible or difficult to reverse. Thus, sintering is more easily prevented than cured.
2.3.2. Factors Affecting Metal Particle Growth and Redispersion in Supported Metals
Temperature, atmosphere, metal type, metal dispersion, promoters/impurities, and support surface
area, texture, and porosity are the principal parameters affecting rates of sintering and redispersion (see
Table 8) [8,103–107]. Sintering rates increase exponentially with temperature. Metals sinter relatively
rapidly in oxygen and relatively slowly in hydrogen, although depending upon the support, metal
redispersion can be facilitated by exposure at high temperature (e.g., 500–550 °C for Pt/Al2O3) to
oxygen and chlorine, followed by reduction. Water vapor also increases the sintering rate of supported
metals, likely through chemical-assisted sintering effects similar to those described in Section 2.4.3.
Normalized dispersion (percentage of metal exposed at any time divided by the initial
percentage exposed) versus time data in Figure 18 show that at temperatures of 650 °C or higher,
rates of metal surface area loss (measured by hydrogen chemisorption) due to sintering of Ni/silica
in hydrogen atmosphere are significant, causing 70% loss of the original metal surface area within
50 h at 750 °C. In reducing atmosphere, metal crystallite stability generally decreases with
decreasing metal melting temperature, i.e., in the order Ru > Ir > Rh > Pt > Pd > Ni > Cu > Ag,
although this order may be affected by relatively stronger metal–support interactions, e.g., the
observed order of decreasing stability of supported platinum in vacuum is Pt/Al2O3 > Pt/SiO2 >
Pt/C. In oxidizing atmospheres, metal crystallite stability depends on the volatility of metal oxides
and the strength of the metal–oxide–support interaction. For noble metals, metal stability in air
decreases in the order Rh > Pt > Ir > Ru; formation of volatile RuO4 accounts for the relative
instability of ruthenium.

28
Table 8. Effects of Important Reaction and Catalyst Variables on Sintering Rates of
Supported Metals Based on General Power-Law Expression (GPLE) Data a.
Variable

Effect

Temperature
Atmosphere
Metal
Support
Promoters
Pore size
a

Sintering rates are exponentially dependent on T; Eact varies from 30 to 150 kJ/mol. Eact
decreases with increasing metal loading; it increases in the following order with
atmosphere: NO < O2 < H2 < N2
Sintering rates are much higher for noble metals in O2 than in H2 and higher for noble
and base metals in H2 relative to N2. Sintering rate decreases for supported Pt in
atmospheres in the following order: NO > O2 > H2 > N2
Observed order of decreasing thermal stability in H2 is Ru > Ir ≅ Rh > Pt; thermal
stability in O2 is a function of (1) volatility of metal oxide and (2) strength of metal
oxide–support interaction
Metal–support interactions are weak (bond strengths of 5–15 kJ/mol); with a few
exceptions, thermal stability for a given metal decreases with support in the following
order: Al2O3 > SiO2 > carbon
Some additives decrease atom mobility, e.g., C, O, CaO, BaO, CeO2, GeO2; others
increase atom mobility, e.g., Pb, Bi, Cl, F, or S. Oxides of Ba, Ca, or Sr are “trapping
agents” that decrease sintering rate
Sintering rates are lower for porous versus nonporous supports; they decrease as
crystallite diameters approach those of the pores
Refs. [8,103–107]. For the definition of a GPLE, see Equation 2 later in this section.

Figure 18. Normalized nickel surface area (based on H2 adsorption) versus time
data during sintering of 13.5% Ni/SiO2 in H2 at 650, 700, and 750 °C. Reproduced
from [108]. Copyright 1983, Elsevier.
The effect of temperature on sintering of metals and oxides can be understood physically in
terms of the driving forces for dissociation and diffusion of surface atoms, which are both
proportional to the fractional approach to the absolute melting point temperature (Tmp). Thus, as
temperature increases, the mean lattice vibration of surface atoms increases; when the Hüttig
temperature (0.3Tmp) is reached, less strongly bound surface atoms at defect sites (e.g., edges and
corner sites) dissociate and diffuse readily over the surface, while at the Tamman temperature
(0.5Tmp), atoms in the bulk become mobile. Accordingly, sintering rates of a metal or metal oxide
are significant above the Hüttig temperature and very high near the Tamman temperature; thus, the
relative thermal stability of metals or metal oxides can be correlated in terms of the Hüttig or
Tamman temperatures [109]. This can be illustrated from values of the melting and Tamman
temperatures for noble and base metals and their compounds listed in Table 9. For example,

29
sintering of copper catalysts for methanol synthesis is promoted by traces of chlorine in the feed,
which react at about 225 °C (500 K) with the active metal/metal oxide surface to produce a highly
mobile copper chloride phase having a Tamman temperature of only 79–174 °C (352–447 K)
relative to 405–527 °C (678–800 K) for copper metal or metal oxides [110].
Table 9. Values of Melting and Tamman Temperatures ( °C) for Common Catalytic
Metals and Their Compounds a.
Compound

Tmp, K

TTamman, K

THüttig, K

Ag
Au
Co
Cu
CuO
Cu2O
CuCl2
Cu2Cl2
Fe
Mo
MoO3
MoS2
Ni
NiO
NiCl2
Ni(CO)4
Rh
Rh2O3
Ru
Pd
PdO
Pt
PtO
PtO2
PtCl2
PtCl4
Zn
ZnO

1233
1336
1753
1356
1599
1508
893
703
1808
2883
1068
1458
1725
2228
1281
254
2258
1373
2723
1828
1023
2028
823
723
854
643
693
2248

617
668
877
678
800
754
447
352
904
1442
534
729
863
1114
641
127
1129
687
1362
914
512
1014
412
362
427
322
347
1124

370
401
526
407
480
452
268
211
542
865
320
437
518
668
384
76
677
412
817
548
307
608
247
217
256
193
208
674

a

Adapted from Ref. [109].

Promoters or impurities affect sintering and redispersion by either increasing (e.g., chlorine and
sulfur) or decreasing (e.g., oxygen, calcium, cesium) metal atom mobility on the support; in the
latter case, this is due to their high resistance to dissociation and migration due to high melting
points, as well as their hindering dissociation and surface diffusion of other atoms. Similarly,
support surface defects or pores impede surface migration of metal particles—especially
micropores and mesopores with pore diameters about the same size as the metal crystallites.
Historically, sintering rate data were fitted to a simple power-law expression (SPLE) of the form

30
D
)
D 0 = ks ( D ) n
dt
D0

d(
−

(1)

where k s is the sintering rate constant, D0 the initial dispersion, and n is the sintering order, which
for typical catalyst systems may vary from 3 to 15; unfortunately, the SPLE is, in general, not valid
for sintering processes because it assumes that surface area or dispersion ultimately reaches zero,
given sufficient time, when in fact, for a given temperature and atmosphere, a nonzero or limiting
dispersion is observed after long sintering times. Moreover, the use of the SPLE is further
questionable because variations in sintering order are observed as a function of time and
temperature for a given catalyst in a fixed atmosphere [105–107]; thus, data obtained for different
samples and different reaction conditions cannot be quantitatively compared. Nevertheless, it has been
shown by Fuentes [111,112] and Bartholomew [104–106] that the effects of temperature, atmosphere,
metal, promoter, and support can be quantitatively determined by fitting sintering kinetic data to the
general power-law expression (GPLE)

െ

ௗቀ

ವ
ቁ
ವబ

ௗ௧

ൌ ݇௦ ቀ

஽

஽బ

െ

஽೐೜ ௠
஽బ

ቁ

(2)

which adds a term −Deq/D0 to account for the observed asymptotic approach of the typical
dispersion versus time curve to a limiting dispersion Deq at infinite time; m, the order of sintering,
is found to be either 1 or 2. A recently compiled, comprehensive quantitative treatment of previous
sintering rate data based on the GPLE with an order m of 2 [104–106] quantitatively addresses the
effects of catalyst properties and reaction conditions on sintering rate. Some of these data are
summarized in Table 10 [108,113–115]. These data show, for example, that the rate constant, and
hence the rate of sintering, is less for Ni/Al2 O3 than for Pt/Al2 O 3 , an unexpected result in view of
the lower heat of vaporization for Ni. This result is possibly explained by a greater metal support
interaction for Ni with alumina.
Table 10. Comparison of Second-Order Sintering Rate Constants and Activation
Energies for Pt, Ni, and Ag Catalysts a .
Catalyst

0.6% Pt/γ-Al2 O3
5% Pt/γ-Al2 O3
15% Ni/γ-Al2 O3
0.6% Pt/γ-Al2 O3
5% Pt/γ-Al2 O3
1.8% Ag/η-Al2 O3

Atm.

D0 b

H2
H2
H2
Air
Air
Air

~0.85
0.10
0.16
~0.85
0.10
0.36

ks c

ks

ks

ks

Eact, d

(400 °C)

(650 °C)

(700 °C)

(750 °C)

kJ/mol

0.007
0.420
0.004
0.024
0.014
0.69

0.310
0.76
0.083
0.29
1.46
-

0.530
0.84
0.13
0.41
2.79
-

1.32
0.97
0.27
0.75
8.51
-

79
13
66
52
97
-

a

Ref.
[113]
[114]
[108]
[113]
[114]

[115]

Refs. [105,106]; b Initial metal dispersion or percentage exposed; c Second-order sintering rate
constant from general power-law expression (GPLE) with units of h −1 ; d Sintering activation energy for
GPLE, −d(D/D0 )/dt = k s [D/D0 − Deq /D0 ]m , where m = 2.

31
Sintering studies of supported metals are generally of two types: (1) studies of commercially
relevant supported metal catalysts and (2) studies of model metal–support systems. The former type
provides useful rate data that can be used to predict sintering rates, while the latter type provides
insights into the mechanisms of metal particle migration and sintering, although the results cannot
be quantitatively extrapolated to predict behavior of commercial catalysts. There is direct evidence
from the previous studies of model-supported catalysts [104,107] for the occurrence of crystallite
migration (mainly in well-dispersed systems early in the sintering process), atomic migration
(mainly at longer sintering times), and spreading of metal crystallites (mainly in oxygen
atmosphere). There is also evidence that under reaction conditions, the surface is dynamic, i.e.,
adsorbates and other adatoms rapidly restructure the surface and slowly bring about faceting;
moreover, thermal treatments cause gradual changes in the distribution of coordination sites to
minimize surface energy. There is a trend in increasing sophistication of spectroscopic tools used to
study sintering and redispersion. In the next decade, we might expect additional insights
into atomic and molecular processes during reaction at the atomic scale using STM, analytical
high resolution transmission electron microscopy (HRTEM), and other such powerful surface
science tools.
2.3.3. Sintering of Catalyst Carriers
Sintering of carriers (supports) has been reviewed by Baker and co-workers [103] and Trimm [116].
Single-phase oxide carriers sinter by one or more of the following processes: (1) surface diffusion,
(2) solid-state diffusion, (3) evaporation/condensation of volatile atoms or molecules, (4) grain
boundary diffusion, and (5) phase transformations. In oxidizing atmospheres, Ȗ-alumina and silica
are the most thermally stable carriers; in reducing atmospheres, carbons are the most thermally
stable carriers. Additives and impurities affect the thermal properties of carriers by occupying
defect sites or forming new phases. Alkali metals, for example, accelerate sintering; while calcium,
barium, nickel, and lanthanum oxides form thermally stable spinel phases with alumina. Steam
accelerates support sintering by forming mobile surface hydroxyl groups that are subsequently
volatilized at higher temperatures. Chlorine also promotes sintering and grain growth in magnesia
and titania during high temperature calcination. This is illustrated in Figure 19 [117]. By contrast,
sulfuric acid treatment of hydrated alumina (gibbsite) followed by two-step calcination, results in a
very stable transitional alumina with needle-like particle morphology [116]. Dispersed metals in
supported metal catalysts can also accelerate support sintering; for example, dispersed nickel
accelerates the loss of Al2O3 surface area in Ni/Al2O3 catalysts.
As an important example of support sintering through phase transformations, Al2O3 has a rich
phase behavior as a function of temperature and preparation. A few among the many important
phases that are stable or metastable, include boehmite, Ȗ-alumina, and Į-alumina [8,118,119].
Other phases are possible and the temperatures at which the phase transitions occur depend on
crystal size and moisture content of the starting material, but as an example, as temperature is
raised, boehmite, which is a hydrated or hydroxyl form of alumina, transforms to Ȗ-alumina
between 300 and 450 °C, then to į-alumina at ~850°C, ș-alumina at ~1000°C, and finally
Į-alumina at ~1125 °C. The corresponding crystal structures for these five phases are

32
orthorhombic, cubic defective spinel, orthorhombic, deformed monoclinic spinel, and hexagonal
close pack (hcp with ABAB stacking) [8,118,119]. The approximate surface areas of these
respective phases, as measured by nitrogen physisorption using Brunauer, Emmett, Teller (BET)
analysis, are approximately 400, 200, 120, 50, and 1 m2/g [8]. The dramatic drop in surface area
during the transition from ș to Į is associated with collapse of the microporous structure and
formation of the dense hcp phase.

Figure 19. BET surface area of titania as a function of thermal treatment and chlorine
content of fresh samples (before pretreatment). Samples were treated at the temperature
indicated for 2 h. Reproduced from [117]. Copyright 1985, Elsevier. Ɣ = Blank TiO2;
Ÿ =TiO2 soaked in H2O; ǻ = TiO2 soaked in HCl/H2O (2.06 wt% Cl); Ŷ = TiO2 soaked
in HCl/H2O (2.40 wt% Cl);ż = TiO2 soaked in HCl/H2O (2.55 wt% Cl); Ƒ = TiO2
soaked in HCl/H2O (2.30 wt% Cl).

33
2.3.4. Effects of Sintering on Catalyst Activity
Baker and co-workers [103] have reviewed the effects of sintering on catalytic activity. Specific
activity (based on catalytic surface area) can either increase or decrease with increasing metal crystallite
size during sintering if the reaction is structure-sensitive, or it can be independent of changes in metal
crystallite size if the reaction is structure-insensitive. Thus, for a structure-sensitive reaction, the impact
of sintering may be either magnified or moderated; while for a structure insensitive-reaction, sintering
has in principle no effect on specific activity (per unit surface area). In the latter case, the decrease in
mass-based activity is proportional to the decrease in metal surface area. Ethane hydrogenolysis and
ethane steam reforming are examples of structure-sensitive reactions, while CO hydrogenation on
supported cobalt, nickel, iron, and ruthenium is largely structure-insensitive in catalysts of moderate
loading and dispersion.
2.4. Gas/Vapor–Solid and Solid-State Reactions
In addition to poisoning, there are a number of chemical routes leading to catalyst deactivation:
(1) reactions of the vapor phase with the catalyst surface to produce (a) inactive bulk and
surface phases (rather than strongly adsorbed species), (b) volatile compounds that exit the
catalyst and reactor in the vapor phase, or (c) sintering due to adsorbate interactions, that we
call chemical-assisted sintering to distinguish it from thermal sintering previously discussed;
(2) catalytic solid-support or catalytic solid-promoter reactions, and (3) solid-state transformations
of the catalytic phases during reaction. Each of these routes is discussed in some detail below.
2.4.1. Gas/Vapor–Solid Reactions
2.4.1.1. Reactions of Gas/Vapor with Solid to Produce Inactive Phases
Dispersed metals, metal oxides, metal sulfides, and metal carbides are typical catalytic phases,
the surfaces of which are similar in composition to the bulk phases. For a given reaction, one of
these catalyst types is generally substantially more active than the others, e.g., only Fe and Ru
metals are active for ammonia synthesis, while the oxides, sulfides, and carbides are inactive. If,
therefore, one of these metal catalysts is oxidized, sulfided, or carbided, it will lose essentially all
of its activity. While these chemical modifications are closely related to poisoning, the distinction
here is that rather than losing activity owing to the presence of an adsorbed species, the loss of
activity is due to the formation of a new phase altogether.
Examples of vapor-induced chemical transformations of catalysts to inactive phases are listed in
Table 11 [8,120–127]. These include the formation of RhAl2O4 in the three-way Pt–Rh/Al2O3
catalyst during high temperature operation in an auto exhaust; oxidation of Fe by low levels of O2
during ammonia synthesis or by H2O during regeneration; dealumination (migration of Al from the
zeolite framework) of Y-zeolite during high temperature catalytic cracking and regeneration in
steam; reaction of SO3 with the alumina support to form aluminum sulfate leading to support
breakdown and catalyst pore plugging in several processes, including CO oxidation in a gas turbine
exhaust, conversion of CO and hydrocarbons in a diesel exhaust converter, and selective catalytic

34
reduction (SCR) of NOx in utility boiler flue gases [8,122–124,127]; oxidation of Fe5C2 to Fe3O4 and of
Co metal supported on alumina or silica to Co surface aluminates or silicates during Fischer–Tropsch
synthesis at high conversions and hence high PH2O; and formation of NiAl2O4 during reaction and steam
regeneration of Ni/Al2O3 in a slightly oxidizing atmosphere above about 500 °C, especially if more
reactive aluminas, e.g., Ȗ, į, or ș forms, are used as supports. Because reaction of SO3 with Ȗ-Al2O3 to
produce Al2(SO4)3 is a serious cause of deactivation of alumina-supported catalysts in several catalytic
processes (e.g., diesel exhaust abatement and SCR), TiO2 or SiO2 carriers are used rather than Al2O3 or
in the diesel or automotive exhaust the alumina catalyst is stabilized by addition of BaO, SrO, or
ZrO2 [8,122–127].
Table 11. Examples of Reactions of Gases/Vapors with Catalytic Solids to Produce
Inactive Phases.
Deactivating chemical
reaction

Catalytic process

Gas/vapor
composition

Catalytic solid

Auto emissions control

N2, O2, HCs, CO,
NO, H2O, SO2

Pt–Rh/Al2O3

2 Rh2O3 +
Ȗ-Al2O3ĺRhAl2O4 + 0.5 [120,121]
O2

Ammonia synthesis and
regeneration

H2, N2

Fe/K/Al2O3

FeĺFeO at >50 ppm O2

Ref.

[8]

FeĺFeO at >0.16 ppm
H2O/H2
H2O induced Al migration
from zeolite framework
[8]
La-Y-zeolite
Catalytic cracking
HCs, H2, H2O
causing zeolite
destruction
N2, O2, 400 ppm
2 SO3 +
CO oxidation, gas turbine
CO, 100–400 ppm
Ȗ-Al2O3ĺAl2(SO4)3
Pt/Al2O3
[8]
exhaust
SO2
which blocks catalyst pores
Formation of Al2(SO4)3 or
N2, O2, HCs (gas
Pt/Al2O3 and
sulfates of Ca, Cu, Fe, or
Diesel HC/soot emissions and liquid), CO, ȕ-zeolite; oxides of V, which block catalysts
[122–124]
NO, H2O, soot,
CaCuFeVK on
control
pores and lower activity
TiO2
for oxidation; Al2O3
SO2
stabilized by BaO
Fe5C2ĺFe3O4 due to
CO, H2, H2O, CO2,
oxidation at high XCO by
Fischer–Tropsch
[125]
Fe/K/Cu/SiO2
HCs
product H2O, CO2
Co + SiO2ĺCoO·SiO2
[126]
Fischer–Tropsch
CO, H2, H2O, HCs
Co/SiO2
and collapse of SiO2 by
product H2O
Selective catalytic
N2, O2, NO, PM a,
Formation of Al2(SO4)3 if
V2O5/WO3/TiO2
[127]
reduction (SCR),
Al2O3 is used
H2O, SO2
stationary
CH4, H2O, CO, H2,
Steam reforming and
Ni + Al2O3ĺNiAl2O4
[8]
Ni/Al2O3
regeneration in H2O
CO2
Traces O2, H2O

a

Particulate matter.

35
2.4.1.2. Reactions of Gas/Vapor with Solid to Produce Volatile Compounds
Metal loss through direct vaporization is generally an insignificant route to catalyst deactivation.
By contrast, metal loss through formation of volatile compounds, e.g., metal carbonyls, oxides,
sulfides, and halides in CO, O2, H2S, and halogen-containing environments, can be significant over
a wide range of conditions, including relatively mild conditions. Classes and examples of volatile
compounds are listed in Table 12. Carbonyls are formed at relatively low temperatures but high
pressures of CO; halides can be formed at relatively low temperatures and low concentration of the
halogens. However, the conditions under which volatile oxides are formed vary considerably with
the metal; for example, RuO3 can be formed at room temperature, while PtO2 is formed at
measurable rates only at temperatures exceeding about 500 °C.
Table 12. Types and Examples of Volatile Compounds Formed in Catalytic Reactions.
Gaseous environment

Compound type

Carbonyls and nitrosyl
carbonyls
Oxides
Sulfides
Halides

CO, NO
O2
H2S
Halogens
a

Example of compound

Ni(CO)4, Fe(CO)5 (0–300 °C) a
RuO3 (25 °C), PbO (>850 °C), PtO2 (>700 °C)
MoS2 (>550 °C)
PdBr2, PtCl4, PtF6, CuCl2, Cu2Cl2

Temperatures of vapor formation are listed in parentheses.

While the chemical properties of volatile metal carbonyls, oxides, and halides are well known,
there is surprisingly little information available on their rates of formation during catalytic
reactions. There have been no reviews on this subject and relatively few reported studies to define
the effects of metal loss on catalytic activity [28,128–141]. Most of the previous work has focused on
volatilization of Ru in automotive converters [128–131]; nickel carbonyl formation in nickel catalysts
during methanation of CO [133,139] or during CO chemisorption at 25 °C [28,135], and formation of
Ru carbonyls during Fischer–Tropsch synthesis [136,137]; volatilization of Pt during ammonia
oxidation on Pt–Rh gauze catalysts [140,141]; and volatilization of Cu from methanol synthesis and
diesel soot oxidation catalysts, leading to sintering in the former and better catalyst–soot contact but
also metal loss in the latter case [109].
Results of selected studies are summarized in Table 13. Bartholomew [131] found evidence of
significant (50%) Ru loss after testing of a Pd–Ru catalyst in an actual reducing automobile exhaust
for 100 h, which he attributed to formation of a volatile ruthenium oxide and which was considered
responsible at least in part for a significant loss (20%) of NO reduction activity.

36

Table 13. Documented Examples of Reactions of Vapor with Solid to Produce
Volatile Compounds.
Catalytic process

Catalytic solid

Vapor formed

Automotive
converter

Pd–Ru/Al2O3

RuO4

Methanation of
CO

Ni/Al2O3

Ni(CO)4

CO chemi-sorption

Ni catalysts

Ni(CO)4

Fischer–Tropsch
synthesis (FTS)

Ru/NaY zeolite,
Ru/Al2O3, Ru/TiO2

Ru(CO)5,
Ru3(CO)12

Ammonia
oxidation

Pt–Rh gauze

PtO2

HCN synthesis

Pt–Rh gauze

PtO2

Methanol
synthesis

CuZnO

CuCl2, Cu2Cl2

Diesel soot
oxidation

Oxides of K, Cu,
Mo, and trace Cl

CuCl2, Cu2Cl2

Comments on deactivation
process

50% loss of Ru during
100-h test in reducing
automotive exhaust
PCO > 20 kPa and T < 425 °C
due to Ni(CO)4 formation,
diffusion and
decomposition on the
support as large crystallites
PCO > 0.4 kPa and T > 0 °C
due to Ni(CO)4 formation;
catalyzed by sulfur
compounds
Loss of Ru during FTS
(H2/CO = 1, 200–250 °C, 1
atm) on Ru/NaY zeolite and
Ru/Al2O3; up to 40% loss
while flowing CO at
175–275 °C over Ru/Al2O3
for 24 h. Rate of Ru loss
less on titania-supported Ru
and for catalysts containing
large metal crystallites
(3 nm) relative to small
metal crystallites (1.3 nm).
Surface carbon lowers loss
Loss: 0.05–0.3 g Pt/ton
HNO3; recovered with Pd
gauze; loss of Pt leads to
surface enrichment with
inactive Rh
Extensive restructuring and
loss of mechanical strength
Mobile copper chloride phase
leads to sintering at reaction
temperature (225 °C)
Mobile copper chloride
improves catalyst–soot
contact; catalyst
evaporation observed

Ref.

[131]

[133]

[134]

[136,137]

[8,142]

[8,143]
[109]

[109]

37
Shen and co-workers [133] found that Ni/Al2O3 methanation catalysts deactivate rapidly during
methanation at high partial pressures of CO (>20 kPa) and temperatures below 425 °C because of
Ni(CO)4 formation, diffusion, and decomposition on the support as large crystallites; under severe
conditions (very high PCO and relatively low reaction temperatures) loss of nickel metal occurs.
Thus, loss of nickel and crystallite growth could be serious problems at the entrance to methanation
reactors where the temperature is low enough and PCO high enough for metal carbonyl formation.
Agnelli and co-workers [139] investigated kinetics and modeling of sintering due to formation and
migration of nickel carbonyl species. They found that the initially sharp crystallite size distribution
evolved during several hours of sintering under low temperature (230 °C) reaction conditions to a
bimodal system consisting of small spherical crystallites and large faceted crystals favoring (111)
planes. The sintering process was modeled in terms of an Ostwald-ripening mechanism coupled
with mass transport of mobile subcarbonyl intermediates. Long-term simulations were found to
predict reasonably well the ultimate state of the catalyst. On the basis of their work, they proposed
two solutions for reducing loss of nickel: (1) increasing reaction temperature and decreasing CO
partial pressure in order to lower the rate of carbonyl formation, and (2) changing catalyst
composition, e.g., alloying nickel with copper or adding alkali to inhibit carbonyl species migration.
Of note, Kuo and Hwang have shown that the particle morphology itself affects the rate of
Ostwald ripening due to different relative chemical potential energies of the surfaces [144]. Using
silver nanoparticles, they found that atoms at sharp edges and corners were removed first, resulting
in more rounded particles for all starting geometries. Thus, initial particle geometry appears to have
an effect in addition to the chemical atmosphere experienced by the particles.
Loss of nickel metal during CO chemisorption on nickel catalysts at temperatures above 0 °C is
also a serious problem; moreover, this loss is catalyzed by sulfur poisoning [28]. In view of the
toxicity of nickel tetracarbonyl, the rapid loss of nickel metal, and the ill-defined adsorption
stoichiometries, researchers are advised to avoid using CO chemisorption for measuring nickel
surface areas; instead, hydrogen chemisorption, an accepted ASTM method with a well-defined
adsorption stoichiometry, is recommended [145]. Figure 20 illustrates a mechanism for the
formation of Ni(CO)4 on a crystallite of nickel in CO atmosphere.
Goodwin and co-workers [136,137] studied the influence of reaction atmosphere, support, and
metal particle size on the loss of Ru due to carbonyl formation. They found that the loss of Ru
during CO hydrogenation (H2/CO = 1, 200–250 °C, 1 atm) on Ru/NaY zeolite and Ru/Al2O3 for
extended periods of time was significant (e.g., up to 40% while flowing CO at 175–275 °C over
Ru/Al2O3 for 24 h). The loss of Ru was significantly less on titania-supported Ru; moreover, the
rate of loss was lower for catalysts containing large metal crystallites (3 nm) relative to those
containing small metal crystallites (1.3 nm). Metal loss was inhibited in part at higher reaction
temperatures as a result of carbon deposition. Thus, while it is clear that loss of ruthenium could be
a serious problem in Fischer–Tropsch synthesis, there are measures in terms of catalyst design and
choice of reaction conditions that can be taken to minimize loss.

38

Figure 20. Formation of volatile tetra-nickel carbonyl at the surface of nickel crystallite
in CO atmosphere. Reproduced from [8]. Copyright 2006, Wiley-Interscience.
One of the most dramatic examples of vapor phase loss of the catalyst occurs during NH3
oxidation on Pt–Rh gauze, an important reaction in the manufacture of nitric acid [8,140,141]. At
the high reaction temperature (~900 °C), formation of a volatile platinum oxide (PtO2) occurs at a
very significant rate; in fact, the rate of loss of 0.05–0.3 g Pt/ton of HNO3 is high enough to provide
a substantial economic incentive for Pt recovery [8]. The most effective recovery process involves
placing a woven Pd-rich alloy gauze immediately below the Pt–Rh gauze to capture the Pt through
formation of a Pd–Pt alloy. Pt loss is also the most significant cause of catalyst deactivation as the
gauze surface becomes enriched in nonvolatile but inactive rhodium oxide [142], requiring
shutdown and catalyst replacement every 3–12 months [8].
Decomposition of volatile platinum oxide species formed during high temperature reaction may
(similar to the previously discussed formation of large crystallites of Ni from Ni(CO)4) lead to
formation of large Pt crystallites and/or substantial restructuring of the metal surface. For example,
Wu and Phillips [146–148] observed surface etching, enhanced sintering, and dramatic surface
restructuring of Pt thin films to faceted particles during ethylene oxidation over a relatively narrow
temperature range (500–700 °C). The substantially higher rate of sintering and restructuring in
O2/C2H4 relative to that in nonreactive atmospheres was attributed to the interaction of free radicals
such as HO2, formed homogeneously in the gas phase, with the metal surface to form metastable
mobile intermediates. Etching of Pt–Rh gauze in a H2/O2 mixture under the same conditions as Pt
surfaces (600 °C, N2/O2/H2 = 90/7.5/2.5) has also been reported [143]. A significant weight loss
was observed in a laminar flow reactor with little change in surface roughness, while in an
impinging jet reactor, there was little weight loss, but substantial restructuring of the surface to
particle-like structures, 1–10 ȝm in diameter; these particles were found to have the same Pt–Rh
composition as the original gauze. The nodular structures of about 10-ȝm diameter formed in these
experiments are strikingly similar to those observed on Pt–Rh gauze after use in production of
HCN at 1100 °C in 15% NH3, 13% CH4, and 72% air (see Figure 21). Moreover, because of the
high space velocities during HCN production, turbulent rather than laminar flow would be

39
expected, as in the impinging jet reactor. While little Pt is volatilized from the Pt–Rh gauze catalyst
during HCN synthesis, the extensive restructuring leads to mechanical weakening of the gauze [8].

Figure 21. (a) SEM of Pt–Rh gauze after etching in N2/O2/H2 = 90/7.5/2.5 at 875 K for
45 h. Reproduced from [143]. Copyright 1992, Elsevier. (b) SEM of Pt–Rh gauze after
use in production of HCN (magnification 1000×). Photograph courtesy of Ted Koch at
DuPont, personal correspondence to the author.
Other examples of catalyst deactivation due to volatile compound formation include (1) loss of the
phosphorus promoter from the VPO catalyst used in the fluidized-bed production of maleic
anhydride, with an attendant loss of catalyst selectivity [8], (2) vapor-phase loss of the potassium
promoter from steam-reforming catalysts in the high temperature, steam-containing environment [8],
and (3) loss of Mo from a 12-Mo-V-heteropolyacid due to formation of a volatile Mo species during
oxydehydrogenation of isobutyric acid to methacrylic acid [138].
While relatively few definitive studies of deactivation by volatile compound formation
have been reported, the previous work does provide the basis for enumerating some general
principles. A generalized mechanism of deactivation by formation of volatile metal compounds can
be postulated (see Figure 22). In addition, the roles of kinetics and thermodynamics can be stated in
general terms:
(1) At low temperatures and partial pressures of the volatilization agent (VA), the overall rate
of the process is limited by the rate of volatile compound formation.
(2) At intermediate temperatures and partial pressures of the VA, the rate of formation of the
volatile compound exceeds the rate of decomposition. Thus, the rate of vaporization is
high, the vapor is stable, and metal loss is high.
(3) At high temperatures and partial pressures of the VA, the rate of formation equals the rate
of decomposition, i.e., equilibrium is achieved. However, the volatile compound may be
too unstable to form or may decompose before there is an opportunity to be transported
from the system. From the previous work, it is also evident that besides temperature and
gas phase composition, catalyst properties (crystallite size and support) can play an
important role in determining the rate of metal loss.

40

Figure 22. Generalized mechanisms and kinetics for deactivation by metal loss.
Reproduced from [8]. Copyright 2006, Wiley-Interscience.
2.4.2. Solid-State Reactions
Catalyst deactivation by solid-state diffusion and reaction appears to be an important mechanism
for degradation of complex multicomponent catalysts in dehydrogenation, synthesis, partial
oxidation, and total oxidation reactions [8,149–160]. However, it is difficult in most of these
reactions to know the extent to which the solid-state processes, such as diffusion and solid-state
reaction, are affected by surface reactions. For example, the rate of diffusion of Al2O3 to the
surface to form an aluminate may be enhanced by the presence of gas-phase oxygen or water or the
nucleation of a different phase may be induced by either reducing or oxidizing conditions.
Recognizing this inherent limitation, the focus here is nevertheless on processes in which formation
of a new bulk phase (and presumably the attendant surface phase) leads to substantially lower
activity. There is probably some overlap with some of the examples given under Gas/Vapor–Solid
Reactions involving reactions of gas/vapor with solid to produce inactive phases.
Examples from the literature of solid-state transformations leading to catalyst deactivation are
summarized in Table 14. They include (1) the formation of KAlO2 during ammonia synthesis at the
Fe/K/Al2O3 catalyst surface, (2) decomposition of the active phase PdO to inactive Pd metal during
catalytic combustion on PdO/Al2O3 and PdO/ZrO2 catalysts, (3) transformation of active carbides
to inactive carbides in Fischer–Tropsch synthesis on Fe/K/Cu catalysts, (4) formation of inactive
V(IV) compounds in SO2 oxidation, and (5) reductive transformation of iron molybdate catalysts
during partial oxidation of benzene, methanol, propene, and isobutene.
Table 14. Examples of Solid-State Transformations Leading to Catalyst Deactivation.
Catalytic process

Catalytic solid

Ammonia synthesis

Fe/K/Al2O3

Catalytic combustion

PdO/Al2O3, PdO/ZrO2

Catalytic combustion

Co/K on MgO, CeO2, or
La2O3

Deactivating chemical reaction

Formation of KAlO2 at catalyst
surface
PdOĺPd at T > 800 °C
Formation of CoO–MgO solid
soln., LaCoO3, or K2O film on
CeO2

Ref.

[159]
[152]
[160]

41
Table 14. Cont.
Catalytic process

Catalytic solid

Deactivating chemical reaction

Dehydrogenation of
ethyl benzene to styrene

Fe2O3/Cr2O3/K2O

Fischer–Tropsch

Fe/K, Fe/K/CuO

Oxidation of SO2 to SO3

V2O5/K2O/Na2O/

Partial oxidation of
benzene to maleic
anhydride

V2O5–MoO3

Partial oxidation of
methanol to
formaldehyde

Fe2(MoO4)3 plus MoO3

Structural reorganization to ȕFeMoO4; reduction of MoO3

[150,156]

Partial oxidation of
propene to acrolein

Fe2(MoO4)3

Reductive transformation of
Mo18O52 to Mo4O11

[153,156]

Partial oxidation of
isobutene to
methacrolein

Fe2(MoO4)3

Reduction to FeMoO4 and MoO3–x

[151,154]

K migration to center of pellet
caused by thermal gradient
Transformation of active carbides to
inactive carbides
Formation of inactive V(IV)
compounds at T < 420–430 °C
Decreased selectivity due to loss of
MoO3 and formation of inactive
vanadium compounds

Ref.

[8]
[157,158]
[155]
[149]

There are basic principles underlying most solid-state reactions in working catalysts that have
been enumerated by Delmon [156]: (1) the active catalytic phase is generally a high-surface-area
defect structure of high surface energy and as such a precursor to more stable, but less active
phases and (2) the basic reaction processes may itself trigger the solid-state conversion of the active
phase to an inactive phase; for example, it may involve a redox process, part of which nucleates the
inactive phase.
A well-documented example of these principles occurs in the partial oxidation of propene to
acrolein on a Fe2(MoO4)3 catalyst [153,156]. This oxidation occurs by the “Mars van Krevelen”
mechanism, i.e., a redox mechanism in which lattice oxygen reacts with the adsorbed hydrocarbon
to produce the partially oxygenated product; the reduced catalyst is restored to its oxidized state
through reaction with gaseous oxygen. In propene oxidation, two atoms of oxygen from the catalyst
are used, one for removing two hydrogen atoms from the olefin and the other one in forming the
unsaturated aldehyde. The fresh, calcined catalyst MoO3 consists of corner-sharing MoO6 octahedra
(with Mo at the center and six oxygen atoms at the corners); but upon reduction to MoO2,
octahedra share edges as shown in Figure 23. However, it has been reported [153,156] that only
slightly reduced (relative to MoO3), open structures such as Mo18O52 and Mo8O23 are the most
active, selective phases; more complete reduction of either of these structures leads to formation of
Mo4O11 (see Figure 24) having substantially lower selectivity. Delmon and co-workers [154,156]
have shown that addition of an oxygen donor such as Sb2O4 facilitates spillover of oxygen and
thereby prevents overreduction and deactivation of the catalyst.

42

Figure 23. Schematic representation of the cyclic reduction/oxidation of twin pairs of
MoO6 octahedra between the corner and the edge-sharing arrangements (boxes
represent MoO6 octahedra with sharing of oxygen atoms at corners for MoO3 or edges
for MoO2). The figure is not completely accurate, because it cannot take into account
the fact that the arrangements are not perpendicular to the main axes of the lattice.
Adapted from [156].

(a)

MoO3

Activation

(b)

Mo18O52

18 Mo columns
(or Mo8O23 with 8 columns)
Aging

(c)

Mo4O11

Figure 24. Schematic representation of the structure of MoO3, Mo18O52, and Mo4O11. The
shear planes in Mo18O52 and Mo4O11 are represented by the oblique arrows (boxes with an
“X” represent MoO5 octahedra). Adapted from [156].

43
2.4.3. Reactions of Gas/Vapor with Solid to Restructure the Surface by Chemical
Assisted Sintering
The surfaces of metallic catalysts can be greatly roughened by interactions with the reactants and/or
products. However, as opposed to forming volatile species that are transported out of the reactor as
discussed in the previous section, these interactions lead to a restructuring of the surface that is similar
to that which occurs during thermal sintering, but at temperatures which are below the Tamman or
Huttig temperatures, respectively defined as 0.5 and 0.3 of the melting point (Tm) of the material, at
which thermal sintering might be expected. Therefore, this surface restructuring must be attributed to
the interaction of the gas phase with the solid. The following three examples from the literature
highlight the chemical-assisted sintering process caused by adsorbate-surface interactions on Ni, Co,
and Pd surfaces.
Chemical sintering of Ni/alumina catalysts in methanation due to formation of volatile Ni(CO)4
followed by its decomposition downstream to large Ni crystallites has been well documented [8,105].
Moreover, deactivation of Ni/alumina by Ni aluminate formation is also observed at the exit of
methanators where temperature is moderately high (T = 450 °C) and steam partial pressure is
maximum [105].
Wilson and de Groot [161] reported that under high pressure (4 bar, H2/CO = 2) and moderate
temperature (523 K) conditions, single crystal Co (0001) surfaces restructured significantly due to
interaction with the CO, which they attributed to an etch-regrowth mechanism. The left hand panel
of Figure 25a shows the scanning tunneling microscopy (STM) image of the single crystal surface,
while Figure 25b shows the same location after exposure to the H2/CO atmosphere for 1 h The
surface restructuring and roughening is profound, with the peaks approximately four atoms high
relative to the previously smooth surface that had only well-defined steps interrupting the (0001)
planar surface.

Figure 25. STM images of the Co (0001) surface (a) before and (b) after 1 h exposure to
4 MPa 2:1 H2:CO atmosphere at 523 K. Reproduced from [161]. Copyright 1995,
American Chemical Society.
More recently, Parkinson et al. [162] have shown that chemical-assisted sintering occurs at
room temperature for palladium supported on magnetite under ultra high vacuum conditions with
CO partial pressures of only 5 × 10í10 mbar. Figure 26 shows four STM images from a movie that

44
demonstrates the surface mobility of the Pd at these low CO partial pressures. Figure 26a is the
surface prior to CO exposure, while Figures 26b–d show the surface as a function of time up to
about an hour of exposure. The authors note that hydroxyl-Pd groups (OH-Pd), identified by
the ×’s in the images, serve as anchoring points for the coalescence of larger Pd clusters. The full
movie, available with the supplementary material for this article [162], is recommended to fully
appreciate the unexpectedly high atomic mobility under these conditions.

Figure 26. “The CO-induced formation of a large Pd cluster. a–d, Four STM images
(14 × 14 nm2, +1 V, 0.2 nA) selected from a 36-frame STM movie (duration 1 h 50 min)
following the deposition of 0.2 ML Pd [on Fe3O4] at RT. Initially (a), isolated Pd atoms
are present, together with hydroxyl groups and one OH–Pd (red cross). After three frames
the background pressure of CO is raised to 5 × 10-10 mbar. Thirty minutes later (frame b),
several mobile ‘fuzzy’ Pd carbonyl species, trapped at other Pd atoms, have formed.
Shortly afterwards (c), three Pd carbonyls and four adatoms have formed a large cluster.
Twenty-five minutes later (d), the cluster has captured another Pd carbonyl, and
diffused to merge with an OH–Pd species.”. Reproduced from [162]. Copyright 2013,
MacMillan Publishers.
2.5. Mechanical Failure of Catalysts
2.5.1. Forms and Mechanisms of Failure
Mechanical failure of catalysts is observed in several different forms that depend on the type of
reactor, including (1) crushing of granular, pellet, or monolithic catalyst forms due to a load in fixed
beds; (2) attrition, the size reduction, and/or breakup of catalyst granules or pellets to produce fines,
especially in fluid or slurry beds; and (3) erosion of catalyst particles or monolith coatings at high
fluid velocities in any reactor design. Attrition is evident by a reduction in the particle size or a
rounding or smoothing of the catalyst particle easily observed under an optical or electron
microscope. Washcoat loss is observed by scanning a cross section of the honeycomb channel with
either an optical or an electron microscope. Large increases in pressure drop in a catalytic process are
often indicative of fouling, masking, or the fracturing and accumulation of attritted catalyst in the
reactor bed.
Commercial catalysts are vulnerable to mechanical failure in large part because of the manner in
which they are formed; that is, catalyst granules, spheres, extrudates, and pellets ranging in diameter

45
from 50 ȝm to several millimeters are in general prepared by agglomeration of 0.02–2 ȝm aggregates
of much smaller primary particles having diameters of 10–100 nm by means of precipitation or gel
formation, followed by spray drying, extrusion, or compaction. These agglomerates have, in general,
considerably lower strengths than the primary particles and aggregates of particles from which they
are formed.
Two principal mechanisms are involved in mechanical failure of catalyst agglomerates:
(1) fracture of agglomerates into smaller agglomerates of approximately 0.2d0–0.8d0 and (2) erosion
(or abrasion) from the surface of the agglomerate of aggregates of primary particles having diameters
ranging from 0.1 to 10 ȝm [163]. While erosion is caused by mechanical stresses, fracture may be
due to mechanical, thermal, and/or chemical stresses. Mechanical stresses leading to fracture or
erosion in fluidized or slurry beds may result from (1) collisions of particles with each other or with
reactor walls or (2) shear forces created by turbulent eddies or collapsing bubbles (cavitation) at high
fluid velocities. Thermal stresses occur as catalyst particles are heated and/or cooled rapidly; they are
magnified by temperature gradients across particles and by differences in thermal expansion
coefficients at the interface of two different materials, e.g., catalyst coating/monolith interfaces; in the
latter case the heating or cooling process can lead to fracture and separation of the catalyst coating.
Chemical stresses occur as phases of different density are formed within a catalyst particle via
chemical reaction; for example, carbiding of primary iron oxide particles increases their specific
volume and micromorphology leading to stresses that break up these particles [164]. A further
example occurs in supported metal catalysts when large quantities of filamentous carbon (according
to reaction mechanisms discussed earlier) overfill catalysts pores, generating enormous stresses that
can fracture primary particles and agglomerates.
2.5.2. Role of Physical and Chemical Properties of Ceramic Agglomerates in Determining Strength
and Attrition Resistance
2.5.2.1. Factors Affecting the Magnitude of Stress Required for Agglomerate Breakage and the
Mechanisms by Which It Occurs
The extent to which a mechanism, i.e., fracture or erosion, participates in agglomerate size
reduction depends upon several factors: (1) the magnitude of a stress, (2) the strength and fracture
toughness of the agglomerate, (3) agglomerate size and surface area, and (4) crack size and radius.
Erosion (abrasion) occurs when the stress (e.g., force per area due to collision or cavitation
pressure) exceeds the agglomerate strength, i.e., the strength of bonding between primary particles.
Erosion rate is reportedly [163] proportional to the external surface area of the catalyst; thus,
erosion rate increases with decreasing agglomerate size.
2.5.2.2. Fracture Toughness of Ceramic Agglomerates
Most heterogeneous catalysts are complex, multiphase materials that consist, in large part, of porous
ceramic materials, i.e., are typically oxides, sulfides, or metals on an oxide carrier or support. When a
tensile stress of a magnitude close to the yield point is applied, ceramics almost always undergo brittle
fracture before plastic deformation can occur. Brittle fracture occurs through formation and propagation

46
of cracks through the cross section of a material in a direction perpendicular to the applied stress.
Agglomerate fracture due to a tensile stress occurs by propagation of internal and surface flaws; these
flaws, created by external stresses or inherent defects, are stress multipliers, i.e., the stress is multiplied
by 2(a/r)0.5, where a is the crack length and r is the radius of curvature of the crack tip; since a/r can
vary from 2 to 1000, the effective stress at the tip of a crack can be 4–60 times the applied stress.
Tensile stress multipliers may be microcracks, internal pores, and grain corners.
The ability of a material to resist fracture is termed fracture toughness. The plane strain fracture
toughness, KIc, is defined as
KIc = Y σ (πα )0.5

(3)

where Y is a dimensionless parameter (often close to 1.0–2.0), the magnitude of which depends upon
both specimen and crack geometries, ı is the applied stress, and a is the length of a surface crack or half
the length of an internal crack. Crack propagation and fracture are likely if the right hand side of
Equation 3 exceeds the experimental value of plane strain fracture toughness (left-hand side of
Equation 3). Plane strain fracture toughness values for ceramic materials are significantly smaller than
for metals and typically below 10 MPa(m)0.5; reported values for nonporous, crystalline alumina
(99.9%), fused silica, and zirconia (3 mol% Y2O3) are 4–6, 0.8, and 7–12 MPa(m)0.5, respectively;
flexural strengths (analogous to yield strengths for metals) for the same materials are 280–550, 100, and
800–1500 MPa [165]. Thus, on the basis of both fracture toughness and flexural strength, nonporous,
crystalline zirconia is much stronger toward fracture than alumina, which in turn is much stronger
than fused silica.
2.5.2.3. Effects of Porosity on Ceramic Agglomerate Strength
The introduction of porosity to crystalline or polycrystalline ceramic materials will, on the basis of
stress amplification, significantly decrease elastic modulus and flexural strength for materials in tension.
This is illustrated by data in Figure 27, showing that elastic modulus and flexural strength of a ceramic
alumina (probably alpha form) are reduced by 75 and 85% respectively as porosity is increased from 0 to
50% [166]. Thus, according to Figure 27b, the flexural strength of typical porous aluminas used as
catalyst supports might lie in the range of 30–40 MPa. However, yield strengths for Ȗ-Al2O3 (shown in the
next section) are factors of 3–50 lower. Nevertheless, the data in Figure 27b suggest that higher strengths
may be possible.
2.5.2.4. Compressive Strengths of Ceramic Materials
Thus far, the discussion has focused mainly on tensile strength, the extent of which is greatly
reduced by the presence of cracks or pores. However, for ceramic materials in compression, there is
no stress amplification due to flaws or pores; thus ceramic materials (including catalytic materials)
in compression are much stronger (approximately a factor of 10) than in tension. In addition, the
strength of ceramic materials can be dramatically enhanced by imposing a residual compressive
stress at the surface through thermal or chemical tempering. Moreover, introduction of binders,

47
such as graphite, enables agglomerates of ceramic powders to undergo significant plastic
deformation before fracture.

Figure 27. The influence of porosity on (a) the modulus of elasticity for aluminum oxide at
room temperature and (b) the flexural strength for aluminum oxide at room temperature.
Reproduced from [166]. Copyright 1956, Wiley.
2.5.3. Tensile Strengths and Attrition Resistance of Catalyst Supports and Catalysts
2.5.3.1. Tensile Strength Data for Catalyst Support Agglomerates
The strengths cited above for nonporous, annealed crystalline or polycrystalline materials do not
necessarily apply to porous catalyst agglomerates, even under compression; rather, agglomerate
strength is dependent upon the strengths of chemical and physical bonds, including the cohesive
energy, between primary particles. Agglomerate strength would depend greatly on the preparation
of the compact. Representative data for catalyst agglomerates (see Table 15) suggest they are
generally substantially weaker than polycrystalline ceramic materials prepared by high temperature
sintering, such as the alumina cited in Figure 27 [163,165,167–171]. For example, Pham and
co-workers [163] found that the breaking strength of a VISTA B alumina agglomerate during
uniaxial compaction is in the range of 5–10 MPa—substantially lower than the reported values for
heat-treated polycrystalline alumina of 280–550 MPa [165]. A large part of this difference (about
85–95%) can be attributed to porosity; however, the remaining 5–15% must be due to differences
in bonding between primary particles. In other words, the bonds between primary particles in
catalyst agglomerates (and some ceramic agglomerates prepared by similar methods) are typically
physical in nature (e.g., involve van der Waals forces) while those in sintered polycrystalline
ceramic agglomerates are principally chemical because of solid bridging of primary particles.
Thus, there appears to be considerable potential for strengthening catalyst agglomerates, since
their strengths are typically factors of 3–50 lower than for conventional, heat-treated ceramics of
similar porosity.

48

Table15. Mechanical Strengths and Attrition Rates of Catalyst Supports Compared to
Those of Sintered Ceramic Agglomerates.
Catalyst support or
ceramic

Preparation/pretreatment/properties

High surface area catalyst supports
Sol–gel granulation/dried 10 h at
Ȗ-Al2O3, 1.2–4.25-mm
40 °C, calcined 3 h at
spheres
450 °C/389 m2/g, dpore = 3.5 nm
Ȗ-Al2O3, 4.25-mm
Alcoa LD-350
spheres
Ȗ-Al2O3, 100 ȝm
VISTA-B-965-500C
Thermal hydrolysis/dried 110 °C,
TiO2 (anatase), 30 ȝm
calcined 2 h at 500 °C/ 92 m2/g, <10nm primary crystallites
Basic precipitation/dried 110 °C,
TiO2 (anatase), 90 ȝm
calcined 2 h at 500 °C/81 m2/g,
10–14-nm primary crystallites
Degussa P25, fumed/4-mm
TiO2 (75% anatase, 25%
extrudates/48 m2/g, Vpore = 0.34 cm3/g,
rutile)
dpore = 21 nm
High surface area catalyst supports (cont.)
Rhone-Poulenc DT51, ppt./4 mm
TiO2 (anatase)
extrudates/92 m2/g, Vpore = 0.40 cm3/g,
dpore = 8, 65 nm
Low surface area ceramics
Spray dried with organic binder;
Al2O3
plastic deformation observed
Al2O3
Heat treated (sintered), 99.9%
TiO2 (Rutile)
Partially sintered
Commercial samples from three
ZrO2 (yttria additive)
companies, spray-dried
ZrO2 (3% Y2O3)
Heat treated (sintered)

Strength,
MPa

Attrition
index,
wt%/h

Ref.

11.6 ± 1.9

0.033

[167]

0.7

0.177

[167]

6.2 ± 1.3

-

[163]

28a

-

[168]

15a

-

[168]

0.9

-

[169]

0.9

-

[169]

2.3

-

[170]

282–551
194

-

[165]
[170]

0.035–0.43

-

[171]

800–1500

-

[165]

a

Rough estimates from break points on relative density versus log[applied pressure] curves; data are
consistent with mass distribution versus pressure curves from ultrasonic tests.

2.5.3.2. Effects of Preparation and Pretreatment on Catalyst Agglomerate Strength
From the data in Table 15 it is evident that even subtle differences in preparation and
pretreatment also affect agglomerate strength. For example, spheres of Ȗ-Al2O3 prepared by sol–gel
granulation are substantially (17 times) stronger than commercial Ȗ-Al2O3 spheres [166]. Moreover,
30- and 90- ȝm diameter particles of TiO2 prepared by thermal hydrolysis or basic precipitation are
30 and 15 times stronger than commercially available 4-mm extrudates [169].
2.5.4. Attrition of Catalyst Agglomerates: Mechanisms, Studies, and Test Methods
Catalyst attrition is a difficult problem in the operation of moving-bed, slurry-bed, or
fluidized-bed reactors. Generally, stronger materials have greater attrition resistance; this

49
conclusion is supported by representative data in Table 15 for Ȗ-Al2O3, showing that the strength
of the alumina prepared by sol–gel granulation is 17 times higher, while its attrition rates is
5 times lower.
The mechanism by which attrition occurs (erosion or fracture) can vary with catalyst or support
preparation, crush strength, and with reactor environment; it can also vary with the mechanical test
method. There is some evidence in the attrition literature, supporting the hypothesis that in the presence
of a large stress, weaker oxide materials are prone to failure by fracture, while stronger materials tend to
erode. For example, in the fluid catalytic cracking process, as new silica–alumina/zeolite catalyst in the
form of 50–150-ȝm spherical agglomerates is added to replace catalyst lost by attrition, the weaker
agglomerates break up fairly rapidly by fracture into smaller subagglomerates, following which the
stronger agglomerates are slowly abraded to produce fine particles of 1–10 ȝm [172]. However, there is
also contrary evidence from Thoma and co-workers [168], showing that fracture may be the preferred
mechanism for strong TiO2 agglomerates, while abrasion is favored for weaker agglomerates. That is,
when subjected to ultrasonic stress, 30-ȝm-diameter agglomerates of amorphous anatase (TiO2)
prepared by thermal hydrolysis were observed to undergo fracture to 5–15-ȝm fragments, while 90-ȝm
agglomerates of polycrystalline anatase prepared by basic precipitation were found to break down by
erosion to 0.1–5-ȝm fragments [168]; in this case, the amorphous anatase was apparently stronger by a
factor of 2 (see Table 15). Supporting a third trend, data from Pham and co-workers [163] show that
attrition mechanism and rate are independent of agglomerate strength, but depend instead on the type of
material. That is, 100-ȝm-diameter agglomerates of precipitated Fe/Cu/K Fischer–Tropsch catalyst
(prepared by United Catalyst Incorporated) and having nearly the same strength shown in Table 15 for
Vista-B Al2O3 (6.3 vs. 6.2 MPa), were found to undergo substantial fracture to 5–30-ȝm fragments (an
increase from 45 to 85%; see Figure 28) as well as substantial erosion to 1 ȝm or less fragments
(increase from 2 to 50%). By comparison, under the same treatment conditions, 90-ȝm-diameter
agglomerates of Vista-B Al2O3 underwent much less attrition, mainly by erosion (20% increase in
0.1–5-ȝm fragments). The very low attrition resistance of the Fe/Cu/K Universal Catalysts, Inc.
(UCI) catalyst is further emphasized by the unsatisfactory outcome of a test of this catalyst by the
U.S. Department of Energy (DOE) in a pilot-scale slurry-phase bubble-column reactor in LaPorte,
TX.; following one day of operation, the filter system was plugged with catalyst fines, preventing
catalyst–wax separation and forcing shutdown of the plant [173].
Thus, based on these three representative examples, it follows that which of the two attrition
mechanisms predominates depends much more on material composition and type than on
agglomerate strength. However, irrespective of mechanism, the rate of attrition is usually greater for
the weaker material.

50

Figure 28. Sedigraph particle size distribution for a United Catalysts, Inc. (UCI)
Fischer-Tropsch catalyst (designated as UCI-LAPI-COMP-DRUMC), used previously in
Department of Energy (DOE) pilot-plant tests. There is considerable particle breakdown
and generation of fine particles after 15 min of ultrasonic irradiation. Reproduced
from [163]. Copyright 1999, Elsevier. —Ƒ— 0 min; – –¸– –5 min; - -ż- - 10 min; – – ǻ –
–15 min.
Figure 29 illustrates the large effect that catalyst preparation method can have on the attrition
resistance of an Fe/Cu Fischer–Tropsch catalyst [174]. This catalyst, prepared by precipitation,
undergoes severe attrition during a 25-min treatment with ultrasonic radiation; indeed, the mass
fraction finer than 0.1–5 ȝm increases from 0 to 65%. However, after a spray drying treatment of
the same catalyst, an increase of only 0 to 10% in the same fractions is evident.

Figure 29. Sedigraph particle size distributions of a precipitated Fe–Cu catalyst, as-prepared
and after spray-drying. The as-prepared catalyst (a) is weak and breaks down easily after 25
min of ultrasonic irradiation, while spray-drying (b) improves its attrition resistance.
Reproduced from [174]. Copyright 2000, Elsevier. —ƶü 0 min; Ƈ 5 min; —ƻü 10 min;
—ǻ— 15 min; —7— 20 min; — ๨— 25 min.

51
In their review of attrition and attrition test methods, Bemrose and Bridgewater [175] discuss
how attrition varies with reactor type, e.g., involves mainly particle–wall impacts in moving pellet
bed reactors and particle–particle impacts in fluidized-bed reactors of high fluid velocity. In fact,
jet attrition of catalyst particles in a gas fluidized-bed involving principally abrasion due to
collision of high-velocity particles has been modeled in some detail [172,176]. Thus, given such
important differences in attrition mechanism, realistic attrition test methods should attempt to
model reactor operation as closely as possible. In addition, the ideal test would require only a small
catalyst sample, a simple, inexpensive apparatus, and a few minutes to complete the test. Relatively
quick, inexpensive single-particle crushing tests have been devised [175]; however, properties of a
single particle are rarely representative of those for the bed; moreover, it is difficult to relate the
results of this crushing test to the actual abrasion process. Realistic tests have been devised for two
reactor types involving a moving catalyst, i.e., an air-jet test for fluidized-bed catalysts [177,178],
and a rotating drum apparatus for moving-bed catalysts [179]; however, the air-jet test requires a
large quantity (e.g., 50 g) of catalyst, an expensive apparatus, and about 20 h to run. In the past
decade, a new jet-cup test has been developed for testing of fluidized-bed catalysts [177,178],
which requires only a 5-g sample and about 1 h to complete; comparisons of results for the
jet-cup and air-jet tests indicate that the two tests give comparable results [177,178]. Nevertheless,
the mechanisms for the two tests are different, i.e., the air-jet (fluid-bed) test is abrasion(erosion-) dominant, while the jet-cup test includes both abrasion and fracture mechanisms [178].
A 30-min, 10-g ultrasonic attrition test based on cavitation has also been developed in the past
decade [168,174,180]; while it likewise involves both abrasion and fracture mechanisms, the
results appear to correlate with other methods. For example, particle size distributions for the same
Co/silica catalyst after ultrasonic, jet-cup, and laboratory-scale, slurry-bed column reactor (SBCR)
tests are very similar (see Figure 30), indicating that both fracture and abrasion mechanisms
operate in the small-scale SBCR. Moreover, the good agreement among the three methods suggests
that both the jet-cup and ultrasonic tests may provide data representative of the attrition process in
laboratory-scale SBCR reactors. It is evident that these two small-scale methods are especially
useful for screening of a series of catalysts to determine relative strength.
Nevertheless, the more realistic large-scale tests are probably needed for accurately determining
design attrition rates of a commercial catalyst to be used in a full-scale process. The observation
that attrition of a fluid catalytic cracking (FCC) catalyst initially involves fracture of weak
agglomerates followed by abrasion of strong agglomerates emphasizes the need to collect and
analyze the particle size distribution of attrited fines as a function of time in order to define which
mechanism (or mechanisms) operates at startup as well as in the steady-state process. Because the
mechanism may be time dependent, rapid, small-scale tests may produce misleading results.
While realistic laboratory-scale tests have been developed for simulating attrition in large
moving-bed and fluidized-bed reactors, no such laboratory test has been developed and demonstrated
yet for simulation of large-scale SBCR reactors, although recent research has focused on the
development of such tests. For example, in laboratory-scale, SBCR tests of supported cobalt catalysts
over several days [180], the attrition resistance decreases in the order Co/Al2O3 > Co/SiO2 >

52
Co/TiO2 (especially the anatase form underwent attrition at a high rate); attrition resistance was
observed to increase with increasing cobalt loading from 10 to 40 wt%.

Figure 30. Particle size distributions of Co/SiO2 catalyst. Adapted from [178]. –––––
Ultrasound 250 W (>10 ȝm);- - - jet cup L/min (>10 ȝm); —Ɣ— Co/SiO2 after SBCR;
Co/SiO2 fresh.
2.5.5. Implications of Mechanistic Knowledge of Attrition for Catalyst Design
The understanding of mechanisms important in attrition of catalyst supports and catalysts, the
relationship between strength and attrition rate for a given material, and test data can be used to
great advantage in the design of attrition resistant catalysts. Several alternatives follow from the
previous discussion for increasing attrition resistance: (1) increasing aggregate/agglomerate
strength by means of advanced preparation methods, e.g., sol–gel granulation, spray drying, and
carefully controlled precipitation methods (see Table 15 and Figure 29 for examples), (2) adding
binders to improve strength and toughness, e.g., the addition of a polyvinylpyrrolidone binder to
agglomerates of quartz sand increases agglomerate strength from 0.1 to 3 MPa [181], (3) coating
aggregates with a porous but very strong material such as ZrO2, e.g., embedding a fluidized-bed
catalyst for partial oxidation of n-butane to maleic anhydride in a strong, amorphous matrix of
zirconium hydrogen phosphate significantly improves its attrition resistance [182], and (4)
chemical or thermal tempering of agglomerates to introduce compressive stresses that increase
strength and attrition resistance, e.g., heating and cooling particles rapidly by passing them through
a low-residence-time, high-temperature furnace to harden the agglomerate exterior, while
preventing significant sintering of or phase changes in the porous interior. The subject of
preventing mechanical degradation and other forms of catalyst deactivation is addressed in greater
detail under Prevention of Catalyst Decay.
2.6. Summary of Deactivation Mechanisms for Solid Catalysts
Causes of solid (heterogeneous) catalyst deactivation are basically threefold: (1) chemical,
(2) mechanical, and (3) thermal. Mechanisms of heterogeneous catalyst deactivation can be
classified into five general areas: (1) chemical degradation including volatilization and leaching,
(2) fouling, (3) mechanical degradation, (4) poisoning, and (5) thermal degradation. Poisoning and
thermal degradation are generally slow processes, while fouling and some forms of chemical and

53
mechanical degradation can lead to rapid, catastrophic catalyst failure. Some forms of poisoning
and many forms of fouling are reversible; hence, reversibly poisoned or fouled catalysts are
relatively easily regenerated. On the other hand, chemical, mechanical, and thermal forms of
catalyst degradation are rarely reversible.
3. Prevention of Catalyst Decay
It is often easier to prevent rather than cure catalyst deactivation. Many poisons and foulants can
be removed from feeds using guard beds, scrubbers, and/or filters. Fouling, thermal degradation,
and chemical degradation can be minimized through careful control of process conditions, e.g.,
lowering temperature to lower sintering rate or adding steam, oxygen, or hydrogen to the feed to
gasify carbon or coke-forming precursors. Mechanical degradation can be minimized by careful
choice of carrier materials, coatings, and/or catalyst particle forming methods.
While treating or preventing catalyst deactivation is facilitated by an understanding of the
mechanisms, additional perspectives are provided by examining the route by which each of the
mechanisms causes loss of catalytic activity, i.e., how it influences reaction rate [109]. Thus,
catalytic activity can be defined in terms of the observed site-based rate constant kobs, which is
equal to the product of the active site density ı (number of sites per area of surface), the site-based
intrinsic rate constant kintr, and the effectiveness factor Ș, i.e.,

kobs = σ kintrη

(4)

Loss of catalytic activity may be due to a decrease in any of the three factors in Equation 4,
whose product leads to kobs. Thus, catalyst deactivation can be caused by (1) a decrease in the site
density ı, (2) a decrease in intrinsic activity (i.e., decrease in kintr), and/or (3) lowered access of
reactants to active sites (decrease in Ș). Poisoning, for example, leads to a loss of active sites, i.e.,
ı = ı0(1 í Į), where Į is the fraction of sites poisoned; sintering causes loss of active sites through
crystallite growth and reduction of active surface area. Fouling can cause both loss of active sites
due to blocking of surface sites as well as plugging of pores, causing a decrease in the effectiveness
Ș. Moreover, poisoning, as discussed earlier, can also lead to a decrease in intrinsic activity by
influencing the electronic structure of neighboring atoms. Thus, each of the deactivation
mechanisms affects one or more of the factors comprising observed activity (see Table 16); all of
the mechanisms, however, can effect a decrease in the number of catalytic sites.

3.1. General Principles of Prevention
The age-old adage that says “an ounce of prevention is worth a pound of cure” applies well to
the deactivation of catalysts in many industrial processes. The catalyst inventory for a large plant may
entail a capital investment of tens of millions of dollars. In such large-scale processes, the economic
return on this investment may depend on the catalyst remaining effective over a period of up to 3–5
years. This is particularly true of those processes involving irreversible or only partially reversible
deactivation (e.g., sulfur poisoning or sintering). Some typical industrial catalysts, approximate
catalyst lifetimes, and factors that determine their life are listed as examples in Table 17. It is
evident that in many processes more than one mechanism limits catalyst life. Moreover, there is a

54
wide variation in catalyst lifetimes among different processes, i.e., from 10í6 to 15 years. While
there is clearly greater interest in extending catalyst lifetimes in processes where life is short, it
should be emphasized that great care must be exercised in protecting the catalyst in any process
from process upsets (e.g., temperature runaway, short-term exposure to impure feeds, or changes in
reactant composition) that might reduce typical catalyst life by orders of magnitude, e.g., from
years to hours.
Table 16. How Deactivation Mechanisms Affect the Rate of a Catalyzed Reaction and
the Rapidity and Reversibility of Deactivation Process.
Effects on reaction rate
Deactivation
mechanism

Decrease in
number of
active sites

Decrease in
intrinisic
activity (kintr)

Decrease in
effectiveness factor
(Ș)

Fast or slow a

Reversible

Chemical degradation
Fouling

×
×

×
×

× b,c
-

Varies
Fast

No
Yes

Mechanical
degradation

×

-

-

Varies

No

Poisoning

×

×

-

Slow

Usually

Thermal
degradation/Sintering

×

× b,d

× b,e

Slow

Sometimes

Vaporization/leaching

×

× b,f

-

Fast

Sometimes

a

Deactivation process

Generally; b In some cases; c Chemical degradation can cause breakdown of support, pore plugging, and loss of

porosity; dIf the reaction is structure-sensitive, sintering could either increase or decrease intrinsic activity; e Sintering
of the support may cause support collapse and loss of porosity; it may also increase average pore diameter. f Leaching
of aluminum or other cations from zeolites can cause buildup of aluminum or other oxides in zeolite pores.

Low temperature
water gas shift
CO + H2OĺCO2 + H2
Hydrocarbon
hydrodesulfurization
R2S + 2 H2ĺH2S + R2
High temperature water
gas shift
CO + H2OĺH2 + CO2
Copper on zinc and
aluminum oxides
Cobalt and molybdenum
sulfides on aluminum oxide
Fe3O4 and chromia

300–400 °C
30 atm

350–500 °C
20–30 atm

Copper on zinc and
aluminum oxides

200–250 °C
10–30 atm

200–300 °C
50–100 atm

Methanol synthesis
CO + 2 H2ĺCH3OH

Supported palladium

30–150 °C
20–30 atm
Vanadium and potassium
sulfates on silica

Supported nickel

250–350 °C
30 atm

420–600 °C
1 atm

Fe with promoters (K2O)
and stabilizer (Al2O3)

Catalyst

450–470 °C
200–300 atm

Operating
conditions

Sulfuric acid manufacturing
2 SO2 + O2ĺ2 SO3

Ammonia synthesis
N2 + 3 H2ĺ2 NH3
Methanation (ammonia and
hydrogen plants)
CO/CO2 + H2ėCH4 + H2O
Acetylene hydrogenation
(“front end”)
C2H2 + H2ĺC2H4

Reaction

1–4

1–10

2–4

2–5

5–10

5–10

5–10

10–15

Typical life
(years)

Activity, mass
transfer, and
pressure drop
Activity and
pressure drop

Slow sintering, pellet
breakage due to steam

Activity

Activity

Activity, pressure
drop, and
mass transfer

Inactive compound
formation; pellet
fracture; plugging
by dust
Slow sintering;
poisoning by S, Cl,
and carbonyls
Slow poisoning and
accelerated sintering
by poisons
Slow coking, poisoning
by metal deposits
in residuum

Activity/selectivity
and temperature

Activity and
pore blockage

Slow poisoning by S,
As, K2CO3 from
plant upsets
Slow sintering

Activity

Catalyst property
affected

Slow sintering

Process affecting life of
catalyst charge

Table 17. Typical Lifetimes and Factors Determining the Life of Some Important Industrial Catalysts a.

55

1–2

Vanadium phosphorus oxide
with transition
metal additives
Copper on zinc oxide

Pt–Rh alloy gauze

Copper chlorides on alumina
(fluidized bed)
Platinum alloys on
treated alumina
Synthetic zeolites
(fluidized bed)

400–520 °C
1–3 atm

220–270 °C
100–300 atm

800–900 °C
1–10 atm

230–270 °C
1–10 atm

460–525 °C
8–50 atm
500–560 °C
2–3 atm

Catalytic cracking of oils

1–3

Silver on Į-alumina with
alkali metal promoters

200–270 °C
10–20 atm

Adapted from Ref. [9].

0.000002

0.01–0.5

0.2–0.5

0.1–0.5

0.5–1

1–3

Ethylene partial oxidation
2 C2H4 + O2ĺ2 C2H4O
Butane oxidation to
maleic anhydride
C4H10 + 3.5 O2ĺC4H2O3 + 4
H2O
Reduction of aldehydes to
alcohols
RCHO + H2ĺRCH2OH
Ammonia oxidation
2 NH3 + 5/2 O2ĺ2 NO + 3
H2O
Oxychlorination of ethylene to
ethylene dichloride
2 C2H4 + 4 HCl + O2 ĺ
2 C2H4Cl2 + 2 H2O
Catalytic hydrocarbon
reforming

Nickel on calcium aluminate
or Į-alumina

500–850 °C
30 atm

Typical life
(years)

Steam reforming, natural gas
CH4 + H2OĺCO + 3 H2

Catalyst

Operating
conditions

Reaction

Table 17. Cont.

Activity and mass
transfer
Activity and
mass transfer

Fluidized state and
activity

Loss by attrition and
other causes resulting
from plant upsets
Coking, frequent
regeneration
Very rapid coking,
continuous regeneration

Selectivity,
fouling by Fe

Activity or
pressure drop

Slow sintering, pellet
breakage (depends on
feedstock)
Surface roughness, loss
of platinum

Activity and
selectivity

Activity and
selectivity

Activity and
pressure drop

Catalyst property
affected

Loss of P; attrition or
pellet breakage; S, Cl
poisoning

Process affecting life of
catalyst charge
Sintering, sulfurpoisoning, carbon
formation, and pellet
breakage due to
plant upsets
Slow sintering,
poisoning by Cl, S

56

57
While complete elimination of catalyst deactivation is not possible, the rate of damage can be
minimized in many cases through understanding of the mechanisms, thereby enabling control of
the deactivation process, i.e., prevention is possible through control of catalyst properties, process
conditions (i.e., temperatures, pressures), feedstock impurities, methods of contacting, and process
design. Figure 31 illustrates general approaches to eliminating or moderating deactivation through
modifications in catalyst and/or process. Examples of how deactivation can be prevented are
discussed below in connection with the most important causes of deactivation: chemical degradation,
fouling by coke and carbon, poisoning, sintering, and mechanical degradation. Principles for
preventing deactivation by these mechanisms are summarized in Table 18, while representative
results from studies focusing on prevention or minimization of catalyst deactivation are summarized
in Table 19.

Figure 31. Approaches to eliminating catalyst deactivation.

Fouling by coke
or carbon

Chemical
degradation

Basic mechanism

Loss of catalytic surface
sites due to formation of
carbon or coke films

Transformation of active
phase to stable,
inactive phase

Oxidation of metal by
contaminant O2 or
reactant/product water

Oxidation of metal
catalysts to inactive
oxides

Formation and growth on
metal surfaces

Free radical reactions at
reactor walls

Free radical reactions in
gas phase

Overreduction of active
oxide phases

Solid-state reaction of active
phase with support or
promoters

Cause

Problem

(1) Avoid accumulation of coke precursors (e.g., atomic carbon,
olefins) through careful choice of reactant conditions or
membranes
(2) Add gasifying agents (e.g., H2, H2O), diluents
(3) Incorporate catalyst additives to increase rate of gasification
or to change ensemble size
(4) Passivate metal surfaces with sulfur
(5) Decrease dispersion
(6) Recycle inerts to flush surface of heavy oligomers and to
moderate temperature

(1) Coat reactor with inert material

(1) Purify feed of oxidants
(2) Minimize reactant/product water by recycle/separation,
staged reactors, and otherwise limiting conversion
(3) Incorporate additives that facilitate resistance to oxidation
(1) Avoid conditions (e.g., oxidizing condition, high steam
pressures, and high temperatures) that favor solid-state reactions
(2) Select combinations of active phase and promoters/supports
that are noninteracting
(1) Stabilize oxidation state using promoters that induce
resistance to reduction or that serve as oxygen donors
(2) Add steam to the reactants to prevent overreduction
(1) Avoid formation of free radicals, lower temperature
(2) Minimize free space
(3) Free radical traps, diluents
(4) Add gasifying agents (e.g., H2, H2O)

Methods of minimization

Table 18. Methods for Preventing Catalyst Decay [8,41].

58

Mechanical
failure

Fouling by coke
or carbon (cont.)

Basic
mechanism

Attrition and/or erosion in
fixed or moving beds

Crushing of granules,
pellets, or monoliths in a
fixed bed

Loss of catalyst
effectiveness; plugging of
pores; destruction of
catalyst

Formation and growth on
metal oxides, sulfides

Loss of catalytic surface
sites due to formation of
carbon or coke films

Abrasion of catalyst coatings
or particles due to
mechanical, thermal, or
chemical stresses

Brittle fracture due to a
mechanical load

Hot spots in pellet or bed

Formation of gas phase
coke, vermicular carbons,
and liquid or solid cokes in
massive quantities

Cause

Problem

Methods of minimization

(1) Minimize formation of free radicals or coke precursors
as above
(2) Use gasifying agents
(3) Incorporate catalyst additives that lower solubility of carbon in
metal or that change ensemble size
(4) Use supports with large pores; large pellets
(1) Use wash coat or small pellets
(2) Use slurry- or fluid-bed reactor, gas diluents
(1) Minimize porosity of pellets or monoliths
(2) Improve bonding of primary particles in agglomerates that
make up pellets or monoliths using advanced forming methods,
e.g., spray drying and controlled thermal treatments
(3) Add binders such as carbon to the support material, which
facilitate plastic deformation and thus protect against
brittle fracture
(4) Chemically or thermally temper agglomerates
(1) Avoid highly turbulent shear flows and/or cavitation, leading
to high erosion rates
(2) Avoid thermal stresses in the preparation and use of catalysts
that lead to fracture or separation of coatings
(3) Avoid formation of chemical phases of substantially different
densities or growth of carbon filaments that cause fracture of
primary particles and agglomerates. Choose supports, support
additives, and coating materials, such as titanates, zirconia, and
zirconates, having high fracture toughness

(1) Utilize measures 1, 2, 3, and 6 for metal surfaces
(2) Design catalyst for optimum pore structure and acidity
(3) Use shape-selective, coke-resistant molecular sieves

Table 18. Cont.

59

Loss of support area

Same as for avoiding loss of metal area

Auto emissions control
Pt– or
Pd–Rh/Al2O3

Deactivation mechanism
Process/Reaction
Catalyst

Method(s) of minimization

Chemical degradation
In three-way catalyst, Rh is very active Place Rh in a separate catalyst layer from Pt or Pd to
prevent alloying; support Rh on ZrO2, which is a
for NO reduction, but it forms a solid
solution with Al2O3 that has no activity noninteracting support for Rh. In general, multilayer
and alloys with Pt or Pd that reduce its strategies (up to 6 layers) are used to prevent undesirable
interactions between different components of the catalyst
activity

Problem/cause

[183–185]

Ref.

Methods of minimization
(1) Purify feed and/or use guard bed to adsorb poison
(2) Employ additives that selectively adsorb poison
(3) Choose reaction conditions that lower adsorption strength
(4) Optimize pore structure and choose mass transfer regimes
that minimize adsorption of poison on active sites
(5) Apply coating that serves as diffusion barrier to poison
(1) Lower or limit reaction temperature while facilitating
heat transfer
(2) Add thermal stabilizers to catalyst; and (3) avoid water

Table 19. Representative Results from Studies Focusing on Prevention/Minimization of Catalyst Deactivation.

Thermal
degradation,
sintering

Metal particle or subparticle
migration at high
temperatures
Crystallization and/or
structural modification
or collapse

Blockage of sites by strong
adsorption of impurity

Loss of catalytic surface
sites

Poisoning

Loss of metal area

Cause

Problem

Basic mechanism

Table 18. Cont.

60

Problem/cause

Method(s) of minimization

Chemical degradation
(1) Employ two- or three-stage process that enables lower
conversion and lower concentrations of steam product in
Fischer–Tropsch
Oxidation of active Co metal crystallites
the first stage. Treat gaseous stream leaving the first or
synthesis
to inactive Co oxides, aluminates, and
second stage to remove water and liquid hydrocarbons
Co supported on Al2O3, silicates and of active iron carbides to
(2) Add noble metal promoters that facilitate and maintain
SiO2, TiO2, and
inactive Fe3O4 or Fe3C in the presence of
high reducibility of the metal or metal carbide phases
Fe/Cu/K/SiO2
high pressure steam at high conversion
(3) Stabilize silica and alumina supports with coatings of
hydrothermally stable materials such as ZrO2 and MgAl2O4
(1) Stabilize reduction state of iron molybdate catalyst
using an oxygen donor such as Į-Sb2O4; the oxygen donor
Partial oxidation of
dissociates molecular oxygen to atomic oxygen that readily
isobutene to methacrolein Overreduction of the catalyst during
spills over to the catalyst
Fe2(MoO4)3,
reaction leads to activity decrease
(2) Mo12BixCeyOz catalyst promoted with Co, Mg, Rb,
Mo12BixCeyOz
and/or Cs oxides is highly resistant to reduction, highly
selective to methacrolein, and long-lived
Steam reforming and
In the absence of steam, PdO is reduced Adding steam to the reactants inhibits oxidation of propane
steam-oxygen conversion
to less active, less thermally stable Pd
at lower reaction temperatures while preventing reduction
of propane
metal
of PdO at higher temperatures (up to 700–900 °C)
Pd/Al2O3

Deactivation
mechanism
Process/Reaction
Catalyst

Table 19. Cont.

[189]

[154,156,188]

[8,126,186,187]

Ref.

61

Catalytic reforming
of naphtha
Pt/Al2O3 promoted with
Re, Sn, Ge, or Ir

Alkylation of isoparaffins
on solid catalysts
Sulfated zirconia, USY a,
Nafion

Alkene oligomerization
Zeolites, esp.
ZSM-5, –22, –23,
beta-zeolite, ferrierite

Deactivation mechanism
Process/Reaction
Catalyst
Method(s) of minimization

Fouling by coke, carbon
(1) Recycle of heavy paraffins flushes the surface of heavy
oligomers while moderating temperature, thereby decreasing
Catalyst fouling by condensation of
the rate of coke formation
heavy oligomers to coke
(2) Addition of steam improves conversion and catalyst
life—probably by cleaning the catalyst surface of
coke precursors
(1) Near critical operation favors desorption and removal of
coke precursors from pores while enabling lower reaction
temperature
Rapid catalyst deactivation due to coke
(2) Remove oxygen, oxygenates, diolefins, and aromatics
formation; unacceptable product
from feed; passivate stainless steel surfaces with silicon
quality, and thermal degradation of
or bases
catalyst during regeneration
(3) Design catalyst for optimum pore structure and acidity
(4) Use stirred-slurry or fluid-bed reactor while minimizing
olefin concentration
(1) Use bimetallic catalyst, e.g., sulfided Pt–Re/Al2O3,
which is substantially more resistant to coke formation and
longer-lived than is Pt/Al2O3. Re sulfide sites break up
large Pt ensembles that produce coke. Sn and Ge have a
Poisoning and fouling by coke produced similar effect; Sn and Ir also improve selectivity
by condensation of aromatics and
(2) Optimize reaction conditions and reactor design, e.g.,
olefins
moving bed and low pressure; maintain optimum Cl and S
contents of catalyst throughout the bed
(3) Near critical reaction mixtures provide an optimum
combination of solvent and transport properties for
maximizing isomerization rates while minimizing coking

Problem/cause

Table 19. Cont.

[8,195–198]

[193,194]

[190–192]

Ref.

62

Hydrocracking of
heavy naphtha
CoMo, NiW, MoW on
Al2O3 or SiO2–Al2O3;
Pt or Pd on Y-zeolite,
mordenite or ZSM-5

Dehydrogenation of
propane and butane
Cr2O3/Al2O3, Cr2O3/ZrO2,
FeO/K/MgO, Pt/Al2O3,
Pt–Sn/Al2O3,
Pt–Sn/KL-zeolite

Deactivation mechanism
Process/Reaction
Catalyst
Method(s) of minimization

Fouling by coke, carbon
(1) Add Sn and alkali metals to Pt/Al2O3—additives reduce
coke coverage of active sites; Sn decreases Pt ensemble
size and enhances reactivity of hydrogen with coke
Catalyst activity is low owing to
(2) Use H2-selective silica membrane to remove product
equilibrium limitations and buildup of
H2, which increases propane conversion; catalyst
product H2; rapid loss of activity occurs
deactivation is slowed and catalyst life increases, probably
owing to coke formation
due to a lowering of surface coverage of reaction
intermediates, including coke precursors, thereby reducing
the rate of coke formation
(1) Optimize metals loading and porosity of catalyst; use
coke-resistant zeolites; incorporate amorphous
silica–alumina, which prevents build up of bulky
compounds in shape-selective zeolites
Loss of activity due to poisoning of
(2) Design process to prevent build up of polynuclear
sites and blocking of small zeolite pores
aromatics, e.g., through distillation, bleeding, flashing,
by coke
precipitation, and adsorption
(3) Decouple aromatics saturation and hydrocracking
reactions to improve selectivity, controllability, and
catalyst life, while decreasing H2 consumption

Problem/cause

Table 19. Cont.

[8,198,204]

[8,199–203]

Ref.

63

Methanol to olefins
or gasoline
Silica–alumina, Y-zeolite,
ZSM-5, other zeolites,
and aluminophosphate
molecular sieves

Methane reforming
CO2/Co/SiO2, Pt/SiO2,
Pt/ZrO2, MgO-supported
noble metals, NiO·MgO
solid solution

Deactivation mechanism
Process/Reaction
Catalyst
Method(s) of minimization

Fouling by coke, carbon
(1) Add MgO or CaO to reduce carbon deposition on Co or
Ni catalysts. CO2 adsorbs strongly on these basic oxides,
possibly providing oxygen atoms that gasify coke
precursors
High rates of carbon formation, which
(2) Adding Sn to Pt catalysts increases stability; ZrO2
rapidly deactivate catalyst
support promotes activity and selectivity by aiding
dissociation of CO2
(3) Add water or H2 or increase pressure to decrease carbon
deposition rate
(1) Maintain a positive methanol concentration through the
reactor (e.g., CSTR) to decrease olefin concentration, favor
olefin–MeOH reaction to higher olefins over olefin–olefin
reactions to coke precursors, substantially decrease coking
and deactivation rates, and thereby greatly improve activity
Severe coking and deactivation of
and selectivity
silica–alumina and Y-zeolite catalysts
(2) Increase concentration of water, which attenuates coke
observed during high conversions of
formation on SAPO-34 by competing with coke precursors
MeOH; also substantial coking of
for active sites
ZSM-5, other zeolites, and
(3) Treat SAPO-34 above 700 °C in steam to lower acidity,
alumino-phosphate molecular sieves
increase catalyst life, and increase selectivity for C2–C3
olefins. Addition of diluent to feed is also beneficial
(4) Silanation decreases activity but improves life of
zeolites, e.g., HY, HZSM-5

Problem/cause

Table 19. Cont.

[209–236]

[205–208]

Ref.

64

Method(s) of minimization

Poisoning of noble metal catalyst by
P and S compounds and large
hydrocarbons from lube oil

100 ppb of HCN and NH3 poisons
cobalt slurry catalyst within 4 days

(1) Poisoning of acid sites by
N-containing compounds.
(2) Deposition of Ni and V metals that
change selectivity and decrease activity

Fischer–Tropsch
synthesis
Co/Al2O3

Fluidized catalytic
cracking (FCC)
USY or REO-Y b in
silica matrix

Optimize pore structure of alumina, deposit noble metals in
layers below the support surface, or provide a diffusion
barrier coating of zeolite or alumina; these measures
prevent access of large poison molecules to catalyst layer
Remove HCN and NH3 to less than 50 ppb total by (1)
catalytic hydrolysis of HCN to NH3, followed by scrubbing
with water or (2) guard bed containing acidic solid
absorbent
(1) FCC matrix serves as a coating to remove N-containing
compounds before they reach zeolites
(2) Add Group 13–15 compounds to passivate metals (Sb
and Bi for Ni and In for V) and/or trap V with MgO or SrO

Fouling by coke, carbon
(1) Use basic supports or oxide promoters, which lower
carbon deposition rate by preventing hydrocracking and by
facilitating adsorption of water, which facilitates
High rates of carbon and coke
gasification of surface carbon
formation, which rapidly deactivate
(2) Promote with S, Cu, or Au, which lower rate of graphite
catalyst
formation on Ni by decreasing ensemble size (since
ensemble size for C-C bond breaking is smaller than for
graphite formation)

Problem/cause

Auto emissions control
Pt–Rh/Al2O3 or Pd/Al2O3

Poisoning

Steam reforming of light
hydrocarbons or naphtha
Ni on MgO,
MgAl2O4 or CaAl2O4
promoted with S, Cu, or
Au

Deactivation mechanism
Process/Reaction
Catalyst

Table 19. Cont.

[8,198]

[213]

[18,212]

[8,60,70,71]

Ref.

65

Catalytic combustion of
methane and LNG
PdO/La2O3, Pr2O3,
CeO2, and ZrO2

(1) Use guard bed or multistage bed to remove metals with
first stage containing large-pore, low-activity catalyst for
removal of metals and subsequent stages containing
progressively smaller-pore, higher-activity catalysts
(2) Use catalysts with bimodal pore distributions

Thermal degradation
(1) Use į- or ș-Al2O3 and Al2O3 spinels having a higher
In close-couple, fast-warm-up
thermal stability than Ȗ-alumina
converters, exhaust temperatures reach
(2) Thermally stabilize Al2O3 with BaO, La2O3, Pr2O3,
1000–1100 °C; conventional Pt–Rh/
CeO2, and ZrO2; stabilize CeO2 with ZrO2 or Pr2O3 and
Ȗ-Al2O3 catalysts sinter rapidly under
ZrO2 with Y
these conditions; CeO2 used as oxygen
(3) Employ PdO that interacts more strongly than Pt with
storage material also sinters rapidly
oxide supports and is hence more stable against sintering
(1) Develop PdO/REO catalysts that resist sintering and
Reaction temperatures ranging up to
decomposition of PdO to Pd up to 1300 °C
1400 °C cause rapid sintering of most
(2) Maintain catalyst temperature below 1000 °C by (a)
catalytic materials. Conversion above
using lean mixtures, followed by post-catalyst injection of
800 °C of PdO to Pd metal is followed
most of the fuel, or (b) employing a metal monolith with
by rapid sintering of Pd and loss of
heat exchange and gradient of catalyst through bed
activity

Pore-mouth poisoning and blockage by
Ni, V, and Fe sulfides present in feed as
organometallics

Hydrotreating of
residuum
Al2O3- supported Mo and
CoMo

Auto emissions control
PdO/į- or ș-Al2O3 doped
with BaO, La2O3, Pr2O3,
CeO2, and ZrO2

Method(s) of minimization
Poisoning
(1) Two-stage operation with removal of H2S
between stages
(2) Split feed into light and heavy streams; desulfurize light
and hydrocrack heavy streams, combine, and conduct deep
hydrogenation/HDS

Noble metal hydrogenation and
high-activity HDS catalysts are
poisoned by H2S

Problem/cause

Hydrotreating of gas oil;
deep HDS
Al2O3- supported CoMo,
noble metals

Deactivation mechanism
Process/Reaction
Catalyst

Table 19. Cont.

[8,224–227]

[8,215–222]

[8,214]

[198]

Ref.

66

Methane steam reforming
Ni on MgAl2O4 or
CaAl2O4

Fluid catalytic cracking
(FCC)
USY, REO-Y

Fischer–Tropsch
synthesis
Co/Al2O3, Co/SiO2,
Co/TiO2

Dehydrogenation of
butene to butadiene
Cr2O3/Al2O3, Pt–Sn/Al2O3

Deactivation mechanism
Process/Reaction
Catalyst
Method(s) of minimization

Thermal degradation
Optimize the operation of staged catalytic reactors (cycle
Permanent loss of catalytic activity by
time between regenerations, temperature, and composition
sintering at high reaction temperatures
of the feed as variables) while placing a limit on the upper
(550–650 °C)
temperature
(1) Employ two-stage process that enables lower
Sintering in hot spots and loss of
conversion, better heat removal, and thereby a smaller
hydrocarbon selectivity at higher
temperature increase in the first reactor
reaction temperatures due to highly
(2) Employ slurry reactor with superior heat transfer
exothermic reaction
efficiency
Dealumination and destruction of
(1) Carry out controlled dealumination or silanization of
zeolite crystallinity and loss of surface
Y-zeolite to produce USY
area/pore volume during
(2) Use of REO-Y to improve thermal stability
high-temperature (650–760 °C, 3 atm)
(3) Limit steam partial pressure during regeneration
regeneration in steam/air
(1) Design relatively low surface area catalyst with rugged,
Sintering of Ni and support during
hydrothermally stable spinel carrier of about 5 m2/g
high-temperature reaction (800–1000 °C)
(2) Form catalyst into rings to facilitate heat transfer and
in high-pressure steam (20 atm)
prevent overheating at the heated tube-wall

Problem/cause

Table 19. Cont.

[8,70]

[8,198]

[186,229]

[8,228]

Ref.

67

[182]

[230–235]

(1) Spray drying improves density and attrition resistance.
(Attrition resistance improves with higher particle density;
attrition resistance decreases in the order Co/Al2O3 >
Co/SiO2 > Co/TiO2)
(2) Addition of SiO2 and/or Al2O3 to TiO2 improves its
attrition resistance; addition of TiO2 or of La2O3 to Al2O3
improves its attrition resistance
(3) Attrition resistance of Co/Al2O3 is improved when the
Ȗ-Al2O3 is formed from synthetic boehmite having a
crystallite diameter of 4–5 nm and is pretreated in acidic
solution having a pH of 1–3

Ref.

Imbed catalyst particles in a strong, amorphous matrix of
zirconium hydrogen phosphate

Method(s) of minimization

USY: ultrastable Y-zeolite. b REO-Y: rare-earth exchanged Y-zeolite.

Attrition in bubble column
slurry reactor

Fischer–Tropsch synthesis
in a bubble-column
slurry reactor
Co/Al2O3, Co/SiO2,
Co/TiO2

a

Attrition in fluidized-bed process

Partial oxidation of
n-butane to maleic
anhydride
VPO

Mechanical degradation

Problem/cause

Deactivation mechanism
Process/Reaction
Catalyst

Table 19. Cont.

68

69

3.2. Prevention of Chemical Degradation (by Vapor–Solid and Solid–Solid Reactions)
The most serious problems of oxidation of metal catalysts, overreduction of oxide catalysts, and
reaction of the active catalytic phase with carrier or promoter, can be minimized or prevented by careful
catalyst and process design (as enumerated in Table 18 and illustrated in Table 19). For example, the
loss of Rh due to solid-state reaction with alumina in the automotive three-way catalyst can be
prevented by supporting Rh on ZrO2 in a separate layer from Pt and/or Pd on alumina [215–222] In
Fischer–Tropsch synthesis, the oxidation of the active cobalt phase in supported cobalt catalysts to
inactive oxides, aluminates, and silicates can be minimized by employing a two- or three-stage process
in which product steam is moderated in the first stage by limiting conversion and in subsequent states
by interstage removal of water [223] It can also be moderated by addition of noble metal promoters that
facilitate and maintain high reducibility of the cobalt and by coating the alumina or silica support with
materials such as ZrO2 that are less likely to react with cobalt to form inactive phases.

3.3. Prevention of Fouling by Coke and Carbon
Rostrup-Nielsen and Trimm [57], Trimm [59], and Bartholomew [60] have discussed principles
and methods for avoiding coke and carbon formation. General methods of preventing coke or
carbon formation are summarized in Table 18. Most of these are based on one important
fundamental principle: carbon or coke results from a balance between the reactions that produce
atomic carbon or coke precursors and the reactions of these species with H2, H2O, or O2 that
remove them from the surface. If the conditions favor formation over gasification, these species
accumulate on the surface and react further to form less active forms of carbon or coke, which
either coat the surface with an inactive film or plug the pores, causing loss of catalyst effectiveness,
pore plugging, or even destruction of the carrier matrix.
Methods to lower rates of formation of carbon or coke precursors relative to their rates of
gasification vary with the mechanism of formation (i.e., gas, surface, or bulk phase) and the nature
of the active catalytic phase (e.g., metal or oxide). For example, gas phase formation can be
minimized by choosing reaction conditions that minimize the formation of free radicals, by using
free-radical traps, by introducing gasifying agents (e.g., H2, H2O) or gas diluents, and by
minimizing the void space available for homogeneous reaction. Similarly, the formation and
growth of carbon or coke species on metal surfaces is minimized by choosing reaction conditions
that minimize the formation of atomic carbon or coke precursors and by introducing gasifying
agents. Selective membranes or supercritical conditions can also be used to lower the gas-phase
and surface concentrations of coke precursors. Since carbon or coke formation on metals
apparently requires a critical ensemble of surface metal atoms and/or dissolution of carbon into
the bulk metal, introduction of modifiers that change ensemble sizes (e.g., Cu or S in Ni or Ru)
or that lower the solubility of carbon (e.g., Pt in Ni) can be effective in minimizing these forms
of deactivation.
For example, in a detailed STM study of submonolayers of Au on Ni(111), Besenbacher and
co-workers [71] found that the electron density of Ni atoms in the vicinity of Au atoms was
increased; from density functional theory (DFT) calculations they concluded that the strength of

70
carbon adsorption (and hence the tendency to form graphite) was decreased on next-nearest
neighbor Ni atoms; from studies of the effects of S adsorption on methane activation and graphite
formation on pure Ni, they were able to infer that the ensemble size needed for methane dissociation
is smaller than that for graphite formation. These fundamental insights were used in the design of a
0.3% Au-promoted 16% Ni/MgAl2O4 catalyst that loses no activity over 4000 h during steam
reforming of n-butane, while the corresponding unpromoted Ni catalyst loses about 5% of its initial
activity (see Figure 32). In contrast to the moderating effects of noble metal additives, addition of
0.5% Sn to cobalt substantially increases the rate of carbon filament formation from ethylene [72],
an effect desirable in the commercial production of carbon filament fibers.

Figure 32. Conversion of n-butane as a function of time during steam reforming in a
3% n–butane–7% hydrogen–3% water in helium mixture at a space velocity of 1.2 hí1.
The dashed curve shows the n-butane conversion for the Ni catalyst (16.8% Ni) and the
solid curve for the Au/Ni catalyst (16.4% Ni/0.3% Au). Reproduced from [71].
Copyright 1998, American Association for the Advancement of Science.
Coke deposition on oxide or sulfide catalysts occurs mainly on strongly acidic sites; accordingly
the rate of coking can be lowered by decreasing the acidity of the support. For example, silanation
of HY and HZSM-5 zeolites decreases their activities but improves catalyst life [236]. In steam
reforming, certain catalyst additives, e.g., MgO, K2O, or U3O8, facilitate H2O or CO2 adsorption
and dissociation to oxygen atoms, which in turn gasify coke precursors [8,60,70].
Similarly, for steam reforming catalysts used for light alcohol and oxygenate conversion, the
addition of partially reducible oxides, like ceria, in nickel perovskite (La1íxCexNiO3) catalysts [237]
or as a support for a cobalt catalyst [238], reduce the rate of carbon deposition. Alternatively, the
reaction atmosphere may be modified to increase the gasification rate by adding oxidizing reactants
(e.g., O2 and/or CO2) to reduce the rate of coke deposition [63]. This process is often described as
autothermal reforming because it tends to balance the endothermic steam reforming reactions with
exothermic reactions that make the process thermally neutral.
As in the case of poisoning (see below), there are certain reactor bed or catalyst geometries that
minimize the effects of coking on the reaction. For example, specific film-mass transport or pore
diffusion regimes favor coke or carbon deposition on either the outside or inside of the catalyst
pellet [239,240]. Choosing supports with relatively large pores minimizes pore plugging; choice of

71
large-diameter, mechanically-strong pellets avoids or delays reactor plugging. However, in view of
the rapidity at which coke and carbon can deposit on, plug, and even destroy catalyst particles, the
importance of preventing the onset of such formation cannot be overemphasized.
Reforming of naphtha provides an interesting case study of catalyst and process designs to avoid
deactivation by coking [8,196–198,241]. The classical Pt/Al2O3 catalyst is bifunctional; that is, the
metal catalyzes dehydrogenation, while the acid sites of the Al2O3 catalyze isomerization and
hydrocracking. Together, the two functions catalyze dehydrocylization and aromatization. Addition
of Re, Sn, or Ge, to Pt and sulfiding of the Pt–Re catalyst substantially reduce coke formation by
diluting large Pt ensembles that would otherwise produce large amounts of coke, while addition of
Sn and Ir improves selectivity for dehydrogenation relative to hydrogenolysis, the latter of which
leads to coke formation. Naphtha reforming processes are designed for (1) high enough H2 pressure
to favor gasification of coke precursors while minimizing hydrocracking, (2) maintenance of Cl
and S contents throughout the bed to ensure optimum acidity and coke levels, and (3) low enough
overall pressure to thermodynamically and kinetically favor dehydrogenation and dehydrocylization.
Accordingly, optimal process conditions are a compromise between case 1 and case 3. The
above-mentioned improvements in catalyst technologies, especially resistance to coking, have
enabled important process improvements, such as optimal operation at lower pressure; thus,
processes have evolved over the past two to three decades from conventional fixed-bed reactors at
high pressure (35 bar) using nonregenerative Pt catalysts to low pressure (3.5 bar), slowly
moving-bed, continuously regenerated units with highly selective Pt/Sn catalysts, resulting in
substantial economic benefits [198,241].

3.4. Prevention of Poisoning
Since poisoning is generally due to strong adsorption of feed impurities and since poisoned
catalysts are generally difficult or impossible to regenerate, it is best prevented by removal of
impurities from the feed to levels that will enable the catalyst to operate at its optimal lifetime. For
example, it is necessary to lower the feed concentration of sulfur compounds in conventional
methanation and Fischer–Tropsch processes involving base metal catalysts to less than 0.1 ppm in
order to ensure a catalyst lifetime of 1–2 years. This is typically accomplished using a guard bed of
porous ZnO at about 200 °C. In cracking or hydrocracking reactions on oxide catalysts, it is
important to remove strongly basic compounds, such as ammonia, amines, and pyridines, from the
feed; ammonia in some feedstocks, for example, can be removed by aqueous scrubbing. The
poisoning of catalysts by metal impurities can be moderated by selective poisoning of the unwanted
metal. For example, in catalytic cracking of nickel-containing petroleum feedstocks, nickel
sites, which would otherwise produce copious amounts of coke, are selectively poisoned by
antimony [242]. The poisoning of hydrotreating catalysts by nickel and vanadium metals can be
minimized by (1) using a guard bed of inexpensive Mo catalyst or graded catalyst bed with
inexpensive, low-activity Mo at the top (bed entrance) and expensive, high-activity catalyst at the
bottom (see Figure 33) and (2) by depositing coke prior to the metals, since these metal deposits
can be physically removed from the catalyst during regeneration [243].

72

Figure 33. Staged reactor system with decreasing pore size strategy for
hydrodemetalization (HDM)/hydrodesulfurization (HDS) of residuum. Reproduced
from [214]. Copyright 1993, Marcel Dekker.
It may be possible to lower the rate of poisoning through careful choice of reaction conditions
that lower the strength of poison adsorption [60] or by choosing mass-transfer-limiting regimes that
limit deposits to the outer shell of the catalyst pellet, while the main reaction occurs uninterrupted
on the interior of the pellet [239]. The manner in which the active catalytic material is deposited on
a pellet (e.g., uniformly or in an eggshell or egg yolk pattern) can significantly influence the life of
the catalyst [17,244].
An example of reducing catalyst poisoning (and oxidation) through process design has
been reported in a process patent for staged hydrocarbon synthesis via the Fischer–Tropsch
reaction [245]. While cobalt catalysts are favored because of their high activities and while it is
desirable to achieve high conversions of CO in the process, the one-pass conversion for cobalt is
limited by (1) its tendency to be oxidized at high partial pressures of product water observed at
high CO conversions and (2) its tendency under these conditions to form the oxygenated products
(e.g., alcohols and aldehydes) that poison or suppress its synthesis activity. One alternative is to
separate products and recycle the unused CO and H2, but this requires costly recompression and
separation of the oxygenates. Costly separation and/or poisoning can be prevented by operating a
first-stage reactor containing a cobalt catalyst to a moderately high conversion followed by reacting
the remaining CO and H2 in a second stage to above 95% conversion on an iron catalyst, which is
not sensitive to the oxygenates and which shifts some of the product water to H2 and CO2, thus
minimizing its hydrothermal degradation.
An example of reducing catalyst poisoning through catalyst design occurs in abatement of
emissions for automotive and motorcycle engines [18,212]. Application of an alumina or zeolite
coating, or alternatively preparing the active phase in a sublayer, provides a diffusion barrier that
prevents or slows the access of poisons from the fuel or oil (e.g., phosphorus and/or zinc from
lubricating oil or corrosion products) to the catalyst surface. The principle is to optimize the pore
size distribution of the diffusion barrier to provide access to the catalytic phase of relatively small
hydrocarbon, CO, NO, and O2 molecules, while preventing access of larger molecules, such as
from lubricating oil and/or particulates.

73
Finally, another strategy that has been employed to reduce the impact of poisoning, particularly
for sulfur, is the inclusion of traps or “getters” as part of the catalyst. These species, including rare
earth oxides of thulium (Tm) [50] or Ce [51] and simple zinc oxide, essentially act as sacrificial
stoichiometric reactants to protect the active metal by preferentially adsorbing the poison. These
traps can extend the catalyst life, but because they are not catalytic as they perform, they are
necessarily temporary agents if the poison remains in the feed to the process.

3.5. Prevention of Sintering
Since most sintering processes are irreversible or are reversed only with great difficulty, it is
important to choose reaction conditions and catalyst properties that avoid such problems.
Metal growth is a highly activated process; thus, by choosing reaction temperatures lower than
0.3–0.5 times the melting point of the metal, rates of metal sintering can be greatly minimized. The
same principle holds true in avoiding recrystallization of metal oxides, sulfides, and supports. Of
course, one approach to lowering reaction temperature is to maximize activity and surface area of
the active catalytic phase.
Although temperature is the most important variable in the sintering process, differences in reaction
atmosphere can also influence the rate of sintering. Water vapor, in particular, accelerates the
crystallization and structural modification of oxide supports. Accordingly, it is vital to minimize the
concentration of water vapor in high temperature reactions on catalysts containing high surface
area supports.
Besides lowering temperature and minimizing water vapor, it is possible to lower sintering rates
through addition of thermal stabilizers to the catalyst. For example, the addition of higher melting
noble metals (such as rhodium or ruthenium) to a base metal (such as nickel) increases the thermal
stability of the base metal [106]. Addition of Ba, Zn, La, Si, and Mn oxide promoters improves
the thermal stability of alumina [246]. These additives can affect product selectivity, but generally
positively toward desired products, and always through extending the productive life of the
catalysts [8].
Designing thermally stable catalysts is a particular challenge in high temperature reactions, such as
automotive emissions control, ammonia oxidation, steam reforming, and catalytic combustion. The
development of thermally stable automotive catalysts has received considerable attention, thus providing
a wealth of scientific and technological information on catalyst design (e.g., Refs. [8,215–222]). The
basic design principles are relatively simple: (1) utilize thermally and hydrothermally stable
supports, e.g., high-temperature į- or ș-aluminas or alkaline-earth or rare-earth oxides that form
ultrastable spinels with Ȗ-alumina; (2) use PdO rather than Pt or Pt–Rh for high temperature
converters, since PdO is considerably more thermally stable in an oxidizing atmosphere because of
its strong interaction with oxide supports; and (3) use multilayer strategies and/or diffusion barriers
to prevent thermally induced solid-state reactions (e.g., formation of Rh aluminate) and to moderate
the rate of highly exothermic CO and hydrocarbon oxidations. For example, a typical three-way
automotive catalyst may contain alkaline-earth metal oxides (e.g., BaO) and rare-earth oxides (e.g.,
La2O3 and CeO2), for stabilizing Pt and/or PdO on alumina, and ZrO2 as a thermal stabilizer for the

74
CeO2 (an oxygen storage material) and as a noninteracting support for Rh in a separate layer or in a
separate phase in a composite layer (see Figure 34).

Figure 34. Conceptual design (by C. H. Bartholomew) of an advanced three-way
catalyst for auto emissions control. Catalyst layer 1 is wash-coated first onto the
monolithic substrate and consists of (a) well-dispersed Pd, which serves to oxidize
CO/hydrocarbons and to reduce NO and (b) CeO2/ZrO2 crystallites (in intimate contact
with Pd), which store/release oxygen respectively, thereby improving the performance
of the Pd. Catalyst layer 2 (added as a second wash coat) is a particle composite of
Rh/ZrO2 (for NO reduction) and Pt/La2O3–BaO/Al2O3 (with high to moderately-high
activity for oxidation of CO and hydrocarbons). A thin (50–80 ȝm) coat of Al2O3,
deposited over catalyst layer 2, acts as a diffusion barrier to foulants and/or poisons.
Both the Al2O3 layer and catalyst layer 2 protect the sulfur-sensitive components of
catalyst layer 1 from poisoning by SO2..
Often, ideal metal dispersions require metal nanoparticles to be distributed closely together, but
these particles are thermodynamically unstable on the surface and undergo rapid sintering, as
described in Section 2.3 above. Recently, in an attempt to reduce sintering rates, researchers have
attempted to stabilize the metal nanoparticles by first dispersing them on a support, encapsulating
them in the same or another metal oxide, and then opening porosity to the particles (e.g., [247,248]).
These approaches have met with varying degrees of success, but point to promising new areas of
synthesis techniques that have the potential to reduce or to eliminate deactivation by sintering.

3.6. Prevention of Mechanical Degradation
While relatively few studies have focused on this topic, there are nevertheless principles that guide
the design of processes and catalysts in preventing or minimizing mechanical degradation (see
Table 19). In terms of catalyst design, it is important to (1) choose supports, support additives, and
coatings that have high fracture toughness, (2) use preparation methods that favor strong bonding
of primary particles and agglomerates in pellets and monolith coatings, (3) minimize (or rather

75
optimize) porosity (thus maximizing density), and (4) use binders, such as carbon, to facilitate
plastic deformation and thus protect against brittle fracture. Processes (and to some extent
preparation procedures) should be designed to minimize (1) highly turbulent shear flows or
cavitation that lead to fracture of particles or separation of coatings, (2) large thermal gradients or
thermal cycling leading to thermal stresses, and (3) formation of chemical phases of substantially
different densities or formation of carbon filaments leading to fracture of primary particles and
agglomerates. Nevertheless, thermal or chemical tempering can be used in a controlled fashion to
strengthen catalyst particles or agglomerates.
Examples of catalyst design to minimize attrition can be found in the recent scientific [230,231]
and patent [232–235] literature focusing on the Fischer–Tropsch synthesis in slurry reactors. These
studies indicate that (1) spray drying of particles improves their density and attrition resistance;
(2) addition of silica and/or alumina into titania improves its attrition resistance, while addition of only
2000–3000 ppm of titania to Ȗ-alumina improves alumina’s attrition resistance; and (3) preformed
alumina spheres promoted with La2O3 provide greater attrition resistance relative to silica. Increasing
attrition resistance is apparently correlated with increasing density [230,231,235]. According to
Singleton and co-workers [235], attrition resistance of Co/Al2O3 is improved when the Ȗ-alumina
support is (1) formed from synthetic boehmite having a crystallite diameter of 4–5 nm and (2) is
pretreated in acidic solution having a pH of 1–3 (see Figure 35); moreover, attrition resistance
decreases in the order Co/Al2O3 > Co/SiO2 > Co/TiO2 and is greater for catalysts prepared by aqueous
versus nonaqueous impregnation.

Figure 35. Effect of solution pH on the attrition resistance of 70-ȝm Ȗ-Al2O3 particles
measured in jet-cup tests [235]. The % increase in fines is defined at the % increase of
particles of less than 11 ȝm.
4. Regeneration of Deactivated Catalysts
Despite our best efforts to prevent it, the loss of catalytic activity in most processes is inevitable.
When the activity has declined to a critical level, a choice must be made among four alternatives:
(1) restore the activity of the catalyst, (2) use it for another application, (3) reclaim and recycle the
important and/or expensive catalytic components, or (4) discard the catalyst. The first alternative

76
(regeneration and reuse) is almost always preferred; catalyst disposal is usually the last resort,
especially in view of environmental considerations.
The ability to reactivate a catalyst depends upon the reversibility of the deactivation process. For
example, carbon and coke formation is relatively easily reversed through gasification with hydrogen,
water, or oxygen. Sintering on the other hand is generally irreversible, although metal redispersion
is possible under certain conditions in selected noble metal systems. Some poisons or foulants
can be selectively removed by chemical washing, mechanical treatments, heat treatments, or
oxidation [249,250]; others cannot be removed without further deactivating or destroying
the catalyst.
The decision to regenerate/recycle or discard the entire catalyst depends largely on the rate of
deactivation. If deactivation is very rapid, as in the coking of cracking catalysts, repeated or
continuous regeneration becomes an economic necessity. Precious metals are almost always
reclaimed where regeneration is not possible. Disposal of catalysts containing nonnoble heavy
metals (e.g., Cr, Pb, or Sn) is environmentally problematic and should be a last resort; if disposal is
necessary, it must be done with great care, probably at great cost. Accordingly, a choice to discard
depends upon a combination of economic and legal factors [250]. Indeed, because of the scarcity of
landfill space and an explosion of environmental legislation, both of which combine to make
waste-disposal prohibitively expensive, there is a growing trend to regenerate or recycle spent
catalysts [251,252]. A sizeable catalyst regeneration industry benefits petroleum refiners by helping
to control catalyst costs and to limit liabilities [253,254]; it provides for ex situ regeneration of
catalyst and recovery/recycling of metals, e.g., of cobalt, molybdenum, nickel, and vanadium from
hydroprocessing catalysts [251].
Consistent with its importance, the scientific literature treating catalyst regeneration is significant
and growing (includes nearly 1000 journal articles since 1990). Regeneration of sulfur-poisoned
catalysts has been reviewed by Bartholomew and co-workers [28]. Removal of coke and carbon from
catalysts has received attention in reviews by Trimm [59,250], Bartholomew [60], and Figueiredo [1].
Redispersion of sintered catalysts has been discussed by Ruckenstein and Dadyburjor [101],
Wanke [102], and Baker and co-workers [103]. Useful case studies of regeneration of
hydrotreating [255] and hydrocarbon-reforming catalysts [256] have also been reported. The
proceedings of the 9th International Symposium on Catalyst Deactivation (2001) contains 12 papers
treating catalyst regeneration [257]. Regeneration, recycling, and disposal of deactivated heterogeneous
catalysts have been reviewed briefly by Trimm [250].
The patent literature treating catalyst regeneration/reactivation is enormous (more than
17,000 patents); the largest fraction of this literature describes processes for regeneration of
catalysts in three important petroleum refining processes, i.e., FCC, catalytic hydrocarbon
reforming, and alkylation. However, a significant number of patents also claim methods for
regenerating absorbents and catalysts used in aromatization, oligomerization, catalytic combustion,
SCR of NO, hydrocracking, hydrotreating, halogenation, hydrogenation, isomerization, partial
oxidation of hydrocarbons, carbonylations, hydroformylation, dehydrogenation, dewaxing,
Fisher–Tropsch synthesis, steam reforming, and polymerization.

77
Conventional methods for regenerating (largely in situ) coked, fouled, poisoned, and/or sintered
catalysts in some of these processes and representative examples thereof [258–297] are
summarized in Table 20, while the basic principles and limitations involved in regeneration of
coked, poisoned, and sintered catalysts are briefly treated in the subsections that follow.

4.1. Regeneration of Catalyst Deactivated by Coke or Carbon
Carbonaceous deposits can be removed by gasification with O2, H2O, CO2, and H2. The temperature
required to gasify these deposits at a reasonable rate varies with the type of gas, the structure and
reactivity of the carbon or coke, and the activity of the catalyst. Walker and co-workers [302] reported
the following order (and relative magnitudes) for rates of uncatalyzed gasification at 10 kN/m3 and
800 °C: O2 (105) > H2O (3) > CO2 (1) > H2 (3 × 10í3). However, this activity pattern does not apply
in general for other conditions and for catalyzed reactions [1]. Nevertheless, the order of decreasing
reaction rate of O2 > H2O > H2 can be generalized.
Rates of gasification of coke or carbon are greatly accelerated by the same metal or metal oxide
catalysts upon which carbon or coke deposits. For example, metal-catalyzed coke removal with H2
or H2O can occur at a temperature as low as 400 °C [1]; ȕ-carbon deposited in methanation can be
removed with H2 over a period of a few hours at 400–450 °C and with oxygen over a period of
15–30 min at 300 °C [60]. However, gasification of more graphitic or less reactive carbons or coke
species in H2 or H2O may require temperatures as high as 700–900 °C [1], conditions, of course,
that result in catalyst sintering.
Because catalyzed removal of carbon with oxygen is generally very rapid at moderate
temperatures (e.g., 400–600 °C), industrial processes typically regenerate catalysts deactivated by
carbon or coke in air. Indeed, air regeneration is used to remove coke from catalysts in catalytic
cracking [81], hydrotreating processes [255], and catalytic reforming [256].
One of the key problems in air regeneration is avoiding hot spots or overtemperatures which
could further deactivate the catalyst. The combustion process is typically controlled by initially
feeding low concentrations of air and by increasing oxygen concentration with increasing carbon
conversion [255,303]; nitrogen gas can be used as a diluent in laboratory-scale tests, while steam is
used as a diluent in full-scale plant operations [303]. For example, in the regeneration of
hydrotreating catalysts, McCulloch [255] recommends keeping the temperature at less than 450 °C
to avoid the Ȗ- to Į-alumina conversion, MoO3 sublimation, and cobalt or nickel aluminate
formation, which occur at 815, 700, and 500–600 °C respectively.
Because coke burn-off is a rapid, exothermic process, the reaction rate is controlled to a large
extent by film heat and mass transfer. Accordingly, burn-off occurs initially at the exterior surface
and then progresses inward, with the reaction occurring mainly in a shrinking shell consistent with
a “shell-progressive” or “shrinking-core” model, as illustrated in Figure 36 [304]. As part of this
same work, Richardson [304] showed how experimental burn-off rate data can be fitted to various
coking transport models, e.g., parallel or series fouling. Burn-off rates for coke deposited on
SiO2/Al2O3 catalysts were reported by Weisz and Goodwin [305]; the burning rate was found to be
independent of initial coke level, coke type, and source of catalyst.

Problem/cause

Method(s) of regeneration/phenomena
studied/conclusions

Deactivation by coke, carbon
(1) ZSM-5 catalyst for light olefin oligomerization
Alkene aromatization
containing 2–3% coke is treated in 8–10% steam/air
oligomerization
Catalyst fouling by condensation of
mixture (1300 kPa, 93 °C inlet) in a fluidized bed
Zeolites, esp. ZSM-5, -22, heavy oligomers to coke
(2) A coked crystalline alumogallosilicate is contacted with
-23, beta-zeolite, ferrierite
oxygen at a concentration of 0.05–10 vol%, 420–580 °C,
and 300–4000 hí1
(1) Coked zeolite is regenerated in liquid phase (P > 3500 kPa)
fluid bed with H2 in two steps: (a) at reaction temperature
(20–50 °C) and (b) at 25 °C above reaction temperature
Alkylation of isoparaffins Rapid catalyst deactivation due to coke
(2) Coked Pd- and Pt/Y-zeolite catalysts containing
on solid catalysts
formation; unacceptable product
10–13% coke are regenerated in either air or H2; H2
a
Sulfated zirconia, USY , quality, and thermal degradation of
treatment enables removal of most of the coke at low to
Nafion, silicalite, ZSM-5 catalyst during regeneration
moderate temperatures; higher temperatures are required
for air (3) USY and other zeolites are regenerated in
supercritical isobutane

Deactivation mechanism
Process/Reaction
Catalyst

[260–263]

[258,259]

Ref.

Table 20. Conventional Methods for and Representative Examples of Catalyst Regeneration from Scientific and Patent Literatures.

78

Dehydrogenation of
propane and butane
Cr2O3/Al2O3, Cr2O3/ZrO2,
FeO/K/MgO, Pt/Al2O3,
Pt–Sn/Al2O3,
Pt–Sn/KL-zeolite

Catalytic reforming
of naphtha
Pt/Al2O3 promoted with
Re, Sn, Ge, or Ir

Deactivation mechanism
Process/Reaction
Catalyst
Method(s) of regeneration/phenomena
studied/conclusions

Deactivation by coke, carbon
(1) Coke on Pt bimetallic reforming catalyst is removed
off-stream in a fixed or moving bed at 300–600 °C,
followed by oxychlorination (350–550 °C)
(2) Coke on Pt/zeolite is removed in halogen-free oxygencontaining gas at T < 415 °C
(3) Sintering during oxidation of coke on Pt–Ir/Al2O3
Poisoning and fouling by coke produced catalyst can be minimized by low regeneration
temperatures
by condensation of aromatics and
(4) Study of influence of heating rate, temperature, and
olefins
time on structural properties of regenerated Pt–Sn/Al2O3
(5) Study of effects of Cl, Sn content, and regeneration
sequence on dispersion and selectivity of Pt–Sn/Al2O3
(6) Regenerated Pt–Re/Al2O3 is more stable than the fresh
catalyst in n-heptane conversion and more selective for
toluene
(1) Temperatures gradients were measured during burn off
of coke formed on a chromia–alumina catalyst during
butene dehydrogenation; data were used in developing a
Catalyst activity is low due to
mathematical model for predicting temperatures and
equilibrium limitations and build-up of
coke profiles
product H2; rapid loss of activity occurs
(2) Coked supported palladium catalyst used in the
due to coke formation
dehydrogenation of dimethyltertrahydronaphthalenes to
dimethylnaphthalenes is reactivated with an organic polar
solvent at a temperature below 200 °C

Problem/cause

Table 20. Cont.

[270,271]

[264–269]

Ref.

79

Hydrocracking of
heavy naphtha
CoMo, NiW, MoW on
Al2O3 or SiO2–Al2O3; Pt
or Pd on Y-zeolite,
mordenite, or ZSM-5

Fluid catalytic cracking
(FCC) of heavy
hydrocarbons
USY or REO-Yb in
silica matrix

Fischer–Tropsch
synthesis
Co/Al2O3

Deactivation mechanism
Process/Reaction
Catalyst
Method(s) of regeneration/phenomena
studied/conclusions

Deactivation by coke, carbon
(1) Carbidic surface carbon deposited on cobalt can be
largely removed in hydrogen at 170–200 °C and in steam at
Loss of activity due to blocking of sites 300–400 °C
by carbon overlayers and heavy
(2) Slurry-phase cobalt catalysts may lose 50% activity
hydrocarbons
during reaction over a period of a few days; the activity can
be rejuvenated in situ by injecting H2 gas into vertical draft
tubes inside the reactor
(1) Process and apparatus for increasing the coke burning
capacity of FCC regenerators; auxiliary regenerator
partially burns off the coke at turbulent or fast fluidizedRapid loss of activity due to poisoning
bed conditions
of acid sites and blocking of small
(2) Multistage fluidized-bed regeneration of spent FCC
zeolite pores by coke
catalyst in a single vessel by incorporating two relatively
dense phase fluidized beds beneath a common dilute
phase region
(1) Regeneration of noble metal/zeolite via progressive
partial removal of carbonaceous deposits under controlled
Loss of activity due to poisoning of acid oxidizing conditions to maximize sorption of a probe
sites and blocking of small zeolite pores molecule while minimizing metal sintering
by coke
(2) Regeneration of noble metal/zeolite in air at about
600 °C, followed by a mild treatment in aqueous ammonia
to improve catalytic activity

Problem/cause

Table 20. Cont.

[277,278]

[275,276]

[272–274]

Ref.

80

Methanol to olefins or
gasoline
Silica– alumina,
Y-zeolite, ZSM-5,
other zeolites, and
aluminophosphate
molecular sieves

Hydrotreating of gas oil
Al2O3-supported Mo and
CoMo, NiMo, NiCoMo,
MoW, NiW

Deactivation mechanism
Process/Reaction
Catalyst
Method(s) of regeneration/phenomena
studied/conclusions

Deactivation by coke, carbon
(1) TPO studies of oxidative regeneration of CoMo and
NiW HDS catalysts; sulfur is removed at 225–325 °C,
carbon at 375–575 °C. Redispersion of NiW was observed
by EXAFS
(2) Physicochemical changes in CoMo and NiCoMo HDS
catalysts during oxidative regeneration, including
redispersion of Co, Ni, and Mo oxides and surface area
loss, were examined
(3) Changes in NiMo catalyst structure and coke
Loss of activity due to formation of
composition during reaction and regeneration were
types I, II, and III coke on metal sulfide
examined and correlated
and alumina surfaces and in pores
(4) Properties of NiMo catalyst deactivated during shale oil
hydrogenation and regenerated in O2 or H2 were examined.
Regeneration in 1.6% O2 was more effective than that in
5% H2. Ni aluminate spinel was observed after burn off
(5) Hard and soft cokes formed on CoMo catalysts during
HDS of gas oil were characterized. At low coke levels,
hard coke was more easily removed in H2 than in O2
(6) Spent catalysts are washed with solvent and contacted
with steam at about 600 °C
(1) Kinetics of coke burnoff from a SAPO-34 used in
converting methanol to olefins were studied; kinetics are
Severe coking and deactivation of
strongly dependent on the nature of the coke. Kinetics are
silica–alumina and Y-zeolite catalysts
slowed by strong binding of coke to acid sites
observed during high conversions of
(2) ZSM-34 catalyst used in conversion of methanol to
MeOH; also substantial coking of
light olefins is effectively regenerated in H2-containing gas;
ZSM-5, other zeolites, and
this approach avoids the formation of catalyst-damaging
aluminophosphate molecular sieves
products such as steam that would be formed during burn
off in air

Problem/cause

Table 20. Cont.

[281,282]

[279,280,
294–297]

Ref.

81

Problem/cause

(1) Poisoning of acid sites by
N- containing compounds.
(2) Deposition of Ni and V metals on
acid sites which change selectivity and
decrease activity

Poisoning of metal sites by arsenic,
sulfur, and other poisons

Pore-mouth poisoning and blockage by
Ni, V, and Fe sulfides present in feed as
organometallics

Deactivation mechanism
Process/Reaction
Catalyst

FCC of residuum
USY or REO-Y
in silica matrix

Hydrogenation or
dechlorination
Ni/SiO2, Pd/Al2O3

Hydrotreating
of residuum
Al2O3-supported Mo
and CoMo

Poisoning
(1) Organometallic solutions of Sb and Bi are added to
process stream to passivate Ni by forming inactive Ni–Sb
and Ni–Bi species
(2) V metal deposits are trapped by reaction with
magnesium orthosilicate to form an unreactive magnesium
vanadium silicate
(3) Spent metal-contaminated catalyst is demetallized by
chlorinating and washing, followed by contacting with
NH4F and one antimony compound
(4) Metal-contaminated catalyst is contacted with an
aqueous solution of a carboxylic acid (e.g., formic, acetic,
citric, or lactic acid)
(5) Metal-contaminated catalyst is contacted with HCl,
HNO3, or H2SO4
(6) Metal-contaminated catalyst is contacted with reducing
CO gas to form gaseous metal carbonyls that are separated
from the catalyst
(1) Regeneration of Ni/SiO2 catalyst poisoned by thiophene
using a sequence of oxidation–reduction treatments at low
PO2 and 1 atm H2 respectively
(2) Regeneration in dilute hypochlorite solution of a
Pd/Al2O3 catalyst deactivated during the aqueous-phase
dechlorination of trichloroethylene in the presence of
sulfite or HSí ions present in ground water
(1) Regeneration of catalysts containing V, Ni, or Fe by
contacting with H2O2 solution and organic acid
(2) Following removal of coke by air or solvent wash,
catalyst is acid leached to remove undesired metals

Method(s) of regeneration/phenomena
studied/conclusions

Table 20. Cont.

[287,288]

[285,286]

[281,282,298–301]

Ref.

82

Hydrocracking of
heavy naphtha
CoMo, NiW, MoW on
Al2O3 or SiO2–Al2O3;
Pt or Pd on Y-zeolite,
mordenite, or ZSM-5
Hydrotreating of gas
oil and residuum
Al2O3-supported Mo
and CoMo

Catalytic reforming
of naphtha
Pt/Al2O3 promoted with
Re, Sn, Ge, or Ir; Pt/KLzeolite

Deactivation mechanism
Process/Reaction
Catalyst
Method(s) of regeneration/phenomena
studied/conclusions

USY: ultrastable Y-zeolite. b REO-Y: rare-earth exchanged Y-zeolite.

[292,293]

(1) Oxidative regeneration of hydroprocessing catalyst at
600 °C optimizes surface area and Mo dispersion
(2) Oxidative regeneration in several steps with a final
oxidation at 500–600 °C to restore residual catalyst activity

Sintering of Mo and Co sulfides causing
formation of large sulfide crystals and
loss of active surface area
a

[291]

Redispersion of noble metals on molecular sieves including
silica-aluminates, ALPOS, SAPOS

[266,269,289,290]

Ref.

Sintering of noble metal causing
formation of large metal crystallites and
loss of active surface area

Thermal degradation
(1) Redispersion of Pt–Ir bimetallic catalysts using a wet
HCl/air treatment, since the conventional oxychlorination is
not effective
Sintering of Pt causing formation of
(2) Redispersion of Pt/KL-zeolite using wet HCl/air
large metal crystallites and loss of
treatment followed by brief calcination and reduction (3)
active surface area
Redispersion of Pt–Re/Al2O3 in Cl2 and O2 (4)
Redispersion of supported Pt, other noble metals, and Ni in
Cl2 and O2

Problem/cause

Table 20. Cont.

83

84

Figure 36. Shell-progressive regeneration of fouled pellet Reproduced from [304].
Copyright 1972, American Chemical Society.
4.2. Regeneration of Poisoned Catalysts
Much of the previous literature has focused on regeneration of sulfur-poisoned catalysts used
in hydrogenations and steam reforming. Studies of regeneration of sulfur-poisoned Ni, Cu, Pt, and Mo
with oxygen/air, steam, hydrogen, and inorganic oxidizing agents have been reported [28].
Rostrup-Nielsen [306] indicates that up to 80% removal of surface sulfur from Mg- and
Ca-promoted Ni, steam reforming catalysts occurs at 700 °C in steam. The presence of both SO2
and H2S in the gaseous effluent suggests that the following reactions occur:
Ni-S + H2OĺNiO + H2S

(5)

H2S + 2H2OĺSO2 + 3H2

(6)

Although this treatment is partially successful in the case of low-surface-area steam
reforming catalysts, the high temperatures required for these reactions would cause sintering of most
high-surface-area nickel catalysts.
Regeneration of sulfur-poisoned catalysts, particularly base metal catalysts, in air or oxygen has
been largely unsuccessful. For example, the treatment of nickel steam-reforming catalysts in steam
and air results in the formation of sulfates, which are subsequently reduced back to nickel sulfide
upon contact with hydrogen. Nevertheless, sulfur can be removed as SO2 at very low oxygen
partial pressures, suggesting that regeneration is possible under carefully controlled oxygen
atmospheres, including those provided by species such as CO2 or NO that dissociate to oxygen.
Apparently, at low oxygen pressures, the oxidation of sulfur to SO2 occurs more rapidly than the
formation of nickel oxide, while at atmospheric pressure the converse is true, i.e., the sulfur or
sulfate layer is rapidly buried in a nickel oxide layer. In the latter circumstance, the sulfur atoms
diffuse to the nickel surface during reduction, thereby restoring the poisoned surface. Regeneration
of sulfur-poisoned noble metals in air is more easily accomplished than with steam, although it is
frequently attended by sintering. Regeneration of sulfur-poisoned nickel catalysts using hydrogen is

85
impractical because (1) adsorption of sulfur is reversible only at high temperatures at which
sintering rates are also high and (2) rates of removal of sulfur in H2 as H2S are slow even at
high temperature.
Inorganic oxidizing agents such as KMnO4 can be used to oxidize liquid phase or adsorbed sulfur to
sulfites or sulfates [16]. These electronically shielded structures are less toxic than the unshielded
sulfides. This approach has somewhat limited application, i.e., in partial regeneration of metal catalysts
used in low temperature liquid-phase hydrogenation reactions or in liquid-phase destruction of
chlorinated organic compounds. For example, Lowrey and Reinhard [286] reported successful
regeneration in dilute hypochlorite solution of a Pd/Al2O3 catalyst deactivated during the aqueous-phase
dechlorination of trichloroethylene (TCE) in the presence of sulfite or HSí ions. These poisons are
formed by sulfate-reducing bacteria present in natural groundwater and are apparently adsorbed on
the alumina or Pd surfaces more strongly than sulfate ions. Figure 37 illustrates how readily the
poisoned catalyst is regenerated by dilute hypochlorite solutions; indeed, it is evident in Figure 37b
that regeneration every 5–10 days successfully maintains the catalytic conversion of TCE around
25% (a value only slightly less than that observed for reaction in distilled water).

Figure 37. Effect of regeneration (R) with hypochlorite of Pd/Al2O3 catalysts used for
aqueous phase dechlorination of trichloroethylene in the presence of HSí/SO32.
Reproduced from [286]. Copyright 1992, American Chemical Society.

86

4.3. Detailed Case Study on Regeneration of Selective Catalytic Reduction (SCR) Catalysts
4.3.1. Introduction to SCR: Key to Abatement of NOx from Coal Utility Boilers
NOx, generally defined as NO and NO2, emissions from coal utility boilers (approximately 30% of
total NOx emissions in the U.S.) contribute substantially to the formation of acid rain and photochemical
smog, which in turn damage human health, property, agriculture, lakes, and forests. Selective
catalytic reduction (SCR) technology has been used in utility boilers since the 1980s in Japan and
Europe in response to stringent NOx removal regulations. By 2000, SCR systems had been installed
in coal-fired boilers totaling roughly 25 and 55 GW in Japan and Europe respectively [307,308].
Equivalent stringent NOx abatement regulations were enacted later in the U.S. by the EPA, including
(1)
(2)
(3)
(4)

the 1990 ARP and OTC mandates, requiring states to reduce NOx emissions by 80%;
the 1995 OTC-Phase 1 requiring Reasonably Available Control Technology (RACT);
the 1998 NOx SIP Call setting up a regional cap-and-trade program for 20 eastern states
based on an equivalent NOx emission rate of 0.15 lb/106-Btu; and
the 2005 Clean Air Interstate Rule (CAIR) requiring all states to meet Best Available
Retrofit Technology (BART) for existing plants, equivalent to emission rates of less
than 0.05–0.10 lb/106-Btu [309,310].

By 2006, about 100 GW of coal-fired steam boilers in the U.S. used SCR. Presently, the U.S.
has about 140 GW [309] of coal-boiler SCR capacity; world-wide, an estimated 300 GW of
coal-boiler SCR is in operation.
Prior to the more recent stringent U.S. emissions regulations, boiler and engine manufacturers
successfully reduced NOx emissions by 30–60% using modifications to combustion processes,
including reducing excess air, adding two-stage combustion features, altering burner design, etc.
However, meeting the new reduction targets of 80–90% is, in general, only possible through
catalytic after-treatment (SCR). Given ever more restrictive NOx emission standards and the fact
that worldwide power production from coal could double or triple in the next decade to an
estimated 1500 GW [311], total installed SCR unit capacity is expected to grow commensurately,
providing continued investment and design challenges in this area.
4.3.2. Selective Catalytic Reduction of NOx
4.3.2.1. Reaction Chemistry and Preferred Catalysts
Selective catalytic reduction (SCR) is a process in which a reducing agent, typically NH3, reacts
selectively with the NOx to produce N2 without consumption of the excess O2 present in the flue
gas. Desirable stoichiometric reactions for SCR of NO and NO2 (Equations 7 and 8) occur with
high activity and selectivity to N2 within a narrow temperature window of 300–400 °C on preferred
commercial catalysts.
4NH3 + 4NO + O2ĺ4N2 + 6H2O

(7)

87
4NH3 + 2NO2 + O2ĺ3N2 + 6H2O

(8)

Undesirable side reactions include oxidations of SO2 (present in the flue gas) and the reducing
agent NH3. While only a small fraction of the SO2 present in the flue gas is catalytically oxidized to
SO3, this acid precursor either corrodes downstream heat-exchange surfaces or reacts with NH3 to
form ammonium sulfates, which in turn can foul catalyst and/or heat exchange surfaces. Oxidation
of NH3 to either NO or N2 may also occur at temperatures above 400 °C.
A typical commercial vanadia catalyst consists of 1 wt% V2O5 and 10 wt% WO3 (alternatively
6 wt% MoO3) supported on high-surface-area TiO2 (mostly anatase, 60–80 m2/g). TiO2 has the
decided advantage over Al2O3 as a support, since the former stabilizes the active vanadia species
and does not form a bulk sulfate in the presence of SO2-containing flue gases; thus TiO2 promotes
activity and extends catalyst life. WO3 and MoO3 prevent the transformation of anatase to rutile;
they reside on basic sites of TiO2, blocking adsorption of SO3, thereby preventing sulfation of the
support. Additionally, WO3 and MoO3 increase Brønsted acidity, promoting NOx reduction while
lowering SO2 oxidation rate. Commercial vanadia-titania catalysts are typically supplied in the
form of extruded monoliths or plates (see Figure 38), forms which minimize pressure drop [8].

a.

b.

Figure 38. SCR catalyst support geometries: (a) extruded ceramic monolith; and (b) plate.
Reproduced from [8]. Copyright 2006, Wiley-Interscience.
4.3.2.2. SCR Process Options
Two process options in terms of SCR reactant placement have found broad use for SCR units
installed in coal-fired plants:
(1)
(2)

the high dust unit (HDU) involving placement of the SCR unit after the economizer
and prior to the air heater, particulate collector, and SO2 scrubber; and
the tail end unit (TEU) involving placement of the SCR unit following the
SO2 scrubber.

The HDU is used more widely in the U.S. and the TEU more frequently in Europe and Japan.
The HDU has the advantage of providing flue gas to the SCR unit at its ideal temperature range
of 300–400 °C and disadvantages of
(1)
(2)

deactivation of the catalyst due to erosion, fouling, and poisoning by fly ash thereby
limiting its useful life to about 3–4 years;
large monolith channel design to limit plugging by fly ash, but which also limits the
amount of active catalyst per reactor volume; and

88
(3)

requirement for a low activity catalyst to limit oxidation of SO2 to SO3 and the attendant
formation of ammonium sulfates which foul and corrode downstream heat exchangers.

The TEU enables use of a smaller volume of high activity catalyst with small diameter channels,
since particulates and SO2 have already been removed upstream; moreover, since deactivation rate
is much lower due to the absence of fly ash and other poisons, catalyst life is substantially extended
(i.e., to 15–20 years). A significant disadvantage is that the outlet scrubber gas, which is only about
120 °C, must be reheated to at least 200–250 °C for the SCR to occur at reasonable rates. The
energy cost of reheating only 100 °C can be as much as 4–6% of the boiler capacity, unless a
regenerative heat exchanger is used. In addition, the SCR catalyst must be designed to operate at
significantly lower temperatures (200–290 °C relative to a typical 300–400 °C for an HDU).
Given the long life of the TEU catalyst, no regeneration is necessary. However, regeneration of
the HDU catalyst is highly desirable, since the regeneration cost is significantly lower than the cost
of a new catalyst. With this background, further discussion focuses on the deactivation and
regeneration of the HDU catalyst.
4.3.3. Catalyst Deactivation, Rejuvenation, and Regeneration
4.3.3.1. Catalyst Deactivation
SCR catalysts have typical process lifetimes around 2–7 years, depending upon their application
and placement in a power plant or other such facility. The principal causes of SCR catalyst
deactivation [8,312] are fourfold:
(1)
(2)
(3)
(4)

fouling/masking of (deposition of solids on) catalyst surfaces, pores, and channels by fly
ash components (e.g., sulfates and phosphates of Ca, K, and Na) or ammonium bisulfate;
chemical poisoning of active sites by elements present in upstream lubricants or
originating in the fuel such as As, Se, and P and alkali and alkaline earth metals;
hydrothermal sintering of the titania, especially as a result of high-temperature
excursions; and
abrasion or erosion by fly ash.

Erosion, fouling, and masking from fly ash and poisoning by As and alkali metals are specific to
SCR catalysts installed near the hot, high-particulate side of a coal-fired boiler, accounting for the
significantly lower catalyst life of 2–4 years for this configuration.
Formation of ammonium bisulfate depends on flue gas temperature, SO3 concentration and NH3
concentration [313]. Deposition of ammonium bisulfate is more likely to occur in catalyst pores at
lower reactor temperatures in low-dust or tail-end (TEU) SCR units and on cooler surfaces of heat
exchangers. Figure 39a shows typical activity loss versus time performance for a set of commercial
V/Ti catalysts tested in a DOE pilot SCR unit installed in a slip-stream near the exit of a coal-fired
boiler (HDU location) using high sulfur, Eastern U.S. coals; 20% of the initial catalyst activity is
lost in about 14,000 h (1.6 years); however, the plant will not shut down until 50–60% of the initial
activity has been lost (around 3–4 years). Activity and NH3 slip are plotted against NH3/NO ratio
for the same catalysts in Figure 39b. To maintain NH3 slip (exit NH3 concentration) below a target

89
maximum of 2–5 ppm (2 is highly preferred), the NH3/NO ratio must be maintained near 0.8; under
these conditions NO conversion is about 88%.

Figure 39. (a) Catalyst activity (k/ko) vs. time; (b) Typical SCR performance.
Reproduced from [313].

Prevention of deactivation requires optimal choices of catalyst design and process conditions.
Abrasion, fouling, and/or poisoning by fly ash can be prevented by installation of a hot-side
electrostatic precipitator or installing an active, low-temperature catalyst at the tail end of the
process. Sintering is minimized by using catalyst promoters that enhance thermal stability and by
maintaining reaction temperatures below critical values. The MoO3 promoter extends catalyst life
(in coal boilers) by preferentially adsorbing vapor-phase As which would otherwise adsorb on
active V 4+ sites. Free CaO in the fly ash (up to 3%) also scavenges As to low levels, forming
calcium arsenide particles which are collected with the fly ash. Many U.S. coals contain adequate
CaO; however, if the CaO content of the coal is too low, it can be added to the boiler or fuel.
However, CaO levels above 3% of the fly ash are undesirable, since CaO reacts with SO2 to form
CaSO4 which masks the exterior surface of the catalyst. Fouling by ammonium bisulfate is
minimized by keeping exit SO3 and NH3 concentrations low and maintaining reaction temperatures
above about 230 °C; SO3 formation is minimized by keeping reaction temperatures below 350 °C

90
or by using lower activity V2O5/TiO2 or zeolite catalysts that have low selectivities for SO3.
Ultimately, however, extra catalyst volume is typically added to SCR reactors to extend periods
between catalyst replacements.
For plants fueled by coal, substantial carry-over of inorganic ash occurs to HDU SCR units, a
small, but significant fraction of which deposits on monolith walls, masks or blocks catalyst
macropores, and plugs flow channels [314]. Extensive fouling necessitates the use of air lancing to
purge the ash out of the catalyst channels. Figure 40 reveals the extent of serious channel plugging
and erosion of an SCR catalyst in a pilot plant following several thousand hours of operation in flue
gas containing coal fly ash. Plugging and excessive pressure drop are avoided by keeping monolith
cell width at or above 7 mm.

Figure 40. Catalyst channel plugging (left) and damage due to erosion (right) during
operation in an SCR facility. Reproduced from [313].
The type and extent of chemical deactivation depends on operating conditions, fuel type,
catalyst geometry, shut-downs for boiler maintenance, etc. Mini-pilot tests and subsequent
full-scale SCR operating experience have provided little evidence of poisoning by basic minerals
from Western United States coals; rather they indicate that deactivation occurs principally by
masking of catalyst layers and plugging of catalyst pores by CaSO4 and other fly ash minerals.
Moreover, laboratory analysis of catalysts exposed to power plant slip streams indicates that
mineral poisons do not penetrate deep into catalyst pores [315,316] nor do they adsorb on Brønsted
acid sites unless plant conditions cause moisture to condense on the catalyst.
4.3.3.2. Plant Operating Strategy to Maximize Catalyst Life
A typical SCR unit consists of a series of two to four catalyst layers (three is most common for
coal boiler cleanup) through which the flue gas usually flows downward (see Figure 41). A layer of
fresh catalyst can be added as catalyst performance declines over time [317]. Two general schemes
are followed for replacing the spent catalyst, both of which take into consideration the relative
activity or design activity level, a parameter that is usually defined as the ratio of NOx conversion at
any time divided by that produced by the fresh catalyst. Once the NOx reduction performance
declines to the minimum design activity level (typically 65–75% of fresh activity), the catalyst can
either be replaced entirely (simultaneous replacement scheme) or one layer can be replaced at a
time (sequential replacement scheme), usually beginning at the top and working down [313,318].

91
The sequential method results in increased overall catalyst life (on a per-volume-replaced
basis), while annual replacement cost would be 60% lower for the simultaneous scheme (see
Figure 42 [319]). Thus, optimal, cost-effective design of an SCR unit requires considering both the
initial capital and annual costs.

Figure 41. Vertical-flow fixed-bed SCR reactor. DOE SCR demonstration facility at
Gulf Power Company’s Plant Crist. Reproduced from [313].
1.00
0.75
0.50
Replace One Layer Every 10,000 hrs

0.25

Replace All Layer Every 20,000 hrs
Design Activity

0.00

0

10000

20000

30000

40000

50000

Time (hours)
Figure 42. SCR replacement strategies: comparison of total replacement on a 20,000 h
cycle relative to sequential replacement on a 10,000 h cycle while maintaining constant
catalyst volume. Reproduced from [319].
Operating experience for commercial SCR installations has been better than anticipated.
Catalyst lifetimes of 3–4 years at overall efficiencies of 75–90% for HDU’s have been observed for
electric boiler installations [312]. The principal contributors to operating cost include catalyst
replacement cost, shutdown cost for catalyst replacement, and plant derating cost associated with

92
catalyst pressure drop. Catalyst replacement or regeneration was typically required within
2–3 years and catalyst replacement times varied from 2–7 days. Pressure drop ranged from
0.8–15 cm of water for the various catalyst configurations and volumes. Pressure drops for plate
type catalysts were significantly lower than for monolithic catalysts.
4.3.3.3. Catalyst Rejuvenation and Regeneration
While high-dust-catalyst life of 2–3 years is acceptable, advances in SCR catalyst regeneration
technologies make it possible to extend life by several additional years. Recent experience indicates
that even after long-term exposure to fly ash, foulants, and poisons, SCR catalysts may be
successfully regenerated to the original performance or better [307,308,320–322].
4.3.3.4. Methods of Renewing Catalysts
Deactivated catalysts may be cleaned, rejuvenated, and/or regenerated. Cleaning commonly
refers to removal of physical restrictions such as monolith channels plugged with fly ash or channel
surfaces covered with a loose dust layer; these restrictions are easily removed in situ using
compressed air, although cleaning will also be done as a first step in the other methods.
Rejuvenation refers to relatively mild treatments that remove catalyst poisons or foulants inside the
catalyst pores and restore part of the catalytic activity; these treatments are often done in situ or
on-site. Rejuvenation involves removal of blinding layers and partial removal of some poisons;
thus, activity is partly recovered, but none is added. Regeneration involves the off-site, complete
restoration of catalytic activity through a series of relatively sophisticated treatments, some of
which remove not only poisons and foulants, but also a part or much of the active catalytic
materials from the support; hence, regeneration also involves restoration of the catalytically active
materials bringing the catalyst to its original state or one of even higher activity. SCR catalysts are
routinely and regularly cleaned or “blown out” during operation, while rejuvenation or regeneration
is typically done after approximate 50–60% of the initial activity of the catalyst has been lost. In
situ rejuvenation (ISR) treatments were practiced early (e.g., 1990s and early 2000s), while off-site
regeneration (OSR) is now the predominant practice because of its greater effectiveness.
4.3.3.5. Rejuvenation or Regeneration?
According to McMahon [322], rejuvenating SCR catalyst may be more cost-effective than
regenerating, if the catalyst is fairly new or the SCR system does not operate year around (as in the
case of plants operating only during high pollutant levels, known as the “ozone season”).
Otherwise, the choice between rejuvenation and regeneration depends largely on economics, i.e.,
(1)
(2)
(3)
(4)
(5)

the plant’s dispatch economics, including transportation costs;
length of catalyst service;
costs of removing and replacing the catalyst;
the impact of the fuels combusted, i.e., coal, oil, or gas; and
the location of the catalyst in the plant, i.e., HDU or TGU.

93
Examples of rejuvenation treatments are found in the scientific and patent literature. For
example, work by Zheng and Johnsson [323] and others (e.g., [324,325]) indicates that activity of
poisoned catalysts might be partially regenerated by washing with water, sulfuric acid, NH4Cl,
and/or catalyst precursor solutions (e.g., ammonium paratungstate and vanadyl sulfate), as well as a
combination of washing and treatment with gaseous SO2. The extent to which these rejuvenation
methods are effective in restoring a significant fraction of the original catalyst activity
varies significantly.
4.3.3.5.1. Rejuvenation

On-site rejuvenation methods generally include the following procedural types: (i) removal of
dust in the monolith channels with compressed air followed by (ii) washing catalyst in a tank
containing agitated, deionized water to remove the CaSO4 coating and alkali metal salts deposited
by fly ash (the solution is generally mildly acidic due to impurities on the catalyst) or acidic
aqueous solution (pH = 1–2 in either case) in a tank; (iii) rinsing vigorously with deionized water
(usually in the same tank) to remove the dissolved and suspended deposits; and (iv) drying slowly
in clean air at room temperature followed by drying gently in hot air. Examples of on-site
regeneration methods include those developed and practiced in the time frame of 1995–2002 by
SCR-Tech, SBW, Saar Energie, Steag, EnBW, HEW, BHK, and Integral [326–328]. The method
described by Schneider and Bastuck [327] provided for adding catalytic materials, i.e., vanadium
and tungsten oxides (via impregnation of the V and W salts) to the cleaned catalyst.
The patent of Budin et al. [328] provides for more sophisticated treatments, including use of
(i) nonionic surfactants and complex-forming or ion-exchange additives, (ii) washing with an acid
or base, (iii) using acoustic radiation to remove fly-ash components, and (iv) addition of catalytic
materials (oxides of V, W, Mo free of alkali and alkaline-earth metals, halogen, and sulfur) to
restore activity, although few details or conditions of use are provided. In fact, no examples are
provided in any of the patents cited directly above; accordingly, it is unclear to what extent and
under what conditions the more sophisticated methods were used for on-site regeneration. The
methods claimed by Budin et al. [328] are clearly more readily applied in off-site regeneration, as
will be clear from the discussion below.
4.3.3.5.2. Regeneration
Bullock & Hartenstein [320], Cooper et al. [329], and McMahon [322] build a strong case for
off-site regeneration and a comprehensive catalyst management program.
4.3.3.6. A Comprehensive Approach to Catalyst Management
The approach [320,322] includes
(1)

strategies for extending catalyst life and reusability and planning for catalyst
removal/rotation to coincide with power plant outages;

94
(2)
(3)

(4)

catalyst inspection and testing before and following regeneration with replacement of
badly damaged catalyst which is unregenerable;
off-site regeneration using a series of robust washing and chemical treatments to remove
channel blockages, deactivated catalyst metals, and poisons, followed by chemical
treatments to restore active catalytic materials; and
gentle drying/calcination in air to high temperatures to produce catalytically active oxides.

4.3.3.7. Common Regeneration Practices
Normal regeneration procedures [307,308,320,322,330–333] are designed to enhance removal
of blockages, deactivated catalyst, and poisons and restore active catalytic material. These typically
include the following steps:
(1)
(2)

(3)

(4)
(5)

pressurized wet and dry treatments to remove channel blockages and outer dust layers;
washing of catalyst units in tanks containing agitated water augmented with surfactants,
dispersants, ion-exchange materials, emulsifiers, acid, base, and/or acoustic radiation to
remove the outer CaSO4 coating, alkali metal salts deposited in the catalyst pores, and
deactivated (e.g., As-poisoned) catalyst;
rinsing repeatedly in deionized water and repeating ultrasonic treatments between or in
concert with chemical treatments, with a final rinse to finish removal of any catalyst or
fouling residue;
reimpregnation of the clean support with salts of the active catalytic materials (V, Mo,
and W); and
drying (calcining) at low heating rates to decompose the salts of the active catalytic
materials to active metal oxides of V, Mo, and W.

4.3.3.8. Regeneration Process Profile: SCR-Tech Regeneration Process
SCR-Tech is the most prominent and experienced off-site regeneration company with 13 years
of experience in the regeneration business and a documented record of research and development,
going back to their German parent company ENVICA, who in 1997 began developing an offsite
regeneration process. SCR-Tech was the first and until 2008 the only company in the U.S. to
perform off-site regeneration. In September 2007, Evonik Energy Services (formerly Steag) opened
an SCR catalyst regeneration facility in the U.S.
The SCR-Tech regeneration process involves a number of different process steps illustrated in
Figure 43. Upon receipt of a shipment of catalyst, catalyst elements from several modules are
inspected and analyzed; results of the analysis provide a basis for determining the precise protocol
for treatment, i.e., the number and order of processing steps [334,335]. A large catalyst module is then
led through a protocol of soaking, washing, ultrasonic treatment, arsenic and/or phosphorus
removal (as needed), replenishment of V and Mo, neutralization, and rinsing in various soaking
pits, as shown in Figure 43; all of these wet chemical steps are performed at controlled pH and
temperature. Finally, the catalyst is dried, inspected, and packaged for shipment. Performance
guarantees are provided for complete removal of blinding layers, catalyst activity (typically higher

95
after regeneration), SO2 to SO3 conversion rate (typically lower), mechanical stability (the same),
and deactivation rate (the same) such that all properties of the regenerated catalyst are as good or
better than the new catalyst.
A comparison of the physical appearances of SCR monoliths and plates before and after
regeneration in Figure 44 reveals the rigor of the SCR-Tech cleaning process. The nearly complete
removal of poisons originally in high concentrations by the regeneration process is demonstrated in
Figure 45. Surface concentrations of CaO, P2O5, SiO2, and SO4 were also substantially reduced.
Table 21 compares the costs of regenerating versus buying a new catalyst [322]. This case is for
a typical 500 MW unit with 650 m3 of catalyst contained in 450 modules (150 modules in each of
3 layers). The purchase cost of new catalyst in 2006 was $3500 to 4500 per m3. The cost to
regenerate the catalyst is approximately 60% of this price. Thus, the purchase cost of one layer is
$758,000 to $975,000 as compared to a regeneration cost of $455,000 to $585,000 resulting in
savings per layer of $303,000 to $390,000 or $910,000 to $1.2 million for three layers. Assuming
the SCR unit runs year around (as most do now) and catalyst life is three years, the annual savings
due to regeneration is in the range of $300,000 to $600,000. The disposal cost for an SCR catalyst
can range from $50 to $2,000/ton, the upper figure based on the cost of treating the vanadium as
hazardous waste. Hence the disposal cost could be as high as $500,000 for a layer of catalyst.
According to McMahon, SCR catalysts can be regenerated from 3 to 7 times.
4.3.4. SCR Catalyst Case Study Summary Observations and Conclusions
1. Off-site regeneration processes are more sophisticated and demanding than on-site rejuvenation
processes; the off-site regeneration processes provide significantly more efficient cleaning and
reconstitution of the catalyst with full recovery of activity—sometimes greater than the fresh
catalyst activity. Rejuvenation provides only partial (up to 85%) recovery of the original activity.
2. The development of offsite processes for regeneration of SCR catalysts is relatively new, having
occurred largely over the past 10–15 years. SCR-Tech was the first and until 2008 the only
company to operate an off-site regeneration facility in the U.S.
3. Because surface deposits are a primary deactivation mechanism, especially in HDU catalysts,
extensive multi-step treatments are required, but rejuventation or regeneration appear to be a
cost-effective method of catalyst management for SCR catalysts.

96

Warehouse Storage

Regeneration Plant

Figure 43. SCR-Tech catalyst regeneration process. Reproduced from [322,335–337].
Reproduced with permission of Electric Power and CoaLogix, Inc.

(a)

(b)

Figure 44. (a) Monolith and (b) plate SCR catalysts before and after SCR-Tech
regenerative treatment. Reproduced from [334]. Courtesy CoaLogix, Inc.

97

Figure 45. Concentration of principle poisons before and after regeneration.
Reproduced from [320]. Courtesy CoaLogix, Inc.
Table 21. Cost per layer (217 m3 or 150 modules) of new versus regenerated SCR
catalyst. Adapted from [322]. Copyright 2006, Electric Power.
Catalyst Handling Step
Removal from SCR
system
Transport out
Purchase price
Shipping
Installation
Net
savings
from
regeneration
Disposal cost

New
Comparable

Regenerated
Comparable

Comparable
Comparable
$758,000–$975,000
$455,000–$585,000
Comparable
Comparable
Comparable
Comparable
$303,000–$390,000 pls disposal cost
$20,000–$500,000

0

4.4. Redispersion of Sintered Catalysts
During catalytic reforming of hydrocarbons on platinum-containing catalysts, growth of 1-nm
platinum metal clusters to 5–20-nm crystallites occurs. An important part of the catalyst regeneration
procedure is the redispersion of the platinum phase by a high temperature treatment in oxygen and
chlorine, generally referred to as “oxychlorination.” A typical oxychlorination treatment involves
exposure of the catalyst to HCl or CCl4 at 450–550 °C in 2–10% oxygen for a period of 1–4 h (see
details in Table 22). During coke burning, some redispersion occurs, e.g., dispersion (D) increases from
0.25 to 0.51, while during oxychlorination the dispersion is further increased, e.g., from 0.51 to
0.81 [256]. A mechanism for platinum redispersion by oxygen and chlorine is shown in Figure 46 [256].
It involves the adsorption of oxygen and chlorine on the surface of a platinum crystallite and
formation of AlCl3, followed by the formation of PtCl2(AlCl3)2 complexes that dissociatively

98
adsorb on alumina to oxychloro-platinum complexes. These latter complexes form monodisperse
platinum clusters upon subsequent reduction.
Table 22. Typical Regeneration Procedure for Reforming Catalysts a.
(1)

Preliminary operations: cool the catalyst to about 200 °C and strip hydrocarbons and H2 with N2

(2) Elimination of coke by combustion: inject dilute air (0.5% O2) at 380 °C and gradually increase
oxygen content to about 2% by volume while maintaining temperature below 450–500 °C to prevent
further sintering of the catalyst. To prevent excessive leaching of Cl2, HCl or CCl4 may be injected during
the combustion step
(3) Restoration of catalyst acidity: restoration of acidity occurs at 500 °C by injection of a chlorinated
compound in the presence of 100–200 ppm water in air
(4) Redispersion of the metallic phase: expose the catalyst to a few Torr of HCl or CCl4 in 2–10% O2
in N2 at 510–530 °C for a period of about 4 h. After redispersion, O2 is purged from the unit and the
catalyst is reduced in H2
a

Ref. [255,256].

Figure 46. Proposed mechanism for redispersion by oxychlorination of alumina-supported
platinum. Reproduced from [256]. Copyright 1982, Brill Nijhoff Publishers.
Some guidelines and principles regarding the redispersion process are worth enumerating:
(1)

(2)

In cases involving a high degree of Pt sintering or poisoning, special regeneration
procedures may be required. If large crystallites have been formed, several successive
oxychlorinations are performed [256].
Introducing oxygen into reactors in parallel rather than in series results in a significant
decrease in regeneration time [101].

99
(3)

(4)

Introduction of hydrocarbons present in the reactor recycle after regeneration is said to
stabilize the catalyst; solvents such as ammonium acetate, dilute nitric acid containing
lead nitrate, and EDTA and its diammonium salt are reported to dissolve out metal
aggregates without leaching out the dispersed metal [101].
The procedures for redispersion of Pt/alumina are not necessarily applicable to Pt on
other supports or to other metals. For example, Pt/silica is redispersed at lower
temperature and higher Cl2 concentration (150–200 °C and 25% Cl2). Pd/alumina can be
redispersed in pure O2 at 500 °C. While Pt–Re/alumina is readily redispersed by
oxychlorination at 500 °C, Pt–Ir/alumina is not redispersed in the presence of O2, unless
the catalyst is pretreated with HCl [266].

An extensive scientific and patent literature of redisperson describes the use of chlorine, oxygen,
nitric oxide, and hydrogen as agents for redispersion of sintered catalysts (summarized in Table 23).
Most of the early literature shows positive effects for chlorine compounds in the presence of
oxygen in redispersing alumina-supported platinum and other noble metals. Recent literature
demonstrates the need for understanding the detailed surface chemistry in order to successfully
develop and improve redispersion processes, especially in more complex catalyst systems such as
alumina-supported bimetallics. For example, on the basis of a fundamental study of the
redispersion surface chemistry, Fung [266] developed a redispersion procedure for Pt–Ir bimetallic
catalysts using a wet HCl/air treatment, since the conventional oxychlorination is not effective for
this catalyst.
Redispersion of alumina-supported platinum and iridium crystallites is also possible in
a chlorine-free oxygen atmosphere, if chlorine is present on the catalyst. The extent of redispersion
depends on the properties of the Pt/Al2O3 catalyst and temperature; for example, the data in
Figure 47 [102] for two different catalysts [catalyst 1 is a commercial Pt/Al2O3 (Engelhard); catalyst 2
is Pt/Al2O3 (Kaiser KA-201) impregnated with chloroplatinic acid] show that the maximum
increases in dispersion occur at about 550 °C. The data also show that redispersion does not occur
in a hydrogen environment. The question whether redispersion of platinum occurs only in oxygen
without chlorine present on the catalyst remains controversial.
Two models, “the thermodynamic redispersion model” and “the crystallite splitting model,”
have been advanced to explain the redispersion in oxygen [101,102,361]. The “thermodynamic”
redispersion model hypothesizes the formation of metal oxide molecules that detach from the
crystallite, migrate to active sites on the support, and form surface complexes with the support.
Upon subsequent reduction, the metal oxide complexes form monodisperse metal clusters. In the
“crystallite splitting” model, exposure of a platinum crystallite to oxygen at 500 °C leads to
formation of a platinum oxide scale on the outer surface of the crystallite, which stresses and
ultimately leads to splitting of the particle [361]. Dadyburjor hypothesizes that the crystallite
splitting model is most applicable to the behavior of large crystallites and to all particles at
relatively small regeneration times, while the thermodynamic migration model is useful for small
particles and most particles after longer regeneration times.

100

Table 23. Representative Patents Prior to 1990 Treating Catalyst Redispersion.
Dispersing agent
class

Dispersing agent

Chlorine-Containing
Cl2, Cl + halogen
Cl, H2O, O2
HCl, Cl–O
Cl, O2
HCl, Cl
Cl, halogen
Cl–H2O
HCl–O–He
Cl, O2
HCl, Cl, H2O, O
HCl, O
Cl, halogen
HCl
Oxygen
O2
Oxygen/N2
O2, N2

Metals/support

Patent No.

Ref.

Pt/zeolite
Pt/zeolite
Ir
Pt–Ir, Ir
Pt–Ir–Re, Pt–Ir/zeolites
Ir, Pt–Ir/Al2O3
Pt–Ir–Se/Al2O3
Pt–Ir–Se/Al2O3
Pt/zeolite
Pt/zeolite
Ir, Pt–Ir/Al2O3
Ir, Pt–Ir/Al2O3
Ir, Pt–Ir/Al2O3

U.S. 4,645,751
U.S. 4,657,874
U.S. 4,491,636
U.S. 4,467,045
U.S. 4,359,400
U.S. 4,480,046
U.S. 4,492,767
U.S. 4,491,635
U.S. 4,855,269
U.S. 4,925,819
U.S. 4,444,896
U.S. 4,444,895
U.S. 4,517,076

[338]
[339]
[340]
[341]
[342]
[343]
[344]
[345]
[346]
[347]
[348]
[349]
[350]

Pt, Re/Al2O3

U.S. 4,482,637

[351]

Cu/Cr, Mn, Ru, Pd, Zn, Si,
Mg, Ca, Sr, Ba

U.S. 4,855,267

[352]

Eu 0,306,170
U.S. 4,891,346
U.S. 4,447,551
U.S. 4,472,514
U.S. 4,473,656

[353]
[354]
[355]
[356]
[357]

U.S. 4,952,543

[358]

U.S. 4,444,897
U.S. 4,472,515

[359]
[360]

Other
NO, NO + halogen
Halogen
Halide
Halide, halogen/H2O
Halogen
NO, NO + halogen, Cl
H2-halides, O2
Halogen, H2O

Pt, Pd/zeolite
Ru, Os, Rh, Pd/Al2O3
Ir, Pt–Ir/Al2O3
Ir, Pt–Ir/Al2O3
Ir, Pt–Ir/Al2O3
Group VIII metals/Al2O3,
SiO2, zeolites
Ir, Pt–Ir/Al2O3
Ir, Pt–Ir/Al2O3

4.4.1. Case Study: Cobalt based Fischer-Tropsch (FT) Catalyst Regeneration
Fischer-Tropsch (FT) synthesis is a catalytic process used to produce long chain hydrocarbons
from synthesis gas consisting of carbon monoxide and hydrogen. Cobalt catalysts were initially
developed by Franz Fischer and Hans Tropsch in the 1920s and similar cobalt-based catalysts are
still in use today [8]. Although more expensive than iron based catalysts that are also used for FTS,
supported cobalt FT catalysts are more active and selective for the desired liquid and wax products.

101

Figure 47. Effects of 1-h treatments in O2 (closed symbols) and H2 (open symbols) on
the dispersion of Pt/Al2O3 catalysts: ż,Ɣ Pt/Al2O3 (Engelhard), Ƒ,Ŷ Pt/KA-201 alumina
(Kaiser). Reproduced from [102]. Copyright 1982, Brill Nijhoff Publishers.
A recent review by the Davis group at the Center for Applied Energy Research at the University
of Kentucky with Bukur at the University of Texas A&M in Qatar [362] focused on the results of
studies using synchrotron radiation to characterize Co FT catalysts. The review includes a detailed
consideration and analysis of the mechanisms and processes of sintering, oxidation, aluminate
formation, and coking and carbide formation and under what operating conditions each is important.
They summarize their and others’ previous findings that oxidation primarily occurs on small (<2 nm)
cobalt crystallites and at high partial pressures of water [362–366]. Further, they highlight
the potentially complicated transformations between CoO and aluminates [362,364,367]. These
complications highlight a complex mechanism that may be related to chemical-assisted sintering of Co
FTS catalysts through a combination of the effect of CoO reduction during the initial activation of the
catalysts and water exposure during operation. First, CoO, present either due to incomplete reduction of
the catalysts [368] or oxidation of the small (<2 nm) crystallites as suggested by Davis’
group [369,370] can apparently increase the sintering rate due to mobility that allows them to aggregate
into larger CoO clusters that are subsequently reduced to metallic Co, as inferred from evidence
presented in a number of studies [79,362,368–371]. Primarily, X-ray absorption near edge (XANES)
analysis shows simultaneous increasing extent of reduction and increasing Co-Co coordination, due
both to removal of oxygen and increases in particle size. Second, water is believed to cause
chemical-assisted sintering [80,367,372–374], especially at high partial pressures that occur at CO
conversions above about 65% [223], although the exact mechanisms are debated. Minor surface
oxidation [373,374] and surface wetting [375] have been proposed, although Saib et al. have
shown that cobalt oxidation is not an important deactivation route [79] in catalysts with Co
particles >~8 nm, which are typical in commercial FTS catalysts.

102
A number of articles by researchers at Sasol, Eindhoven University of Technology, and the
University of South Africa detailed the causes of deactivation and demonstrated the regenerability
of alumina-supported cobalt FT catalysts [79,368,371,376–382]. Through a combination of studies
on single crystal [377] and actual catalysts from pilot plants operated under industrial FT
conditions [368,371], they concluded that contrary to prior hypotheses, neither formation of cobalt
aluminates nor oxidation of the cobalt were significant deactivation mechanisms. In fact, extent of
Co oxidation actually decreased with time on stream [371]. However, Co sintering and carbon
deposition were identified as the primary means of deactivation. In unpublished presentations by these
authors, the relative contributions of carbon deposition and sintering to the deactivation were reported
as roughly equal. More interestingly, both of these deactivation mechanisms could be largely reversed
through high pressure oxidation treatment [376,378], which removes both inactive carbon and
redisperses the cobalt. Through high resolution transmission electron micrographs (HRTEM), the
mechanism of redispersion of the cobalt was identified as the Kirkendall effect, which results in the
formation of spherical shells of cobalt oxide that during subsequent reduction disperse into smaller
crystallites of cobalt (see Figure 48). Bezemer et al. have previously shown that unpromoted Co FT
catalysts require Co crystallites of at least 6 nm in diameter to achieve maximum turnover frequency,
but this is the optimum size because larger crystallites display the same surface activity as the 6 nm
particles [383]. The oxidative regeneration and reduction process described by Hauman et al. [376] and
Weststrate et al. [377] recovers ~95% of the fresh catalyst activity by removing the carbon deposits and
returning the sintered cobalt particles to near the optimum 6 nm size. While the rate per mass of catalyst
is nearly constant following regeneration, some smaller particles are produced on model catalysts
because the rate on a turnover frequency basis decreases by roughly 1/3 compared to the fresh
catalysts [376].

Figure 48. Bright field TEM images showing redispersion of cobalt particles supported
on a flat model silica by oxidative treatment. The center image shows hollow spheres
created by the Kirkendall effect, which form dispersed smaller particles upon re-reduction
in the right hand image. Reproduced from [378]. Copyright 2011, Springer.
These results are significant because they show the power of careful evaluation of the root
causes of deactivation in an important catalytic system and then show how proper choice of
regeneration conditions can extend the life of the catalysts by redispersion of the active metal.
However, promoters may not be redispersed as completely as the cobalt during repeated
regeneration. Although traditional promoters, like Pt and Ru, appear to remain with the Co and

103
maintain their effect, some promoters like Au tend to segregate and lose their promotion effect, as
indicated by TPR peaks shifting to higher temperatures [384].
5. Summary
This article focuses on the causes, mechanisms, prevention, modeling, and treatment
(experimental and theoretical) of deactivation. Several general, fundamental principles are evident:
(1)

(2)

(3)

The causes of deactivation are basically of three kinds: chemical, mechanical, and
thermal. The five intrinsic mechanisms of catalyst decay, (a) poisoning, (b) fouling, (c)
thermal degradation, (d) chemical degradation, and (e) mechanical failure, vary in their
reversibility and rates of occurrence. Poisoning and thermal degradation are generally
slow, irreversible processes, while fouling with coke and carbon is generally rapid and
reversible by regeneration with O2 or H2.
Catalyst deactivation is more easily prevented than cured. Poisoning by impurities can
be prevented through careful purification of reactants or mitigated to some extent by
adding traps or “getters” as components of the catalyst. Carbon deposition and coking
can be prevented by minimizing the formation of carbon or coke precursors through
gasification, careful design of catalysts and process conditions, and by controlling
reaction rate regimes, e.g., mass transfer regimes, to minimize effects of carbon and
coke formation on activity. Sintering is best avoided by minimizing and controlling the
temperature of reaction, although recent developments have focused on encapsulating
metal crystallites to eliminate mobility, while still allowing access for reactants
and products.
Catalyst regeneration is feasible in some circumstances, especially to recover activity
loss due to rapid coking or longer term deactivation associated with loss of active metal
dispersion. Typically, regeneration or rejuvenation strategies are dictated by process or
economic necessity to obtain desired process run lengths. Life cycle operating strategies
are important considerations when evaluating catalyst regeneration/rejuvenation versus
replacement decisions. Rejuvenation treatments can extend the useful life of catalysts.
Selective catalytic reduction catalysts provide an example of rejuvenation practiced in
a commercial process.

Acknowledgments
The authors wish to acknowledge Brigham Young University for its support and the employees
of MDPI, for their untiring assistance in publishing this review.
Author Contributions
Calvin H. Bartholomew was the primary author of this review. Morris D. Argyle provided
assistance in writing and revising the case studies and updating the article in response to the
reviewers’ comments.

104
Conflicts of Interest
The authors declare no conflict of interest.
References
1.

2.
3.
4.
5.

6.
7.

8.
9.
10.
11.
12.
13.
14.
15.
16.
17.

18.

Figueiredo, J.L. Carbon formation and gasification on nickel. In Progress in Catalyst
Deactivation; (NATO Advanced Study Institute Series E, No. 54); Martinus Nijhoff
Publishers: Boston, MA, USA, 1982; pp. 45–63.
Hughes, R. Deactivation of Catalysts; Academic Press, London, UK, 1984.
Oudar, J.; Wise, H. Deactivation and Poisoning of Catalysts; Marcel Dekker: New York, NY,
USA, 1985.
Butt, J.B.; Petersen, E.E. Activation, Deactivation, and Poisoning of Catalysts; Academic
Press: San Diego, CA, USA, 1988.
Denny, P.J.; Twigg, M.V. Factors determining the life of industrial heterogeneous catalysts.
In Catalyst Deactivation 1980 (Studies in Surface Science and Catalysis); Delmon, B.,
Froment, G.F., Eds.; Elsevier: Amsterdam, The Nerthelands, 1980; Volume 6; pp. 577–599.
Bartholomew, C.H. Catalyst Deactivation. Chem. Eng. 1984, 91, 96–112.
Butt, J.B. Catalyst deactivation and regeneration. In Catalysis—Science and Technology;
Anderson, J.R., Boudart, M., Eds.; Springer-Verlag: New York, NY, USA, 1984; Volume 6,
pp. 1–63.
Bartholomew, C.H.; Farrauto, R.J. Fundamentals of Industrial Catalytic Processes, 2nd ed.;
Wiley-Interscience: Hoboken, NJ, USA, 2006.
Delmon, B.; Froment, G.F. Catalyst Deactivation 1980 (Studies in Surface Science and
Catalysis); Elsevier: Amsterdam, The Nerthelands, 1980; Volume 6.
Delmon, B.; Froment, G.F. Catalyst Deactivation 1987 (Studies in Surface Science and
Catalysis); Elsevier: Amsterdam, The Nerthelands, 1987; Volume 34.
Bartholomew, C.H.; Butt, J.B. Catalyst Deactivation 1991 (Studies in Surface Science and
Catalysis); Elsevier: Amsterdam, The Nerthelands, 1991; Volume 68.
Delmon, B.; Froment, G.F. Catalyst Deactivation 1994 (Studies in Surface Science and
Catalysis); Elsevier: Amsterdam, The Nerthelands, 1994; Volume 88.
Bartholomew, C.H.; Fuentes, G.A. Catalyst Deactivation 1997 (Studies in Surface Science
and Catalysis); Elsevier: Amsterdam, The Nethelands, 1997; Volume 111.
Delmon, B.; Froment, G.F. Catalyst Deactivation 1999 (Studies in Surface Science and
Catalysis); Elsevier: Amsterdam, The Nethelands, 1999; Volume 126.
Moulijn, J.A. Catalyst Deactivation. Appl. Catal. A 2001, 212, 1–255.
Maxted, E.B. The poisoning of metallic catalysts. Adv. Catal. 1951, 3, 129–177.
Hegedus, L.L.; McCabe, R.W. Catalyst poisoning. In Catalyst Deactivation 1980 (Studies in
Surface Science and Catalysis); Delmon, B., Froment, G.F., Eds.; Elsevier: Amsterdam,
The Nethelands, 1980; Volume 6, pp. 471–505.
Hegedus, L.L.; McCabe, R.W. Catalyst Poisoning; Marcel Dekker: New York, NY,
USA, 1984.

105
19. Butt, J.B. Catalyst poisoning and chemical process dynamics. In Progress in Catalyst
Deactivation (NATO Advanced Study Institute Series E, No. 54); Figueiredo, J.L., Ed.;
Martinus Nijhoff Publishers: Boston, MA, USA, 1982; pp. 153–208.
20. Barbier, J. Effect of poisons on the activity and selectivity of metallic catalysts. In
Deactivation and Poisoning of Catalysts; Oudar, J., Wise, H., Eds.; Marcel Dekker: New
York, NY, USA, 1985; pp. 109–150.
21. Bartholomew, C.H. Mechanisms of nickel catalyst poisoning. In Catalyst Deactivation 1987
(Studies in Surface Science and Catalysis); Delmon, B., Froment, G.F., Eds.; Elsevier:
Amsterdam, The Nethelands, 1987; Volume 34, pp. 81–104.
22. Rostrup-Nielsen, J.R. Promotion by poisoning. In Catalyst Deactivation 1991 (Studies in
Surface Science and Catalysis); Bartholomew, C.H., Butt, J.B., Eds.; Elsevier: Amsterdam, The
Nethelands, 1991; Volume 68, pp. 85–101.
23. Inga, J.; Kennedy, P.; Leviness, S. A1 Fischer-tropsch process in the presence of nitrogen
contaminants. WIPO Patent WO 2005/071044, 4 August 2005.
24. Völiter, V.J.; Hermann, M. Katalytische Wirksamkeit van reinem und von CO-vergiftetem
Platin bei der p-H2-Umwandlung. Z. Anorg. Allg. Chem. 1974, 405, 315.
25. Baron, K. Carbon monoxide oxidation on platinum-lead films. Thin Solid Films 1978, 55,
449–462.
26. Clay, R.D.; Petersen, E.E. Catalytic activity of an evaporated platinum film progressively
poisoned with arsine. J. Catal. 1970, 16, 32–43.
27. Madon, R.J.; Seaw, H. Effect of Sulfur on the Fischer-Tropsch Synthesis. Catal. Rev.—Sci.
Eng. 1977, 15, 69–106.
28. Bartholomew, C.H.; Agrawal, P.K.; Katzer, J.R. Sulfur poisoning of metals. Adv. Catal.
1982, 31, 135–242.
29. Rostrup-Nielsen, J.R. Sulfur Poisoning. In Progress in Catalyst Deactivation (NATO
Advanced Study Institute Series E, No. 54); Figueiredo, J.L., Ed.; Martinus Nijhoff
Publishers: Boston, MA, USA, 1982; pp. 209–227.
30. Wise, H.; McCarty, J.; Oudar, J. Sulfur and carbon interactions with metal surfaces. In
Deactivation and Poisoning of Catalysts; Oudar, J., Wise, H., Eds.; Marcel Dekker: New
York, NY, USA, 1985; pp. 1–50.
31. Rostrup-Nielsen, J.R.; Nielsen, P.E.H. Catalyst deactivation in synthesis gas production, and
important syntheses. In Deactivation and Poisoning of Catalysts; Oudar, J., Wise, H., Eds.;
Marcel Dekker: New York, NY, USA, 1985; pp. 259–323.
32. Grossmann, A.; Erley, W.; Ibach, H. Adsorbate-induced surface stress and surface
reconstruction: oxygen, sulfur and carbon on Ni(111). Surf. Sci. 1995, 337, 183–189.
33. Ruan, L.; Stensgaard, I.; Besenbacher, F.; Lægsgaard, E. Observation of a missing-row
structure on an fcc (111) surface: The (5 ¥3 ×2)S phase on Ni(111) studied by scanning
tunneling microscopy. Phys. Rev. Lett. 1993, 71, 2963–2966.
34. Kitajima, Y.; Yokoyama, T.; Ohta, T.; Funabashi, M.; Kosugi, N.; Kuroda, H. Surface
EXAFS and XANES studies of (5v3× 2)S/Ni(111). Surf. Sci. 1989, 214, L261–L269.

106
35. Perdereau, M.; Oudar, J. Structure, mécanisme de formation et stabilité de la couche
d'adsorption du soufre sur le nickel. Surf. Sci. 1970, 20, 80–98.
36. Oudar, J. Sulfur adsorption and poisoning of metallic catalysts. Catal. Rev.—Sci. Eng. 1980,
22, 171–195.
37. McCarroll, J.J.; Edmonds, T.; Pitkethly, R.C. Interpretation of a complex low energy electron
diffraction pattern: Carbonaceous and sulphur-containing structures on Ni(111). Nature 1969,
223, 1260–1262.
38. Edmonds, T.; McCarroll, J.J.; Pitkethly, R.C. Surface structures formed during the interaction
of sulphur compounds with the (111) face of nickel. J. Vac. Sci. Technol. 1971, 8, 68–74.
39. Bartholomew, C.H. Mechanisms of catalyst deactivation. Appl. Catal. A 2001, 212, 17–60.
40. Ruan, L.; Besenbacher, F.; Stensgaard, I.; Lægsgaard, E. Atom-resolved studies of the
reaction between H2S and O on Ni(110). Phys. Rev. Lett. 1992, 69, 3523–3526.
41. Hepola, J.; McCarty, J.; Krishnan, G.; Wong, V. Elucidation of behavior of sulfur on nickel-based
hot gas cleaning catalysts. Appl. Catal. B 1999, 20, 191–203.
42. Erley, W.; Wagner, H. Sulfur poisoning of carbon monoxide adsorption on Ni(111). J. Catal.
1978, 53, 287–294.
43. Rendulic, K.D.; Winkler, A. The initial sticking coefficient of hydrogen on sulfur- and
oxygen-covered polycrystalline nickel surfaces. Surf. Sci. 1978, 74, 318–320.
44. Goodman, D.W.; Kiskinova, M. Chemisorption and reactivity studies of H2 and CO on
sulfided Ni(100). Surf. Sci. 1981, 105, L265–L270.
45. Kiskinova, M.; Goodman, D.W. Modification of chemisorption properties by electronegative
adatoms: H2 and CO on chlorided, sulfided, and phosphided Ni(100). Surf. Sci. 1981, 108,
64–76.
46. Johnson, S.; Madix, R.S. Desorption of hydrogen and carbon monoxide from Ni(100),
Ni(100)p(2 × 2)S, and Ni(100)c(2 × 2)S surfaces. Surf. Sci. 1981, 108, 77–98.
47. Madix, R.J.; Thornberg, M.; Lee, S.B. CO-sulfur interaction on Ni(110); evidence for local
interactions, not long range electronic effects. Surf. Sci.1983, 133, L447–L451.
48. Hardegree, E.L.; Ho, P.; White, J.M. Sulfur adsorption on Ni(100) and its effect on CO
chemisorption: I. TDS, AES and work function results. Surf. Sci. 1986, 165, 488–506.
49. Erekson, E.J.; Bartholomew, C.H. Sulfur poisoning of nickel methanation catalysts: II.
Effects of H2S concentration, CO and H2O partial pressures and temperature on reactivation
rates. Appl. Catal. 1983, 5, 323–336.
50. Jacobs, G.; Ghadiali, F.; Pisanu, A.; Padro, C.L.; Borgna, A.; Alvarez, W.E.; Resasco, D.E.
Increased Sulfur Tolerance of Pt/KL Catalysts Prepared by Vapor-Phase Impregnation and
Containing a Tm Promoter. J. Catal. 2000, 191, 116–127.
51. Jongpatiwut, S.; Sackamduang, P.; Rirksomboon, T.; Osuwan, S.; Alvarez, W.E.; Resasco,
D.E. Sulfur- and water-tolerance of Pt/KL aromatization catalysts promoted with Ce and Yb.
Appl. Catal. A 2002, 230, 177–193.

107
52. Jacobs, G.; Ghadiali, F.; Pisanu, A.; Borgna, A.; Alvarez, W.E.; Resasco, D.E.
Characterization of the morphology of Pt clusters incorporated in a KL zeolite by vapor phase
and incipient wetness impregnation. Influence of Pt particle morphology on aromatization
activity and deactivation. Appl. Catal. A 1999, 188, 79–98.
53. McVicker, G.B.; Kao, J.L.; Ziemiak, T.J.J.; Gates, W.E.; Robbins, J.L.; Treacy, M.M.J.;
Rice, S.B.; Vanderspurt, T.H.; Cross, V.R.; Ghosh, A.K. Effect of Sulfur on the Performance
and on the Particle Size and Location of Platinum in Pt/KL Hexane Aromatization Catalysts.
J. Catal. 1993, 139, 48–61.
54. Farbenindustrie, I.G. Improvements in the Manufacture and Production of Unsaturated
Hydrocarbons of Low Boiling Point. British Patent 322,284, 5 December 1929.
55. Kritzinger, J.A. The role of sulfur in commercial iron-based Fischer–Tropsch catalysis with
focus on C2-product selectivity and yield. Catal. Today 2002, 71, 307–318.
56. Sparks, D.E.; Jacobs, G.; Gnanamani, M.K.; Pendyala, V.R.R.; Ma, W.; Kang, J.;
Shafer, W.D.; Keogh, R.A.; Graham, U.M.; Gao, P.; et al. Poisoning of cobalt catalyst used
for Fischer–Tropsch synthesis. Catal. Today 2013, 215, 67–72.
57. Rostrup-Nielsen, J.R.; Trimm, D.L. Mechanisms of carbon formation on nickel-containing
catalysts. J. Catal. 1977, 48, 155–165.
58. Trimm, D.L. The Formation and Removal of Coke from Nickel Catalyst. Catal. Rev.—Sci.
Eng. 1977, 16, 155–189.
59. Trimm, D.L. Catalyst design for reduced coking (review). Appl. Catal. 1983, 5, 263–290.
60. Bartholomew, C.H. Carbon deposition in steam reforming and methanation. Catal. Rev.—Sci.
Eng. 1982, 24, 67–112.
61. Albright, L.F.; Baker, R.T.K. Coke Formation on Metal Surfaces (ACS Symposium Series
202); American Chemical Society: Washington, DC, USA, 1982.
62. Menon, P.G. Coke on catalysts—harmful, harmless, invisible and beneficial types. J. Mol.
Catal. 1990, 59, 207–220.
63. Rostrup-Nielsen, J.R. Conversion of hydrocarbons and alcohols for fuel cells. Phys. Chem.
Chem. Phys. 2001, 3, 283–288.
64. Trane-Restrup, R.; Resasco, D.E.; Jensen, A.D. Steam reforming of light oxygenates. Catal.
Sci. Technol. 2013, 3, 3292–3302.
65. De Lima, S.M.; da Silva, A.M.; da Costa, L.O.O.; Assaf, J.M.; Jacobs, G.; Davis, B.H.;
Mattos, L.V.; Noronha, F.B. Evaluation of the performance of Ni/La2O3 catalyst prepared
from LaNiO3 perovskite-type oxides for the production of hydrogen through steam reforming
and oxidative steam reforming of ethanol. Appl. Catal. A 2010, 377, 181–190.
66. Deken, J.D.; Menon, P.G.; Froment, G.F.; Haemers, G.; On the nature of carbon in NiĮ-Al2O3
catalyst deactivated by the methane-steam reforming reaction. J. Catal. 1981, 70, 225–229.
67. Durer, W.G.; Craig, J.H., Jr.; Lozano, J. Surface carbon and its effects on hydrogen
adsorption on Rh(100). Appl. Surf. Sci. 1990, 45, 275–277.
68. Moeller, A.D.; Bartholomew, C.H. Deactivation by carbon of nickel and nickel-molybdenum
methanation catalysts. Prepr.—Am. Chem. Soc., Div. Fuel Chem. 1980, 25, 54–70.

108
69. Marschall, K.-J.; Mleczko, L. Short-contact-time reactor for catalytic partial oxidation of
methane. Ind. Eng. Chem. Res. 1999, 38, 1813–1821.
70. Rostrup-Nielsen, J.R. Catalytic steam reforming. In Catalysis—Science and Technology;
Anderson, J.R., Boudart, M., Eds.; Springer-Verlag: New York, NY, USA, 1984; Volume 5,
pp. 1–117.
71. Besenbacher, F.; Chorkendorff, I.; Clausen, B.S.; Hammer, B.; Molenbroek, A.M.;
Nørskov, J.K.; Stensgaard, I. Design of a surface alloy catalyst for steam reforming. Science
1998, 279, 1913–1915.
72. Nemes, T.; Chambers, A.; Baker, R.T.K. Characteristics of carbon filament formation from
the interaction of CobaltíTin particles with ethylene. J. Phys. Chem. B 1998, 102, 6323-6330.
73. Bartholomew, C.H.; Strasburg, M.V.; Hsieh, H. Effects of support on carbon formation and
gasification on nickel during carbon monoxide hydrogenation. Appl. Catal. 1988, 36,
147–162.
74. Vance, C.K.; Bartholomew, C.H. Hydrogenation of carbon dioxide on group viii metals: III,
Effects of support on activity/selectivity and adsorption properties of nickel. Appl. Catal.
1983, 7, 169–177.
75. Baker, R.T.K.; Chludzinski, J.J., Jr. Filamentous carbon growth on nickel-iron surfaces:
The effect of various oxide additives. J. Catal. 1980, 64, 464–468.
76. Brown, D.E.; Clark, J.T.K.; Foster, A.I.; McCarroll, J.J.; Sims, M.L. Inhibition of coke
formation in ethylene steam cracking. In Coke Formation on Metal Surfaces (ACS
Symposium Series 202); Albright, L.F., Baker, R.T.K., Eds.; American Chemical Society:
Washington, DC, USA, 1982; pp. 23–43.
77. Bitter, J.H.; Seshan, K.; Lercher, J.A. Deactivation and coke accumulation during CO2/CH4
reforming over Pt catalysts. J. Catal. 1999, 183, 336–343.
78. Rostrup-Nielsen, J.R. Coking on nickel catalysts for steam reforming of hydrocarbons. J.
Catal. 1974, 33, 184–201.
79. Saib, A.M.; Moodley, D.J.; Ciobica, I.M.; Hauman, M.M.; Sigwebela, B.H.; Weststrate, C.J.;
Niemanstsverdriet, J.W.; van de Loosdrecht, J. Fundamental understanding of deactivation
and regeneration of cobalt Fischer–Tropsch synthesis catalysts. Catal. Today 2010, 154,
271–282.
80. Tsakoumis, N.E.; Rønning, M.; Borg, O.; Rytter, E.; Holmen, A. Deactivation of cobalt based
Fischer–Tropsch catalysts: A review. Catal. Today 2010, 154, 162–182.
81. Gates, B.C.; Katzer, J.R.; Schuit, G.C.A. Chemistry of Catalytic Processes; McGraw-Hill:
New York, NY, USA, 1979.
82. Naccache, C. Deactivation of acid catalysts. In Deactivation and Poisoning of Catalysts;
Oudar, J., Wise, H., Eds.; Marcel Dekker: New York, NY, USA, 1985; pp. 185–203.
83. Appleby, W.G.; Gibson, J.W.; Good, G.M. Coke Formation in Catalytic Cracking. Ind. Eng.
Chem. Process Des. Dev. 1962, 1, 102–110.
84. Beuther, H.; Larson, O.H.; Perrotta, A.J. The mechanism of coke formation on catalysts.
In Catalyst Deactivation 1980 (Studies in Surface Science and Catalysis); Delmon, B.,
Froment, G.F., Eds.; Elsevier: Amsterdam, The Nerthelands, 1980; Volume 6, pp. 271–282.

109
85. Gayubo, A.G.; Arandes, J.M.; Aguayo, A.T.; Olazar, M.; Bilbao, J. Deactivation and acidity
deterioration of a silica/alumina catalyst in the isomerization of cis-butene. Ind. Eng. Chem.
Res. 1993, 32, 588–593.
86. Augustine, S.M.; Alameddin, G.N.; Sachtler, W.M.H. The effect of Re, S, and Cl on the
deactivation of Pt Ȗ-Al2O3 reforming catalysts. J. Catal. 1989, 115, 217–232.
87. Guisnet, M.; Magnoux, P. Coking and deactivation of zeolites: Influence of the Pore
Structure. Appl. Catal. 1989, 54, 1–27.
88. Bauer, F.; Kanazirev, V.; Vlaev, C.; Hanisch, R.; Weiss, W. Koksbildung in
ZSM-5-Katalysatoren. Chem. Tech. 1989, 41, 297–301.
89. Grotten, W.A.; Wojciechowski, B.W.; Hunter, B.K. On the relationship between coke
formation chemistry and catalyst deactivation. J. Catal. 1992, 138, 343–350.
90. Bellare, A.; Dadyburjor, D.B. Evaluation of modes of catalyst deactivation by coking for
cumene cracking over zeolites. J. Catal. 1993, 140, 510–525.
91. Uguina, M.A.; Serrano, D.P.; Grieken, R.V.; Vènes, S. Adsorption, acid and catalytic changes
induced in ZSM-5 by coking with different hydrocarbons. Appl. Catal. A 1993, 99, 97–113.
92. Li, C.; Chen, Y.-W.; Yang, S.-J.; Yen, R.-B. In-situ FTIR investigation of coke formation on
USY zeolite. Appl. Surf. Sci. 1994, 81, 465–468.
93. Buglass, J.G.; de Jong, K.P.; Mooiweer, H.H. Analytical studies of the coking of the zeolite
ferrierite. J. Chem. Soc. Abstr. 1995, 210, 105-PETR.
94. Chen, D.; Rebo, H.P.; Moljord, K.; Holmen, A. Effect of coke deposition on transport and
adsorption in zeolites studied by a new microbalance reactor. Chem. Eng. Sci. 1996, 51,
2687–2692.
95. Guisnet, M.; Magnoux, P.; Martin, D. Roles of acidity and pore structure in the deactivation
of zeolites by carbonaceous deposits. In Catalyst Deactivation 1997 (Studies in Surface
Science and Catalysis); Bartholomew, C.H., Fuentes, G.A., Eds.; Elsevier: Amsterdam, The
Nethelands, 1997; Volume 111, pp. 1–19.
96. Masuda, T.; Tomita, P.; Fujikata, Y.; Hashimoto, K. Deactivation of HY-type zeolite catalyst
due to coke deposition during gas-oil cracking. In Catalyst Deactivation 1999 (Studies in
Surface Science and Catalysis); Delmon, B., Froment, G.F. Eds.; Elsevier: Amsterdam, The
Nethelands, 1999; Volume 126, pp. 89–96.
97. Cerqueira, H.S.; Magnoux, P.; Martin, D.; Guisnet, M. Effect of contact time on the nature
and location of coke during methylcyclohexane transformation over a USHY zeolite. In
Catalyst Deactivation 1999 (Studies in Surface Science and Catalysis); Delmon, B., Froment,
G.F., Eds.; Elsevier: Amsterdam, The Nethelands, 1999; Volume 126, pp. 105–112.
98. Wanke, S.E.; Flynn, P.C. The sintering of supported metal catalysts. Catal. Rev.—Sci. Eng.
1975, 12, 93–135.
99. Wynblatt, P.; Gjostein, N.A. Supported metal crystallites. Prog. Solid State Chem. 1975, 9,
21–58.
100. Ruckenstein, E.; Pulvermacher, B. Kinetics of crystallite sintering during heat treatment of
supported metal catalysts. AIChE J. 1973, 19, 356–364.

110
101. Ruckenstein, E.; Dadyburjor, D.B. Sintering and redispersion in supported metal catalysts.
Rev. Chem. Eng. 1983, 1, 251–356.
102. Wanke, S.E. Sintering of commercial supported platinum group metal catalysts. In Progress
in Catalyst Deactivation (NATO Advanced Study Institute Series E, No. 54); Figueiredo, J.L.,
Ed.; Martin Nijhoff Publishers: Boston, MA, USA, 1982; pp. 315–328.
103. Baker, R.T.; Bartholomew, C.H.; Dadyburjor, D.B. Sintering and Redispersion: Mechanisms
and Kinetics. In Stability of Supported Catalysts: Sintering and Redispersion; Horsley, J.A.,
Ed.; Catalytica: Mountain View, CA, USA, 1991; pp. 169-225.
104. Bartholomew, C.H. Model catalyst studies of supported metal sintering and redispersion
kinetics. In Catalysis (Specialist Periodical Report); Spivey, J.J., Agarwal, S.K., Eds.; Royal
Society of Chemistry, Cambridge, UK, 1993; Volume 10, pp. 41–82.
105. Bartholomew, C.H. Sintering kinetics of supported metals: new perspectives from a unifying
GPLE treatment. Appl. Catal. A 1993, 107, 1–57.
106. Bartholomew, C.H. Sintering kinetics of supported metals: perspectives from a generalized
power law approach. In Catalyst Deactivation 1994 (Studies in Surface Science and
Catalysis); Delmon, B., Froment, G.F., Eds.; Elsevier: Amsterdam, The Nethelands, 1994;
Volume 88, pp. 1–18.
107. Bartholomew, C.H. Sintering and redispersion of supported metals: Perspectives from the
literature of the past decade. In Catalyst Deactivation 1997 (Studies in Surface Science and
Catalysis); Bartholomew, C.H., Fuentes, G.A., Eds.; Elsevier: Amsterdam, The Nethelands,
1997; Volume 111, pp. 585–592.
108. Bartholomew, C.H.; Sorenson, W. Sintering kinetics of silica- and alumina-supported nickel
in hydrogen atmosphere. J. Catal. 1983, 81, 131–141.
109. Moulijn, J.A.; van Diepen, A.E.; Kapteijn, F. Catalyst deactivation: is it predictable?: What to
do? Appl. Catal. A 2001, 212, 3–16.
110. Bridger, G.W.; Spencer, M.S. Methanol synthesis. In Catalyst Handbook, 2nd ed.; Twigg,
M.V., Ed.; Manson Publishing: London, UK, 1996; pp. 441–468.
111. Fuentes, G.A. Catalyst deactivation and steady-state activity: A generalized power-law
equation model. Appl. Catal. 1985, 15, 33–40.
112. Fuentes, G.A.; Ruiz-Trevino, F.A. Towards a better understanding of sintering phenomena in
catalysis. In Catalyst Deactivation 1991 (Studies in Surface Science and Catalysis);
Bartholomew, C.H., Butt, J.B., Eds., Elsevier: Amsterdam, The Nethelands, 1991; Volume 68,
pp. 637–644.
113. Bournonville, J.P.; Martino, G. Sintering of Alumina Supported Platinum. In Catalyst
Deactivation 1980 (Studies in Surface Science and Catalysis); Delmon, B., Froment, G.F.,
Eds., Elsevier: Amsterdam, The Nethelands, 1980; Volume 6, pp. 159–166.
114. Somorjai, G.A. Small-Angle X-Ray Scattering and Low Energy Electron Diffraction Studies
on Catalyst Surfaces, In X-ray and Electron Methods of Analysis (Progress in Analytical
Chemistry); Van Olphen, H., Parrish, W., Eds., Plenum Press; New York, NY, USA, 1968;
Volume 1, pp. 101–126.

111
115. Seyedmonir, S.R.; Strohmayer, D.E.; Guskey, G.J.; Geoffroy, G.L.; Vannice, M.A.
Characterization of supported silver catalysts: III. Effects of support, pretreatment, and
gaseous environment on the dispersion of Ag. J. Catal. 1985, 93, 288–302.
116. Trimm, D.L. Thermal stability of catalyst supports. In Catalyst Deactivation 1991 (Studies in
Surface Science and Catalysis); Bartholomew, C.H., Butt, J.B. Eds.; Elsevier: Amsterdam,
The Nethelands, 1991; Volume 68, pp. 29–51.
117. Shastri, A.G.; Datye, A.K.; Schwank, J. Influence of chlorine on the surface area and
morphology of TiO2. Appl. Catal. 1985, 14, 119–131.
118. Oberlander, R.K. Aluminas for catalysts: Their preparations and properties. In Applied
Industrial Catalysis; Leach, B.E., Ed.; Academic Press: Orlando, FL, USA, 1984; Volume 3,
pp. 64–112.
119. Wefers, K.; Misra, C. Oxides and Hydroxides of Aluminum; Alcoa Technical Paper No. 19;
Alcoa Laboratories: Pittsburg, PA, USA, 1987.
120. Hegedus, L.L.; Baron, K. Phosphorus accumulation in automotive catalysts. J. Catal. 1978,
54, 115–119.
121. Summers, J.; Hegedus, L.L. Modes of catalyst deactivation in stoichiometric automobile
exhaust. Ind. Eng. Chem. Prod. Res. Dev. 1979, 18, 318–324.
122. Peter-Hoblyn, J.D.; Valentine, J.M.; Sprague, B.N.; Epperly, W.R. Methods for reducing
harmful emissions from a diesel engine. U.S. Patent 6,003,303, 21 December 1999.
123. Manson, I. Self-regenerating diesel exhaust particulate filter and material. U.S. Patent
6,013,599, 11 January 2000.
124. Deeba, M.; Lui, Y.K.; Dettling, J.C. Four-way diesel exhaust catalyst and method of use. U.S.
Patent 6,093,378, 25 July 2000.
125. Dry, M.E. The fischer-tropsch synthesis. In Catalysis—Science and Technology; Anderson,
J., Boudart, M., Eds.; Springer-Verlag: New York, NY, USA, 1981; pp. 159–218.
126. Huber, G.W.; Guymon, C.G.; Stephenson, B.C.; Bartholomew, C.H. Hydrothermal stability
of Co/SiO2 Fischer-Tropsch synthesis catalysts. In Catalyst Deactivation 2001 (Studies in
Surface Science and Catalysis); Spivey, J.J., Roberts, G.W., Davis, B.H., Eds.; Elsevier:
Amsterdam, The Netheland, 2001; Volume 139, pp. 423–430.
127. Busca, G.; Lietti, L.; Ramis, G.; Berti, F. Chemical and mechanistic aspects of the selective
catalytic reduction of NOx by ammonia over oxide catalysts: A review. Appl. Catal. B 1998,
18, 1–36.
128. Kobylinski, T.P.; Taylor, B.W.; Yong, J.E. Stabilized ruthenium catalysts For NOx reduction.
Proc. SAE, 1974, Paper 740250.
129. Shelef, M.; Gandhi, H.S. The reduction of nitric oxide in automobile emissions: Stabilisation
of catalysts containing ruthenium. Platinum Met. Rev. 1974, 18, 2–14.
130. Gandhi, H.S.; Stepien, H.K.; Shelef, M. Optimization of ruthenium-containing, stabilized,
nitric oxide reduction catalysts. Mat. Res. Bull. 1975, 10, 837–845.
131. Bartholomew, C.H. Reduction of nitric oxide by monolithic-supported palladium-nickel and
palladium-ruthenium alloys. Ind. Eng. Chem. Prod. Res. Dev. 1975, 14, 29–33.

112
132. Clark, R.W.; Tien, J.K.; Wynblatt, P. Loss of palladium from model platinum-palladium
supported catalysts during annealing. J. Catal. 1980, 61, 15–18.
133. Shen, W.M.; Dumesic, J.A.; Hill, C.G. Criteria for stable Ni particle size under methanation
reaction conditions: Nickel transport and particle size growth via nickel carbonyl. J. Catal.
1981, 68, 152–165.
134. Pannell, R.B.; Chung, K.S.; Bartholomew, C.H. The stoichiometry and poisoning by sulfur of
hydrogen, oxygen and carbon monoxide chemisorption on unsupported nickel. J. Catal. 1977,
46, 340–347.
135. Lohrengel, G.; Baerns, M. Determination of the metallic surface area of nickel and its
dispersion on a silica support by means of a microbalance. Appl. Catal. 1981, 1, 3–7.
136. Qamar, I.; Goodwin, J.G. Fischer-Tropsch Synthesis over Composite Ru Catalysts In
Proceedings of 8th North American Catalysis Society Meeting, Philadelphia, PA, USA, 1983,
Paper C-22.
137. Goodwin, J.G.; Goa, D.O.; Erdal, S.; Rogan, F.H. Reactive metal volatilization from
Ru/Al2O3 as a result of ruthenium carbonyl formation. Appl. Catal. 1986, 24, 199–209.
138. Watzenberger, O.; Haeberle, T.; Lynch, D.T.; Emig, G. Deactivation of Heteropolyacid
Catalysts. In Catalyst Deactivation 1991 (Studies in Surface Science and Catalysis);
Bartholomew, C.H., Butt, J.B., Eds.; Elsevier: Amsterdam, The Nethelands, 1991; Volume
68, pp. 441–448.
139. Agnelli, M.; Kolb, M.; Mirodatos, C. Co hydrogenation on a nickel catalyst: 1. Kinetics and
modeling of a low-temperature sintering process. J. Catal. 1994, 148, 9–21.
140. Lee, H.C.; Farrauto, R.J. Catalyst deactivation due to transient behavior in nitric acid
production. Ind. Eng. Chem. Res. 1989, 28, 1–5.
141. Farrauto, R.J.; Lee, H.C. Ammonia oxidation catalysts with enhanced activity. Ind. Eng.
Chem. Res. 1990, 29, 1125–1129.
142. Sperner, F.; Hohmann, W. Rhodium-platinum gauzes for ammonia oxidation. Platinum Met.
Rev. 1976, 20, 12–20.
143. Hess, J.M.; Phillips, J. Catalytic etching of Pt/Rh gauzes. J. Catal. 1992, 136, 149–160.
144. Kuo, C.L.; Hwang, K.C. Does morphology of a metal nanoparticle play a role in ostwald
ripening processes? Chem. Mater. 2013, 25, 365–371.
145. Bartholomew, C.H. Hydrogen adsorption on supported cobalt, iron, and nickel. Catal. Lett.
1990, 7, 27–51.
146. Wu, N.L.; Phillips, J. Catalytic etching of platinum during ethylene oxidation. J. Phys. Chem.
1985, 89, 591–600.
147. Wu, N.L.; Phillips, J. ReactionǦenhanced sintering of platinum thin films during ethylene
oxidation. J. Appl. Phys. 1986, 59, 769–779.
148. Wu, N.L.; Phillips, J. Sintering of silica-supported platinum catalysts during ethylene
oxidation. J. Catal. 1988, 113, 129–143.

113
149. Bielanski, A.; Najbar, M.; Chrzaoszcz, J.; Wal, W. Deactivation of the V2O5-MoO3, catalysts
in the selective oxidation of eenzene to maleic anhydride and the changes in its morphology
and chemical composition. In Catalyst Deactivation 1980 (Studies in Surface Science and
Catalysis); Delmon, B., Froment, G.F., Eds.; Elsevier: Amsterdam, The Nethelands, 1980;
Volume 6, pp. 127–140.
150. Burriesci, N.; Garbassi, F.; Petrera, M.; Petrini, G.; Pernicone, N. Solid state reactions in
Fe-Mo oxide catalysts for methanol oxidation during aging in industrial plants. In Catalyst
Deactivation 1980 (Studies in Surface Science and Catalysis); Delmon, B., Froment, G.F.,
Eds.; Elsevier: Amsterdam, The Nethelands, 1980; Volume 6, pp. 115–126.
151. Xiong, Y.L.; Castillo, R.; Papadopoulou, C.; Dada, L.; Ladriere, J.; Ruiz, P.; Delmon, B.
The protecting role of antimony oxide against deactivation of iron molybdate in oxidation
catalysts. In Catalyst Deactivation 1991 (Studies in Surface Science and Catalysis);
Bartholomew, C.H., Butt, J.B., Eds.; Elsevier: Amsterdam, The Nethelands, 1991; Volume 68,
pp. 425–432.
152. Farrauto, R.J.; Hobson, M.; Kennelly, T.; Waterman, E. Catalytic chemistry of supported
palladium for combustion of methane. Appl. Catal. A 1992, 81, 227–237.
153. Gai-Boyes, P.L. Defects in Oxide Catalysts: Fundamental Studies of Catalysis in Action.
Catal. Rev.—Sci. Eng. 1992, 34, 1–54.
154. Delmon, B. Solid-state reactions in catalysts during ageing: Beneficial role of spillover. In
Catalyst Deactivation 1994 (Studies in Surface Science and Catalysis, Volume 88); Delmon,
B., Froment, G.F.; Eds., Elsevier: Amsterdam, The Nethelands, 1994, pp. 113–128.
155. Erickson, K.M.; Karydis, D.A.; Boghosian, S.; Fehrmann, R. Deactivation and compound
formation in sulfuric-acid catalysts and model systems. J. Catal. 1995, 155, 32–42.
156. Delmon, B. Solid state reactions in catalysts: An approach to real active systems and their
deactivation. In Catalyst Deactivation 1997 (Studies in Surface Science and Catalysis);
Bartholomew, C.H., Fuentes, G.A., Eds.; Elsevier: Amsterdam, The Nethelands, 1997;
Volume 111, pp. 39–51.
157. Jackson, N.B.; Datye, A.K.; Mansker, L.; O’Brien, R.J.; Davis, B.H. Deactivation and
attrition of iron catalysts in synthesis gas. In Catalyst Deactivation 1997 (Studies in Surface
Science and Catalysis); Bartholomew, C.H., Fuentes, G.A., Eds.; Elsevier: Amsterdam, The
Nethelands, 1997; Volume 111, pp. 501–516.
158. Eliason, S.A.; Bartholomew, C.H. Temperature-programmed reaction study of carbon
transformations on iron fischer-tropsch catalysts during steady-state synthesis. In Catalyst
Deactivation 1997 (Studies in Surface Science and Catalysis); Bartholomew, C.H.,
Fuentes, G.A., Eds.; Elsevier: Amsterdam, The Nethelands, 1997; Volume 111, pp. 517–526.
159. Baranski, A.; Dziembaj, R.; Kotarba, A.; Golebiowski, A.; Janecki, Z.; Pettersson, J.B.C.
Deactivation of iron catalyst by water-potassium thermal desorption studies. In Catalyst
Deactivation 1999 (Studies in Surface Science and Catalysis); Delmon, B., Froment, G.F.,
Eds.; Elsevier: Amsterdam, The Nethelands, 1999; Volume 126, pp. 229–236.

114
160. Querini, C.A.; Ravelli, F.; Ulla, M.; Cornaglia, L.; Miró, E. Deactivation of Co, K catalysts
during catalytic combustion of diesel soot: Influence of the support. In Catalyst Deactivation
1999 (Studies in Surface Science and Catalysis); Delmon, B., Froment, G.F., Eds.; Elsevier:
Amsterdam, The Nethelands, 1999; Volume 126, pp. 257–264.
161. Wilson, J.; de Groot, C. Atomic-scale restructuring in high-pressure catalysis. J. Phys. Chem.
1995, 99, 7860–7866.
162. Parkinson, G.S.; Novotny, Z.; Argentero, G.; Schmid, M.; Pavelec, J.; Kosak, R.; Blaha, P.;
Diebold, U. Carbon monoxide-induced adatom sintering in a Pd–Fe3O4 model catalyst.
Nat. Mater. 2013, 12, 724–728.
163. Pham, H.N.; Reardon, J.; Datye, A.K. Measuring the strength of slurry phase heterogeneous
catalysts. Powder Technol. 1999, 103, 95–102.
164. Kalakkad, D.S.; Shroff, M.D.; Kohler, S.; Jackson, N.; Datye, A.K. Attrition of precipitated
iron Fischer-Tropsch catalysts. Appl. Catal. A 1995, 133, 335–350.
165. Callister, W.D. Materials Science and Engineering: An Introduction; John Wiley, & Sons,
Inc.: New York, NY, USA, 2000.
166. Coble, R.L.; Kingery, W.D. Effect of porosity on physical properties of sintered alumina.
J. Am. Ceram. Soc. 1956, 39, 377–385.
167. Deng, S.G.; Lin, Y.S. Granulation of sol-gel-derived nanostructured alumina. AIChE J. 1997,
43, 505–514.
168. Thoma, S.G.; Ciftcioglu, M.; Smith, D.M. Determination of agglomerate strength
distributions: Part 1. Calibration via ultrasonic forces. Powder Technol. 1991, 68, 53–61.
169. Bankmann, M.; Brand, R.; Engler, B.H.; Ohmer, J. Forming of high surface area TiO2 to
catalyst supports. Catal. Today 1992, 14, 225–242.
170. Kenkre, V.M.; Endicott, M.R. A theoretical model for compaction of granular materials.
J. Am. Ceram. Soc. 1996, 79, 3045–3054.
171. Song, H.; Evans, J.R.G. A die pressing test for the estimation of agglomerate strength. J. Am.
Ceram. Soc. 1994, 77, 806–814.
172. Werther, J.; Xi, W. Jet attrition of catalyst particles in gas fluidized beds. Powder Technol.
1993, 76, 39–46.
173. Bhatt, B.L.; Schaub, E.S.; Hedorn, E.C.; Herron, D.M.; Studer, D.W.; Brown, D.M. Liquid
phase Fischer-Tropsch synthesis in a bubble column. In Proceedings of Liquefaction
Contractors Review Conference; Stiegel, G.J., Srivastava, R.D., Eds.; U.S. Department of
Energy: Pittsburgh, PA, USA, 1992; pp. 402–423.
174. Pham, H.N.; Datye, A.K. The synthesis of attrition resistant slurry phase iron Fischer-Tropsch
catalysts. Catal. Today 2000, 58, 233–240.
175. Bemrose, C.R.; Bridgewater, J. A review of attrition and attrition test methods. Powder
Technol. 1987, 49, 97–126.
176. Ghadiri, M.; Cleaver, J.A.S.; Tuponogov, V.G.; Werther, J. Attrition of FCC powder in the
jetting region of a fluidized bed. Powder Technol. 1994, 80, 175–178.
177. Weeks, S.A.; Dumbill, P. Method speeds FCC catalyst attrition resistance determinations.
Oil Gas J. 1990, 88, 38–40.

115
178. Zhao, R.; Goodwin, J.G.; Jothimurugesan, K.; Spivey, J.J.; Gangwal, S.K. Comparison of
attrition test methods: ASTM standard fluidized bed vs jet cup. Ind. Eng. Chem. Res. 2000,
39, 1155–1158.
179. Doolin, P.K.; Gainer, D.M.; Hoffman, J.F. Laboratory testing procedure for evaluation of
moving bed catalyst attrition. J. Test. Eval. 1993, 21, 481.
180. Oukaci, R.; Singleton, A.H.; Wei, D.; Goodwin, J.G. Attrition resistance of Al2O3-supported
cobalt F-T catalysts. In Proceedings of the 217th National Meeting, ACS Division of
Petroleum Chemistry, Anaheim, CA, USA, 21–25 March 1999; p. 91.
181. Adams, M.J.; Mullier, M.A.; Seville, J.P.K. Agglomerate strength measurement using a
uniaxial confined compression test. Powder Technol. 1994, 78, 5–13.
182. Emig, G.; Martin, F.G. Development of a fluidized bed catalyst for the oxidation of n-butane
to maleic anhydride. Ind. Eng. Chem. Res. 1991, 30, 1110–1116.
183. Silver, R.G.; Summers, J.C.; Williamson, W.B. Catalysis and Automotive Pollution Control
II, Elsevier: Amsterdam, The Nethelands, 1991; p. 167.
184. Fisher, G.B.; Zammit, M.G.; LaBarge, J. Investigation of Catalytic Alternatives to Rhodium
in Emissions Control; SAE Report 920846; SAE International: Warrendale, PA, USA, 1992.
185. Farrauto, R.J.; Heck, R.M. Catalytic converters: state of the art and perspectives. Catal.
Today 1999, 51, 351–360.
186. Beer, G.L. Extended catalyst life two stage hydrocarbon synthesis process. U.S. Patent
6,169,120, 2 January 2001.
187. Bartholomew, C.H.; Stoker, M.W.; Mansker, L.; Datye, A. Effects of pretreatment, reaction,
and promoter on microphase structure and Fischer-Tropsch activity of precipitated iron
catalysts. In Catalyst Deactivation 1999 (Studies in Surface Science and Catalysis);
Delmon, B., Froment, G.F., Eds.; Elsevier: Amsterdam, The Nethelands, 1999; Volume 126,
pp. 265–272.
188. Nagano, O.; Watanabe, T. Method for producing methacrolein. U.S. Patent 5,728,894,
17 March 1998.
189. Maillet, T.; Barbier, J.; Duprez, D. Reactivity of steam in exhaust gas catalysis III. Steam and
oxygen/steam conversions of propane on a Pd/Al2O3 catalyst. Appl. Catal. B 1996, 9,
251–266.
190. Mathys, G.M.K.; Martens, L.R.M.; Baes, M.A.; Verduijn, J.P.; Huybrechts, D.R.C. Alkene
oligomerization. WIPO Patent WO 1993/016020 (A3), 16 September 1993.
191. Mathys, G.M.K.; Martens, L.R.M.; Baes, M.A.; Verduijn, J.P.; Huybrechts, D.R.C. Alkene
oligomerization. U.S. Patent 5,672,800, 30 September 1997.
192. Stine, L.O.; Muldoon, B.S.; Gimre, S.C.; Frame, R.R. Process for oligomer production and
saturation. U.S. Patent 6,080,903, 26 June 2000.
193. Subramaniam, B.; Arunajatesan, V.; Lyon, C.J. Coking of solid acid catalysts and strategies
for enhancing their activity. In Catalyst Deactivation 1999 (Studies in Surface Science and
Catalysis); Delmon, B., Froment, G.F., Eds.; Elsevier: Amsterdam, The Nethelands, 1999;
Volume 126, pp. 63–77.

116
194. Ginosar, D.M.; Fox, R.V.; Kong, P.C. Solid catalyzed isoparaffin alkylation at supercritical
fluid and near-supercritical fluid conditions. WIPO Patent WO 1999/033769 (A1), 8 July 1999.
195. Ribeiro, F.H.; Bonivardi, A.L.; Kim, C. Transformation of platinum into a stable,
high-temperature, dehydrogenation-hydrogenation catalyst by ensemble size reduction with
rhenium and sulfur. J. Catal. 1994, 150, 186–198.
196. Ginosar, D.; Subramaniam, B. Coking and activity of a reforming catalyst in near-critical and
dense supercritical reaction mixtures. In Catalyst Deactivation 1994 (Studies in Surface
Science and Catalysis); Delmon, B., Froment, G.F., Eds.; Elsevier: Amsterdam, The
Nethelands, 1994; Volume 88, pp. 327–334.
197. Petersen, E.E. Catalyst deactivation: Opportunity amidst woe. In Catalyst Deactivation 1997
(Studies in Surface Science and Catalysis); Bartholomew, C.H., Fuentes, G.A., Eds.; Elsevier:
Amsterdam, The Nethelands, 1997; Volume 111, pp. 87–98.
198. Gosselink, J.W.; Veen, J.A.R.V. Coping with catalyst deactivation in hydrocarbon
processing. In Catalyst Deactivation 1999 (Studies in Surface Science and Catalysis);
Delmon, B., Froment, G.F., Eds.; Elsevier: Amsterdam, The Nethelands, 1999; Volume 126,
pp. 3–16.
199. Lin, L.; Zao, T.; Zang, J.; Xu, Z. Dynamic process of carbon deposition on Pt and Pt–Sn
catalysts for alkane dehydrogenation. Appl. Catal. 1990, 67, 11–23.
200. Resasco, D.E.; Haller, G.L. Catalytic dehydrogenation of lower alkanes. In Catalysis
(Specialist Periodiodical Report; Spivey, J.J., Agarwal, S.K., Eds.; Royal Society of
Chemistry, Cambridge, UK, 1994; Volume 11, pp. 379–411.
201. Cortright, R.D.; Dumesic, J.A. Microcalorimetric, Spectroscopic, and Kinetic Studies of
Silica Supported Pt and Pt/Sn Catalysts for Isobutane Dehydrogenation. J. Catal. 1994, 148,
771–778.
202. Weyten, H.; Keizer, K.; Kinoo, A.; Luyten, J.; Leysen, R. Dehydrogenation of propane using
a packed-bed catalytic membrane reactor. AIChE J. 1997, 43, 1819–1827.
203. Praserthdam, P.; Mongkhonsi, T.; Kunatippapong, S.; Jaikaew, B.; Lim, N. Determination of
coke deposition on metal active sites of propane dehydrogenation catalysts. In Catalyst
Deactivation 1997 (Studies in Surface Science and Catalysis); Bartholomew, C.H.,
Fuentes, G.A., Eds.; Elsevier: Amsterdam, The Nethelands, 1997; Volume 111, pp. 153–158.
204. Rose, B.H.; Kiliany, T.R. Dual catalyst system. WIPO Pat. App. WO 2000/069993 (A1),
12 May 2000.
205. Guerrero-Ruiz, A.; Sepulveda-Escribano, A.; Rodriguez-Ramos, I. Cooperative action of
cobalt and MgO for the catalysed reforming of CH4 with CO2. Catal. Today 1994, 21,
545–550.
206. Qin, D.; Lapszewicz, J. Study of mixed steam and CO2 reforming of CH4 to syngas on
MgO-supported metals. Catal. Today 1994, 21, 551–560.
207. Stagg, S.; Resasco, D. Effects of promoters and supports on coke formation on Pt catalysts
during CH4 reforming with CO2. In Catalyst Deactivation 1997 (Studies in Surface Science
and Catalysis); Bartholomew, C.H., Fuentes, G.A., Eds.; Elsevier: Amsterdam, The
Nethelands, 1997; Volume 111, pp. 543–550.

117
208. Fujimoto, K.; Tomishige, K.; Yamazaki, O.; Chen, Y.; Li, X. Development of catalysts for
natural gas reforming: Nickel-magnesia solid solution catalyst. Res. Chem. Intermed. 1998,
24, 259–271.
209. Marshall, C.L.; Miller, J.T. Process for converting methanol to olefins or gasoline. U.S.
Patent 5,191,142, 2 March 1993.
210. Gayubo, A.G.; Aguayo, A.T.; Campo, A.E.S.D.; Benito, P.L.; Bilbao, J. The role of water on
the attenuation of coke deactivation of a SAPO-34 catalyst in the transformation of methanol
into olefins. In Catalyst Deactivation 1999 (Studies in Surface Science and Catalysis);
Delmon, B., Froment, G.F., Eds.; Elsevier: Amsterdam, The Nethelands, 1999; Volume 126,
pp. 129–136.
211. Barger, P.T. SAPO catalysts and use thereof in methanol conversion processes. U.S. Patent
5,248,647, 28 Sepetember 1993.
212. Cox, J.P.; Evans, J.M. Exhaust gas catalyst for two-stroke engines. WIPO Patent WO
1999/042202 (A1), 20 February 1998.
213. Leviness, S.C.; Mart, C.J.; Behrmann, W.C.; Hsia, S.J.; Neskora, D.R. Slurry hydrocarbon
synthesis process with increased catalyst life. WIPO Patent WO 1998/050487 (A1),
2 May 1997.
214. Bartholomew, C.H. Catalyst deactivation in hydrotreating of residua: A review. In Catalytic
Hydroprocessing of Petroleum and Distillates; Oballa, M., Shih, S. Eds.; Marcel Dekker:
New York, NY, USA, 1993; pp. 1–32.
215. Summers, J.; Williamson, W.B. Palladium-only catalysts for closed-loop control. In
Environmental Catalysis 1993; Armor, J.N., Ed.; American Chemical Society: Washington,
DC, USA, 1993; Volume 552, pp. 94–113.
216. Dettling, J.; Hu, Z.; Lui, Y.K.; Smaling, R.; Wan, C.Z.; Punke, A. Smart Pd TWC technology
to meet stringent standards. In Studies in Surface Science and Catalysis; Elsevier:
Amsterdam, The Nethelands, 1995; Volume 96, pp. 461–472.
217. Berndt, M.; Ksinsik, D. Catalyst and process for its preparation. U.S. Patent 4,910,180, 20
March 1990.
218. Prigent, M.; Blanchard, G.; Phillippe, P. Catalyst supports/catalysts for the treatment of
vehicular exhaust gases. U.S. Patent 4,985,387, 15 January 1991.
219. Williamson, W.B.; Linden, D.G.; Summers, J.C., II. High-temperature three-way catalyst for
treating automotive exhaust gases. U.S. Patent 5,041,407, 20 August 1991.
220. Williamson, W.B.; Linden, D.G.; Summers, J.C., II. High durability and exhuast catalyst with
low hydrogen sulfide emissions. U.S. Patent 5,116,800, 26 May 1992.
221. Narula, C.K.; Watkins, W.L.H.; Chattha, M.S. Binary La-Pd oxide catalyst and method of
making the catalyst. U.S. Patent 5,234,881, 10 August 1993.
222. Wan, C.-Z.; Tauster, S.J.; Rabinowitz, H.N. Catalyst composition containing platinum and
rhodium components. U.S. Patent 5,254,519, 19 October 1993.
223. Argyle, M.D.; Frost, T.S.; Bartholomew, C.H. Modeling cobalt fischer tropsch catalyst
deactivation with generalized power law expressions. Top. Catal. 2014, 57, 415–429.

118
224. Furuya, T.; Yamanaka, S.; Hayata, T.; Koezuka, J.; Yoshine, T.; Ohkoshi, A. Hybrid catalytic
combustion for stationary gas turbine—Concept and small scale test results. In Proceedings of
Gas Turbine Conference and Exhibition, Anaheim, CA, USA, 31 May–4 June, 1987; ASME
Paper 87-GT-99.
225. Kawakami, T.; Furuya, T.; Sasaki, Y.; Yoshine, T.; Furuse, Y.; Hoshino, M. Feasibility study
on honeycomb ceramics for catalytic combustor. In Proceedings of Gas Turbine and
Aeroengine Congress and Exposition, Toronto, ON, Canada, 4–8 June 1989; ASME Paper
89-GT-41.
226. Betta, R.D.; Ribeiro, F.; Shoji, T.; Tsurumi, K.; Ezawa, N.; Nickolas, S. Catalyst structure
having integral heat exchange. U.S. Patent 5,250,489, 5 October 1993.
227. Fujii, T.; Ozawa, Y.; Kikumoto, S. High pressure test results of a catalytic combustor for gas
turbine. J. Eng. Gas Turbines Power 1998, 120, 509–513.
228. Borio, D.O.; Schbib, N.S. Simulation and optimization of a set of catalytic reactors used for
dehydrogenation of butene into butadiene. Comput. Chem. Eng.1995, 19, S345–S350.
229. Fiato, R.A. Two stage process for hydrocarbon synthesis. U.S. Patent 5,028,634, 2 July 1991.
230. Zhao, R.; Goodwin, J.G., Jr.; Jothimurugesan, K.; Gangwal, S.K.; Spivey, J.J. Spray-dried
Iron Fischer-Tropsch catalysts. 1. Effect of structure on the attrition resistance of the catalysts
in the calcined state. Ind. Eng. Chem. Res. 2001, 40, 1065–1075.
231. Zhao, R.; Goodwin, J.G., Jr.; Jothimurugesan, K.; Gangwal, S.K.; Spivey, J.J. Spray-dried
iron Fischer-Tropsch catalysts. 2. Effect of carburization on catalyst attrition resistance. Ind.
Eng. Chem. Res. 2001, 40, 1320–1328.
232. Singleton, A.H.; Oukaci, R.; Goodwin, J.G. Processes and catalysts for conducting
fischer-tropsch synthesis in a slurry bubble column reactor. U.S. Patent 5,939,350, 17
August 1999.
233. Plecha, S.; Mauldin, C.H.; Pedrick, L.E. Titania catalysts, their preparation and use in
Fischer-Tropsch synthesis. U.S. Patent 6,087,405, 11 July 2000.
234. Plecha, S.; Mauldin, C.H.; Pedrick, L.E. Titania catalysts, their preparation and use in
fischer-tropsch synthesis. U.S. Patent 6,124,367, 26 September 2000.
235. Singleton, A.H.; Oukaci, R.; Goodwin, J.G. Reducing fischer-tropsch catalyst attrition losses
in high agitation reaction systems. WIPO Patent WO 2000/071253 (A3), 30 November 2000.
236. Seitz, M.; Klemm, E.; Emig, G. Poisoning and regeneration of NOx adsorbing catalysts for
automotive applications. In Catalyst Deactivation 1999 (Studies in Surface Science and
Catalysis); Delmon, B., Froment, G.F. Eds.; Elsevier: Amsterdam, The Nethelands, 1999;
Volume 126, pp. 211–218.
237. de Lima, S.M.; da Silva, A.M.; da Costa, L.O.O.; Assaf, J.M.; Mattos, L.V.; Sarkari, R.;
Venugopal, A.; Noronha, F.B. Hydrogen production through oxidative steam reforming
of ethanol over Ni-based catalysts derived from La1íxCexNiO3 perovskite-type oxides.
Appl. Catal. B 2012, 121–122, 1–9.
238. Song, H.; Ozkan, U.S. Ethanol steam reforming over Co-based catalysts: Role of oxygen
mobility. J. Catal. 2009, 261, 66–74.

119
239. Masamune, S.; Smith, J.M. Performance of fouled catalyst pellets. AIChE J. 1966, 12,
384–394.
240. Murakami, Y.; Kobayashi, T.; Hattori, T.; Masuda, M. Effect of intraparticle diffusion on
catalyst fouling. Ind. Eng. Chem. Fundam. 1968, 7, 599–605.
241. Stork, W.H.J. Molecules, catalysts and reactors in hydroprocessing of oil fractions. In
Hydrotreatment and Hydrocracking of Oil Fractions (Studies in Surface Science and
Catalysis); Froment, G.F., Delmon, B., Grange, P., Eds.; Elsevier: New York, NY, USA,
1997; Volume 106, pp. 41–67.
242. Parks, G.D.; Schaffer, A.M.; Dreiling, M.J.; Shiblom, C.M. Surface Studies on the Interaction
of Nickel and Antimony on Cracking Catalysts. Prepr.—Am. Chem. Soc., Div. Petr. Chem.
1980, 25, 335.
243. Trimm, D.L. Introduction to Catalyst Deactivation. In Progress in Catalyst Deactivation
(NATO Advanced Study Institute Series E, No. 54); Figueiredo, J.L., Ed.; Martinus Nijhoff
Publishers: Boston, MA, USA,1982; pp. 3–22.
244. Becker, E.R.; Wei, J.J. Nonuniform distribution of catalysts on supports: I. Bimolecular
Langmuir reactions. J. Catal. 1977, 46, 365–381.
245. Long, D.C. Staged hydrocarbon synthesis process. U.S. Patent 5,498,638, 12 March 1996.
246. Powell, B.R., Jr. Stabilitzation of high surface area aluminas. In Proceedings of the Materials
Research Society Annual Meeting, Boston, MA, USA, 16–21 November 1980; Paper H9.
247. Dai, Y.; Lim, B.; Yang, Y.; Cobley, C.M.; Li, W.; Cho, E.C.; Grayson, B.; Fanson, P.T.;
Campbell, C.T.; Sun, Y.; et al. A sinter-resistant catalytic system based on platinum
nanoparticles supported on TiO2 nanofibers and covered by porous silica. Angew. Chem.
Int. Ed. 2010, 49, 8165–8168.
248. Shang, L.; Bian, T.; Zhang, B.; Zhang, D.; Wu, L.-Z.; Tung, C.-H.; Yin, Y.; Zhang, T.
Graphene-supported ultrafine metal nanoparticles encapsulated by mesoporous silica: Robust
catalysts for oxidation and reduction reactions. Angew. Chem. Int. Ed. 2014, 53, 250–254.
249. Heck, R.; Farrauto, R. Catalytic Air Pollution Control: Commercial Technology;
Van Nostrand Reinhold: New York, NY, USA, 1995.
250. Trimm, D.L. The regeneration or disposal of deactivated heterogeneous catalysts.
Appl. Catal. A 2001, 212, 153–160.
251. Berrebi, G.; Dufresne, P.; Jacquier, Y. Recycling of spent hydroprocessing catalysts:
EURECAT technology. Environ. Prog. 1993, 12, 97–100.
252. D’Aquino, R.L. For catalyst regeneration, incentives anew. Chem. Eng. 2000, 107, 32–33, 35.
253. Chang, T. Regeneration industry helps refiners control costs, limit liabilities. Oil Gas J. 1998,
96, 49.
254. Blashka, S.R.; Duhon, W. Catalyst regeneration. Int. J. Hydrocarbon Eng.1998, 4, 60–64.
255. McCulloch, D.C. Catalytic hydrotreating in petroleum refining. In Applied Industrial
Catalysis; Leach, B.E., Ed.; Academic Press: New York, NY, USA, 1983; pp. 69–121.
256. Franck, J.P.; Martino, G. Deactivation and Regeneration of Catalytic-Reforming Catalysts.
In Progress in Catalyst Deactivation (NATO Advanced Study Institute Series E., No. 54);
Figueiredo, J.L., Ed.; Martinus Nijhoff Publishers: Boston, MA, USA, 1982; pp. 355–397.

120
257. Spivey, J.J.; Roberts, G.W.; Davis, B.H. Catalyst Deactivation 2001 (Studies in Surface
Science and Catalysis; Elsevier: Amsterdam, The Nethelands, 2001; Volume 139.
258. Haddad, J.H.; Harandi, M.N.; Owen, H. Upgrading light olefin fuel gas in a fluidized bed
catalyst reactor and regeneration of the catalyst. U.S. Patent 5,043,517, 27 August 1991.
259. Ueda, M.; Murakami, T.; Shibata, S.; Hirabayashi, K.; Kondoh, T.; Adachi, K.; Hoshino, N.;
Inoue, S. Process for the regeneration of coke-deposited, crystalline silicate catalyst. U.S.
Patent 5,306,682, 26 April 1994.
260. Zhang, S.Y.-F.; Gosling, C.D.; Sechrist, P.A.; Funk, G.A. Dual regeneration zone solid
catalyst alkylation process. U.S. Patent 5,675,048, 7 October 1997.
261. Panattoni, G.; Querini, C.A. Isobutane alkylation with C4 olefins: regeneration of
metal-containing catalysts. In Catalyst Deactivation 2001 (Studies in Surface Science
and Catalysis); Spivey, J.J., Roberts, G.W., Davis, B.H., Eds.; Elsevier: Amsterdam,
The Nethelands, 2001; Volume 139, pp. 181.
262. Thompson, D.N.; Ginosar, D.M.; Burch, K.C. Regeneration of a deactivated USY alkylation
catalyst using supercritical isobutene. Appl. Catal. A 2005, 279, 109–116.
263. Petkovic, L.M.; Ginosar, D.M.; Burch, K.C. Supercritical fluid removal of hydrocarbons
adsorbed on wide-pore zeolite catalysts. J. Catal. 2005, 234, 328–339.
264. Dufresne, P.; Brahma, N. Offsite regeneration process for a catalyst containing at least one
precious metal. U.S. Patent 5,854,162, 29 December 1998.
265. Innes, R.A.; Holtermann, D.L.; Mulaskey, B.F. Low temperature regeneration of coke
deactivated reforming catalysts. U.S. Patent 5,883,031, 16 March 1999.
266. Fung, S.C. Regenerating a reforming catalyst—this bimetallic noble-metal catalyst is used in
a refining process to convert saturated-hydrocarbons to aromatics - simply burning off the
carbon was not enough—a procedure had to be found for redispersing the metals on the
alumina support as well. Chemtech 1994, 24, 40–44.
267. Alfonso, J.C.; Aranda, D.A.G.; Schmal, M.; Frety, R. Regeneration of a Pt-Sn/Al2O3 catalyst:
Influence of heating rate, temperature and time of regeneration. Fuel Proc. Technol. 1997, 50,
35–48.
268. Arteaga, G.J.; Anderson, J.A.; Rochester, C.H. Effects of Catalyst Regeneration with and
without Chlorine on Heptane Reforming Reactions over Pt/Al2O3 and Pt–Sn/Al2O3. J. Catal.
1999, 187, 219–229.
269. Pieck, C.L.; Vera, C.R.; Parera, J.M. Study of industrial and laboratory regeneration of
Pt-Re/Al2O3 catalysts. In Catalyst Deactivation 2001 (Studies in Surface Science and
Catalysis); Spivey, J.J., Roberts, G.W., Davis, B.H., Eds.; Elsevier: Amsterdam, The
Nethelands, 2001; Volume 139, pp. 279–286.
270. Acharya, D.R.; Hughes, R.; Kennard, M.A.; Liu, Y.P. Regeneration of fixed beds of coked
chromia—alumina catalyst. Chem. Eng. Sci. 1992, 47, 1687–1693.
271. Sikkenga, D.L.; Zaenger, I.C.; Williams, G.S. Palladium catalyst reactivation. U.S. Patent
4,999,326, 12 March 1991.
272. Ekstrom, A.; Lapszewicz, J. On the role of metal carbides in the mechanism of the
Fischer-Tropsch reaction. J. Phys. Chem. 1984, 88, 4577–4580.

121
273. Ekstrom, A.; Lapszewicz, J. The reactions of cobalt surface carbides with water and their
implications for the mechanism of the Fischer-Tropsch reaction. J. Phys. Chem. 1987, 91,
4514–4619.
274. Pedrick, L.E.; Mauldin, C.H.; Behrmann, W.C. Draft tube for catalyst rejuvenation and
distribution. U.S. Patent 5,268,344, 7 December 1993.
275. Owen, H.; Schipper, P.H. Process and apparatus for regeneration of FCC catalyst with
reduced NOx and or dust emissions. U.S. Patent 5,338,439,16 August 1994.
276. Raterman, M.F. Two-stage fluid bed regeneration of catalyst with shared dilute phase.
U.S. Patent 5,198,397, 30 March 1993.
277. Apelian, M.R.; Fung, A.S.; Hatzikos, G.H.; Kennedy, C.R.; Lee, C.-H.; Kiliany, T.R.;
Ng, P.K.; Pappal, D.A. Regeneration of noble metal containing zeolite catalysts via partial
removal of carbonaceous deposits. U.S. Patent 5,393,717, 28 February 1995.
278. Clark, D.E. Hydrocracking process using a reactivated catalyst. U.S. Patent 5,340,957, 23
August 1994.
279. Yoshimura, Y.; Sato, T.; Shimada, H.; Matsubayashi, N.; Imamura, M.; Nishijima, A.;
Yoshitomi, S.; Kameoka, T.; Yanase, H. Oxidative regeneration of spent molybdate and
tungstate hydrotreating catalysts. Energy Fuels 1994, 8, 435–445.
280. Oh, E.-S.; Park, Y.-C.; Lee, I.-C.; Rhee, H.-K. Physicochemical changes in
hydrodesulfurization catalysts during oxidative regeneration. J. Catal. 1997, 172, 314–321.
281. Aquavo, A.T.; Gayubo, A.G.; Atutxa, A.; Olazar, M.; Bilbao, J. Regeneration of a catalyst
based on a SAPO-34 used in the transformation of methanol into olefins. J. Chem. Tech.
Biotech. 1999, 74, 1082–1088.
282. Forbus, N.P.; Wu, M.M.-S. Regeneration of methanol/methyl ether conversion catalysts. U.S.
Patent 4,777,156, 11 October 1988.
283. Krishna, A.; Hsieh, C.; English, A.E.; Pecoraro, T.; Kuehler, C. Additives improve FCC
process; increase catalyst life, lower regeneration temperatures and improve yield and quality
of products cost effectively. Hydrocarbon Processing 1991, 59–66.
284. Altomare, C.; Koermer, G.; Schubert, P.; Suib, S.; Willis, W. A designed fluid cracking
catalyst with vanadium tolerance. Chem. Mater. 1989, 1, 459–463.
285. Aguinaga, A.; Montes, M. Regeneration of a nickel/silica catalyst poisoned by thiophene.
Appl. Catal. A 1992, 90, 131–144.
286. Lowry, G.V.; Reinhard, M. Pd-Catalyzed TCE Dechlorination in Groundwater: Solute
Effects, Biological Control, and Oxidative Catalyst Regeneration. Environ. Sci. Technol.
2000, 34, 3217–3223.
287. Trinh, D.C.; Desvard, A.; Martino, G. Process for regenerating used catalysts by means of
hydrogen peroxide aqueous solution stabilized with an organic compound. U.S. Patent
4,830,997, 16 May 1989.
288. Sherwood, D.E. Process for the reactivation of spent alumina-supported hydrotreating
catalysts. U.S. Patent 5,230,791, 27 July 1993.

122
289. Fung, S.C. Deactivation and regeneration/redispersion chemistry of Pt/KL-zeolite. In Catalyst
Deactivation 2001 (Studies in Surface Science and Catalysis); Spivey, J.J., Roberts, G.W.,
Davis, B.H., Eds.; Elsevier: Amsterdam, The Nethelands, 2001; Volume 139, pp. 399–406.
290. Didillon, B. Catalyst regeneration process and use of the catalyst in hydrocarbon conversion
processes. U.S. Patent 5,672,801, 30 September 1997.
291. Tsao, Y.-Y.P.; von Ballmoos, R. Reactivating catalysts containing noble metals on molecular
sieves containing oxides of aluminum and phosphorus. U.S. Patent 4,929,576, 29 May 1990.
292. Dufresne, P.; Brahma, N.; Girardier, F. Off-site regeneration of hydroprocessing catalysts.
Rev. Inst. Fr. Pet. 1995, 50, 283–293.
293. Clark, F.T.; Hensley, A.L., Jr. Process for regenerating a spent resid hydroprocessing catalyst
using a group IIA metal. U.S. Patent 5,275,990, 4 January 1994.
294. Brito, A.; Arvelo, R.; Gonzalez, A.R. Variation in structural characteristics of
a hydrotreatment catalyst with deactivation/regeneration cycles. Ind. Eng. Chem. Res. 1998,
37, 374–379.
295. Teixeira-da-Silva, V.L.S.; Lima, F.P.; Dieguez, L.C. Regeneration of a deactivated
hydrotreating catalyst. Ind. Eng. Chem. Res. 1998, 37, 882–886.
296. Snape, C.E.; Diaz, M.C.; Tyagi, Y.R.; Martin, S.C.; Hughes, R. Characterisation of coke on
deactivated hydrodesulfurisation catalysts and a novel approach to catalyst regeneration. In
Catalyst Deactivation 2001 (Studies in Surface Science and Catalysis); Spivey, J.J.,
Roberts, G.W., Davis, B.H., Eds.; Elsevier: Amsterdam, The Nethelands, 2001; Volume 139,
pp. 359–365.
297. Sherwood, D.E., Jr.; Hardee, J.R., Jr. Method for the reactivation of spent alumina-supported
hydrotreating catalysts. U.S. Patent 5,445,728, 29 August 1995.
298. Maholland, M.K.; Fu, C.-M.; Lowery, R.E.; Kubicek, D.H.; Bertus, B.J. Demetallization and
passivation of spent cracking catalysts. U.S. Patent 5,021,377, 4 June 1991.
299. Kubicek, D.H.; Fu, C.-M.; Lowery, R.E.; Maholland, K.M. Reactivation of spent cracking
catalysts. U.S. Patent 5,141,904, 25 August 1992.
300. Fu, C.-M.; Maholland, M.; Lowery, R.E. Reactivation of spent, metal-containing cracking
catalysts. U.S. Patent 5,151,391, 29 September 1992.
301. Hu, Y.; Luo, B.; Sun, K.; Yang, Q.; Gong, M.; Hu, J.; Fang, G.; Li, Y. Dry
regeneration-demetalization technique for catalyst for residuum and/or heavy oil catalytic
cracking. U.S. Patent 6,063,721, 16 May 2000.
302. Walker, P.L., Jr.; Rusinko, F., Jr.; Austin, L.G. Gas Reactions of Carbon. Adv. Catal. 1959,
11, 133–221.
303. Fulton, J.W. Regenerating Spent Catalyst. Chem. Eng. 1988, 96, 111–114.
304. Richardson, J.T. Experimental determination of catalyst fouling parameters. Carbon profiles.
Ind. Eng. Chem. Process Des. Dev. 1972, 11, 8–11.
305. Weisz, P.B.; Goodwin, R.B. Combustion of carbonaceous deposits within porous catalyst
particles: II. Intrinsic burning rate. J. Catal. 1966, 6, 227–236.
306. Rostrup-Nielsen, J.R. Some principles relating to the regeneration of sulfur-poisoned
nickel catalyst. J. Catal. 1971, 21, 171–178.

123
307. Hartenstein, H. Feasibility of SCR Technology for NOx Control Technology for the Milton R.
Young Station, Center, North Dakota; Expert Report on behalf of the United States
Department of Justice: Washington, DC, USA, July 2008.
308. Hartenstein, H.; Ehrnschwender, M.; Sibley, A.F. SCR Regeneration-10 Years of R&D and
Commercial Application. In Proceedings of Power Plant Air Pollution Control Mega
Symposium, Baltimore, MD, USA, 24–28 August 2008; Paper No. 104.
309. EPA; Office of Air and Radiation. Update on Cap and Trade Programs for SO2 and NOx.
Presented at the Enviornmental Markets Association, 12 Annual Spring Conference, Miami,
FL, USA, 29April–2 May 2008.
310. EPA. NOx Budget Trading Program—Basic Information. Available online: http://www.epa.
gov/airmarkt/progsregs/nox/docs/NBPbasicinfo.pdf (accessed on 31 October 2014).
311. McIlvaine Company. News Release: Coal to Play Bigger Role in Future Power Generation
With World Capacity of 1450 GW by 2012. Available online: http://www.mcilvainecompany.
com/brochures/Alerts_for_Internet/s9%20moni/1104coal%20to%20play%20bigger%20role
%20in%20future%20power%20generation.htm (accessed on 31 October 2014).
312. Hjalmarsson, A.-K. NOx Control Technologies for Coal Combustion; IEACR/24; IEA Coal
Research: Lexington, KY, USA, 1990.
313. U.S. Department of Energy; Southern Company Services. Control of Nitrogen Oxide
Emissions: Selective Catalytic Reduction. Clean Coal Technology; Topical Report 9; U.S.
Department of Energy: Pittsburgh PA, USA; Southern Company Services: Birmingham AL,
USA; 1997.
314. Janssen, F.; Meijer, R. Quality control of DeNOx catalysts: Performance testing, surface
analysis and characterization of DeNOx catalysts. Catal. Today 1993, 16, 157–185.
315. Ashton, J.; Nackos, A.; Bartholomew, C.H.; Hecker, W.C.; Baxter, L. Poisoning/Deactivation
of Vanadia/Titanium Dioxide SCR Catalysts in Coal Systems. Presented at the 19th Annual
Western States Catalysis Club Symposium, Albuquerque, NM, USA, 25 February 2005.
316. Guo, X.; Nackos, A.; Ashton, J.; Bartholomew, C.H.; Hecker, W.C.; Baxter, L. Poisoning
Study of V2O5-WO3/TiO2 Catalysts by Na, K, and Ca. Presented at 19th Annual Western
States Catalysis Club Symposium, Albuquerque, NM, USA, 25 February 2005.
317. Forzatti, P.; Lietti, L.; Tronconi, E. Nitrogen Oxides Removal–Industrial. In Encyclopedia of
Catalysis; Horváth, T.I., Ed.; Wiley, Hoboken, NJ, USA, 2003.
318. Rubin, E.S., Kalagnanam, J.R., Frey, C.H., Berkenpas, M.B. Integrated Environmental
Control Modeling of Coal-Fired Power Systems. Journal of the Air and Waste Management
Association 1997, 47, 1180-1188.
319. Cochran, J.R.; Ferguson, A.W. Selective Catalytic Reduction for NOx Emission Control. In
Proceedings of the First International Conference on Air Pollution Zannetti, P., Brebbia, C.A.,
Garcia Gardea, J.E., Ayala Milian, G., Eds., Monterrey, Mexico, 16-18 Feb. 1993, Volume 1,
pp. 703–718.

124
320. Bullock, P.E.; Hartenstein, H.U. O&M Cost Reduction of a Coal-Fired US Merchant Plant
Through an Optimized SCR Management Strategy Involving Catalyst Regeneration. In
Proceedings of the Conference on Selective Catalytic Reduction and Selective Non-Catalytic
Reduction for NOx Control, Pittsburgh, PA, USA, 15–17 May 2002.
321. Babcock Hitachi’s Rejuvenation of Mehrum SCR Catalyst. FGD and DeNOx Newsletter;
No. 315; The McIlvaine Company: Northfield, IL, USA; 2004.
322. McMahon, B. Catalyst regeneration: The business case. Power 2006, 150, 36-39.
323. Zheng, Y.; Jensen, A.; Johnsson, J. Laboratory investigation of selective catalytic reduction
catalysts: Deactivation by potassium compounds and catalyst regeneration. Ind. Eng.
Chem. Res. 2004, 43, 941–947.
324. Khodayari, R.; Odenbrand, C. Regeneration of commercial SCR catalysts by washing and
sulphation: effect of sulphate groups on the activity. Appl. Catal. B 2001, 33, 277–291.
325. Khodayari, R.; Odenbrand, C. Regeneration of commercial TiO2-V2O5-WO3 SCR catalysts
used in bio fuel plants. Appl. Catal. B 2001, 30, 87–89.
326. Dörr, H.-K.; Koch, G.; Bastuck, W. Method for renewed activation of honeycomb-shaped
catalyst elements for denitrating flue gases. U.S. Patent, 6,387,836 (B1), 14 May 2002.
327. Schneider, P.; Bastuck, W. Method to re-enable honeycomb catalyst elements constructed for
the denitrification of flue gases. DE Patent 10222915 (B4), 15 January 2004.
328. Budin, R.; Krotla, K.; Rabitsch, H. Process for regenerating used deNOx or dedioxin catalytic
converters. U.S. Patent, 6,484,733 (B2), 26 November 2002.
329. Cooper, M.; Harrison, K.; Lin, C. An Experimental Program to Optimize SCR Catalyst
Regeneration for Lower Oxidation of SO2 to SO3. Proceedings of the 2006 Environmental
Controls Conference, Pittsburgh, PA, USA, 16–18 May 2006.
330. Patel, N.; Hartenstein, H.-U.; Wenz, F. Process for decoating a washcoat catalyst substrate.
U.S. Patent 7,559,993 (B1), 14 July 2009.
331. Patel, N.; Hartenstein, H.-U.; Wenz, F. Process for decoating a washcoat catalyst substrate.
U.S. Patent 6,929,701, 16 August 2005.
332. Tate, A.; Skipper, J.; Wenz, F. Environmentally sound handling of deactivated SCR catalyst.
Coal Power, 31 July 2008.
333. Hartenstein, H.U.; Hoffman, T. Method of regeneration of SCR catalyst. U.S. Patent
7,723,251, 25 May 2010.
334. Servatius, P.; Schluttig, A. Lifetime extension of SCR-DeNOx catalysts using SCR-Tech’s
high efficiency ultrasonic regeneration process. In Proceedings of 2001 Conference of
Selective Catalytic Reduction and Selective Non-Catalytic Reduction for NOx Control,
Pittsburgh, PA, USA, 16–18 May 2001.
335. Hartenstein, H.; Gutberlet, H. Catalyst Regeneration—An Integral Part of Proper Catalyst
Management. Proceedings of the 2001 Workshop on Selective Catalytic Reduction,
Baltimore, MD, USA, 13–15 November 2001.

125
336. Wenz, F.; Deneault, R.; Franklin, H.N. The Goals, Challenges and Successes of Regenerating
Selective Catalytic Reduction Catalyst. In Proceedings of the Electric Power 2006
Conference–NOx Control III–SCR Technologies, Atlanta, GA, USA, 2–4 May 2006.
337. CoaLogix. Available online: http://www.coalogix.com (accessed on 31 October 2014).
338. McCullen, S.B.; Wong, S.S.; Huang, T.J. Regeneration of noble metal-highly siliceous zeolite
with sequential hydrogen halide and halogen or organic-halogen compound treatment. U.S.
Patent 4,645,751, 24 February 1987.
339. Borghard, W.S.; Huang, T.J.; McCullen, S.B.; Schoennagel, H.J.; Tsao, Y.-.P.; Wong, S.S.
Redispersion of agglomerated noble metals on zeolite catalysts. U.S. Patent 4,657,874, 14
April 1987.
340. Fung, S.C.; Rice, R.W. Process using halogen/oxygen for reactivating iridium and selenium
containing catalysts. U.S. Patent 4,491,636, 1 January 1985.
341. Fung, S.C. Redispersion of Ir catalysts by low temperature reduction step. U.S. Patent
4,467,045, 21 August 1984.
342. Landolt, G.R.; McHale, W.D.; Schoennagel, H.J. Catalyst regeneration procedure. U.S.
Patent 4,359,400, 16 November 1982.
343. Fung, S.C.; Weissman, W.; Carter, J.L. Reactivation process for iridium-containing catalysts
using low halogen flow rates. U.S. Patent 4,480,046, 30 October 1984.
344. Fung, S.C. Low temperature decoking process for reactivating iridium and selenium
containing catalysts. U.S. Patent 4,492,767, 8 January 1985.
345. Fung, S.C.; Weissman, W.; Carter, J.L.; Kmak, W.S. Reactivating iridium and selenium
containing catalysts with hydrogen halide and oxygen. U.S. Patent 4,491,635, 1 January 1985.
346. Mohr, D.H. Process for regenerating a monofunctional large-pore zeolite catalyst having high
selectivity for paraffin dehydrocyclization. U.S. Patent 4,855,269, 8 August 1989.
347. Fung, S.C.; Tauster, S.J.; Koo, J.Y. Method of regenerating a deactivated catalyst. U.S. Patent
4,925,819, 15 May 1990.
348. Fung, S.C.; Kmak, W.S. Reactivation of iridium-containing catalysts by halide pretreat and
oxygen redispersion. U.S. Patent 4,444,896, 24 April 1984.
349. Fung, S.C.; Weissman, W.; Carter, J.L. Reactivation process for iridium-containing catalysts
using low halogen flow rates. U.S. Patent 4,444,895, 24 April 1984.
350. Boyle, J.P.; Gilbert, J.B. Reactivation of iridium-containing catalysts. U.S. Patent 4,517,076,
14 May 1985.
351. Buss, W.D.; Hughes, T.R. In situ hydrocarbon conversion catalyst regeneration and sulfur
decontamination of vessels communicating with catalyst reactor. U.S. Patent 4,482,637,
13 November 1984.
352. Cheng, W.-H. Regeneration of methanol dissociation catalysts. U.S. Patent 4,855,267, 8
August 1989.
353. Huang, Y.-Y.; LaPierre, R.B.; McHale, W.D. Process for dispersing or redispersing a group
VIII noble metal species on a porous inorganic support. European Patent 0,306,170 (B1), 8
March 1989.

126
354. Hucul, D.A. Redispersal of Ru, Os, Rh and Pd catalysts and processes therewith. U.S. Patent
4,891,346, 2 January, 1990.
355. Fung, S.C.; Rice, R.W. Process for reactivating iridium-containing catalysts. U.S. Patent
4,447,551, 8 May 1984.
356. Fung, S.C. Process for reactivating iridium-containing catalysts in series. U.S. Patent
4,472,514, 18 September 1984.
357. S.C. Fung; R.W. Rice Process for reactivating iridium-containing catalysts. U.S. Patent
4,473,656, 25 September 1984.
358. Huang, Y.-Y.; LaPierre, R.B.; McHale, W.D. Process for dispersing or redispersing
a Group VIII noble metal species on a porous inorganic support. U.S. Patent 4,952,543,
28 August 1990.
359. Fung, S.C.; Weissman, W.; Carter, J.L.; Kmak, W.S. Reactivating iridium-containing
catalysts with hydrogen halide and oxygen. U.S. Patent 4,444,897, 24 April 1984.
360. Fung, S.C. Low temperature decoking process for reactivating iridium containing catalysts.
U.S. Patent 4,472,515 18 September 1984.
361. Dadyburjor, D.B. Regions of validity for models of regeneration of sintered supported metal
catalysts. In Catalyst Deactivation (Studies in Surface Science and Catalysis); Delmon, B.,
Froment, G.F., Eds.; Elsevier: Amsterdam, The Nethelands, 1980; Volume 6, pp. 341–351.
362. Jacobs, G.; Ma, W.; Gao, P.; Todic, B.; Bhatelia, T.; Bukur, D.B.; Davis, B.H. The
application of synchrotron methods in characterizing iron and cobalt Fischer–Tropsch
synthesis catalysts. Catal. Today 2013, 214, 100–139.
363. Van Steen, E.; Claeys, M.; Dry, M.E.; van de Loosdrecht, J.; Viljoen, E.L.; Visagie, J.L.
Stability of nanocrystals:ௗ Thermodynamic analysis of oxidation and re-reduction of cobalt in
water/hydrogen mixtures. J. Phys. Chem. B 2005, 109, 3575–3577.
364. Ma, W.; Jacobs, G.; Ji, Y.; Bhatelia, T.; Bukur, D.B.; Khalid, S.; Davis, B.H. Fischer–Tropsch
synthesis: Influence of CO conversion on selectivities, H2/CO usage ratios, and catalyst
stability for a Ru promoted Co/Al2O3 catalyst using a slurry phase reactor. Top. Catal. 2011,
54, 757–767.
365. Jacobs, G.; Sarkar, A.; Ji, Y.; Patterson, P.M.; Das, T.K.; Luo, M.; Davis, B.H. Fischer-Tropsch
synthesis: Characterization of interactions between reduction promoters and Co for
Co/Al2O3–based GTL catalysts. In Proceedings of AIChE Annual Meeting, San Francisco,
CA, USA, 12–17 November 2006.
366. Das, T.K.; Jacobs, G.; Patterson, P.M.; Conner, W.A.; Davis, B.H. Fischer–Tropsch
synthesis: characterization and catalytic properties of rhenium promoted cobalt alumina
catalysts Fuel 2003, 82, 805–815.
367. Jacobs, G.; Das, T.K.; Patterson, P.M.; Luo, M.; Conner, W.A.; Davis, B.H. Fischer–Tropsch
synthesis: Effect of water on Co/Al2O3 catalysts and XAFS characterization of reoxidation
phenomena. Appl. Catal. A 2004, 270, 65–76.
368. Moodley, D.J.; Saib, A.M.; van de Loosdrecht, J.; Welker-Nieuwoudt, C.A.; Sigwebela, B.H.;
Niemanstsverdriet, J.W. The impact of cobalt aluminate formation on the deactivation of
cobalt-based Fischer–Tropsch synthesis catalysts. Catal. Today 2011, 171, 192–200.

127
369. Jermwongratanachai, T.; Jacobs, G.; Shafer, W.D.; Ma, W.; Pendyala, V.R.R.; Davis, B.H.;
Kitiyanan, B.; Khalid, S.; Cronauer, D.C.; Kropf, A.J.; Marshall, C.L. Fischer–Tropsch
synthesis: Oxidation of a fraction of cobalt crystallites in research catalysts at the onset of FT
at partial pressures mimicking 50% CO conversion. Top. Catal. 2014, 57, 479–490.
370. Azzam, K.; Jacobs, G.; Ma, W.; Davis, B.H. Effect of cobalt particle size on the catalyst
intrinsic activity for Fischer–Tropsch synthesis. Catal. Lett. 2014, 144, 389–394.
371. Saib, A.M.; Borgna, A.; van de Loosdrecht, J.; van Berge, P.J.; Niemanstsverdriet, J.W.
XANES study of the susceptibility of nano-sized cobalt crystallites to oxidation during
realistic Fischer–Tropsch synthesis. Appl. Catal. A 2006, 312, 12–19.
372. Dalai, A.K.; Davis, B.H. Fischer–Tropsch synthesis: A review of water effects on the
performances of unsupported and supported Co catalysts. Appl. Catal. A 2008, 348, 1–15.
373. Sadeqzadeh, M.; Chambrey, S.; Piché, S.; Fongarland, P.; Luck, F.; Curulla-Ferré, D.;
Schweich, D.; Bousquet, J.; Khodakov, A.Y. Deactivation of a Co/Al2O3 Fischer–Tropsch
catalyst by water-induced sintering in slurry reactor: Modeling and experimental
investigations. Catal. Today 2013, 215, 52–59.
374. Sadeqzadeh, M.; Hong, J.; Fongarland, P.; Curulla-Ferre, D.; Luck, F.; Bousquet, J.;
Schweich, D.; Khodakov, A.Y. Mechanistic modeling of cobalt based catalyst sintering in
a fixed bed reactor under different conditions of Fischer–Tropsch synthesis. Ind. Eng.
Chem. Res. 2012, 51, 11955–11964.
375. Soled, S.L.; Kiss, G.; Kliewer, C.; Baumgartner, J.; El-Malki, E.-M. Learnings from Co
Fischer-Tropsch catalyst studies. Present at 245th ACS National Meeting & Exposition,
New Orleans, LA, USA, 7–11 April 2013.
376. Hauman, M.M.; Saib, A.M.; Moodley, D.J.; Plessis, E.D.; Claeys, M.; van Steen, E.
Re-dispersion of cobalt on a model fischer–tropsch catalyst during reduction–oxidation–
reduction cycles. Chem. Catal. Chem. 2012, 4, 1411–1419.
377. Weststrate, C.J.; Kizilkaya, A.C.; Rossen, E.T.R.; Verhoeven, M.W.G.M.; Ciobica, I.M.;
Saib, A.M.; Niemanstsverdriet, J.W. Atomic and Polymeric Carbon on Co(0001): Surface
Reconstruction, Graphene Formation, and Catalyst Poisoning. J. Phys. Chem. C 2012, 116,
11575–11583.
378. Weststrate, C.J.; Hauman, M.M.; Moodley, D.J.; Saib, A.M.; van Steen, E.;
Niemanstsverdriet, J.W. Cobalt Fischer–Tropsch catalyst regeneration: The crucial role of the
kirkendall effect for cobalt redispersion. Top. Catal. 2011, 54, 811–816.
379. Moodley, D.J.; van de Loosdrecht, J.; Saib, A.M.; Overett, M.J.; Datye, A.K.;
Niemanstsverdriet, J.W. Carbon deposition as a deactivation mechanism of cobalt-based
Fischer–Tropsch synthesis catalysts under realistic conditions. Appl. Catal. A 2009, 354,
102–110.
380. van de Loosdrecht, J.; Balzhinimaev, B.; Dalmon, J.-A.; Niemanstsverdriet, J.W.;
Tsybulya, S.V.; Saib, A.M.; van Berge, P.J.; Visagie, J.L. Cobalt Fischer-Tropsch synthesis:
Deactivation by oxidation?. Catal. Today 2007, 123, 293–302.

128
381. Saib, A.M.; Borgna, A.; van de Loosdrecht, J.; van Berge, P.J.; Niemanstsverdriet, J.W. In
Situ surface oxidation study of a planar Co/SiO2/Si(100) model catalyst with nanosized cobalt
crystallites under model FischeríTropsch synthesis conditions. J. Phys. Chem. B 2006, 110,
8657–8664.
382. Moodley, D.J.; van de Loosdrecht, J.; Saib, A.M.; Niemanstsverdriet, J.W. The formation and
influence of carbon on cobalt-based Fischer-Tropsch synthesis catalysts: An integrated
review. In Advances in Fischer-Tropsch Synthesis, Catalysts, and Catalysis; Davis, B.H.,
Occelli, M.L., Eds.; CRC Press Taylor & Francis Group: Boca Raton, FL, USA, 2010;
pp. 49–81.
383. Bezemer, G.L.; Herman, J.H.; Kuipers, P.C.E.; Oosterbeek, H.; Holewign, J.E.; Xu, X.;
Kapteijn, F.; van Dillen, A.J.; de Jong, K.P. Cobalt particle size effects in the
FischeríTropsch reaction studied with carbon nanofiber supported catalysts. J. Am. Chem.
Soc. 2006, 128, 3956–3964.
384. Jacobs, G.; Ribeiro, M.C.; Ma, W.; Ji, Y.; Khalid, S.; Sumodjo, P.T.A.; Davis, B.H. Group 11
(Cu, Ag, Au) promotion of 15%Co/Al2O3 Fischer–Tropsch synthesis catalysts. Appl. Catal. A
2009, 361, 137–151.

129

Deactivation and Regeneration of Commercial Type
Fischer-Tropsch Co-Catalysts—A Mini-Review
Erling Rytter and Anders Holmen
Abstract: Deactivation of commercially relevant cobalt catalysts for Low Temperature
Fischer-Tropsch (LTFT) synthesis is discussed with a focus on the two main long-term
deactivation mechanisms proposed: Carbon deposits covering the catalytic surface and re-oxidation
of the cobalt metal. There is a great variety in commercial, demonstration or pilot LTFT operations
in terms of reactor systems employed, catalyst formulations and process conditions. Lack of
sufficient data makes it difficult to correlate the deactivation mechanism with the actual process
and catalyst design. It is well known that long term catalyst deactivation is sensitive to the
conditions the actual catalyst experiences in the reactor. Therefore, great care should be taken
during start-up, shutdown and upsets to monitor and control process variables such as reactant
concentrations, pressure and temperature which greatly affect deactivation mechanism and rate.
Nevertheless, evidence so far shows that carbon deposition is the main long-term deactivation
mechanism for most LTFT operations. It is intriguing that some reports indicate a low deactivation
rate for multi-channel micro-reactors. In situ rejuvenation and regeneration of Co catalysts are
economically necessary for extending their life to several years. The review covers information
from open sources, but with a particular focus on patent literature.
Reprinted from Catalysts. Cite as: Rytter, E.; Holmen, A. Applying the Behavioural Family Therapy
Model in Complex Family Situations. Catalysts 2015, 5, 478-499.
1. Introduction
In a gas-to-liquid (GTL) plant the high H2/CO ratio obtained from reforming of natural gas to
synthesis gas (syngas) obviates the need for shifting CO with steam to yield more hydrogen (and
CO2) for the FT unit. This is one main reason for using a cobalt catalyst instead of the much
cheaper iron alternative as catalytic metal for the FT reaction. In addition, the cobalt catalyst is
more active and has a simpler product slate of mainly paraffins and some Į–olefins. However, both
cobalt metal in itself, precious metal promoters as well as advanced overall formulations, make the
catalyst inherently costly. Further, Co catalysts typically lose about half their activity within a few
months. Assuming an economically acceptable catalyst lifetime of 2–3 years, this means that the
catalyst cost will add several USD to the price per bbl of produced synthetic crude. Therefore,
improving catalyst stability is a major focus among technology providers and plant operators. It
follows logically that a basic understanding of the mechanisms involved in the deactivation process is
vital to improving catalyst stability. Fortunately, it appears that, at least for most commercial Co
catalysts, rejuvenation of catalyst activity is possible.
A comprehensive review of deactivation of Co FT catalysts appeared a few years ago [1]. This
review discusses a wide variety of deactivation mechanisms comprising sintering; re-oxidation of
cobalt, including surface oxidation; formation of stable compounds between cobalt and the support,

130
e.g., cobalt aluminate; surface reconstruction; formation of carbon species on the cobalt surface;
carbiding; and poisoning. However, less focus is given in the review to long-term deactivation under
commercial conditions.
Historically, details of deactivation mechanisms and rates have been scarce particularly as only
a few plants are operated commercially. Nevertheless, some data can be found in the patent
literature, mainly based on operation of pilot or demonstration plants, and in conference
presentations. A complicating factor is that an industrial process is typically operated at constant
global production, i.e., deactivation is counteracted by a steady increase in operating temperature.
Only a few reports on deactivation under these most relevant conditions could be found.
Fortunately, Sasol and its collaborator, Eindhoven University of Technology, have published
extensive data on mechanisms and rates of deactivation at industrial conditions for their
Co/Pt/alumina catalyst. Several of their papers have focused on long term deactivation due to
polycarbon deposition [2].
We would also like to draw attention to an extensive report to US-DOE where long term
experiments are reported focusing mainly on the effect of water [3]. Catalysts with the formulations
Co(15 wt.%)/Re(0.2; 0.5; 1.0 wt.%)/Ȗ-Al2O3, Co(10 wt.%)/Ru(0.2 wt.%)/TiO2, Co(15 wt.%)/Pt(0;0.5
wt.%)/Ȗ-Al2O3, Co(15 wt.%)/Ru(0.2;0.5;1.0 wt.%)/ Ȗ-Al2O3 and Co(12.4 wt.%)/SiO2 were tested
for up to 3500 h in a 1L autoclave (CSTR-reactor). The authors claim that carbon deposition may
be minimized by careful temperature control, and that deactivation caused by sintering and
oxidation are the major concerns. These conclusions are controversial and have been disputed by
several investigations; see later in this review and in our previous review [1].
Argyle et al. have fitted previously published activity versus time data to first or second order
general power law rate expressions incorporating a limiting activity and have shown how parallel
routes, e.g., sintering and carbon deposition deactivation, can be modeled separately. For example,
their model predicts that during a 60 day run under typical FTS conditions a commercial Co
catalyst loses about 30% activity within 10–15 days due to rapid sintering, while an additional 30%
activity is lost gradually over the 60 day period due to carbon [4].
Causes of deactivation may depend on catalyst material and properties, e.g., support, promoters,
dispersion, extent of reduction, etc.; reactor type; and especially operating conditions. It appears
that after an initial break-in period during which cobalt is equilibrated with its reactor environment
in terms of crystallite size, possibly crystal structure, and degree of reduction, a slow long term
deactivation is observed. The origin of this latter deactivation period is discussed in terms of carbon
formation and/or re-oxidation of the metal.
2. Catalyst Activity
To understand catalyst deactivation, it is first necessary to understand the factors that determine
initial catalyst activity. Activity is largely dependent on the degree of reduction of the cobalt metal
precursor and the average size of the cobalt crystallites, which together determine the surface
density of catalytically active sites, i.e., the dispersion of the metal. It has been verified that the
turn-over-frequency (TOF) is rather constant for Co crystallites larger than 6–8 nm [5]. As activity
falls off rapidly below this threshold, methods for making very high dispersion catalysts have

131
limited relevance for FT-synthesis. Cobalt crystallite size and degree of reduction depend on
several factors, including cobalt precursor, support material and its pretreatment; pore diameter,
pore volume and available total surface area; method of impregnation or deposition; drying and
calcination conditions; reduction conditions; etc. It is especially important to calcine and reduce
the catalyst at optimum conditions, i.e., optimal gas flow rates, temperature ramp and final
temperature [6]. For example, overly high calcination and reduction temperatures result in large
cobalt oxide and cobalt metal crystallites and, therefore, undesirably low dispersion due to
over-sintering. For a given degree of reduction and crystallite size, the activity per kg catalyst is
proportional to the cobalt loading. The loading of a commercial type catalyst may vary from 12 to
30 wt.% and is a compromise between several catalyst properties. For instance, a lower surface
area and pore volume support will be able to accommodate less cobalt, but might be considerably
more attrition resistant.
To optimize cobalt crystallite size is not particularly challenging as long as one is able to control
preparation conditions. It is industrial practice to add a metal promoter that enhances the degree of
reduction and maintains a targeted dispersion [7,8]. The literature provides no solid evidence that
such metal promoters are able to enhance reaction rates or surface concentration of intermediates.
Promoters used at a commercial or demo scale include platinum, rhenium and ruthenium. The
promoter will add significantly to the cost of the catalyst; catalyst grade Re is today priced at ca.
3000 USD/kg and Pt at 45,000 USD/kg [9]. Typical loadings are up to 0.5 and 0.05 wt.% for Re
and Pt, respectively. It can also be mentioned that a possible effect of cobalt being in the fcc or hcp
crystallographic phase, with the latter being more active, has been reported [10]. However, studies
on the actual configuration of an active cobalt crystal are needed to be able to correlate activity
with atomic arrangement on the surface of a working catalyst. It is well documented that the
catalyst support has a strong influence on the selectivity to C5+ of the process, but as long as known
impurities like alkali, alkaline earth metals and sulfur are eliminated it is less clear how
support chemistry and pore structure influence activity of a catalyst for a given cobalt crystallite
size. However, different supports have varying interactions with cobalt oxide and therefore
influence reducibility.
It is well known that reaction rate greatly depends on process conditions. Generally, rate
increases with increasing temperature and overall pressure. Furthermore, indications are that the
rate increases with H2/CO ratio, possibly due to higher methane make, and decreases with
increasing conversion as the partial pressure of syngas is reduced and a high level of product water
may block active sites [11].
Of the CHx monomers generated on the surface of a catalyst, CH2 is probably the most abundant
intermediate and is probably readily incorporated in the chain during polymerization. A smaller
portion of the monomer will be hydrogenated all the way to methane. The growing chain can
terminate by ȕ–hydrogen abstraction and leave the surface as an olefin or be hydrogenated to an
alkane. Olefins can also be hydrogenated in a secondary reaction. There is evidence from experiments
at low conversion and small catalyst particle size that the primary product is dominated by olefins,
but for practical purposes the olefin to paraffin ratio is well above two for C3 and then diminishes
rapidly with chain length [12].

132
3. Fischer-Tropsch Reactors
Apart from the type of FT-catalyst, selection of the FT-reactor, as well as how it is operated and
incorporation in the XTL flowsheet, is the principal factor influencing catalyst deactivation.
Comparison of properties of the main reactor types for low-temperature Fischer-Tropsch synthesis
is given in Table 1. Conversion per path will vary, mainly with the propensity for heat removal and
temperature control, whereas the temperature typically is between 200 and 250 °C and the pressure
will be in the range 15–30 bar. H2/CO ratio in the make-up gas from the syngas generator will be
slightly below 2, whereas the outlet ratio of the FT-reactor will be considerable lower; down to 1.2,
in some cases even below 1.0. Evidently, high temperature and low H2/CO ratio are expected to
promote deactivation, but reports on these effects are not available.
Table 1. Properties for different reactor types for Low-Temperature Fischer-Tropsch
synthesis *.
Reactor

Conversion
per path
(%)

Capacity
per reactor
(bbl/day)

Tubular fixed-bed

30–35

”6000

Slurry bubble
column

55–65

”25,000

Micro-channel

65–75

”1000

Characteristics
” 30,000 tubes with catalysts
pellets or extrudates.
Internal heat exchanger and
optional product filter
Metal block with
< 2mm diameter channels

* Based on open literature and patents for commercial Low Temperature Fischer-Tropsch (LTFT)
synthesis with cobalt catalysts [13,14]. See also references in Table 2.

In a comparative study of reactor types for LTFT synthesis Guettel and Turek conclude that the
productivity per reactor volume of a slurry bubble-column reactor or monolith reactor is up to one
order of magnitude higher than for a fixed-bed or micro-channel reactor [15]. However, this
conclusion is at variance with other work, reported for micro-channel reactors for which superior
productivities have been reported [16].
Fixed-bed reactor. Due to the necessity of controlling the heat evolved during reaction, the
design of a fixed-bed FT reactor is based on a multi-tube heat-exchange type of reactor where
catalyst pellets are loaded into the tube bundles and the shell contains evaporating water. In this
type of reactor axial temperature typically increases through a maximum of 5–20°°C, and it is
imperative to avoid hot-spots which cause sintering. In order to minimize deactivation due to
temperature effects several measures are taken. Once-through CO conversion is limited to 30–35%.
Tube diameter is typically 2.5–5 cm and the size of the catalyst pellets or extrudates is relatively
small, in the range of 1–3 mm. Small extrudates size is also important to secure good radial mixing
and minimize diffusion limitations, thus maintaining high selectivity to liquids. It has been shown
that above ca. 200 μm particle size the higher effective H2/CO ratio in the inner part of the pellets
significantly reduces C5+ yield [12]. Therefore, an egg-shell catalyst design is preferred where only
the outer parts are impregnated with active metal. Another factor limiting the applicable superficial

133
gas velocity and the global rate is pressure drop. In general, a challenge related to deactivation in a
fixed-bed reactor is variations in the partial pressures of reactants and products along the tube
length and within the catalyst particles.
H2/CO ratio depends on several factors, but the make-up gas typically has a ratio slightly below
2 for maximizing C5+ selectivity. With recycle of the product gas to the FT-section the feed ratio to
the reactor may be significantly lower, possibly 1.6–1.7, and there may be a gradual reduction from
the inlet to around 1.4–1.5 at the outlet. Compared to a slurry reactor, the average gas composition
may be richer in hydrogen and due to less efficient temperature control, the overall operating
temperature usually lower. A consequence is lower reaction rates, but this is at least in part
compensated for by a lower average partial pressure of product water and thereby a higher syngas
pressure. The overall effect on deactivation, in particular carbon deposition, is complex and
challenging to predict.
A distinct advantage of the tubular fixed-bed reactor is a well proven commercial design.
Several tens of thousands of tubes can be incorporated within the reactor shell. Scale-up is
comparatively easy, and optimization can be done in a single tube laboratory reactor. Operational
experience with catalyst fouling or attrition and resulting difficulties with loading and unloading tubes
are trade secrets, but it can be expected that an experienced operator is able to control these factors.
Minimizing catalyst deactivation or being able to perform in situ regeneration is critical in order to
reduce catalyst consumption and avoid an extensive unloading-reloading sequence. The liquid product
is inherently separated from the catalyst and any need for removing residual particles and metal
components will be low.
Slurry bubble-column reactor (SBCR). Catalyst particles are suspended in the liquid
hydrocarbon product of the FT process and synthesis gas is bubbled through the slurry. Depending
on the density of the catalysts particles, their diameter and the superficial gas velocity, there is a
profile of solid concentration diminishing from bottom to top of the reactor. Gaseous components
leave from the top of the reactor. Higher boiling products have to be removed from the reactor as
liquid, and separated from the catalyst. Several methods for this purpose have been patented, both
in situ and ex situ techniques. Broadly they can be classified as employment of filters, settling
devices, magnetic separation and hydrocyclones. Sasol uses internal filters combined with
secondary polishing filters of the product [14].
Settling of the catalyst should be avoided as overheating and consequently catalyst deactivation
will occur. Particularly critical are the gas distribution system and, depending on design measures
to prevent particle settling in stagnant regions below the nozzles. One way to improve overall
liquid circulation, and thereby avoid settling, is to install so-called internal down-comers. A serious
threat to the catalyst in a slurry operation is any upset in production, like a sudden reduction in
syngas flow. Without adequate back-up systems such events will lead to settling and serious
overheating in the catalyst mass due to continuous FT-synthesis with residual syngas. Similar
conditions may occur in slurry separation devices like filters, but no public information is available
on any effects on catalyst deactivation in these devices.
An SBCR operates preferentially in the churn turbulent flow regime for best distribution of
catalyst particles as well as minimizing mass and heat transport restrictions. In the churn turbulent

134
flow regime there is a mixture of smaller and larger bubbles that undergo frequent beak-up and
coalescence. This mechanism prevents serious film transport restrictions on the catalyst slurry
interphase, and with catalyst particles below 200 ȝm the H2/CO ratio as well as water vapor
pressure can be assumed constant over the entire reactor volume. Thus, deactivation should be
more easily controlled compared to other reactor configurations. Further details on operation of
SBCRs can be found in the book on Fischer-Tropsch technology by Steinberg and Dry [17].
Reasons for selection of a slurry bubble column reactor for Fischer-Tropsch synthesis include:
(1) a comparatively simple construction; (2) high space-time-yield and catalyst efficiencies;
(3) high heat transfer coefficients; and (4) isothermal conditions. Continuous catalyst regeneration
of a slip stream is a viable option. Challenges include minimizing catalyst particle attrition and
efficiently separating catalyst from the products. Efficient liquid and gas back-mixing and a high
exit water concentration lead to high selectivity; the high exit water concentration is beneficial in
reducing coke deposition. On the other hand reactant concentrations are lower than the average of a
fixed-bed reactor resulting in comparatively lower global rates. Single pass conversion is typically
in the range of 55–65%, significantly higher than for fixed-bed. Conversion is limited by the
feasible height of the reactor, but there is also an upper conversion limit above about 75–80% for
which the water-gas-shift activity will lead to possible catalyst oxidation and a steep increase in
CO2 yield [18]. Naturally, extensive recycle of syngas in the FT-section of the plant is necessary to
obtain a very high overall CO conversion.
Micro-channel reactor. Micro-channel reactor technology holds great promise for process
intensification due to outstanding heat and mass transfer rates [19,20]. Combined with highly
active and stable catalysts, micro-channel reactors can achieve high volume based productivities. In
some cases very high conversions ( ࡱ 90%) can be obtained while maintaining high C5+ selectivity.
Detailed studies of flow and temperature behavior have shown that a micro-channel reactor can
operate isothermally and with very low pressure drop [19]. Except for the normal initial
deactivation, the catalyst in the micro-channel reactor is remarkably stable even at very high
conversions [21]. Observed rates of deactivation appear to be lower in the micro-channel reactor
compared with the fixed bed laboratory reactor at similar conditions. Velocys and CompactGTL
are operating microchannel demonstration plants [20].
3. Commercial Catalyst Formulations
Scale-up to commercial catalyst production is demanding, and little public information on the
industrial manufacturing processes is available. Great care must be taken to obtain a homogeneous
catalyst material, but the targeted distribution of cobalt on the support depends on the actual
process. For slurry catalysts with diameters typically in the range 40–120 μm pore diffusion
resistance of the syngas is negligible ensuring full utilization of the available surface area [12]. For
micro-channel reactors the catalyst either will resemble a slurry catalyst or be impregnated onto
special trays that are inserted into the channels. On the other hand, a fixed-bed catalyst is typically
designed as an egg-shell or rim catalyst in which only the outer few hundred micrometers contain
the active phase. As to the degree of reduction it has been shown by Sasol that the initial value is

135
not critical, as the syngas will reduce the catalyst further under the first months of operation,
thereby increasing CO conversion [2].
Cobalt FT catalysts can be classified according to supports and promoters used. Table 2 lists
several commercial type catalysts. Precious metal promoters like Pt and Ru may facilitate
hydrogenation activity and hence reduce propensity for carbon deposition as indicated, for
example, by Exxon Mobil for ruthenium.
Some relative activities have been estimated and are included in the table. The values are based
on fixed-bed data reported in the patent literature. As the process conditions vary considerably, an
attempt to normalize the activities is made by using a simplified kinetic expression with activation
energy of 110 kJ/mol, and partial pressures are average pressures in the reactor [22]. It is recognized
that this comparison is only approximate, but still a guide. Activities are generally lower for a given
catalyst operated in slurry compared to fixed-bed in spite of limited apparent diffusion limitations.
The origin of this effect is not understood. Particularly low activities have been reported by Nippon
and ENI/IFP, perhaps an indication of large cobalt crystals in the working catalyst.
It appears that Shell favors a titania based support over their previous zirconia modified silica
system. Promoters are either Mn or V, the latter claimed to lower CO2 make [23]. Titania has
relatively large pores and moderately low surface areas, but is known to facilitate high selectivities
to C5+ products. Fixed-bed catalysts contain modifiers like citric acid added prior to the forming
step. In addition to Shell, it should be mentioned that BP is promoting its fixed-bed technology and
claims that a CO2 resistant support is vital [24].
Velocys/Oxford Catalyst and Compact GTL are offering micro-channel fixed-bed technologies
with catalyst diameter or thickness of 0.1–0.5 mm. Velocys’ carbon combustion preparation
technique may very well take the edge off initial sintering during FT synthesis by optimizing cobalt
crystallite size already at the catalyst manufacturing stage. Loading and off-loading the catalyst can
be particular challenge for these systems, but efficient methods for this purpose are claimed by
these companies.
The platinum promoted Sasol catalyst is prepared on Ȗ-alumina (Puralox SCCa-2/150 or -5/150:
pore volume 0.5 mL/g; surface area 150 m2/g) and stabilized by impregnation with tetra-ethoxysilane (TEOS) followed by calcination to give a surface concentration of ca. 2.5 Si atoms/nm2. This
procedure apparently modifies the surface so that the support becomes less prone to dissolution in
the acidic product water. The GTL.F1/Statoil catalyst is based on a larger pore diameter Ȗ-alumina
modified with nickel and fired at high temperature to produce a nickel-aluminate (spinel)/Į-alumina
mixture. The pore properties resemble titania-based supports, but with very high attrition resistance.
Also the ENI/IFP catalyst is supported on Si-modified Ȗ-alumina, but probably strengthened by
silanation and calcination giving a final SiO2 content of 6–7 wt.%. Other support modification
methods have been described by ENI/IFP in earlier patents, including formation of spinel
compounds. It is unclear whether the catalyst formulation contains a reduction or other type of
promoter. In their slurry catalyst development, Nippon has apparently adopted a silica based
support formulation similar to that of Shell’s previous fixed-bed catalyst, but using ruthenium as a
promoter. No information on attrition resistance has been revealed, as is the case for most other
slurry catalysts as well. Syntroleum used a catalyst similar to Sasol’s, but also with a ruthenium

136
promoter. Exxon Mobil, that pioneered titania as a support with an alumina binder, BP, ConocoPhillips and Syntroleum have terminated their developments in FT-technology. However, BP is still
licensing their FT technology and Exxon Mobil has announced that the technology is ready for
commercialization should the right project be prioritized.
Table 2. Formulation of commercial type cobalt catalysts and their application *.
Technology
provider
Sasol
Shell
GTL.F1
ENI/IFP/
Axens
Nippon Oil
Syntroleum
BP
Exxon
Mobil
ConocoPhillips
Compact
GTL
Oxford cat.
/Velocys

Promoter

Reactor
type

Reactor
scale
(bbl/d)

Pt

Slurry

16,000

Mn; V

Fixed

6000

0.3

[26]

Re

Slurry

1000

0.9

[27]

?

Slurry

20

0.19

[28]

Ru

Slurry

500

0.16

[29]

Ru

Slurry

80

?

Fixed

Re

Slurry

200

Ru/Pt/Re

Slurry

400

Alumina?

Re?

Micro

500

[33]

Titaniasilica

Re

Micro

1000

[34]

Support/
modifier
Ȗ-Alumina/
Si **
Titania
Nialuminate/
Į-Alumina
Ȗ-alumina/
SiO2;spinel
Silica/
Zirconia
Ȗ-Alumina/
Si **; La
ZnO
Titania/
Ȗ-Alumina
Ȗ-Alumina/
Boron

Relative
activity

Reference

[25]

[30]
[24]
[31]
0.68

[32]

* Deduced from open literature and patents. Actual commercial formulation may vary. ** Si from TEOS;
tetraethoxy silane.

4. Causes of Deactivation
From our previous review on deactivation during LTFT synthesis the main causes of
deactivation are sintering, re-oxidation of cobalt, formation of stable compounds between cobalt
and the support, surface reconstruction, formation of carbon species on the cobalt surface,
carbiding and poisoning [1]. In addition there can be a loss of catalyst material from the reaction
zone due to attrition. The chemical environment is challenging with a number of reactive chemical
species generated including significant amounts of water. In addition, the exothermicity of the
reaction may lead to hot spots in the catalyst.
There appear to be two main “schools” for describing long-term deactivation mechanisms based
on demo slurry operations, one favoring re-oxidation [35,36], and one poly-carbon formation on the
surface [2]. It should be realized that both catalyst system and process conditions can affect the results.

137
In addition, an initial sintering stage may be observed if the fresh catalyst contains crystallites in the
range of 6–12 nm. The consequences of severe deactivation can be a significant decline in the
activity over a typical design period for catalyst life of two years to an estimated 25–30% of the
initial value. In addition, slurry catalyst loss due to attrition can be significant. All previous
experience considered catalyst replacement due to deactivation will contribute significantly to the
operational costs of an FT plant.
In a study on the effect of impurities it was found [37], by impregnating 400 ppm of the impurity
element from nitrate precursors, a poisoning effect which decreases in the following order:
Na > Ca > K > Mg > P,
Mn, Fe and Cl showing minimal effects. The latter is surprising as chlorine causes a 25%
reduction in hydrogen chemisorption. Even alkali concentrations of less than 100 ppm have a large
effect on the rate (site time yield) [38]. However, no effect of alkali on the hydrogen chemisorption
was observed. The impact of alkali and alkaline earth elements is far stronger than any stoichiometric
blocking of surface sites, and might be related to the strong electronegativity of the elements leading
to blocking of steps on cobalt thus preventing CO dissociation [39]. Particular care must be taken to
avoid alkali and alkaline earth elements in impregnation fluids and washing water as well as
contaminants in the catalyst support. Sulfur as H2S or (CH3)2S added to the syngas gives a
deactivation consistent with stoichiometric blocking of cobalt surface sites as in situ measurement of
cobalt dispersion by H2S is consistent with hydrogen derived dispersion of a fresh catalyst. The effect
of ammonia appears to be strongly catalyst dependent, and reports vary from negligible influence to
strong negative consequences.
Contributions to attrition of catalysts for three-phase slurry bubble column operation include
mechanical abrasion and breakage of catalyst particles; chemical dissolution; and synergisms
between these mechanisms. It appears that Sasol focuses more on avoiding chemical attack on their
Ȗ-alumina based support [40], whereas Statoil/GTL.F1 [41], IFP/ENI/Axens [28], and probably
Exxon [42], have developed more mechanically robust slurry catalysts.
Sasol has reported deactivation profiles in several publications for their Co/Pt/modified Ȗ-alumina
catalyst [2]. We refer to the section below on carbon deposition and to the previous review for
further details on the Sasol work and their extensive documentation of carbon deposits [1].
Statoil/GTL.F1 have disclosed deactivation profiles at several conferences for their attrition
resistant catalyst of Co/Re/aluminate spinel catalyst. In a CSTR slurry reactor test over 3000 h the
temperature was adjusted regularly, typically in 2–3 weeks intervals, to keep conversion reasonable
constant [43]. Somewhat surprisingly, temperature was decreased from 222 °C to 215 °C during
the operation meaning that the catalyst activity increased regularly. This is in line with the reported
enhanced reduction during first months of slurry FT-synthesis. More surprisingly, the C5+ selectivity
increased simultaneously by ca. 5%, significantly more than expected given the reduction in
temperature. In another presentation on CSTR results, the rate of hydrocarbon formation increased
up to 800 h time-on-stream (TOS) followed by decline towards end of test at 1600 h [44].
Characterization of a commercial catalyst after ca. one month operation in a commercial scale

138
slurry-bubble column confirms an increase in degree of reduction [45]. This is concurrent with
sintering probably facilitated by a high steam partial pressure [46].
5. Deactivation by Carbon Deposition
Higher hydrocarbons (waxes) are desired products from LTFT synthesis on Co catalysts. The
hydrocarbons can accumulate on the surface and can slowly be converted to carbon or coke that
blocks the active sites. By using TPR Lee et al. [47] could distinguish between several forms of
carbon on the catalyst surface from CO disproportionation; they suggested that carbon was
present in two forms: atomic and polymeric carbon. Support for stabilization of graphitic carbon on
an fcc-Co(111) surface has been obtained by quantum-mechanical calculations [48]. DFT data
indicate chemical bonding between graphene and cobalt, as also supported by other studies [49].
Direct STM evidence for formation of graphene on Co(0001) has been demonstrated through
decomposition of ethylene [50]. It was found that carbon on the surface induces cobalt reconstruction
and weakens CO and H2 adsorptions.
Thus, there are ample investigations showing that carbon in different forms can interact with
and block cobalt surfaces. Support for this deactivation mechanism comes from a few long term
studies using commercial catalysts in pilot reactors. Sasol studied catalyst deactivation by
periodically removing samples from a pilot slurry bubble column reactor operated for 6 months [2].
Wax was removed by inert solvent extraction before the catalyst samples were characterized by
temperature programmed hydrogenation and oxidation, chemisorption, TEM and LEIS. Polymeric
carbon was found both on the alumina support and on cobalt. This carbon is resistant to hydrogen
treatment at temperatures above the FT synthesis temperature. The amount of polymeric carbon
correlated well with observed long term deactivation. From XANES data they ruled out oxidation of
cobalt during the run, but there was significant sintering taking place during the first 10–15 days on
stream. Moodley et al. [2] concluded that accumulation of polymeric carbon was responsible for at
least a part of long term catalyst deactivation.
Build-up of graphitic or polymeric carbon as deactivation mechanism was recognized by
Syntroleum [51]. By TGA-MS they estimated that about 1% carbon was deposited on the catalyst.
Carbon deposition on cobalt/ZnO has also been proposed by BP based on results from a
demonstration plant and laboratory studies [52].
6. Deactivation by Re-Oxidation
Schanke et al. investigated the influence of water on deactivation of unpromoted or Re
promoted alumina supported cobalt catalysts. Adding 20–28% steam to 50% syngas in the feed of a
lab-scale fixed-bed reactor resulted in significant deactivation due to oxidation of highly dispersed
cobalt crystals and surface cobalt atoms [53]. Although these experiments clearly show oxidation
of cobalt, the conditions represent very high conversion levels (> 80%). Oxidation takes place within
the stability range for bulk cobalt metal, and is presumed to be a consequence of surface reactivity
of small crystallites. It is also evident that the effect of water depends critically on the support
material used, e.g., samples of Ȗ–alumina from different sources behave very differently. Even a

139
positive effect of water on activity has been reported; see [1,3] for further discussion. From activity
tests over a range of conversions in a slurry CSTR reactor Co/Re/Ȗ-alumina catalyst activity is
observed to slightly increase with conversion up to 85%; above 85% conversion it drops rapidly [18].
At high conversions high partial pressures of H2O may oxidize small cobalt crystallites and promote
aluminate formation thus enhancing WGS activity as shown by a significant CO2 make. Similar
results were found for a Pt promoted catalyst [54].
Very recently, the group of A.Y. Kodakov in cooperation with Total published evidence for
surface oxidation of cobalt nano-crystallites in alumina supported, Pt promoted catalysts during FT
synthesis [55]. However, surface oxidation was only clearly evident by STEM-EELS after an
excursion to 340 °C and 100% CO conversion. B.H. Davis and coworkers exposed a freshly
reduced catalyst directly to a water vapor pressure equivalent to 50% conversion [56]. They
observed oxidation of a fraction of the smaller cobalt crystallites when supported on alumina or
activated carbon, and recommended that careful crystallite size management is required for a
commercial catalyst. Kliewer et al. studied redox transformations of cobalt catalysts by TEM in
terms of agglomeration of the metal, mixed-oxide formation with the support and reversible
oxidation followed by reduction under mild hydrogen treatment [57]. They claim that reactor and
TEM studies show that nanoscale Co crystallites can oxidize to CoO during commercially relevant
FT synthesis conditions in spite of bulk thermodynamic data that suggest otherwise [58]. The
propensity for oxidation is enhanced by small Co crystallites and high CO conversion with
attendant high water vapor pressure and a high H2O to CO ratio in the reactor. The oxide
crystallites thus formed can be fully reduced by hydrogen at standard FT temperature and pressure.
Reported TEM images from Exxon show that the oxidized cobalt metal wets the support surface
and thereby facilitates contact between nearby crystallites. This has also been illustrated in a
presentation from Statoil in Figure 1 [59]. Both images show a thickness of a cobalt oxide
crystallite of ca. 5 nm wetting the surface. Further, the TEM image indicates mixed orientation of
several cobalt oxide crystallites and an amorphous layer at the support interphase. Upon
re-reduction metal crystallites may agglomerate depending on the initial spatial distribution on the
catalyst support.
That cobalt distribution can vary significantly with catalyst preparation procedure is illustrated
in Figure 2 for a Co/Re/Ȗ-alumina catalyst [43]. Cobalt dispersions are comparable, but we see that
the distribution of clusters of the oxide varies significantly. Although it can be imagined that small
well dispersed clusters are less prone to deactivation, this needs to be verified.
It is well known that mixed-oxides can be formed between cobalt and silica, alumina and titania.
In part, a surface layer of mixed-oxide is formed during catalyst preparation from water solution
followed by drying and calcination. It has also been claimed that enhanced mixed oxide formation
takes place during FT-reaction at high conversion levels (> 70%) concurrent with oxidation of
cobalt to Co2+. In the case of silica supported catalysts, well defined crystalline needles of cobalt
silicate are rapidly formed at higher conversions [60].

140

Figure 1. Wetting of cobalt oxide on a support. Left: EELS spectra; alumina: red;
cobalt: green. Right: TEM.

Figure 2. Cobalt oxide clusters on a Ȗ-alumina. Bar shows 1 μm for all images.
Redox reactions of cobalt have been used deliberately to enhance the catalytic properties of FT
catalysts by employing a reduction-oxidation-reduction (ROR) procedure [61,62]. An ROR
treatment can increase Co dispersion presumably by forming hollow oxide domes during controlled
oxidation that break up into smaller Co crystallites during re-reduction [63]. Whether the resulting
Co crystals that will be in close proximity with each other experience agglomeration during the
first months of commercial FT synthesis has apparently not been reported.
7. Catalyst Rejuvenation and Regeneration
Regeneration of cobalt LTFT synthesis catalysts is largely described in the patent literature. The
options involve treatment of the catalyst with air (oxygen), hydrogen and/or CO and variations
thereof in addition to procedures for removing produced wax. Therefore, regeneration addresses
reversing the main deactivation processes of carbon deposition, metal oxidation and sintering by
combustion, reduction and re-dispersion, respectively. A review of early reports on regeneration
covering 1930–1952 has been presented [64]. It was concluded that there is no universal process for
regenerating Co FT-catalysts; the art over this relatively short time period comprising conflicting
results with patents covering a wide range of processes involving oxidation, reduction, combined
oxidation-reduction, steam-reduction, operating at elevated temperatures and solvent extraction.

141
Commercial regeneration processes are either in situ in the FT-reactor itself or ex situ after
removal of part of or the entire catalyst inventory. Indications are that Shell successfully
regenerates their catalyst regularly, but it is undisclosed whether this takes place inside the tubes of
the fixed-bed reactor or in a separate unit, whereas Sasol apparently removes part of the slurry from
the reactor continuously for regeneration and re-deployment of the catalyst into the reactor. The
latter approach allows continuous operation of the LTFT synthesis. Micro-channel reactors pose
special challenges depending on the catalyst configuration. In situ regeneration is an option, or the
catalyst can be removed for external treatment either by unloading the catalyst particles or
removing multi-channel trays with catalyst attached.
A summary of regeneration procedures from some of the main industrial companies that are or
recently have been involved in Fischer-Tropsch technology development is given in Table 3. The
table is representative of available information but should not be assumed to be complete or
up-to-date; nevertheless it illustrates what is probably the preferred approach of each individual
company. In the first column the type of commercial regeneration process and primary FT-reactor
type are listed and whether the regeneration is intermittent or continuous, while data in the other
columns refer to the actual test protocol and results described in the patent literature. Note that
reactor type can be different in columns one and three.
It is clearly possible to regenerate a deactivated catalyst to a level close to the original activity
by different combinations of wax removal, hydrogenation and combustion of carbonaceous
deposits; followed by re-reduction if needed. There is unfortunately little information available on
the long-term performance of regenerated catalysts.
Sasol has published a brief summary of their procedure for removing most of the wax, followed
by combustion of the remaining carbonaceous species and complete reactivation [65]. Specifically,
Sasol describes dewaxing by hydrogenolysis with pure hydrogen for 2 h at 220 °C and reduction
for 2 h at 350 °C [66]. After passivation with CO2, the catalyst is subjected to oxidation with air in
a fluidized-bed calcination unit at 250 °C for 6 h under a pressure preferably of ca. 10 bar.
Re-reduction is performed at 425 °C and 98% of the original activity is regained. In cooperation
with Eindhoven University, Sasol has investigated mechanisms of deactivation and regeneration for
both model and commercial catalysts [67]. After a heptane wash and reoxidation, several FT cycles
were demonstrated with no apparent permanent loss in activity. The oxidation step is described in
terms of the Kirkendall effect involving spreading of a Co oxide film during oxidation followed by
breaking up of the film to form small re-dispersed Co crystallites.
It appears that Shell is aiming at in situ regeneration in the tubes of their fixed-bed FT reactor,
although an external post FT-reaction step is part of their described procedure [69]. Regeneration is
based on a procedure that most likely comes from their Bintulu plant in Malaysia. After wax
removal a mild hydrogenation and oxidation is conducted. The catalyst is then taken out of the FTreactor, treated with a concentrated aqueous ammonia solution and subsequently with CO2, giving
Co ammonium carbonate. Most likely the latter procedure gives cobalt ammine carbonate
complexes suitable for re-dispersing cobalt [76].

142
Table 3. Summary of regeneration concepts and procedures based on patents and
presentations from LTFT technology companies.
Technology owner
Regeneration
configuration*

Catalyst

FT test

Wax removal

Primary

TOS

reactor

step

hydrogenation

Sasol [66,68]

Co/Pt/

Slurry; continuous

alumina

Ex situ

-

Shell [69]

Co/Mn/

Fixed-bed;

Titania

intermittent in situ

- Years

GTL.F1 [70]

Co/Re/

Slurry; continuous

aluminate

Ex situ

-

ExxonMobil [71]
Slurry; continuous
Ex situ

Co/Re/titania
- A few days

H2 strip at
Slurry**

Fixed-bed
Full scale

Slurry**

cyclohexane

No

wash
Fixed-bed
(lab)

Co/
zirconia-silica

Ex situ

-

ConocoPhillips [73]

Co/Ru or Re/

Fixed-bed

Slurry; continuous

(F-)alumina

(lab) or

Ex situ

- 1014 h

slurry.**

Co/alumina

Slurry**

Ex situ

Diluted H2

Filtration + H2
strip

No

Regeneration

Oxidation/

effect

Re-dispersion

(activity)

Air in fluidbed
Diluted O2 at

Gas oil wash
Draining or

Nippon Oil [72]

Syntroleum [74]

H2 at 350 °C

xylene wash

Slurry; continuous

Slurry; continuous

220 °C or

Calcination/

98%
> 100%

270 °C;

< 100%

NH3/CO2

(800 h TOS)

Diluted air in
fluid-bed
Diluted O2 in
fixed-bed

98–115%

ࡱ 100%

Steaming at
Slurry (lab)

“de-oiling”

No

200 °C; 25 bar

95%

in fixed-bed
7% H2/steam at
No

No

80–95%

3 bar
N2 at
316–343 °C

- 4000 h

300 °C/

No

Diluted O2 in

“Good

fluid-bed

performance”

O2 in situ

ࡱ 100%

Oxford Catalyst/
Velocys [75]

Co/Re/ titania-

Micro-

Micro-channel; in

silica?

channel

N2 flow?

H2 in situ

situ intermittent
* Estimated. ** Size of slurry reactor is undisclosed, but the technology providers have had both lab and pilot/demo units in operation.

Regeneration may re-disperse cobalt, possibly even to a level higher than for a freshly reduced
catalyst [70]. A promoted cobalt catalyst on a spinel support was subjected to an extended run in a
slurry-bubble column reactor as described by Schanke et al. [14]. Actual time on stream was not
reported, but samples from four different TOS’s were analyzed and regenerated. Data from the
patent have been plotted in Figure 3. All activity data are from standard runs in a laboratory
fixed-bed reactor at 210 °C, 20 bar pressure and H2/CO ratio of 2:1, and after a conventional
reduction protocol with hydrogen at 350 °C. Fresh and used catalysts, all embedded in wax, were
drained to remove excessive wax at 85 °C before activity testing. A successive reduction in activity
to ca. 50% was experienced. If the spent catalyst was calcined at 300 °C to burn off excessive
wax and carbon deposits before testing, the activity is only slightly reduced compared to the fresh
catalyst. Interestingly, even an enhanced activity was experienced if the wax was removed by
cyclohexane/n-heptane solvent extraction before calcination. Pore volumes and surface areas of freshly

143
calcined catalysts and used catalysts subject to wax draining and oxidation are unchanged. From these
data it appears that hydrogenation by itself has moderate or no effect on regeneration, but that burning
off coke deposits completely rejuvenates the catalyst.

Relative catalyst activity

1,2
1
0,8
0,6
0,4
0,2

Wax draining; No calcination
Wax draining + calcination
Solvent extraction + calcination

0
Fresh

Used-1

Used-2

Used-3

Used-4

Successive time-on-stream
Figure 3. Catalyst activity before and after oxidation of carbon deposits of a catalyst
used in a slurry bubble column reactor (data from ref. [70]).
In an early patent, Exxon researchers describe successful rejuvenation by atmospheric hydrogen
treatment at typical FT synthesis temperatures of 200–230 °C of spent Co/Ru/titania catalysts [77].
The promoter is needed to facilitate the rejuvenation and is claimed also to inhibit carbon deposits [78].
However, examples are only for FT runs of 10–30 days at 50–60% CO conversion in lab-scale,
fixed-bed reactors, and therefore may only addresses hydrogenolysis of very heavy waxes that
partially block pores and possibly hydrogenation of oligomeric carbons on the surface of the catalyst.
In a later patent, as shown in the table above, Exxon Mobil has demonstrated that reduction with
hydrogen after oxidative regeneration can be performed in the slurry FT-reactor itself at mild process
conditions of 200 °C and 20 bar, thus resembling the FT synthesis conditions. The company has a
number of patents describing regeneration procedures, including adding more active metal after
combustion of carbon deposits [79].
As pointed out above, - deposits of heavy hydrocarbon waxes reside in the pores of a used
catalyst that should be largely removed before regeneration. Hydrogen treatment may reduce part
of the wax through hydrogenation, but may also leave residual components at the surface.
ConocoPhillips claims to have designed a suitable reactivation procedure that both removes heavy
hydrocarbons and reduces the active metal [73]. For a 19 wt.% Co/0.1 wt.% Ru on alumina catalyst
run for 40 days at standard FT-conditions, best regeneration results were obtained in a fixed-bed

144
using a mixture of 93% steam/7% H2 at 300 °C and 3 bar where 95% of initial activity was
regained. Similar experiments for a catalyst composition of 19 wt.% Co/0.1 wt.% Re on fluorinated
alumina in a slurry reactor were less successful as 71% of initial activity was obtained, up from
33% after the FT period.
In their continuous regeneration of a spent slurry catalyst, Syntroleum focuses on removing as
much wax as possible by nitrogen stripping at 300–350 °C and 3 bar, followed by calcination.
Improved cobalt dispersion and reducibility are claimed, but no actual performance data are
reported for the regenerated catalyst [74]. Nippon Oil has followed an alternative approach where
they hydrothermally treat the spent catalyst with steam at elevated pressure [72]. Much of the
activity is regained, although it appears that the TOS between regenerations is relatively short in
view of the high activity level of ca. 75% before regeneration, thus moderating build-up of
polymeric carbonaceous deposits.
Oxford Catalyst/Velocys have presented interesting long-term operation and regeneration
performances for their micro-channel reactor [75]. Velocys’ catalyst and reactor operate at much
higher space velocities and productivities than conventional catalyst/large-scale FT reactors. To
compensate for deactivation, temperature is increased from 205 °C to 232 °C after ca. 650 days
TOS, accompanied by a slight increase in CO conversion from 71.2 to 73.6%. Naturally, the
temperature increase is accompanied by a reduction in C5+ selectivity, in this case by as much as
from 87.9% to 75.8%. Both activity and selectivity are fully restored after in situ hydrogen
treatment followed by calcination with oxygen and re-reduction. The first step of hydrogenation
probably re-reduces smaller cobalt crystallites and partly removes some wax and deposits. The
hydrogenation step can be carried out under conditions resembling FT-synthesis and can thus be
carried out at the plant location and with the catalyst loaded in the reactor. However, burning off
the carbonaceous deposits is far more demanding and probably requires removing the catalyst from
the reactor. These results are qualitatively in agreement with what can be expected from large scale
fixed-bed and slurry bubble column operations although the rate of deactivation for the Velocys’
catalyst is quantitatively smaller than those reported in available literature for large scale reactors.
The alternative micro-channel LTFT provider, CompactGTL, appears to have only a regeneration
patent directed at removing accumulated ammonia deposits [80].
8. Conclusions
An attempt to visualize the main catalyst deactivation mechanisms that impact FT catalyst
activity during commercial operation is shown in Figure 4. The regions shown are approximate
with respect to TOS and conversion, and there certainly will be some overlap. The following
takes place:
•

In the initial phase, (a few weeks) there will be driving forces towards both an increase in
activity by reduction as well as a decrease due to sintering. The net effect can be positive or
negative and will to a large degree depend on the catalyst formulation, pretreatment and FT
reactor environment. These mechanisms are well documented for slurry operations, but are
less evident for fixed-bed and micro-reactors.

145
•
•

•

Sintering is favored by small Co crystallites and high conversion levels due to enhanced
water activity, whereas reduction is facilitated by low conversion and water vapor pressure.
To a large extent oxidation of cobalt depends on operating conditions and may proceed at
high conversion conditions, particularly for smaller cobalt crystals and may involve
interaction with the support material.
Polycarbon deposition is the principal long-term cause of deactivation for all LTFT
catalysts and reactor types since commercial catalysts are stabilized against oxidation by
operation at realistic conversions, while reduction and sintering are short term phenomena
which determine steady-state activity after just 1–2 months of TOS.

Figure 4. Main reactions taking place during equilibration and deactivation of Co
Fischer-Tropsch catalysts.
Although deactivation during the first 6–18 months of operation can cause 30–60% loss in
activity in low-temperature Fischer-Tropsch synthesis with cobalt based catalysts, fortunately it
appears possible to regenerate the catalyst to approach the activity of a freshly equilibrated catalyst.
In a slurry reactor system a slip stream can be taken from the reactor for continuous rejuvenation.
For a fixed-bed reactor, including micro-channel reactors, an intermittent in situ rejuvenation and
in situ regeneration procedures can presumably be implemented as long as removal of wax and
combustion of carbonaceous deposits can be controlled to prevent large temperature excursions.
Nevertheless, based on patent literature, ex situ regeneration procedures appear to be the norm for
most FT processes.
Acknowledgement
The financial support from inGAP centre of research-based innovation, supported by the
Norwegian research council under contract no. 174893, is acknowledged with gratitude.

146
References and Notes
1.
2.

3.

4.
5.

6.
7.
8.

9.
10.
11.

12.

13.
14.

15.

Tsakoumis, N.E.; Rønning, M.; Borg, Ø.; Rytter, E.; Holmen, A. Deactivation of cobalt based
Fischer-Tropsch catalysts. A review. Catal. Today 2010, 154, 162–182e.
Moodley, D.J.; van de Loosdrecht, J.; Saib, A.M.; Overett, M.J.; Datye, A.K.;
Niemantsverdriet, J.W. Carbon deposition as a deactivation mechanism of cobalt-based
Fischer–Tropsch synthesis catalysts under realistic conditions. Appl. Catal.2009, 354, 102–110.
Davis, B.H.; Iglesia, E. Technology Development for Iron and Cobalt Fischer-Tropsch
Catalysts, Final Technical Report. Available online: http://www.fischer-tropsch.org/
DOE/DOE_reports/40308/FC26-98FT40308-f/FC26-98FT40308-f_toc.htm (accessed on
13 March 2015).
Argyle, M.D.; Frost, T.S.; Bartholomew, C.H. Cobalt Fischer–Tropsch Catalyst Deactivation
Modeled Using Generalized Power Law Expressions. Top. Catal. 2014, 57, 415–429.
Bezemer, G.L.; Bitter, J.H.; Kuipers, H.P.C.E.; Oosterbeek, H.; Holewijn, J.E.; Xu, X.;
Kapteijn, F.; van Dillen, A.J.; de Jong, K.P. Cobalt particle size effects in the Fischer-Tropsch
reaction studied with carbon nanofiber supported catalysts. J. Am. Chem. Soc. 2006, 128,
3956–3964.
Borg, Ø. Role of Alumina Support in Cobalt Fischer-Tropsch Synthesis. Ph.D. Theses,
NTNU, Norway, April 2007.
Diehl, F.; Khodakov, A.Y. Promotion of Cobalt Fischer-Tropsch Catalysts with Noble Metals:
A Review. Oil Gas Sci. Technol. Rev. IFP. 2009, 64, 11–24.
Ma, W.; Jacobs, G.; Keogh, R.A.; Bukur, D.B.; Davis, B.H. Fischer-Tropsch synthesis: Effect
of Pd, Pt, Re, and Ru noble metal promoters on the activity and selectivity of a 25%Co/Al2O3
catalyst. Appl. Catal. A 2012, 437–438, 1–9.
Information could be found at http://www.metal-pages.com/metalprices/rhenium/ and
http://www.metal-pages.com/metals/platinum/metal-prices-news-information/.
Ducreux, O.; Rebours, B.; Lynch, J.; Roy-Auberger, M.; Bazin, D. Microstructure of
Supported Cobalt Fischer-Tropsch Catalysts. Oil Gas Sci. Technol. Rev. IFP. 2009, 64, 49–62.
Blekkan, E.A.; Borg, Ø.; Frøseth, V.; Holmen, A. Fischer-Tropsch synthesis on cobalt
catalysts: The effect of water. In Catalysis; Royal Society of Chemistry: Cambridge, UK,
2007; Volume 20, pp. 1–21e.
Rytter, E.; Eri, S.; Skagseth, T.H.; Schanke, D.; Bergene, E.; Myrstad, R.; Lindvåg, A.
Catalyst Particle Size of Cobalt/Rhenium on Porous Alumina and the Effect on
FischeríTropsch Catalytic Performance. Ind. Eng. Chem. Res. 2007, 46, 9032–9036
Zennaro, R. The ENI-IFP/Axens GTL technology for the conversion of natural gas.
Proceedings of 6th World GTL Summit, London, UK, 17–19 May 2006.
Schanke, D.; Wagner, M.; Taylor, P. Proceedings of 1st annual gas processing symposium,
Doha, Qatar, 10–12 January 2009; Alfadala, H., Reklaitis, G.V.R., El-Halwagi, M.M., Eds.;
Elsevier: Amsterdam, The Netherlands, 2009.
Guettel, R.; Turek, T. Comparison of different reactor types for low temperature
Fischer–Tropsch synthesis: A simulation. Chem. Eng. Sci. 2009, 64, 955–964.

147
16. LeViness, S. presentation, Energy Frontiers International, 22nd October 2012. Available
online: http://www.velocys.com/arcv/press/ppt/EFI%202012%20Presentation%20121019_1_
rev2.pdf (accessed on 3 December 2014).
17. Steynberg, A.P.; Dry, M.E.; Davis, B.H.; Breman, B.B. Fischer-Tropsch Reactors. Stud. Surf.
Sci. Catal. 2004, 152, 64–195.
18. Schanke, D.; Lian, P.; Eri, S.; Rytter, E.; Sannæs, B.H.; Kinnari, K.J. Optimisation of
Fischer-Tropsch Reactor Design and Operation in GTL Plants. Stud. Surf. Sci. Catal. 2001,
136, 239–244.
19. Myrstad, R.; Eri, S.; Pfeifer, P.; Rytter, E.; Holmen, A. Fischer-Tropsch synthesis in a
microstructured reactor. Catal. Today 2009, 147S, S301–S304.
20. Holmen, A.; Myrstad, R.; Venvik, H.J.; Zhu, J.; Chen, D. Monolithic, microchannel and
carbon nanofibers/carbon felt reactors for syngas conversion by Fischer-Tropsch synthesis.
Catal. Today 2013, 216, 150–157.
21. Yang, J.; Eiras, S.B.; Myrstad, R.; Pfeifer, P.; Venvik, H.J.; Holmen, A. Fischer-Tropsch
synthesis over high-loading Co-based catalysts in a microreactor. Prep. ACS Div. Energy
Fuels 2014, 59, 828–830.
22. Rytter, E.; Ochoa-Fernández, E.; Fahmi, A. In Catalytic Process Development for Renewable
Materials; Imhof, P., van der Waal, J.-K., Eds.; Wiley: Weinheim, Germany, 2013; pp. 265–308.
23. Creyghton, E.J.; Mesters, C.M.A.M.; Reynhout, M.J.; Verbist, G.L.M.M. Process for
preparing a catalyst. Patent WO2008071640 A2, June 2008.
24. Freide, J.J.H.M.F.; Collins, J.P.; Nay, B.; Sharp, C. A history of the BP Fischer-Tropsch
catalyst from laboratory to full scale demonstration in Alaska and beyond. Stud. Surf.
Sci. Catal. 2007, 163, 37–44.
25. Van Berge, P.J.; Barradas, S. Catalysts. U.S. Patent 7365040 B2, March 2008.
26. Dogterom, R.J.; Mesters, C.M.A.M.; Reynhout, M.J. Process for preparing a hydrocarbon
synthesis catalyst. Patent WO2007068731 A1, June 2007.
27. Rytter, E.; Skagseth, T.H.; Wigum, H.; Sincadu, N. Enhanced strength of alumina based Co
Fischer-Tropsch catalyst. Patent WO 2005072866, August 2005.
28. Bellussi, G.; Carluccio, L.C.; Zennaro, R.; del Piero, G. Process for the preparation of
Fischer-Tropsch catalysts with a high mechanical, thermal and chemical stability.
Patent WO2007009680 A1, January 2007.
29. Ikeda, M.; Waku, T.; Aoki, N. Catalyst for Fisher-Tropsch synthesis and method for
producing hydrocarbon. Patent WO2005099897 A1, October 2005.
30. Inga, J.; Kennedy, P.; LeViness, S. Fischer-Tropsch process in the presence of nitrogen
contaminants. Patent WO2005071044, August 2005.
31. Soled, S.L.; Fiato, R.A.; Iglesia, E. Cobalt-ruthenium catalyst for Fischer-Tropsch synthesis.
Eur. Pat. Appl. 87310896, June1989.
32. Srinivasan, N.; Espinoza, R.L.; Coy, K.L.; Jothimurugesan, K. Fischer-Tropsch catalysts using
multiple precursors. U.S. Patent 6822008 B2, November 2004.
33. Compact GTL. The modular gas solution. Available online: http://www.compactgtl.com
(accessed on 24 February 2015).

148
34. Velocys.
Think
Smaller.
Available
online:
http://www.velocys.com/resources/
Velocys_Booklet.pdf (accessed on 24 February 2015).
35. Kliewer, C.E.; Soled, S.L.; Kiss, G. Characterizing Intrinsic Deactivation in Cobalt-Catalyzed
Fischer-Tropsch Synthesis. Available online: http:///www.nycsweb.org/files/37637807.pdf
(accessed on 24 February 2015).
36. Kliewer, C.E.; Kiss, G.; Soled, S.L. Characterizing Intrinsic Deactivation in Cobalt-Catalyzed
Fischer-Tropsch Synthesis. Microsc. Microanal. 2010, 16, 1258–1259.
37. Borg, Ø.; Hammer, N.; Enger, B.C.; Myrstad, R.; Lindvåg, O.A.; Eri, S.; Skagseth, T.H.;
Rytter, E. Effect of biomass-derived synthesis gas impurity elements on cobalt Fischer–Tropsch
catalyst performance including in situ sulphur and nitrogen addition. J. Catal. 2011, 279,
163–173.
38. Balonek, C.M.; Rane, A.H.L.S.; Schmidt, E.R.L.; Holmen, A. Effect of alkali metal impurities
on Co-Re catalysts for Fischer-Tropsch synthesis from biomass-derived syngas. Catal. Lett.
2010, 138, 8–13.
39. Lillebø, A.H. Conversion of biomass derived synthesis gas into liquid fuels via the
Fischer-Tropsch synthesis process: Effect of alkali and alkaline earth metal impurities and CO
conversion levels on cobalt based catalysts. Ph.D. Thesis, NTNU, Norway, September 2014.
40. Van Berge, P.J.; van de Loosdrecht, J.; Barradas, S. Method of treating an untreated catalyst
support and forming a catalyst precursor and catalyst from the teated support. U.S. Patent
6875720 B2, October 2005.
41. Rytter, E.; Eri, S.; Schanke, D.; Wigum, H.; Skagseth, T.H.; Borg, Ø.; Bergene, E.
Development of an Attrition Resistant Fischer–Tropsch Catalyst for Slurry Operation. Topics
Catal. 2011, 54, 801–810.
42. Behrmann, W.C.; Davis, S.M.; Mauldin, C.H. Catalyst preparation. Eur. Patent 535790 A1,
August 1992.
43. Rytter, E.; Eri, S.; Schanke, D.; Wigum, H.; Skagseth, T.H.; Sincadu, N.; Bergene, E.
Proceedings of 9th Novel Gas Conversion Symposium, Lyon, France, 30 May–3 June 2010.
44. Rytter, E.; Schanke, D.; Eri, S.; Wigum, H.; Skagseth, T.H.; Lian, P.; Sincadu, N.
Optimization of Statoil's Fischer-Tropsch Co/Re/alumina catalyst. Prep. ACS Div. Petro. Chem.
2004, 49, 182–183.
45. Tsakoumis, N.E.; Dehghan-Niri, R.; Rønning, M.; Walmsley, J.C.; Borg, Ø.; Rytter, E.;
Holmen, A. X-ray absorption, X-ray diffraction and electron microscopy study of spent cobalt
based catalyst in semi-commercial scale Fischer–Tropsch synthesis. Appl. Catal. A 2014, 479,
59–89.
46. Bartholomew, C.H. Sintering kinetics of supported metals: New perspectives from a unifying
GPLE treatment. Appl. Catal. A 1993, 107, 1–57.
47. Lee, D.; Lee, J.-H.; Ihm, A.-K. Effect of Carbon Deposits on Carbon Monoxide
Hydrogenation over Alumina Supported Cobalt Catalyst. Appl. Catal. 1988, 36, 199–207.
48. Swart, J.C.W.; van Steen, E.; Ciobícã, I.M.; van Santen, R.A. Interaction of graphene with
FCC-Co(111). Phys. Chem. Chem. Phys. 2009, 11, 803–807.

149
49. Tan, K.F.; Xu, J.; Chang, J.; Borgna, A.; Saeys, M.A. Carbon deposition on Co catalysts
during Fischer–Tropsch synthesis: A computational and experimental study. J. Catal. 2010,
274, 121–129
50. Weststrate, C.J.; KÕzÕlkaya, A.C.; Rossen, E.T.R.; Verhoeven, M.W.G.M.; Ciobîcă, I.M.;
Saib, A.M.; Niemantsverdriet, J.W. Atomic and Polymeric Carbon on Co(0001): Surface
Reconstruction, Graphene Formation, and Catalyst Poisoning. J. Phys. Chem. C 2012, 116,
11575–11583.
51. Gruver, V.; Young, R.; Engman, J.; Robota, H.J. The Role of Accumulated Carbon in
Deactivating Cobalt Catalysts during FT Synthesis in a Slurry-Bubble-Column Reactor.
Prep. Am. Chem. Soc., Div. Petrol. Chem. 2005, 50, 164–169.
52. Freide, J.J.H.M.F.; Gamlin, T.D.; Hensman, J.R.; Nay, B.; Sharp, C. Deactivation of
Cobalt-Based Catalyst in the F-T Process. J. Nat. Gas. Chem. 2004, 13, 1–9.
53. Schanke, D.; Hilmen, A.M.; Bergene, E.; Kinnari, K.; Rytter, E.; Ådnanes, E.; Holmen, A.
Study of the deactivation mechanism of Al2O3-supported cobalt Fischer-Tropsch catalyst.
Catal. Lett. 1995, 34, 269–284.
54. Li, J.; Zhan, X.; Zhang, Y.; Jacobs, G.; Das, T.; Davis, B.H. Fischer–Tropsch synthesis: effect
of water on the deactivation of Pt promoted Co/Al2O3 catalysts. Appl. Catal. A 2002, 228,
203–212.
55. Lancelot, C.; Ordomsky, V.V.; Stéphan, O.; Sadeqzadeh, M.; Karaca, H.; Lecroix, M.;
Curulla-Ferré, D.; Luck, F.; Fongarland, P.; Griboval-Constant, A.; Kodakov, A.Y. Direct
Evidence of Surface Oxidation of Cobalt Nanoparticles in Alumina-Supported Catalysts for
Fischer-Tropsch Synthesis. ACS Catal. 2014, 4, 4510–4515.
56. Jermwongratanachai, T.; Jacobs, G.; Shafer, W.D.; Ma, W.; Pendyala, V.R.R.; Davis, B.H.;
Kitiyanan, B.; Khalid, S.; Cronauer, D.C.; Kropf, A.J.; Marshall, C.L. Fischer–Tropsch
Synthesis: Oxidation of a Fraction of Cobalt Crystallites in Research Catalysts at the Onset of
FT at Partial Pressures Mimicking 50% CO Conversion. Top. Catal. 2014, 57, 479–490.
57. Kliewer, C.E.; Kiss, G.; DeMartin, G.J. Ex situ transmission electron microscopy: A fixed-bed
approach. Microsc. Microanal. 2006, 12, 135–144.
58. Kubaschewski, O.; Alcock, C.B. Metallurgical Thermochemistry, 5th ed.; Pergaman: New
York, NY, USA, 1979.
59. Rytter, E.; Schanke, D.; Eri, S.; Wigum, H.; Skagseth, T.H.; Bergene, E. Some QA and
optimization issues during development of the Statoil FT-catalyst. Available online:
http://www.globalspec.com/reference/43550/203279/qa-and-optimization-issues-duringdevelopment-of-the-statoil-ft-catalyst (accessed on 24 February 2015).
60. Kiss, G.; Kliever, C.E.; DeMartin, G.J.; Culross, C.C.; Baumgartner, J.E. Hydrothermal
deactivation of silica-supported cobalt catalysts in Fischer–Tropsch synthesis. J. Catal. 2003,
217, 127–140.
61. Eri, S.; Goodwin, J.G., Jr.; Marcelin, G.; Riis, T. Catalyst for Production of Hydrocarbons.
U.S. Patent 4801573, January 1987.
62. Kobylinski, T.P.; Kibby, C.L.; Pannell, R.B.; Leigh, E.L. Synthesis Gas Conversion using
ROR Activated Catalyst. Eur. Patent 253924, July 1988.

150
63. Ernst, B.; Bensaddik, A.; Hilaire, L.; Chaumette, P.; Kiennemann, A. Study on a cobalt silica
catalyst during reduction and Fischer-Tropsch reaction: In situ EXAFS compared to XPS and
XRD. Catal. Today 1998, 39, 329–341.
64. Arcuri, K.B.; LeViness, S.C. The Regeneration of Hydrocarbon Synthesis Catalyst—A Partial
Review of the Related Art Published during 1930 to 1952. Available online:
http://www.fischer-tropsch.org/primary_documents/presentations/AIChE%202003%20Spring
%20National%20Meeting/Paper%2086a%20Arcuri%20Regeneration.pdf (accessed on 24
February 2015).
65. Weststrate, C.; Hauman, M.; Moodley, D.; Saib, A.; Steen, E.; Niemantsverdriet, J. Cobalt
Fischer–Tropsch Catalyst Regeneration: The Crucial Role of the Kirkendall Effect for Cobalt
Redispersion. Topics Catal. 2011, 54, 811–815.
66. Van de Loosdrecht, J.; Saib, A.M. Catalysts. Patent WO 2008/139407, November 2008.
67. Saib, A.M.; Gauché, J.L.; Weststrate, C.J.; Gibson, P.; Boshoff, J.H.; Moodley, D.J.
Fundamental Science of Cobalt Catalyst Oxidation and Reduction Applied to the
Development of a Commercial Fischer-Tropsch Regeneration Process. Ind. Eng. Chem. Res.
2014, 53, 1816–1824.
68. Van de Loosdrecht, J.; Booysen, W.A. Catalysts. U.S. Patent 2011/0245355, June 2011.
69. Bezemer, G.L.; Nkrumah, S.; Smits, J.T.M. Process for Regenerating a Catalyst. U.S. Patent
2012/0165417, December 2011.
70. Rytter, E.; Eri, S.; Skagseth, T.H.; Borg, Ø. Fischer-Tropsch Catalyst Regeneration. U.S.
Patent 0210939, August 2013.
71. Soled, S.L.; Baumgartner, J.E.; Kiss, G. In situ catalyst regeneration/activation process.
U.S. Patent 6900151, November 2003
72. Ono, H.; Nagayasu, Y.; Hayasaka, K. Method for Manufacturing a Regenerated Fischer-Tropsch
Synthesis Catalyst and Hydrocarbon Manufacturing Method. U.S. Patent 2012/0322899,
December 2012.
73. Wright, H.A.; Raje, A.P.; Espinoza, R.L. Pressure Swing Catalyst Reneration Procedure for
Fischer-Tropsch Catalyst. U.S. Patent 6962947, November 2005.
74. Huang, J.-R.; Arcuri, K.; Agee, K.; Schubert, P.F. Process for Regenerating a Slurry
Fischer-Tropsch Catalyst. U.S. Patent 6989403, April 2004.
75. LeViness, S. 245th ACS National Meeting, April 2013, New Orleans. Available online:
http://www.velocys.com/arcv/press/ppt/AIChE%202013%20Presentation%2004–30–13%20%20LeViness%20v1_rev1.pdf (accessed on 10 March 2105).
76. Bailey, S.; Gray, G.; Lok, M.C. Methods for the production of cobalt catalysts supported on
silicon dioxide and their use. Patent WO 2001/062381, August 2001.
77. Iglesia, E.; Soled, S.L.; Fiato, R.A. Cobalt-ruthenium catalysts for Fischer-Tropsch synthesis
and process for their preparation. U.S. Patent 4738948, April 1988.
78. Iglesia, E.; Soled, S.L.; Fiato, R.A.; Via, G.H. Bimetallic Synergy in Cobalt Ruthenium
Fischer-Tropsch Synthesis Catalysts. J. Catal. 1993, 143, 345–368.

151
79. Daage, M.; Koveal, R.J.; Chang, M. Catalyst Regeneration. U.S. Patent 6800579,
January 2002.
80. Minnie, O.R.; Ndifor, E.N.; Kavian, S.; Smith, B.; Knowles, R. Process for the regeneration
of a Fischer-Tropsch Catalyst. Patent WO 2013/093423, June 2013.

152

Influence of Reduction Promoters on Stability of
Cobalt/γ-Alumina Fischer-Tropsch Synthesis Catalysts
Gary Jacobs, Wenping Ma and Burtron H. Davis
Abstract: This focused review article underscores how metal reduction promoters can impact
deactivation phenomena associated with cobalt Fischer-Tropsch synthesis catalysts. Promoters can
exacerbate sintering if the additional cobalt metal clusters, formed as a result of the promoting
effect, are in close proximity at the nanoscale to other cobalt particles on the surface. Recent efforts
have shown that when promoters are used to facilitate the reduction of small crystallites with the
aim of increasing surface Co0 site densities (e.g., in research catalysts), ultra-small crystallites
(e.g., <2–4.4 nm) formed are more susceptible to oxidation at high conversion relative to larger
ones. The choice of promoter is important, as certain metals (e.g., Au) that promote cobalt oxide
reduction can separate from cobalt during oxidation-reduction (regeneration) cycles. Finally, some
elements have been identified to promote reduction but either poison the surface of Co0 (e.g., Cu),
or produce excessive light gas selectivity (e.g., Cu and Pd, or Au at high loading). Computational
studies indicate that certain promoters may inhibit polymeric C formation by hindering C-C coupling.
Reprinted from Catalysts. Cite as: Jacobs, G.; Ma, W.; Davis, B.H. Influence of Reduction
Promoters on Stability of Cobalt/γ-Alumina Fischer-Tropsch Synthesis Catalysts. Catalysts 2014,
4, 49-76.
1. Introduction
Fischer-Tropsch synthesis (FTS) making use of cobalt catalysts is the core of the gas-to-liquids
(GTL) process [1,2]. Due to the high H2/CO syngas ratio derived from reforming of natural gas,
additional water gas shift is not required to adjust the ratio upward, and internal water-gas shift
(WGS) activity is undesirable. This is one benefit of cobalt catalysts relative to iron catalysts for
GTL, as the former typically possess low intrinsic WGS activity. Because cobalt is much more
expensive than its iron counterpart, and because the reaction occurs on the surface, it is important
to disperse the cobalt metal particles in order to improve usage efficiency.
A typical cost effective way to do so is to impregnate the pre-calcined support with a cobalt nitrate
solution by wet or dry (incipient wetness) impregnation followed by drying, air calcination to
decompose the cobalt nitrate precursor to cobalt oxide, and reduction (e.g., 10 h in hydrogen gas at
350 °C) to Co0 crystallites (typically in the range of 5 to 20 nm). The surface of Co particles
provides the catalytically active sites for Fischer-Tropsch synthesis.
However, with typical reduction of supported cobalt at low temperatures (e.g., 350–400 °C)
appropriate for obtaining active small crystallites of 6–15 nm, a sizeable fraction (typically in the
range of 15–70%) of the cobalt remains in the oxide form, mainly as CoO. The fraction of
unreduced cobalt is larger for supports such as alumina which interact strongly with cobalt oxide
and for low cobalt loadings, e.g., less than 10–15% on such supports. The extent of the interaction
increases with decreasing loading of cobalt. At low loadings (e.g., < 5%), 60 to 80% of the cobalt is

153
present as CoO strongly bound to the support surface, i.e., a surface cobalt aluminate, CoO*Al2O3,
which requires very high temperatures to reduce [3]. At high loadings (e.g., 15%–30%Co), cobalt
will be present primarily as Co3O4, which reduces in two steps: Co3O4 + H2 = 3CoO + H2O and
3CoO + 3CoO = Co + 3H2O, for which maximum rates of reduction occur at about 300–350 and
500–650 °C, respectively [4]. Thus, following reduction of an unpromoted 15–30% Co/alumina at
350–400 °C for 5–15 h, a significant fraction (30–60%) of CoO typically remains [5–8]. Since
higher extents of reduction (80–90%) of cobalt are highly desirable, i.e., correlated with higher
activity on a per g catalyst basis, as well as improved C5+ selectivity, there is considerable interest
and widespread application of noble metal promoters, which facilitate the reduction of cobalt
oxides and increase the surface density of cobalt active sites.
This article reviews a number of stability issues associated with the application of promoters for
cobalt FTS catalysts. Examples are provided to demonstrate a number of considerations for
selecting a noble metal for Co catalysts. The main point of the article is that each promoter has its
own advantages and set of issues that must be addressed and, in some cases, still defined.
2. Results and Discussion
Figure 1 compares temperature programmed reduction (TPR) TPR profiles of a number of noble
metal and Group IB-promoted 15%Co/γ-Al2O3 catalysts pertinent to this review. The commonly
used promoters are Pt, Re, and Ru and the solid line profiles are at close to atomically equivalent
loadings. Pt and Ru facilitate the reduction of both steps of cobalt oxide reduction, while Re
catalyzes the reduction of primarily the second step. This has been explained by Re oxide reducing
at a higher temperature than Pt and Ru and that a reduced form of the promoter is required to obtain
the promoting effect [7]; however, further confirmation of this is needed as the oxidation state
remains in question [9]. Both Pt and Re appear to be more effective at facilitating reduction relative
to Ru, but higher loadings of Ru can be used to further facilitate reduction, as indicated by the
dashed profile at 0.5%Ru loading. Similar trends were reported in the TPR peak locations in a
recent investigation of promoter characteristics by Cook et al. [9], as shown in Figure 2.
Cu, Ag, and Au (Group IB) promoters are also effective at promoting the reduction of cobalt
oxides, as described in Figure 1 (top) [10]. However, note that the loadings indicated by the solid
lines are approximately three times higher on an atomic basis than those of the commonly used
promoters (Pt, Re, Ru) shown in the lower part of the figure. The costs of Ag and Cu are, whether
on a weight or atomic basis, much lower than any of the other promoters shown. Therefore, it was
of interest to explore their ability to facilitate reduction at even higher loadings. Increasing the
loading by a factor of 3.3 resulted in further and important shifts of both TPR peaks of cobalt oxide
reduction to lower temperatures (Figure 1, top).

154
Figure 1. TPR profiles demonstrate the effectiveness of Cu for facilitating reduction
of cobalt oxides. Curve labels: unpromoted 15%Co/Al2O3 (thick solid) and
Cu-promoted with 0.033%Cu (thick dashed), 0.49%Cu (thin solid), and 1.63%Cu
(thin dash-dotted) by weight. (Reproduced with permission from [7] and [10] Copyright
2002, 2009, Elsevier).

155
Figure 2. TPR profiles by Cook et al. [9] show that Pt and Re are more efficient than
Ru in facilitating reduction of cobalt oxides over 25%Co/γ-Al2O3 catalysts with
equivalent atomic loading (i.e., noble metal / Co ratio was 0.007).

The choice of promoter metal and its loading may influence the stability of cobalt catalysts in a
number of ways. The first section examines how reduction promoter type and loading influence the
activity and selectivity of cobalt catalysts, while the second section discusses how promoters may
exacerbate deactivation rates through oxidation and a possible complex sintering mechanism, the
two of which are not mutually exclusive. A brief summary of the application of computational
methods is also provided, which discusses the location of promoter relative to cobalt, the resistance
or sensitivity of cobalt to oxidation depending on size, and how promoters may hinder carbon
formation pathways. Finally, in adding a second catalytic metal to the catalyst, the ability to
regenerate the catalyst in a simple and effective manner becomes an important concern.
2.1. Influence of Promoter Choice and Loading on Catalyst Activity and Selectivity
2.1.1. Example #1—Copper
The first example demonstrates a relatively inexpensive metal that is highly effective for
promoting the reduction of cobalt oxides: Cu, which is a common promoter in iron carbide FTS
catalysts [11,12]. As of this writing, Cu is approximately 0.015% of the cost of Pt on a mass basis
and would seem to be an ideal candidate as a promoter.
With increases in extent of reduction of cobalt (from 49.8% for 15%Co/Al2O3 to 53.2, 69.4, and
93.3% for 0.033%, 0.49%, and 1.63% Cu-promoted 15%Co/Al2O3 catalysts, respectively) the
active metal site densities with Cu addition increased also, and hydrogen chemisorption uptakes
measured by TPD increased from 72 μmol H2/gcat for the unpromoted 15%Co/Al2O3 catalyst to
82, 140, and 172 μm H2/gcat for the 0.033%, 0.49%, and 1.63% Cu-promoted 15%Co/Al2O3
catalysts [10]. However, the metal dispersions do not account for the partitioning of metal type on

156
the surface of Co particles, or the influence of the presence of Cu on the ensembles of Co required
for conducting the synthesis. Surface enrichment by Cu has been detected in bimetallic Cu-Co
catalysts before [13].
Table 1 shows a comparison of XCO at the same weight hourly space velocity of two
Cu-promoted 15%Co/Al2O3 catalysts relative to the unpromoted 15%Co/Al2O3 catalyst. Despite
increases in metal site densities as measured by hydrogen TPD, a decrease in XCO is observed,
which is exacerbated at higher Cu loading [10]. These results suggest a poisoning of surface Co
sites, likely due to enrichment of Cu at the surface. This finding is further supported by the changes
in selectivity that occur when comparing the catalysts to the unpromoted catalyst at a similar
conversion level. Table 2 shows that at the lower Cu promoter loading, the methane is slightly
increased, C5+ is slightly decreased [10]. However, increasing the Cu promoter loading further
leads to a prohibitive increase in methane selectivity (21.6% versus 9.2%) and a precipitous drop in
C5+ selectivity (47.7% versus 81.6%) [10].
Table 1. XCO for two Cu promoted 15%Co/Al2O3 catalysts at a SV of 4.2 NL/gcat/h
relative to the unpromoted catalyst. Conditions: 220 °C, 1.6 MPa, H2/CO = 2.0 (adapted
with permission from [10] Copyright 2009, Elsevier).
Catalyst

TOS (h)

XCO (%)

15%Co/Al2O3
0.49%Cu-15%Co/Al2O3
1.63%Cu-15%Co/Al2O3

26–98
30–99
25–104

28.7
27.9
14.2

SV
(NL/gcat/h)
4.2
4.2
4.2

Table 2. Two comparisons of product selectivity at similar XCO levels for two Cu
promoted 15%Co/Al2O3 catalysts relative to the unpromoted catalyst. Conditions:
220 °C, 1.6 MPa, H2/CO = 2.0 (adapted with permission from [10] Copyright
2009, Elsevier).
Catalyst

XCO (%)

15%Co/Al2O3
0.49%Cu-15%Co/Al2O3
15%Co/Al2O3
1.63%Cu-15%Co/Al2O3 *

47.8
50.6
28.7
29.9

SV
(NL/gcat/h)
2.0
1.7
4.2
1.0

S(CH4)

S(C5+)

S(CO2)

8.9
9.9
9.2
21.6

80.6
76.6
81.6
47.7

0.82
0.83
0.67
1.51

* Due to the low activity of the 1.63%Cu promoted catalyst, a separate comparison was made at lower
XCO, as it was not possible to decrease SV further.

2.1.2. Example #2—Silver and Gold
A comparison between Ag and Au shows that, in the case of Ag promoted 15%Co/Al2O3, the
catalyst achieves higher activity (Table 3) and C5+ selectivity (Table 4) than the unpromoted
catalyst at both high and low loadings of promoter. A 15% Co/Al2O3 catalyst promoted with 1.51%
Au also performs better than the unpromoted catalyst with both an increase in productivity and a
slight improvement in selectivities [10]. However, at a higher Au loading (5.05%) the catalyst

157
performed poorly with a steep drop in productivity (from XCO of 51.7 at 1.51%Au to an XCO of
14.1 at 5.05%Au at SV = 4.2, Table 3) and adverse impacts on selectivity (C5+ is 60.1% compared
to 81.6% for the unpromoted catalyst at XCO of ~28%, Table 4) [10]. Thus, noble metal loading of
promoter is important, not only from the standpoint of cost.
In a recent detailed kinetic investigation [14], which was a collaboration between CAER and
Texas A&M University in Qatar, modeling results point to the presence of two kinds of sites on the
Co FTS catalyst for the production of methane—FTS sites from standard Anderson-Schulz-Flory
kinetics and additional sites for methanation. The results of the Au promoted catalyst at the lower
loading, and the Ag promoted catalysts at both low and high loadings, suggest that these Group IB
promoters assist in either blocking methanation sites or controlling the relative surface fugacity of
hydrogen relative to adsorbed CO and intermediates on the surface of the cobalt catalyst.
Table 3. XCO for two Ag and Au-promoted 15%Co/Al2O3 catalysts at a SV of
4.2 NL/gcat/h relative to to the unpromoted catalyst. Conditions: 220 °C, 1.6 MPa,
H2/CO = 2.0 (adapted with permission from [10] Copyright 2009, Elsevier).
Catalyst
15%Co/Al2O3
1.51%Au-15%Co/Al2O3
5.05%Au-15%Co/Al2O3
0.83%Ag-15%Co/Al2O3
2.76%Ag-15%Co/Al2O3

TOS (h)
26–98
26–57
30–84
20–47
22–92

XCO (%)
28.7
51.7
14.1
50.4
46.9

SV (NL/gcat/h)
4.2
4.2
4.2
4.2
4.2

Table 4. Product selectivities for Ag and Au-promoted 15%Co/Al2O3 catalysts at
XCO values comparable to the unpromoted catalyst. Conditions: 220 °C, 1.6 MPa,
H2/CO = 2.0 (adapted with permission from [10] Copyright 2009, Elsevier).
Catalyst
15%Co/Al2O3
1.51%Au-15%Co/Al2O3
0.83%Ag-15%Co/Al2O3
2.76%Ag-15%Co/Al2O3
15%Co/Al2O3
5.05%Au15%Co/Al2O3*

XCO (%)
47.8
50.0
50.4
46.9
28.7

SV (NL/gcat/h)
2.0
4.2
4.2
4.2
4.2

S(CH4)
8.9
8.0
7.7
7.6
9.2

S(C5+)
80.6
83.7
83.6
85.0
81.6

S(CO2)
0.82
0.83
0.94
0.87
0.67

27.1

1.0

18.0

60.1

1.68

* Due to the low activity of the 5.05%Au promoted catalyst, a separate comparison was made at lower
XCO, as it was not possible to decrease SV further.

2.1.3. Example #3—Common Promoters (Pt, Re, Ru)
Slight differences in selectivity can also be achieved with the commonly used reduction
promoters, which are Pt, Re, and Ru [15], as compiled in Table 5. Ruthenium itself is catalytically
active for the FTS reaction, and higher alpha values have been measured in the hydrocarbon
distribution [16]. Therefore, it is not surprising that lower methane and higher C5+ were achieved
with a 0.26%Ru-25%Co/Al2O3 catalyst relative to an unpromoted one. Re promoter was also found

158
to give slightly better selectivities, in agreement with the work of Borg et al. [17] (Table 6).
However, an atomically equivalent amount of Pt slightly (though not prohibitively) worsened the
selectivities, and an attempt to use Pd to replace Pt resulted in a significantly poorer
product distribution.
Table 5. Product selectivities* for 25%Co/Al2O3 catalysts containing commonly used
promoters (Pt, Re, Ru) or Pd at XCO values comparable an unpromoted reference
catalyst. Conditions: 220 °C, 2.2 MPa, H2/CO = 2.1 (adapted with permission from [15]
Copyright 2012, Elsevier).
Catalyst
25%Co/Al2O3
0.26%Ru-25%Co/Al2O3
0.48%Re-25%Co/Al2O3
0.50%Pt-25%Co/Al2O3
0.27%Pd-25%Co/Al2O3

XCO (%)
49.4
51.3
49.6
48.0
50.3

SV (NL/gcat/h)
4.3
7.6
8.0
5.6
4.9

S(CH4)
7.9
7.0
7.2
8.3
11.5

S(C5+)
83.4
86.8
86.0
83.0
75.9

* S(CO2) ranged from 0.35–0.75% in all catalysts. All data taken within first 81 h on-stream.

Table 6. Product selectivities from data taken at 210 °C, 2.0 MPa, H2/CO = 2, and XCO
of 43–44% (adapted with permission from [17] Copyright 2009, Elsevier) over
20%Co/γ-Al2O3 catalysts without or with 0.5%Re using narrow pore (7.4 nm), medium
pore (12.3 nm), and wide pore (16.7 nm) supports.
Catalyst
Co/NPA
CoRe/NPA
Co/MPA
CoRe/MPA
Co/WPA
CoRe/WPA

Hydrocarbon selectivity (%)
C1
C2-C4
C 5+
9.0
9.9
81.1
8.8
9.5
81.7
8.6
8.7
82.8
8.4
8.3
83.4
8.0
7.5
84.5
8.0
7.2
84.9

In terms of differences in catalyst structure, the three commonly used promoters have, in a
number of cases, been observed to be in atomic contact with Co (e.g., as an alloy), with no
presence of promoter-promoter coordination at relatively low loadings (Re [18–20], Ru [21],
Pt [22–24]). This is not always the case (e.g., Ru [9,16]), indicating that loading and preparation
method are also factors that influence coordination environment. Pretreatment is also a factor. For
example, Iglesia et al. [25] noted with Ru-Co/TiO2 catalysts that, with increasing calcination
temperature, total coordination of Ru with neighbors increases to suggest sintering, but that most of
the increase in coordination is due to Ru taking on coordination with Co; thus, calcination
promoted mixing of the two metals.
Moreover, Chonco et al. [26] have recently demonstrated with physical mixtures of Pt/Al2O3
and Co/Al2O3 that atomic coordination of the promoter to cobalt is not always required to obtain a
reduction promoting effect. In our work, unlike the Re, Ru, and Pt promoted Co/alumina catalysts
at low promoter loadings, Pd promoter exhibited some promoter-promoter (i.e., Pd-Pd) coordination,

159
suggesting the presence of well dispersed islands of Pd that likely gave rise to excessive
hydrogenation activity [15] and rapid deactivation relative to the other three promoted catalysts [15].
Considering commercial research, a patent by Sasol researchers [1] examined Ru and Re
promoters (of a catalyst containing 30 g Co and 100 g Al2O3) at 0.41 g and 3.0 g levels,
respectively, versus a catalyst containing just 0.05 g of Pt and reported slightly higher productivity
with the Pt promoted catalyst (0.349 kg HC/kg cat/h at XCO = 87 vol.% with Pt versus 0.307 and
0.281 kg HC/kg cat/h for Ru and Re at XCO = 77 and 70%, respectively). The conditions were
220 °C, 2.0 MPa, H2/CO = 2/1; space velocity of 2.0 mn3/h/kg catalyst. Under the same conditions,
similar productivity (~0.29 kg HC/kg cat/h at XCO = 73 vol.%) was observed with a 0.28 g Ru as
with 0.05 g Pt with catalysts of lower loading (20 g Co and 100 g Al2O3). The results appear to
indicate that Pt is a very efficient promoter.
In a patent by Conoco researchers [27], the benefit of adding Re on selectivity was highlighted.
Examples 27 through 31 in that patent compare 1%Re promoted 20%Co/Al2O3 catalysts with an
unpromoted catalyst. Higher conversions (77–100%) and C5+ productivities (240–270 g/h/kgcat) for
the Re promoted catalysts relative to the unpromoted catalyst (XCO of 65% and C5+ productivity of
170 g/h/kgcat) were reported. In addition, improvements in selectivities were observed as well,
including improvements in alpha (0.89–0.90 with Re promotion versus 0.88 for the unpromoted
catalyst) and decreases in methane (9–15 wt.% for Re promoted versus 18 wt.% for the
unpromoted catalyst). Many other examples of Re promotion only or in combination with other
elements are also highlighted in the patent report.
2.1.4. Example #4—Impact of Loading for Pt and Ag Promoted Catalysts
The final example is provided to show that, just because the promoter forms promoter-promoter
bonds (i.e., as in the case of Pd described in the previous subsection), it should not immediately be
ruled out. Ag, by itself, is a catalyst that is only weakly active for hydrogenation, and its addition as
a promoter does result in significant Ag-Ag coordination, but the resulting activity and selectivity
of the Co catalyst is improved. Figure 3 (left) compares the Ag K-edge EXAFS Fourier transform
magnitude spectra of Ag-promoted 25%Co/Al2O3 catalysts as a function of Ag loading. While
a lower distance peak for Ag-Co coordination is suggested (and confirmed by EXAFS fittings),
with increasing loading of Ag, the general trend in EXAFS fittings for Ag-promoted 25%Co/Al2O3
catalysts was increasing Ag-Ag coordination (higher distance peak), such that the NAg-Co/NAg-Ag
ratio decreased from 0.59 at 0.276%Ag loading to 0.16 at 2.76%Ag loading [28].
Figure 3 (right) compares the Pt LIII-edge EXAFS Fourier transform magnitude spectra of
Pt-promoted 25%Co/Al2O3 catalysts. A single low distance peak indicates primarily Pt-Co
coordination, with no visible Pt-Pt coordination being evident. As shown in Table 7, as loading was
increased for the Pt promoted catalyst, a slight negative impact on selectivity occurred, with
a slight decrease in C5+ and a slight increase in the WGS rate [28]. For all Pt-promoted catalysts,
slightly lower C5+ selectivities were observed compared to the unpromoted catalyst. With the
Ag-promoted catalysts, C1 and C5+ selectivities were slightly improved at all loadings [28]. Thus,
the presence of a weakly hydrogenating metal [29], Ag, did not adversely affect selectivity to a
significant degree, even when excessive amounts of promoter were added [28].

160
In summary, the above examples show that (1) metals that facilitate reduction of cobalt oxides
do not automatically increase XCO on a per gram of catalyst basis; (2) type and loading of promoter
influence activity and selectivity such that a metal that may promote XCO (on a per g basis) at lower
loading may or may not poison or adversely impact surface fugacities (and selectivity) at higher
loadings; and (3) the intrinsic hydrogenation rate of the promoter is an important factor to consider,
as it may adversely or beneficially influence selectivity.
2.2. Influence of Promoter Addition on Oxidation and Complex Sintering of Cobalt
2.2.1. Reoxidation of Small Cobalt Crystallites at the Onset of FTS at Realistic Conversions
The primary aims of adding a metal promoters are to (1) lower reduction temperature thereby
increasing extent of reduction to Co metal and (2) boost active site densities by facilitating the
reduction of cobalt oxide crystallites in strong interaction with the alumina support, such that
clusters of cobalt metal crystallites can be formed to provide the active surface for carrying out the
FTS reaction. Thus, when a promoter is added, the additional gain in active site density will be due
in large part to the reduction of smaller cobalt oxide species having stronger interactions with the
support. Depending on the loading of cobalt and method of preparation, if smaller cobalt metal
crystallites (e.g., <2–4.4 nm [30,31]) within cobalt clusters are formed, they may be susceptible to
reoxidation [30,32]. Some researchers have recently indicated that Co clusters less than 6–8 nm
have lower intrinsic activity [33,34]. Additional investigations are needed in this area. At
commercially relevant FTS conditions, a problem was identified by us in defining intrinsic activity;
as chemisorption is conducted on freshly activated catalysts, any oxidation of such small Co
clusters that occurs at commercially relevant conditions can mask a measurement of intrinsic
activity at the level of the active site [35]. Therefore, it is important to take into account the
oxidation state of Co in the working FTS catalyst.
To probe the role of oxidation, a recent XANES study was made whereby freshly
reduced unpromoted and Pt-promoted cobalt/alumina catalysts were subjected to H2:CO:H2O
mixtures typical of the 50% conversion condition of a slurry phase reactor [36] for one hour. A
lower-than-commercial loading of 10% cobalt was utilized in order to favor the formation of small
cobalt crystallites after activation that fall in the range of being susceptible to reoxidation. The
average cobalt cluster size (i.e., cluster of crystallites) was ~5 nm [36]. Even though the
10%Co/Al2O3 catalyst was reduced at 550 °C as opposed to 400 °C for the 0.5%Pt-10%Co/Al2O3
catalyst, the extent of reduction from XANES indicated that the Pt-promoted catalyst had a higher
extent of reduction, as defined by the intensity of the white line. When switching to conditions to
mimic 50% conversion, the white line intensity in the XANES spectra of both catalysts increased
significantly (Figure 4), but the change was more severe for the Pt-promoted catalyst [36]. This
reoxidation occurred rapidly, is confined to a fraction of cobalt, and is not associated with the
initial decay period commonly observed in FTS reaction tests, which may take on the order of days
to establish. Reoxidation of the small cobalt crystallites (<2–4.4 nm) has been recently verified by
in-situ XRD and magnetometer investigations [37].

161
Figure 3. (left) Ag promoted 25%Co/Al2O3 catalysts displayed EXAFS peaks that
could only be fitted well by including both Ag-Co and Ag-Ag coordination; (right) Pt
promoted 25%Co/Al2O3 catalysts displayed a single peak in the first coordination shell
that could be fitted well by only including Pt-Co coordination (adapted with permission
from [28] Copyright 2013, Elsevier).

Table 7. Product selectivity * of 25%Co/Al2O3 catalysts having different loadings of
Ag and Pt (adapted with permission from [28] Copyright 2013, Elsevier).
Catalyst
25%Co/Al2O3
0.5%Pt-25%Co/Al2O3
2.0%Pt-25%Co/Al2O3
5.0%Pt-25%Co/Al2O3
0.276%Ag-25%Co/Al2O3
1.11 %Ag-25%Co/Al2O3
2.76%Ag-25%Co/Al2O3

XCO (%)
51.0
52.0
45.0
52.5
46.4
48.1
44.5

SV (NL/gcat/h)
3.4–4.2
1.7–12
9.0–12
10–16
8.8–12
9.3–12
7.0–12

S(CH4)
8.3
9.1
9.1
9.5
7.4
7.3
7.6

S(C5+)
82.5
81.2
81.9
80.7
84.1
83.7
84.1

S(CO2)
0.8
1.1
1.1
3.2
0.4
0.4
0.6

* All data taken within first 58 h on-stream. T = 220 °C; P = 2.2 MPa; H2/CO = 2.1.

An extreme case in terms of small Co cluster size was also recently examined where Co
particles were infused in the pores of KL-zeolite support by a CVD method [35] to produce a
0.5%Pt-5%Co/KL catalyst with 1 nm cobalt particles. The catalyst, following activation in
hydrogen, exhibited an extent of reduction of 75%. However, after exposure to FTS conditions
(220 °C, 1.8 MPa, H2/CO 1.95, SV of 3.0 NL/gcath), the extent of reduction fell to 33%, as
measured by XANES spectroscopy. A loss in Co-Co coordination and growth of Co-O

162
coordination was quantified by EXAFS. XCO fell to 3.0%, and the resulting product selectivities
were very poor (CH4 selectivity of 29.6%, C5+ selectivity of 49.8%, and CO2 selectivity of 3.4%).
Figure 4. (left) Pt promoted and (right) unpromoted 10%Co/Al2O3 catalyst after (solid)
hydrogen activation (550 °C for unpromoted and 400 °C for Pt-promoted) and (dashed)
after exposure to H2, CO, and H2O partial pressure ratios mimicking XCO = 50% at
20.7 bar (adapted with permission from [36] Copyright 2014, Springer).

2.2.2. Sintering and Co Support Compound Formation during Initial Deactivation Period Prior to
Leveling-off Period
The mechanisms of sintering during FTS on Co catalysts during typical commercial operation
are not well understood and hence will not be discussed in detail in this review. There is evidence
to indicate that H2O accelerates sintering of metallic particles [38–41]; on this aspect, a mechanism
involving surface oxidation, coalescence, and formation of larger clusters has been postulated by
Sadeqzadeh et al. [39]. Further research is needed in this area.
Unreduced cobalt oxide in the working FT catalyst could be problematic since it may coalesce,
reduce, and provide a driving force for the sintering of cobalt metal particles; this leads to net
reduction with time onstream in the initial catalyst decay period (i.e., which follows the onset
period of ~1 h as described in the previous section) [32,42]. Moreover, small cobalt oxide species
can react with the support and contribute to the formation of cobalt aluminates [42,43]. Since a

163
promoter can facilitate the reduction of smaller species that may, depending on size, be susceptible
to reoxidation at high conversion (as previously described), any cobalt oxide formed at the onset
can contribute to either net reduction/sintering with time on-stream or cobalt aluminate formation.
Net reduction and changes in Co-Co coordination consistent with sintering were observed for a
0.2%Re-15%Co/Al2O3 research catalyst at CAER during 2000 h of operation (Figure 5) [44,45].
Sintering, i.e., growth of Co metal crystallites, and carbon formation were observed by
Saib et al. [31] during a 56-day commercial test of a 0.05%Pt/20%Co catalyst. Sintering was
determined to be rapid, reaching completion within the initial 15–16 days based on analysis by
synchrotron-XRD of samples withdrawn from the reactor during this period; an increase in average
crystallite diameter from 9 to 14 nm was observed. Formation in the catalyst of unreactive surface
carbons, which restructured or poisoned the catalysts, occurred relatively more slowly over the
56-day period. Carbon deposits were analyzed by TPR during the latter part of this run [46]. During a
second 140-day run [32], extent of reduction (EOR) was determined by periodically removing catalyst
samples and analyzing them by XANES (Figure 6). EOR was determined to increase from 53 to 89%
over the 140-day run. A small amount of cobalt aluminate formation was also observed, as it was
detected in a used commercial 20%Co/0.05%Pt/Al2O3 catalyst [31]. Cobalt aluminate was also
observed in used 0.2%Re-15%Co/Al2O3 research catalyst samples [44,47] (Figure 7).
Figure 5. (left) XANES and (right) EXAFS spectra as a function of time for
a 0.2%Re-15%Co/Al2O3 catalyst. T = 220 °C, 2.0 MPa, SV = 5 SL/h/gcat, H2/CO = 2:1.
Adapted with permission from [44] (Copyright 2003, Elsevier) and [45] (2006).

164
The impact of the promoter on cobalt aluminate formation from initially reduced small Co
crystallites is difficult to assess, because unpromoted cobalt/alumina catalysts contain more
residual unreduced CoO after activation, and this residual oxide, which is inactive for FTS, can also
react with the support to form cobalt aluminate. Thus, there is a question as to how much cobalt
aluminate is formed from the oxidation of very small (e.g., <2 – 4.4 nm) crystallites of cobalt metal
at the onset of CoO formation (e.g., within the first hour, as shown in Figure 7 [44,45]) and its
subsequent reaction with alumina, and how much is formed from the reaction of residual CoO (i.e.,
leftover due to incomplete reduction) with the support, as described by Sasol researchers [43]. The
cobalt aluminate was detected by them in XANES derivative spectra, as shown in Figure 8. The
results from the previous section suggest that the former is dependent on Co crystallite size and
P(H2O)/P(syngas) ratio, which is in turn influenced by conversion level. In a recent kinetic
investigation, excursions of a Ru-promoted 25%Co/Al2O3 catalyst (average cluster size of 5.0 nm)
to high CO conversion levels (e.g., XCO > 75%) resulted in the oxidation of a fraction of cobalt
clusters [21] (as demonstrated by changes in the lineshape of the XANES derivative spectra) and
increases in CO2 and CH4 selectivities (Figure 9). The oxidized cobalt is active for WGS, and the
additional H2 produced therefrom tends to increase the C1 product. Thus, there appears to be an
unfavorable synergism in the selectivities of CO2 and C1 when this threshold is surpassed
(Figure 9) [21].
Crystallite size sensitivity for cobalt aluminate formation was also suggested in water co-feeding
studies. Although H2O co-feeding can lead to improvements in activity and selectivity for certain
cobalt catalysts [48], when a Pt promoter was utilized to facilitate the reduction of Co oxides in a
15%Co/Al2O3 catalyst (average cluster diameter = 5.6 nm), at 28 vol.% added H2O the catalyst
underwent significant cobalt aluminate formation, as demonstrated in Figure 10 (left) and (right) [49,50]
along with catastrophic deactivation (75% drop in XCO). An unpromoted 25%Co/Al2O3 catalyst
with larger cluster size (11.8 nm average diameter) [40] was more resistant to this phenomenon.
Thus, on the one hand, a promoter is very useful for boosting Co° site densities during activation
when the support interaction with Co oxides is high. On the other hand, if the strongly interacting
Co oxides are reduced and form tiny Co0 crystallites on the surface, they are more sensitive to H2O.
Higher loadings can help to make the catalyst more robust by increasing cobalt size, a technique
that has been implemented commercially, and thereby reoxidation and subsequent Co aluminate
formation may be largely avoided. With a commercial catalyst stabilized against these processes,
only up to ~3% cobalt aluminate was formed during realistic FTS conditions [31]. However, it
should be noted that when exposed to 1.0 MPa H2O an increase was observed to 10% cobalt
aluminate [43]; thus, water co-feeding or operating at high conversions may have drawbacks,
depending on catalyst type and conditions utilized. A schematic of the structural changes discussed
for research catalysts, including reoxidation of tiny (<2–4.4 nm) cobalt crystallites at the startup of
Fischer-Tropsch synthesis at industrially relevant conversions, is shown in Figure 11 [42].

165
Figure 6. XANES analysis of a series of used Co/Pt/Al2O3 catalyst samples retrieved
from a 100 barrel/day slurry bubble column reactor operated at 220 °C, 2.0 MPa, (H2 +
CO) and conversions between 50% and 70%, feed gas composition = 50 vol.% H2 and
25 vol.% CO, P(H2O)/P(H2) = 1–1.5, P(H2O) = 0.4–0.6 MPa. Reproduced with
permission from [32]. Copyright (2006) Elsevier.

Figure 7. From the run shown in Figure 5, XANES derivative spectra of (left) freshly
reduced/passivated catalysts, which could be fit with Co0 and CoO, and (right) used
catalyst samples, which could only be fitted with Co0, CoO, and CoAl2O4. Adapted with
permission from [44] (Copyright 2003, Elsevier) and [45] (2006).

166
Figure 8. Formation of a minor cobalt aluminate component at 1.0 MPa H2O by
increasing conversion in a continuously stirred tank reactor (CSTR) reactor run at
230 °C, 2.0 MPa, 50 vol.% H2, 25 vol.% CO and 25 vol.% inerts. Reproduced with
permission from [43]. Copyright (2011) Elsevier.

Figure 9. Changes in CO2 and CH4 selectivities as a function of conversion over
0.27%Ru-25%Co/Al2O3 catalyst (Co cluster size of ‫׽‬5 nm by hydrogen
chemisorption/pulse reoxidation) at 220 °C, 1.5 MPa, H2/CO = 2.1 and SV = 0.3–15
NL/gcath (reproduced with permission from [21] Copyright 2011, Springer).

167
Figure 10. (left) Co-feeding of H2O over 0.5%Pt-15%Co/Al2O3 at T = 210 °C, P = 2.0
MPa, H2:CO = 2:1, SV = 8 SL/gcath reveals that irreversible deactivation (and a minor
reversible effect) is observed at 28% H2O addition; (right) XANES analysis of the
used catalyst reveals formation of cobalt aluminate through reaction of the CoO formed
with the support. Reprinted with permission from [40,49,50]. Copyright (2002, 2003,
2004) Elsevier.

Figure 11. Proposed explanation for the deactivation of alumina-supported cobalt
nanoparticles in research catalysts as a function of time on-stream. Adapted with
permission from [42]. Copyright (2013) Elsevier [42]. The figure emphasizes why
commercial catalysts adopt larger crystallite diameters for the purpose of stability.

168
2.3. Regeneration
In recent years, there has been interest in carrying out oxidation-reduction cycles in order to
explore the potential for regenerating metal promoted cobalt catalysts. In an earlier work [16], we
simulated regeneration with 2%Ru-15%Co/alumina catalysts having relatively high Ru loadings.
Part of the aim was to see if any promoter mixing or separation occurred at the atomic level, and
the resulting influence on selectivity. To do so, TPR, EXAFS/XANES, and transmission electron
microscopy with elemental mapping were applied. It was difficult to detect atomic mixing at the
promoter loadings utilized, though elemental mapping showed that the metals were in close
proximity to one another at the scale of nanometers. With reduction-oxidation cycles, TPR profiles
revealed that while the first step of the TPR shifted to slightly higher temperatures, the second peak
(i.e., CoO to Co0) shifted to slightly lower temperatures [16] for these heavily Ru-loaded catalysts.
Westrate et al. [51] compared Pt and Ru promoted Co/alumina catalysts subjected to
oxidation-reduction treatments. Following oxidation treatment both the cobalt and promoter phases
are well mixed and in an oxidized state. This finding is in agreement with our EXAFS results for
a Ag-promoted Co/alumina catalyst [10]. Upon reduction, noble metals form bonds with cobalt
metal, again in agreement with our earlier findings for Pt, Re, and Ru promoted catalysts at low
promoter loadings [42]. During simulated regeneration by oxidation, the promoter separates from
the cobalt phase and is found inside a ring of Co3O4, in “Kirkendall voids” as shown in Figure 12.
Re-reduction of this state leads to a decrease in the promoter concentration at the surface of Co
particles, as observed by XPS.
Figure 12. Pt promoter (oxide form) is separated from Co3O4 and found within
Kirkendall voids within the Co3O4 particle. Reproduced with permission from [51].
Copyright (2013) Elsevier.

A key point regarding regeneration is whether or not the promoter continues to facilitate cobalt
oxide reduction once an oxidation (e.g., carbon burn-off) step has been conducted. The proximity
of the promoter to cobalt is important; although this will likely vary with promoter chemistry (see
earlier comment regarding van Steen’s group’s use of physical mixtures to demonstrate that atomic
contact may not be necessary in all cases [26]). To probe this attribute further, a TPR and XANES

169
investigation was carried out to screen a number of promoter metals [52]. An example comparing
two promoters—Pt and Au—is provided in Figure 13 to demonstrate the methodology. The
oxidation-reduction cycle involved a 4 h calcination in flowing air at 350 °C followed by reduction
for 10 h in hydrogen at 350 °C. It is evident (Figure 13) that Pt improves its effectiveness after
three oxidation-reduction cycles, while Au is no longer effective after the first cycle. In the case of
Au, the TPR profiles move to higher temperatures with increasing reduction-oxidation cycle
number, suggesting that Au is separating from the Co (e.g., by sintering). At the same time,
XANES shows that the oxidation state of Co following oxidation reduction moves toward higher
oxidation state in the activated Au-promoted catalyst. The TPR profiles for Pt-promoted
Co/alumina shift slightly to lower temperature. Corresponding XANES spectra demonstrate that
the cobalt is largely reduced after activation following several simulated regeneration cycles. Thus,
Pt is a more effective promoter for long-term use, although a different regeneration method at
different conditions might be more effective in the case of Au. Both XANES and TPR data
revealed that Re retains its ability to facilitate reduction even after 3 oxidation-reduction cycles.
With Ru, the XANES results indicated that Ru was also effective after 3 ORcycles, although a
slight shift to higher temperature was observed for the CoO to Co0 TPR peak position in the
preliminary study [52].
Figure 13. (Top) TPR profiles after RO cycles and (Bottom) XANES profiles after RO
cycling demonstrate that Pt is more effective at continuing to facilitate reduction after
simulate regeneration. Reproduced with permission from [52]. Copyright (2014) Elsevier.

170
Figure 13. Cont.

2.4. Modeling
2.4.1. Modeling of Site Suppression and Deactivation
Although this manuscript is focused on catalyst structure and its influence on FTS stability, a
brief word should also be made regarding modeling. One puzzling aspect about the water effect
during co-feeding and kinetic investigations is that it can be either positive or negative depending
on the nature of the cobalt catalyst. Figure 10 displays results for a 0.5%Pt promoted 15%Co/Al2O3
catalyst having an average cobalt cluster size of 5.6 nm (i.e., crystallites must be equal to below
this value), and the water effect is negative but exhibits a significant degree of reversibility at levels
below 28 vol.% added H2O. The reversibility was, based on EXAFS/XANES investigations,
suggested to be due to oxidation and re-reduction of small cobalt crystallites. Because oxygen was
bound to the cobalt sites, the behavior could be modeled in terms of a kinetic parameter based on
adsorption inhibition of reactants. On the other hand, Co/silica catalysts having larger cobalt clusters
exhibited a positive effect [53], and positive effects on cobalt catalysts have been suggested to be
due to water increasing the concentration of surface active carbon species (e.g., unsupported Co and
Co/titania [38]) or removing heavy wax from catalyst pores leading to a higher available site
density (e.g., Co/silica with varying pore size [54]). Returning to cobalt/alumina catalysts,
interestingly, by aging the catalyst sufficiently (i.e., the catalyst is significantly deactivated from its
initial condition) [55] or utilizing catalysts with 10+ nm size [56], the deactivation rate becomes
low, and the positive kinetic effect occurring on metallic cobalt particles can be observed. The
main point is that the water effect can be modeled using a simple power law expression with a
water effect parameter, m, and the magnitude and sign (i.e., positive [57] or negative [55,57]) of m
provides valuable information about the structure of the catalyst.

171
r(FT) = kP(CO)aP(H2)b/[1 + m P(H2O)/P(H2)]

(1)

A number of the phenomena associated with the stability and deactivation of cobalt catalysts
have been discussed herein, and these and other aspects of stability (e.g., carbon deposition and
carbide formation) are discussed in the Editor’s book on Catalyst Deactivation [58] and a review
article [59]. At this point in time, a number of important aspects and trends regarding the structure
sensitivity of cobalt catalysts, and especially experimental research catalysts, are known. However,
moving forward, there is a great deal of focus on developing catalysts that make the most efficient
use of cobalt. Co crystallites should be small enough to maximize active site density, but also large
enough that crystallites will be stabilized against reoxidation and sintering. Researchers are also
exploring preparation methods to disperse and adequately position metal clusters in spatially
favorable ways (i.e., as far apart as possible). Some methods include freeze-drying [60], vapor
phase impregnation [61], coating the support with carbon [62], using solvents such as ethylene
glycol [63], optimizing drying temperature or calcination chemistry with dilute NO/N2 [64],
bypassing calcination altogether [53], or locking metal particles onto the support so that they
cannot find one another and undergo agglomeration [65]. Thus, using advanced preparation
methods to obtain spatially uniform distribution of cobalt crystallites is critical, since in
conventional cobalt/alumina catalysts prepared by impregnation, cobalt clusters can be within close
vicinity to one another, or in grapelike clusters, which have been described as “graveyards” for
cobalt active sites during reaction testing [66]. Thus, thinking toward the future, it will be of
increasing importance to model deactivation mechanisms and quantify how much each mechanism
contributes to overall deactivation of the catalyst. Robust models which address the chemistry of
promoters could enable the performance of new research catalysts to be compared to current
commercial catalyst formulations. For example, a recently published forward thinking article by
Argyle et al. [67] addresses modeling of the contributions of several deactivation mechanisms to
overall deactivation rate in Co catalyzed FTS.
2.4.2. Computational Methods Based on First Principles
Related to promoters of FTS catalysts, computational methods based on first principles have
been useful in describing the location of promoter with respect to cobalt, for providing insight into
the requirements for cobalt oxidation by H2O, and for determining how the promoter may influence
carbon deposition.
Computational methods are being utilized to elucidate the preferential location for the occupancy
of the promoter with respect to the cobalt atoms that make up the cluster. For example, a combined
study making using of Low Energy Ion Scattering (i.e., on a 1%Re-12%Co/Al2O3 catalyst) and
computational DFT modeling (i.e., on a Co13Re cluster) determined that there is a preference of Re
promoter to occupy subsurface sites, where it coordinates with a maximum number of cobalt
atoms [68]. This is in agreement with the results of some EXAFS investigations, where direct
Re-Co atomic contact has been observed [19,20].
Molecular modeling has also been conducted to examine surface oxidation of larger cobalt
particles (e.g., as utilized in commercial catalysts) by H2O, and the pathway was ruled out as a

172
significant chemical transformation mechanism for deactivating sites under commercial FTS
conditions [31].
Computational methods have also been used extensively to gain insight into the role of carbon
deposits in catalyst deactivation. For example [69], HR-TEM and computational DFT modeling
were applied to elucidate the preferred occupancies of carbon over different cobalt surfaces. The
stability of various carbon species under reaction conditions was evaluated. Extended graphene
islands and a surface carbide were found to be 99 and 79 kJ/mol more stable than surface CH2
groups. Both carbon phases were suggested to initiate and grow from step sites. Saib et al. [31]
have recently reviewed carbon formation in detail on cobalt FTS catalysts, including the
application of computational methods. They also analyzed used catalysts from a commercial slurry
bubble column reactor and, following wax extraction by THF, carried out temperature programmed
hydrogenation and oxidation measurements to characterize the carbons. The least reactive species
toward hydrogenation, which reacted at 430 °C, was assigned to polymeric carbons. A model [70]
showing the location of small carbon oligomers of the fcc Co(111) surface, the precursors of
polymeric carbon, was described. Note that polymeric carbon was found on both cobalt and the
support. The authors [31] also reviewed the role of subsurface carbon [69,71], where theoretical
modeling has indicated that subsurface carbon hinders CO adsorption and dissociation processes on
associated Co atoms, and the requirements under which carbon induces clock surface
reconstruction [72]. This, in turn, may cause deactivation of sites via shape changes or, on the other
hand, induce the formation of active sites (e.g., B5 sites [31]: 3-fold sites that more easily
dissociate CO; or triangular nanoscale islands having step edges similar to C7 sites [73], as
observed by scanning tunneling microscopy). The restructuring of cobalt by strong CO
chemisorption (i.e., roughening—leading to more active sites) was described by Schulz et al. [74]
as to be in competition with sintering.
Some computational studies have focused on defining how promoters of cobalt catalysts may
impact carbon formation. Recently, the mechanisms for carbon compound formation on
unpromoted and Pt or Ru promoted Co surfaces were investigated [75]. The activation energies for
carbon-carbon and carbon-carbon-carbon coupling reactions were found to be larger on Pt or Ru
promoted Co surfaces relative to unpromoted Co surface. The results suggest that carbon formation
and thus, carbon compound (e.g., polymeric carbon formation, may be inhibited by the presence of
the promoters. The authors also found that the promoters did not change the activation energy of
C diffusion to the subsurface.
3. Experimental Section
Typical catalyst preparation method: the support used was Sasol Catalox-150 γ-Al2O3. It was
first calcined at 400 °C in a muffle furnace for 4 h. A slurry impregnation method was performed,
whereby the ratio of the volume of loading solution used to the weight of alumina was 1:1, such
that approximately 2.5 times the pore volume of solution was used to prepare the catalyst in two
steps [7]. Due to the solubility limit of cobalt nitrate, multiple impregnation steps were used. After
each impregnation step, the catalyst was dried under vacuum in a rotary evaporator from 80 to
100 °C. Promoter precursors used were tetraammine palladium (II) nitrate, tetraammine platinum

173
(II) nitrate, rhenium oxide (Re2O7), ruthenium nitrosyl nitrate, silver nitrate, HAuCl4, and copper
nitrate. The promoters were added dropwise to achieve incipient wetness impregnation. After final
drying at 80–100 °C, the final catalysts were calcined at 350 °C under flowing air for 4 h.
Typical CSTR reaction test: the catalyst (15 g) was ground and sieved to 170–325 mesh before
loading into a fixed-bed reactor for 10–15 h of ex situ reduction at 350 °C and atmospheric
pressure using a gas mixture of H2/He with a molar ratio of 1:3. The reduced catalyst was then
transferred to a 1-L continuously stirred tank reactor (CSTR) which was previously charged with
315 g of melted Polywax 3000, under the protection of a N2 inert gas. The transferred catalyst was
further reduced in situ at 230 °C at atmospheric pressure using pure hydrogen for another 10 h
before starting the FT reaction. In this study, the FT conditions were 220 °C, 1.5–2.2 MPa,
H2/CO = 2.0–2.1. The space velocity varied between 1.0 and 16 NL/g-cat/h. in order to give about
50% CO conversion in different tests. This allowed us to fairly compare the differences in
hydrocarbon selectivity data resulting from the promoter effect.
4. Conclusions
There are a number of stability issues that must be considered when selecting metal reduction
promoters for use in Fischer-Tropsch synthesis catalysts. If tiny cobalt crystallites (<2–4.4 nm) are
formed by facilitating the reduction of cobalt oxides that are strongly interacting with the support,
they may undergo reoxidation at the onset of FTS at high conversion. Any cobalt oxide either left
unreduced or formed from reoxidation of tiny cobalt crystallites can participate in a complex
sintering mechanism that involves agglomeration of cobalt oxides, re-reduction, and sintering of
the metal. Promoters can also exacerbate sintering if the cobalt metal clusters formed as a result of
the promoting effect are in close proximity to other cobalt particles on the surface. Not all metals that
facilitate cobalt reduction promote activity on a per gram catalyst basis; some will poison the surface
(e.g., Cu). A poor choice of promoter (or poor choice in loading) can also lead to excessive
hydrogenation activity and raise the light gas selectivity (e.g., Pd or Cu; Au at high loading).
Furthermore, certain metals (e.g., Au) that promote cobalt oxide reduction can separate from cobalt
during oxidation-reduction (regeneration) cycles. Therefore, they may not be effective for long-term
use, or they may require non-standard regeneration treatments. Computational studies suggest
that certain promoters (e.g., Pt or Ru) may hinder deactivation by carbon by increasing the
energy barrier for carbon-carbon coupling reactions, while subsurface C formation was not found
to be affected.
Acknowledgments
The authors would like to acknowledge the support of the Commonwealth of Kentucky. We
would also like to thank the Editor for helpful comments.
Conflicts of Interest
The authors declare no conflict of interest.

174
References
1.
2.
3.
4.

5.
6.
7.

8.
9.

10.

11.

12.

13.
14.

Espinoza, R.L.; Visagie, J.L.; van Berge, P.J.; Bolder, F.H. Fischer-Tropsch catalysts
containing iron and cobalt. US Patent No. 5,733,839, March 1998.
Van Berge, P.J.; Barradas, S.; van de Loosdrecht, J.; Visagie, J.L. Advances in the cobalt
catalyzed Fischer-Tropsch synthesis. Erdöl Erdgas Kohle 2001, 117, 138–142.
Wang, W.-J.; Chen, Y.-W. Influence of metal loading on the reducibility and hydrogenation
activity of cobalt/alumina catalysts. Appl. Catal. 1991, 77, 223–233.
Jacobs, G.; Ji, Y.; Davis, B.H.; Cronauer, D.C.; Kropf, A.J.; Marshall, C.L. Fischer-Tropsch
synthesis: Temperature programmed EXAFS/XANES investigation of the influence of support
type, cobalt loading, and noble metal promoter addition to the reduction behavior of cobalt
oxide particles. Appl. Catal. A 2007, 333, 177–191.
Reuel, R.C.; Bartholomew, C.H. The stoichiometries of H2 and CO adsorptions on cobalt:
Effects of support and preparation. J. Catal. 1984, 85, 63–77.
Vada, S.; Hoff, A.; Adnanes, E.; Schanke, D.; Holmen, A. Fischer-Tropsch synthesis on
supported cobalt catalysts promoted by platinum and rhenium. Top. Catal. 1995, 2, 155–162.
Jacobs, G.; Das, T.K.; Zhang, Y.-Q.; Li, J.; Racoillet, G.; Davis, B.H. Fischer-Tropsch
synthesis: support, loading, and promoter effects on the reducibility of cobalt catalysts.
Appl. Catal. A 2002, 233, 263–281.
Farrauto, R.J.; Bartholomew, C.H. Fundamentals of Industrial Catalytic Processes;
John Wiley & Sons: Chichester, UK, 2003; p. 700.
Cook, K.M.; Poudyal, S.; Miller, J.T.; Bartholomew, C.H.; Hecker, W.C. Reducibility of
alumina-supported cobalt Fischer-Tropsch catalysts: Effects of noble metal type, distribution,
retention, chemical state, bonding, and influence on cobalt crystallite size. Appl. Catal. A
2012, 449, 69–80.
Jacobs, G.; Ribeiro, M.C.; Ma, W.; Ji, Y.; Khalid, S.; Sumodjo, P.T.A.; Davis, B.H. Group 11
(Cu, Ag, Au) promotion of 15%Co/Al2O3 Fischer-Tropsch synthesis catalysts. Appl. Catal. A
2009, 361, 137–151.
Aldossary, M.A.; Sharma, P.; Ojeda, M.; Gupta, M.; Fierro, J.L.; Spivey, J.J. Effect of
different Cu loading on Fe-Mg catalyst for Fischer-Tropsch synthesis. In Proceedings of ACS
National Meeting & Exposition, Philadelphia, PA, USA, August 19–23, 2012. Curran
Associates, Inc.: Red Hook, NY, USA, 2012; Volume 57, p. 944.
Li, S.; Li, A.; Krishnamoorthy, S.; Iglesia, E. Effects of Zn, Cu, and K promoters on the
structure and on the reduction, carburization, and catalytic behavior of iron-based
Fischer-Tropsch synthesis catalysts. Catal. Lett. 2001, 77, 197–205.
Cesar, D.V.; Perez, C.A. Quantitative XPS analysis of bimetallic Cu-Co catalysts. Phys. Status
Solidi A 2001, 187, 321–326.
Todic, B.; Bhatelia, T.; Froment, G.F.; Ma, W.; Jacobs, G.; Davis, B.H.; Bukur, D.B. Kinetic
model of Fischer-Tropsch synthesis in a slurry reactor on Co/Re/Al2O3 catalyst. Ind. Eng.
Chem. Res. 2013, 52, 669–679.

175
15. Jacobs, G.; Ma, W.; Gao, P.; Todic, B.; Bhatelia, T.; Bukur, D.B.; Khalid, S.; Davis, B.H.
Fischer-Tropsch synthesis: differences observed in local atomic structure and selectivity with
Pd compared to typical promoters (Pt, Re, Ru) of Co/Al2O3 catalysts. (Special Issue in honor
of the late Prof. Laszlo Guczi 1932–2012). Top. Catal. 2012, 55, 811–817.
16. Jacobs, G.; Sarkar, A.; Ji, Y.; Luo, M.-S.; Dozier, A.; Davis, B.H. Fischer-Tropsch synthesis:
assessment of the ripening of cobalt clusters and mixing between Co and Ru promoter via
oxidation-reduction cycles over lower Co-loaded Ru-Co/Al2O3 catalysts. Ind. Eng. Chem. Res.
2008, 47, 672–680.
17. Borg, O.; Hammer, N.; Eri, S.; Lindvag, O.A.; Myrstad, R.; Blekkan, E.A.; Ronning, M.;
Rytter, E.; Holmen, A. Fischer-Tropsch synthesis over un-promoted and Re-promoted
gamma-Al2O3 supported cobalt catalysts with different pore sizes. Catal. Today 2009, 142,
70–77.
18. Bazin, D.; Borko, L.; Koppany, Zs.; Kovacs, I.; Stefler, G.; Sajo, L.I.; Schay, Z.; Guczi, L.
Re-Co/NaY and Re-Co/Al2O3 bimetallic catalysts: In situ EXAFS and catalytic activity.
Catal. Lett. 2002, 84, 169–182.
19. Jacobs, G.; Chaney, J.A.; Patterson, P.M.; Das, T.K.; Davis, B.H. Fischer-Tropsch synthesis:
Study of the promotion of Re on the reduction property of Co/Al2O3 catalysts by in situ
EXAFS/XANES of Co K and Re LIII edges and XPS. Appl. Catal. A 2004, 264,
203–212.
20. Ronning, M.; Nicholson, D.G.; Holmen, A. In situ EXAFS study of the bimetallic interaction
in a rhenium-promoted alumina-supported cobalt Fischer-Tropsch catalyst. Catal. Lett. 2001,
72, 141–146.
21. Ma, W.; Jacobs, G.; Ji, Y.; Bhatelia, T.; Bukur, D.B.; Khalid, S.; Davis, B.H. Fischer-Tropsch
synthesis: Influence of CO conversion on selectivities, H2/CO usage ratios, and catalyst
stability for a Ru promoted Co/Al2O3 catalyst using a slurry phase reactor. Top. Catal. 2011,
54, 757–767.
22. Guczi, L.; Bazin, D.; Kovacs, I.; Borko, L.; Schay, Z.; Lynch, J.; Parent, P.; Lafon, C.; Stefler,
G.; Koppany, Z.; Sajo, I. Structure of Pt-Co/Al2O3 and Pt-Co/NaY bimetallic catalysts:
Characterization by in situ EXAFS, TPR, XPS and by activity in CO (Carbon Monoxide)
Hydrogenation. Top. Catal. 2002, 20, 129–139.
23. Jacobs, G.; Chaney, J.A.; Patterson, P.M.; Das, T.K.; Maillot, J.C.; Davis, B.H.
Fischer-Tropsch synthesis: Study of the promotion of Pt on the reduction property of Co/Al2O3
catalysts by in situ EXAFS of Co K and Pt LIII edges and XPS. J. Synch. Rad. 2004, 11,
414–422.
24. Sadeqzadeh, M.; Karaca, H.; Safonova, O.V.; Fongarland, P.; Chambrey, S.; Roussel, P.;
Griboval-Constant, A.; Lacroix, M.; Curulla-Ferré, D.; Luck, F.; et al. Identification of the
active species in the working alumina-supported cobalt catalyst under various conditions of
Fischer–Tropsch synthesis. Catal. Today 2011, 164, 62–67.
25. Iglesia, E.; Soled, S.L.; Fiato, R.A.; Via, G.H. Bimetallic synergy in cobalt-ruthenium
Fischer-Tropsch synthesis catalysts. J. Catal. 1993, 143, 345–368.

176
26. Chonco, Z.H.; Nabaho, D.; Claeys, M.; van Steen, E. The role of reduction promoters in
Fischer-Tropsch catalysts for the production of liquid fuels. In Proceedings of 23rd Meeting of
the North American Catalysis Society, 2–7 June 2013, Louisville, KY, USA.
27. Ionkina, O.; Subramanian, M.A.; Chao, W.; Makar, K.M.; Manzer, L.E. Fischer-Tropsch
processes and catalysts with promoters. US Patent 20020010221A1, January 2002.
28. Jermwongratanachai, T.; Jacobs, G.; Ma, W.; Shafer, W.D.; Gnanamani, M.K.; Gao, P.;
Kitiyanan, B.; Davis, B.H.; Klettlinger, J.L.S.; Yen, C.H.; et al. Fischer-Tropsch synthesis:
Comparisons between Pt and Ag promoted Co/Al2O3 catalysts for reducibility, local atomic
structure, catalytic activity, and oxidation-reduction (OR) cycles. Appl. Catal. 2013, 464–465,
165–180.
29. Redjala, T.; Remita, H.; Apostolescu, G.; Mostafavi, M.; Thomazeau, C.; Uzio, D. Bimetallic
Au-Pd and Ag-Pd clusters synthesized by gamma or electron beam radiolysis and study of the
reactivity/structure relationships in the selective hydrogenation of but-1,3-diene. Oil Gas
Sci. Technol. 2006, 61, 789–797.
30. Van Steen, E.; Claeys, M.; Dry, M.E.; van de Loosdrecht, J.; Viljoen, E.L.; Visagie, J.L.
Stability of nanocrystals: thermodynamic analysis of oxidation and re-reduction of cobalt in
water/hydrogen mixtures. J. Phys. Chem. B 2005, 109, 3575–3577.
31. Saib, A.M.; Moodley, D.J.; Ciobica, I.M.; Hauman, M.M.; Sigwebela, B.H.; Weststrate, C.J.;
Niemantsverdriet, J.W.; van de Loosdrecht, J. Fundamental understanding of deactivation and
regeneration of cobalt Fischer-Tropsch synthesis catalysts. Catal. Today 2010, 154, 271–282.
32. Saib, A.M.; Borgna, A.; van de Loosdrecht, J.; van Berge, P.J.; Niemantsverdriet, J.W.
XANES study of the susceptibility of nano-sized cobalt crystallites to oxidation during
realistic Fischer-Tropsch synthesis. Appl. Catal. 2006, 312, 12–19.
33. Bezemer G.L, Bitter J.H.; Kuipers H.P.C.E.; Oosterbeek H.; Holewijn J.E.; Xu X.D.;
Kapteijn, F.; van Dillen, A.J.; de Jong, K.P. Cobalt particle size effects in the Fischer-Tropsch
reaction studied with carbon nanofiber supported catalysts. J. Am. Chem. Soc. 2006, 128,
3956–3964.
34. Borg, Ø.; Dietzel, P.D.C.; Spjelkavik, A.I.; Tveten, E.Z.; Walmsley, J.C.; Diplas, S.; Eri, S.;
Holmen, A.; Rytter, E. Fischer-Tropsch synthesis: cobalt particle size and support effects on
intrinsic activity and product distribution. J. Catal. 2008, 259, 161–164.
35. Azzam, K.; Jacobs, G.; Ma, W.; Davis, B.H. Effect of cobalt particle size on the catalyst
intrinsic activity for Fischer-Tropsch synthesis. Catal. Lett. 2014, 144, 389–394.
36. Jermwongratanachai, T.; Jacobs, G.; Shafer, W.D.; Ma, W.; Pendyala, V.R.R.; Davis, B.H.;
Kitiyanan, B.; Khalid, S.; Cronauer, D.C.; Kropf, A.J.; Marshall, C.L. Fischer-Tropsch
synthesis: Oxidation of a fraction of cobalt crystallites in research catalysts at the onset of FT
at partial pressures mimicking 50% CO conversion. Top. Catal. in press.
37. Fischer, N.; Clapham, B.; Feltes, T.E.; van Steen, E.; Claeys, M. The reoxidation of cobalt
Fischer-Tropsch catalysts. In Proceedings of Syngas 2012 Convention, Cape Town, South
Africa, 1–4 April 2012.
38. Bertole, C.J.; Mims, C.A.; Kiss, G. The effect of water on the cobalt-catalyzed
Fischer-Tropsch synthesis. J. Catal. 2002, 210, 84–96.

177
39. Sadeqzadeh, M.; Hong, J.; Fongarland, P.; Curulla-Ferre, D.; Luck, F.; Bousquet, J.;
Schweich, D.; Khodakov, A.Y. Mechanistic modeling of cobalt based catalyst sintering in a
fixed bed reactor under different conditions of Fischer-Tropsch synthesis. Ind. Eng. Chem. Res.
2012, 51, 11955–11964.
40. Jacobs, G.; Das, T.K.; Patterson, P.M.; Luo, M.; Conner, W.A.; Davis, B.H. Fischer-Tropsch
synthesis: Effect of water on Co/Al2O3 catalysts and XAFS characterization of reoxidation
phenomena. Appl. Catal. A 2004, 270, 65–76.
41. Soled, S.; Kliewer, C.; Kiss, G.; Baumgartner, J. Reversible and irreversible changes in Co
Fischer-Tropsch catalysts during synthesis. In Proceedings of 21st Meeting of the North
American Catalysis Society, San Francisco, CA, USA, 7–12 June 2009.
42. Jacobs, G.; Ma, W.; Gao, P.; Todic, B.; Bhatelia, T.; Bukur, D.B.; Davis, B.H. The application
of synchrotron methods in characterizing iron and cobalt Fischer-Tropsch synthesis catalysts.
Catal. Today 2013, 214, 100–139.
43. Moodley, D.J.; Saib, A.M.; van de Loosdrecht, J.; Welker-Nieuwoudt, C.A.; Sigwebela, B.H.;
Niemantsverdriet, J.W. The impact of cobalt aluminate formation on the deactivation of
cobalt-based Fischer-Tropsch synthesis catalysts. Catal. Today 2011, 171, 192–200.
44. Das, T.K.; Jacobs, G.; Patterson, P.M.; Conner, W.A.; Davis, B.H. Fischer-Tropsch synthesis:
Characterization and catalytic properties of rhenium promoted cobalt alumina catalysts. Fuel
2003, 82, 805–815.
45. Jacobs, G.; Sarkar, A.; Ji, Y.; Patterson, P.M.; Das, T.K.; Luo, M.; Davis, B.H. Fischer-Tropsch
synthesis: characterization of interactions between reduction promoters and Co for
Co/Al2O3–based GTL catalysts. In Proceedings of AIChE Annual Meeting, San Francisco, CA,
USA, 12–17 November 2006.
46. Karaca, H.; Hong, J.; Fongarland, P.; Roussel, P.; Griboval-Constant, A.; Lacroix, M.;
Hortmann, K.; Safonova, O.V.; Khodakov, A.Y. In situ XRD investigation of the evolution of
alumina-supported cobalt catalysts under realistic conditions of Fischer-Tropsch synthesis.
Chem. Commun. 2010, 46, 788–790.
47. Jacobs, G.; Patterson, P.M.; Zhang, Y.-Q.; Das, T.K.; Li, J.; Davis, B.H. Fischer-Tropsch
synthesis: Deactivation of noble metal-promoted Co/Al2O3 catalysts. Appl. Catal. A 2002, 233,
215–226.
48. Kim, C.J. Water addition for increased CO/H2 hydrocarbon activity over catalysts comprising
* e There is virtually no difference. carried out for thousands of hours, the time ranges
specified are very close to each othecobalt, ruthenium, and mixtures thereof which may
include a promoter metal. U.S. Patent 5,227,407, July 1993.
49. Li, J.; Zhan, X.; Zhang, Y.-Q.; Jacobs, G.; Das, T.K.; Davis, B.H. Fischer-Tropsch synthesis:
Effect of water on the deactivation of Pt promoted Co/Al2O3 catalysts. Appl. Catal. A 2002,
228, 203–212.
50. Jacobs, G.; Das, T.K.; Patterson, P.M.; Li, J.; Sanchez, L.; Davis, B.H. Fischer-Tropsch
synthesis: XAFS studies of the effect of water on a Pt-promoted Co/Al2O3 catalyst.
Appl. Catal. A 2003, 247, 335–343.

178
51. Weststrate, C.J.; Saib, A.M.; Niemantsverdriet, J.W. Promoter segregation in Pt and Ru
promoted cobalt model catalysts during oxidation-reduction treatments. Catal. Today 2013,
215, 2–7.
52. Jermwongratanachai, T.; Jacobs, G.; Shafer, W.D.; Pendyala, V.R.R.; Ma, W.; Gnanamani,
M.K.; Hopps, S.; Thomas, G.A.; Kitiyanan, B.; Khalid, S.; et al. Fischer-Tropsch synthesis:
TPR and XANES analysis of the impact of oxidation-reduction (OR) cycles on the reducibility
of Co/alumina catalysts with different promoters (Pt, Ru, Re, Ag, Au, Rh, Ir). Catal. Today
2014, in press.
53. Li, J.; Jacobs, G.; Das, T.K.; Zhang, Y.-Q.; Davis, B.H. Fischer-Tropsch synthesis: Effect of
water on the catalytic properties of a Co/SiO2 catalyst. Appl. Catal. A 2002, 236, 67–76.
54. Dalai, A.K.; Das, T.K.; Chaudhari, K.V.; Jacobs, G.; and Davis, B.H. Fischer-Tropsch
synthesis: Water effects on Co supported on wide and narrow-pore silica. Appl. Catal. A 2005,
289, 135–142.
55. Ma, W.; Jacobs, G.; Sparks, D.E.; Spicer, R.L.; Davis, B.H.; Klettlinger, J.L.S.; Yen, C.H.
Fischer-Tropsch synthesis: Kinetics and water effect study over 25%Co/Al2O3 catalysts.
Catal. Today 2014, in press.
56. Logdberg, S.; Boutonnet, M.; Walmsley, J.C.; Jaras, S.; Holmen, A.; Blekkan, E.A. Effect of
water on the space-time yield of different supported cobalt catalysts during Fischer-Tropsch
synthesis. Appl. Catal. A 2011, 393, 109–121.
57. Ma, W.; Jacobs, G.; Sparks, D.E.; Gnanamani, M.K.; Pendyala, V.R.R.; Yen, C.H.;
Klettlinger, J.L.S.; Tomsik, T.M.; Davis, B.H. Fischer-Tropsch synthesis: support and cobalt
cluster size effects on kinetics over Co/Al2O3 and Co/SiO2 catalysts. Fuel 2011, 90, 756–765.
58. Studies in Surface Science and Catalysis. Catalyst Deactivation 1991: Proceedings of the 5th
International Symposium; Bartholomew, C.H., Butt, J.B., Eds.; Elsevier: Amsterdam, The
Netherlands, 1991; Volume 68.
59. Bartholomew, C.H. Mechanisms of catalyst deactivation. Appl. Catal. A 2001, 212, 17–60.
60. Eggenhuisen, T.M.; Munnik, P.; Talsma, H.; de Jongh, P.E.; de Jong, K.P. Freeze-drying for
controlled nanoparticle distribution in Co/SiO2 Fischer-Tropsch catalyst. J. Catal. 2013, 297,
306–313.
61. Graham, U.M.; Jacobs, G.; Gnanamani, M.; Lipka, S.; Shafer, W.D.; Swartz, C.;
Jermwongratanachai, T.; Chen, R.; Rogers, F.; Davis, B.H. Fischer Tropsch Synthesis: High
Oxygenate-Selectivity of Cobalt Catalysts supported on Hydrothermal Carbons. ACS Catal.
submitted for publication, 2014.
62. van de Loosdrecht, J.; Datt, M.; Visagie, J.L. Carbon coated supports for cobalt based
Fischer-Tropsch catalysts. Top. Catal. 2014, in press.
63. Rane, S.; Borg, O.; Yang, J.; Rytter, E.; Holmen, A. Effect of alumina phases on hydrocarbon
selectivity in Fischer-Tropsch synthesis. Appl. Catal. A 2010, 388, 160–167.
64. Sietsma, J.R.A.; den Breejen, J.P.; de Jongh, P.E.; van Dillen, J.; Bitter, J.H.; de Jong, K.P.
Highly active cobalt-on-silica catalysts for the Fischer-Tropsch synthesis obtained via a novel
calcination. Stud. Surf. Sci. Catal. 2007, 167, 85–90.

179
65. Lu, J.; Elam, J.W.; Stair, P.C. Synthesis and stabilization of supported metal catalysts by
atomic layer deposition. Accts. Chem. Res. 2013, 46, 1806–1815.
66. Soled, S.L.; Kiss, G.; Kliewer, C.; Baumgartner, J.; El-Malki, E.-M. Learnings from Co
Fischer-Tropsch catalyst studies. Abstracts of Papers, ENFL-412, 245th ACS National
Meeting & Exposition, New Orleans, LA, USA, 7–11 April 2013.
67. Argyle, M.D.; Frost, T.S.; Bartholomew, C.H. Cobalt Fischer Tropsch catalyst deactivation
modeled using generalized power law expressions. Top. Catal. 2014, in press.
68. Bakken, V.; Bergene, E.; Rytter, E.; Swang, O. Bimetallic cobalt/rhenium systems: Preferred
position of rhenium through an interdisciplinary approach. Catal. Lett. 2010, 135, 21–25.
69. Tan, K.F.; Xu, J.; Chang, J.; Borgna, A.; Saeys, M. Carbon deposition on Co catalysts during
Fischer-Tropsch synthesis: A computational and experimental study. J. Catal. 2010, 274,
121–129.
70. Swart, J.C.W.; Ciobica, I.M.; van Santen, R.A.; van Steen, E. Intermediates in the formation
of graphitic carbon on a flat FCC-Co(111) surface. J. Phys. Chem. C 2008, 112, 12899–12904.
71. Burghgraef, H. A quantum chemical study of CH and CC bond activation on transition metals.
PhD Thesis, Eindhoven University of Technology: Eindhoven, The Netherlands, June 1995.
72. Ciobica, I.M.; van Santen, R.A.; van Berge, P.J.;. van de Loosdrecht, J. Adsorbate induced
reconstruction of cobalt surfaces. Surf. Sci. 2008, 602, 17–27.
73. Wilson, J.; de Groot, C. Atomic-scale restructuring in high-pressure catalysis. J. Phys. Chem.
1995, 99, 7860–7866.
74. Schulz, H.; Nie, Z.; Ousmanov, F. Construction of the Fischer–Tropsch regime with cobalt
catalysts. Catal. Today 2001, 71, 351–360.
75. Balakrishnan, N.; Joseph, B.; Bhethanabotla, V.R. Effect of Pt and Ru promoters on
deactivation of Co catalysts by C deposition during Fischer-Tropsch synthesis: A DFT study.
Appl. Catal. A 2013, 462–463, 107–115.

180

Deactivation of Pd Catalysts by Water during Low
Temperature Methane Oxidation Relevant to Natural Gas
Vehicle Converters
Rahman Gholami, Mina Alyani and Kevin J. Smith
Abstract: Effects of H2O on the activity and deactivation of Pd catalysts used for the oxidation of
unburned CH4 present in the exhaust gas of natural-gas vehicles (NGVs) are reviewed. CH4
oxidation in a catalytic converter is limited by low exhaust gas temperatures (500–550 °C) and low
concentrations of CH4 (400–1500 ppmv) that must be reacted in the presence of large quantities of
H2O (10–15%) and CO2 (15%), under transient exhaust gas flows, temperatures, and compositions.
Although Pd catalysts have the highest known activity for CH4 oxidation, water-induced sintering
and reaction inhibition by H2O deactivate these catalysts. Recent studies have shown the reversible
inhibition by H2O adsorption causes a significant drop in catalyst activity at lower reaction
temperatures (below 450 °C), but its effect decreases (water adsorption becomes more reversible)
with increasing reaction temperature. Thus above 500 °C H2O inhibition is negligible, while Pd
sintering and occlusion by support species become more important. H2O inhibition is postulated to
occur by either formation of relatively stable Pd(OH)2 and/or partial blocking by OH groups of the
O exchange between the support and Pd active sites thereby suppressing catalytic activity.
Evidence from FTIR and isotopic labeling favors the latter route. Pd catalyst design, including
incorporation of a second noble metal (Rh or Pt) and supports high O mobility (e.g., CeO2) are
known to improve catalyst activity and stability. Kinetic studies of CH4 oxidation at conditions
relevant to natural gas vehicles have quantified the thermodynamics and kinetics of competitive
H2O adsorption and Pd(OH)2 formation, but none have addressed effects of H2O on O mobility.
Reprinted from Catalysts. Cite as: Gholami, R.; Alyani, M.; Smith, K.J. Deactivation of Pd
Catalysts by Water during Low Temperature Methane Oxidation Relevant to Natural Gas Vehicle
Converters. Catalysts 2015, 5, 561-594.
1. Introduction
Natural gas, an abundant energy resource with worldwide proven reserves of over
204.7 trillion m3 [1], is used primarily for electricity generation and heating. The composition of natural
gas (NG) is highly variable, but CH4 typically accounts for 80–90% of the components of NG. CH4 has
the highest H/C ratio among all hydrocarbon fuels and during combustion, generates the lowest amount
of CO2 per unit of energy. The amount of SO2 generated during NG combustion is also relatively low
because the S content of NG is significantly lower than that of gasoline or diesel fuels. These
environmental benefits, together with a relatively low cost of NG, have resulted in an increased interest
in its use as a transportation fuel. Currently there are >16 million natural gas vehicles (NGVs) in
operation around the world, and their numbers are growing at about 20% annually [2]. However, a
significant concern for the wide-spread implementation of NG as a fuel for combustion engines is that

181
unburned CH4, expelled in the engine exhaust, is a significant greenhouse gas with potency more than
25xs that of CO2.
The transportation sector is a major contributor to air pollution through the combustion of
gasoline and diesel fuels, accounting for ~77% of CO emissions, ~47% of hydrocarbon emissions
and ~60% of NOx emissions in the USA [3]. The exhaust gas of a conventional gasoline powered
spark-ignition internal combustion engine (SI-ICE) consists mostly of N2 (70–75%), CO2
(11–13%) and water (10-12%) with about 1–2% of pollutants, specifically unburned hydrocarbons,
CO and NOx [4,5]. The pollutants must be removed before the exhaust gas is emitted to the
atmosphere so as to meet increasingly stringent worldwide emission standards. The pollutants
present in the engine exhaust are dependent on the engine air/fuel (A/F) ratio. For example, if the
A/F ratio is above the stoichiometric value for complete combustion (A/F = 14.6), the
concentration of reducing agents (hydrocarbons and CO) in the exhaust gas decreases whereas the
concentration of oxidizing agents (O2 and NOx) increases. Consequently, several different
strategies have been developed to control engine emissions, depending on the operating conditions
and the target emission levels [5]. Typically, a gasoline engine management system controls the
A/F ratio or the exhaust gas composition (using an oxygen sensor connected to a secondary air
supply) near the stoichiometric value. A single three-way catalyst (TWC) bed, placed in the
exhaust gas flow, ensures simultaneous oxidation of the CO and hydrocarbons and the reduction of
the NOx. Alternatively, dual-bed systems combine a NOx reduction catalyst bed with a separate
oxidation catalyst and secondary air to remove the CO and hydrocarbons. Under lean-burn
conditions a gasoline engine may operate with sufficiently high A/F ratios so as to obtain a
significant reduction in CO and NOx emissions and improved fuel efficiency. The function of the
catalyst in this case is limited to the oxidation of mainly hydrocarbons, while the NOx emissions are
captured using a NOx trap followed by desorption and reduction in a TWC during an occasional
near stoichiometric excursion of the engine. Although lean-burn engines improve fuel efficiency,
the exhaust gas temperature is significantly lower than from conventional gasoline powered
engines, and consequently, catalysts with high oxidation activity at relatively low temperatures are
needed for this application [5].
Modern TWC converters used in gasoline ICEs contain Pt, Rh and Pd, dispersed on a washcoat
applied to a cordierite ceramic monolith or metal monolith [3,5]. The monolith usually has a
honeycomb structure with 1 mm square channels to accommodate the high gas throughputs from
the exhaust with minimal pressure drop. The washcoat, a mix of several metal oxides (Al2O3,
CeO2, ZrO2), is applied to increase the metal support surface area (Al2O3), to improve thermal
stability (ZrO2) and to provide enhanced oxygen storage capacity (CeO2) that widens the operating
range for optimal oxidation and reduction by the catalyst. The metal composition of the converter
varies with application but typically contains 5–20:1 of Pt:Rh with a total metal loading of
0.9–2.2 g Lí1. Pd may be used to replace all or part of the Pt for cost savings [5].
Exhaust gas emissions from NGVs are difficult to control because low concentrations of CH4
(400–1500 ppmv) must be oxidized in the presence of high concentrations of H2O (10-15 vol.%)
and CO2 (15 vol.%) at relatively low exhaust gas temperatures (450–550 °C). The greater strength
of the C-H bond in CH4 (450 kJ/mol) relative to other hydrocarbons [6] implies that catalysts with

182
high CH4 oxidation activity must be used. NGVs operate near the stoichiometric point or under
lean-burn conditions [7,8]. Stoichiometric NGV engines are primarily used in light-duty passenger
cars, whereas lean-burn engines are more common in heavy-duty vehicles such as buses. Over the
past ~20 years, conventional converter technologies have been adapted for NGVs using Pd
catalysts (which have the highest activity for CH4 oxidation [7,9,10]) to adequately reduce (by
50–60%) the CH4 content in NGV exhausts at <500 °C in the presence of high H2O concentrations.
Commercial catalysts for SI-NG engines also typically incorporate a CeO2/ZrO2 solid solution for
high O2 adsorption capacity, which serves to buffer O2 concentration during the rapid switching
between slightly oxidizing and reducing conditions close to a stoichiometric mixture (e.g., [7,11]).
Several papers and reviews have assessed the activity and deactivation of Pd catalysts for CH4
oxidation, supported on Al2O3, SiO2, ZrO2, CeO2, and zeolites; promoted with noble metals, e.g., Pt
and Rh, and with transition metal oxides, e.g., oxides of Co, Ni, and Sn [6,7,10–20]. Studies have
largely focused on CH4 oxidation on supported Pd catalysts containing 0.5 to 5% Pd (typical Pd
loadings in commercial SI-NG monolithic coated catalysts are about 3–7 g.Lí1, equivalent to about
1.5–4 wt.% loading in a monolith washcoat) at temperatures ranging from 450 to 600 °C and at
CH4 concentrations of 0.04 to 1 vol.% (0.04 to 0.15 vol.% for commercially representative tests).
High activity for CH4 oxidation appears to be favored by Pd loadings of 3–5 g Lí1 and dispersions
lower than about 0.12–0.15 [7]. Pd-O sites associated with Pd/CeO2 surfaces appear to have the
highest activity for CH4 oxidation [21,22].
Mechanisms and kinetics of CH4 oxidation over Pd/PdO catalysts have elicited continued debate
in the literature [6,13,14,23], for which data interpretation is complicated by the transitions that Pd
catalysts undergo during thermal pre-treatment and reaction [24]. Furthermore, the high
concentration of H2O in the NGV exhaust and the typically transient reaction conditions that result
from cycling between oxidizing and reducing conditions in the NG engine [6,11] are known to
significantly impact catalyst activity and stability.
The present review is focused on the inhibition and deactivation effects of H2O, especially at the
relatively low temperatures representative of CH4 oxidation over Pd catalysts in a NG engine.
Although previous reviews have addressed the issue of Pd catalyst stability in the presence or
absence of H2O [4,12,20,25], and several catalyst deactivation mechanisms are possible at the
exhaust gas conditions [26], several unresolved issues remain. More recent studies of the past
decade have provided new insights into the effects of H2O, especially at lower temperatures, and
these are the focus of the present review. Note, however, that in many cases, fresh catalysts in
powder form have been evaluated using ideal fixed-bed micro-reactors and simulated exhaust gas
under steady state operating conditions. Tests of monolith catalysts with promoters suitably aged
and operated with A/F frequency and amplitude modulation that occur in a vehicle are few [7,11].
Nonetheless, interpretation of data from ideal catalyst studies allows direct links to be drawn
between fundamental catalyst properties and catalyst performance for CH4 oxidation, whereas in
real systems this may be more difficult to achieve.

183
2. Effects of H2O on CH4 Oxidation over Pd Catalysts
Water is a major component of the engine exhaust and is also a product of the combustion that
occurs in the catalytic converter. In TWCs, H2O acts as an oxidizing agent for CO conversion by
the water-gas-shift reaction and for steam reforming of hydrocarbons [4]. H2O also significantly
affects the thermal stability of the metals (Pt, Rh and Pd) present in the TWC as well as the
support, mostly through sintering mechanisms [4,27,28] and by changes in the Pd oxidation
state during hydrothermal aging [29]. Water may also act as a reaction inhibitor by adsorption onto
the catalyst.
Bounechada et al. [11] reported on the activity of a Pd-Rh (Pd/Rh = 39/1) TWC converter
supported on stabilized Al2O3, promoted with Ce-Zr (Zr/Ce = 3.5) and wash coated on a ceramic
honeycomb monolith, tested under fuel-lean (Ȝ > 1), stoichiometric (Ȝ = 1.00), and fuel-rich (Ȝ < 1)
conditions (gas composition: 0.15 vol.% CH4, 0.6% CO, 0.1% H2, 10% H2O, 10.7% CO2, 0.13%
NO, 0–1.14% O2; Ȝ was varied by changing feed O2 concentration; GHSV = 50,000 hí1). At
stationary conditions (constant Ȝ; steady-state experiment), the CH4 conversion was observed to
continuously decrease under both stoichiometric (52 to 43% after 0.5 h reaction) and fuel-lean
(from 62 to 59% after 0.5 h reaction) conditions, even though injecting a fuel-rich pulse during
fuel-lean stationary operation increased the CH4 conversion to its initial value at the onset of
reaction. The authors attributed the deactivation under fuel-lean conditions to the inhibition effect
of H2O on the CH4 oxidation reaction, whereas under stoichiometric conditions, partial reduction of
PdO due to the lack of oxygen, may lead to a loss in PdO active sites for CH4 oxidation. The
authors also claimed that the presence of high oxygen capacity metals (Ce and Zr) in the catalyst
made the reduction of PdO improbable under stoichiometric conditions. Under fuel-rich conditions,
H2O acts as an oxidant through water-gas shift and steam reforming reactions.
2.1. Water Concentration and Reaction Temperature Effects on CH4 Oxidation Activity of
Pd Catalysts
With the growing interest in NGVs, recent studies have focused on effects of H2O on Pd
catalysts during CH4 combustion [16,18,30–38]. Deactivation or inhibition effects of H2O are
dependent upon several factors including catalyst formulation, reaction temperature, catalyst
time-on-stream history, and H2O concentration. Table 1 summarizes selected data that show effects
of H2O added to the feed gas during CH4 light-off experiments over Pd catalysts. The light-off
temperature (here reported as the temperature corresponding to 30% CH4 conversion during
temperature programmed reaction, T30) increases as the H2O concentration increases, showing a
clear inhibition effect that increases in magnitude with increasing H2O concentration.
In several cases the effects of H2O have been examined by measuring the CH4 conversion at
steady-state, with and without H2O added to the feed gas. A typical set of data, reported by
Persson et al. [35], is shown in Figure 1 using several Pd/Al2O3 catalysts reacted at 500 °C. These
data also show that added H2O significantly suppresses CH4 conversion, but the effect is at least
partially reversible. Similar effects of H2O addition have been reported in the literature, as
summarized in Table 2. These reports confirm that H2O acts as an inhibitor of CH4 oxidation over

184
Pd catalysts and that upon removal of the H2O from the CH4/O2 reactant, the inhibition is partially
reversible [31,33].

Figure 1. Effect of water vapor on the activity for CH4 combustion over Pd/Al2O3 (Ŷ);
2:1 PdPt/Al2O3 (¨); and 1:1 PdPt/Al2O3 (o) at 500 °C; 5 vol.% of steam was added to
the 1.5% CH4/air feed gas, GHSV = 100,000 hí1, for 5 h. [35] Copyright©
2007 Elsevier.
Reaction temperature is another key variable affecting the role of H2O addition. Although the
data of Table 2 cannot be compared directly because of the different operating conditions, they do
show that at 600 °C, the decrease in CH4 conversion with H2O addition is much less significant
than at lower temperatures (400 °C). Several authors have proposed that the deactivation is related
to the reaction of H2O with active PdO sites [16,18,31,40,41], PdO + H2OĺPd(OH)2, resulting in
the formation of inactive Pd(OH)2, as first proposed by Cullis et al. [40]. Burch et al. [31] also
reported a strong inhibitory effect of water on Pd catalysts up to 450 °C. However, at higher
temperatures the negative influence of water on the activity was very small, suggesting that above
450 °C the reverse reaction (Pd(OH)2ĺPdO + H2O) occurs. Eriksson et al. [41] observed a
significant decrease in CH4 conversion over a much wider range of temperatures (200–800 °C)
after adding 18% H2O to a CH4/O2 feed over a Pd/ZrO2 catalyst, which was likely due to
the relatively high H2O concentration used in this study. Different results were reported by
Kikuchi et al. [16] when adding 1 vol.% H2O during CH4 oxidation over a Pd/Al2O3 catalyst, i.e., a
decrease in activity was observed up to about 450 °C and no H2O inhibition was observed at higher
temperatures. However, during addition of 20 vol.% H2O, the inhibiting effect could be observed
up to 600 °C, in qualitative agreement with Eriksson et al. [41].

0.1 wt.% Pd/H-beta

Catalyst

400

Temp
°C

Added H2O, vol.%
0
1
5
10
20
a

120,000

GHSV
hí1

372
372
420
425
445

Calcined at 873K for 6 h; b Calcined at 1273 K for 6 h.

290
315
335
360
365

Reference [16]
1.1% Pd/SnO2
1.1% Pd/Al2O3-36NiO
48,000
1% CH4/20% O2 in N2
T30, °C
360
410
-

0.2% CH4/10%O2
in N2

Feed Gas
mol %

10

Conc.
vol %

100

Period a
min

75

Before
H2O
addition

15

During
H2O
addition

58

After
H2O
addition

CH4 Conversion, %

Water addition

Conversion after 400 min
TOS with periodic water
addition

Comments

300
350
-

[38]

Refs.

Reference [39]
0.9% Pd/ZrO2 b
50,000
0.4% CH4/0.05% CO/5% CO2/10% O2 in N2
T30, °C

0.9% Pd/ZrO2 a

Table 2. H2O inhibition over Pd catalysts during CH4 oxidation.

345
400
430
460
510

Reaction conditions

GHSV, hí1
Dry feed gas composition, vol.%

1.1% Pd/Al2O3

Table 1. Effect of H2O addition on T30 (temperature at 30% CH4 conversion) during temperature-programmed CH4 oxidation over
Pd catalysts.

185

550

600

2 wt.% Pd/Al2O3

1 wt.% Pd/Al2O3

160,000

160,000

100,000

120,000

GHSV
hí1

0.4% CH4 in air

0.4% CH4 in air

1.5% CH4 in air

1% CH4/4%O2 in
N2

Feed Gas
mol %

8

10

5

5

Conc.
vol %

60

60

300

900

Period a

95

95

95

43

90

79

13

8

93

92

30

40

Water addition
CH4 Conversion, %
During
After
Before
H2O
H2O
H2O
addition addition addition

Conversion after 400 min
TOS with periodic water
addition
Conversion after 300 min
TOS with periodic water
addition

Initial activity

Conversion after 35 h TOS
with periodic water addition

Comments

Period: Time length of water addition period. b TOS: Time-on-stream. c HTNU-10 is the H-form of a medium pore zeolite with Si/Al = 7.1.

500

5 wt.% Pd/Al2O3

a

400

Temp
°C

1.3 wt.% Pd/HTNU-10c

Catalyst

Reaction conditions

Table 2. Cont.

[36]

[32]

[35]

[33]

Refs
.

186

187

Figure 2. Delay in the H2O peak with respect to other products obtained by passing
pulses of CH4/O2/He (closed square) and 1 vol.% O2/3.45 vol.% H2O/He (open square)
over Pd/ZrO2 at different temperatures. Reproduced with permission from [42].
Copyright © 2001 Elsevier.
Further insight into the H2O adsorption/desorption phenomena on Pd/ZrO2 catalysts has been
obtained using pulsed-flow experiments [42,43]. Accordingly, pulses of CH4/O2/He (1:4:95 vol %)
were passed over a Pd/ZrO2 catalyst at various temperatures and the products monitored by mass
spectrometer. The time at which the peak maximum for H2O appeared in each spectrum, compared
to other products, was reported as the delay in the H2O peak. The data (Figure 2) show that the H2O
generated during CH4 oxidation lags other products, suggesting a slow H2O adsorption/desorption
equilibrium which might include spillover to the support. As the temperature increases above
450 °C (723 K), the desorption rate of H2O increases and the delay in the H2O peak compared to
the other products is insignificant. This behavior is in agreement with observations from other
studies [30,31,44] that the desorption rate of H2O produced during CH4 oxidation is slow and on
the order of seconds below 450 °C, even though CO2, the other product of reaction, desorbs very
quickly. Increasing temperature above 450 °C removes the desorption time gap between CO2 and
H2O, and thus, no inhibition by H2O occurs. Ciuparu et al. [42] also pulsed gas containing 3.45
vol.% H2O/O2/He but no CH4 (and hence no reaction) through the same catalyst bed (Figure 2),
showing that the H2O generated from CH4 oxidation lags the H2O added to the feed. These data
demonstrate that the adsorption/desorption of H2O from the Pd catalyst surface is temperature
dependent and reaches equilibrium at temperatures above ~450 °C (723 K), even for H2O added in
the gas phase.
Figure 3 compares temperature-programmed-reaction (TPR) profiles for CH4 oxidation obtained
over a Pd/ZrO2 catalyst, from both pulsed and continuous flow experiments with or without H2O
added [42,43]. The pulsed flow TPR profile was obtained by injecting pulses of the reaction
mixture (1/4/95:CH4/O2/He for the “dry” feed and 1/4/95:CH4/O2/He saturated with ~2% H2O for
the “wet” feed) into a He stream every 3 min while ramping the temperature at 0.5 K miní1.
Between consecutive pulses the catalyst was purged in flowing He. The pulsed flow data of Figure

188
3 show that at temperatures above 450 °C (723 K), there is no H2O inhibition, since the
conversions of “dry” and “wet” reaction mixtures are essentially the same. At <450 °C, inhibition
is observed due to a low H2O desorption rate. When H2O is added to the gas phase, the H2O
adsorption rate is enhanced and the rate of desorption is further decreased. With continuous flow of
reactants and a higher H2O concentration, H2O inhibition occurs at high temperatures due to
re-adsorption. The addition of H2O to the feed directs the equilibrium towards more H2O
adsorption on the surface and hence a greater decrease in catalyst activity during CH4 oxidation.

Figure 3. Temperature-programmed reactions during pulsed or continuous flow of
reactants over Pd/ZrO2 with or without H2O in the feed. Reproduced with permission
from [42]. Copyright © 2001 Elsevier.
The above observations are consistent with the following hypotheses: (1) product inhibition of
CH4 oxidation by H2O on PdO catalysts occurs at temperatures below 450 °C; (2) product
inhibition by H2O is enhanced by its slow rate of desorption from the PdO catalyst relative to a
higher rate of CH4 oxidation; (3) PdO and H2O may interact via the reversible reaction: PdO +
H2OļPd(OH)2 yielding inactive Pd(OH)2 and thus reversibly deactivating PdO as first proposed
by Cullis et al. [40]; and (4) the extent of the CH4 oxidation reaction increases with increasing
temperature but is reduced with increasing H2O concentration in the gas phase.
2.2. Catalyst Sintering by H2O
The possibility that addition of H2O may degrade Pd catalysts through a sintering mechanism [26]
has also been investigated. According to Hansen et al. [45], the sintering rate of metal nanoparticles
depends on their size. For nanoparticles <3 nm in diameter, Ostwald ripening is the most likely
sintering mechanism. For larger particles (3–10 nm), both Ostwald ripening and particle migration
and coalescence may occur, but the sintering rate is much slower than for the smaller particles [45].
The particle sintering rate has also been shown to correlate with the vapor pressure of the surface
species [4]. Pd is unique among the PGMs in that the oxide (PdO) has a much lower vapor pressure

189
than the metal (Pd), and consequently, one would expect a very low sintering rate of PdO by Ostwald
ripening [4]. The rate of sintering is also dependent on the support. Lamber et al. [46] suggested that on
SiO2 in the presence of H2O, the formation of silanol (Si-OH) groups favors the migration and
coalescence of Pd, whereas in the absence of H2O, Ostwald ripening is favored. Sintering suppression
has been demonstrated for Pt catalysts using supports that enhance metal-support interactions [28].
Nagai et al. [47] demonstrated a correlation between the O electron density of the support, the
strength of the Pt-O interaction and the resulting crystallite size. Thus, supports with a stronger
metal-support interaction have a higher O electron density and yield smaller Pt crystallites in the
order SiO2 < Al2O3 < ZrO2 < TiO2 < CeO2 [28,47].
Xu et al. [48] reported that the main deactivation mechanism of Pd/Al2O3 catalysts following
exposure to 10 (v/v)% H2O/N2 at 900 °C for up to 200 h is Pd sintering. A substantial decrease in
Pd dispersion from 3.7% to 0.9% over 7 wt.% Pd/Al2O3 and similar decreases at other Pd loadings
after 96 h hydrothermal aging, were observed. As noted by Xu et al. [48], aging the catalyst at 900
°C ensures that PdO decomposition to Pd0 is complete and consequently the more rapid sintering
observed is relevant to the behavior of Pd0 rather than PdO.
Escandon et al. [49] examined effects of hydrothermal aging at lower temperatures, where PdO
is thermodynamically stable [6]. A 1 wt.% Pd/ZrO2-Ce catalyst was hydrothermally aged at 300,
425, and 550 °C in 2% H2O/Air for 30 h, before being evaluated for CH4 oxidation under lean-burn
conditions (5000 ppmv CH4 in dry air). The results, shown in Figure 4, are compared with the same
catalyst, thermally aged at 550 °C in dry air for 30 h (identified as Pd/ZrO2-Ce-550 in
Figure 4) [49]. A significant irreversible decrease in CH4 conversion occurs and the extent of
catalyst deactivation increases with aging temperature (Figure 4). The T50% increases from 375 °C
for the fresh oxidized catalyst (identified as Pd/ZrO2-Ce in Figure 4), to 450 °C for the air-aged
catalyst and to > 550 °C for the hydrothermally aged catalyst. Pd dispersion and BET surface area
of the aged catalysts did not change [49]. Comparing the activity results of the catalyst thermally
aged in air (Pd/ZrO2-Ce-550) with that aged in 2% H2O/air at 550 °C (Pd/ZrO2-Ce-550h), confirms
that catalyst deactivation rate increases in the presence of H2O. The stability of the hydrothermally
aged catalysts during reaction was also evaluated, using both isothermal deactivation experiments
at 500 °C and light-off measurements made after 50 h reaction with 5000 ppmv CH4 in air. The
catalysts aged in the presence of H2O at 300 °C underwent a significant deactivation whereas the
catalyst aged in the presence of H2O at 425 °C was much more resistant to deactivation, and after
25 h time-on-stream was the most active of all the catalysts examined. XRD analysis of the
catalysts showed that the more stable catalysts are associated with the most stable supports [49].
In another study of CH4 oxidation at low temperature (250–450 °C), a change in PdO
dispersion was suggested as the main cause of deactivation of 0.5% Pd/Al2O3 and 0.5% Pd/SiO2
catalysts [50]. Dispersion decreased from 10% for the unused 0.5% Pd/SiO2 catalyst to 5.6% for the
catalyst reacted in 1% CH4/air feed at 450 °C for 7 h, whereas for the 0.5% Pd/Al2O3 catalyst the
corresponding changes in dispersion were 67% to 6.3%, respectively. These observations are in
good agreement with that of Narui et al. [51], in which the PdO dispersion of a 0.5% Pd/Al2O3
catalyst decreased from 14% to 11% after 6 h reaction at 350 °C. Zhang et al. [52] investigated Pd
catalysts supported on ZSM-5 and reported that catalyst stability is improved when CH4 oxidation

190
is carried out in the presence of H2O at 430–480 °C, compared to the reaction in a dry feed. In both
cases, the loss in catalyst activity could be related to reduced PdO dispersion, as determined by the
Pd/Si ratio measured by XPS, but the loss in dispersion is smaller in the presence of H2O [52]. By
contrast, Araya et al. [53] reported an insignificant drop in PdO dispersion (from 31.7% to 28.2%)
of a Pd/SiO2 catalyst after 96 h of reaction at 325 °C in 1.5% CH4/6% O2 in He, despite a
significant decrease in CH4 conversion from 32% to 22%. The extent of catalyst deactivation was
found to further increase in the presence of 3% H2O added to the feed.

Figure 4. CH4 conversion over fresh 1 wt.% Pd/ZrO2-Ce catalyst compared to 1 wt.%
Pd/ZrO2-Ce thermally aged in air at 550 °C (Pd/ZrO2-Ce-550) and hydrothermally aged
at different temperatures in 2% H2O/air (identified as Pd/ZrO2-Ce-TTTh where TTT is
the aging temperature in °C). Reproduced with permission from [49]. Copyright©
2008 Elsevier.
Several studies have demonstrated that catalyst sintering can be reduced by encapsulating
Pd/PdO nanoparticles in support materials. Sinter-resistant Pd catalysts have been prepared by
atomic layer deposition of Al2O3 overlayers on Pd [54], as well as by the synthesis of Pd/SiO2
core-shell structures [55,56]. Cargnello et al. [22] reported a Pd/CeO2 core-shell catalyst supported
on Al2O3 for CH4 oxidation that is about 200xs more active than an equivalent Pd-CeO2/Al2O3
catalyst prepared by wet impregnation. The authors demonstrated that the Pd cores remain isolated
even after heating the catalyst to 850 °C and that the CH4 light-off curves (measured at GHSV of
200,000 hí1 in a feed gas of 0.5% CH4, 2% O2 in Ar) are the same for the fresh catalyst and one
that has been aged at 850 °C for 12 hours. The Pd nanoparticles encapsulated by CeO2 enhance the
metal-support interaction that leads to exceptionally high CH4 oxidation activity and good thermal
stability [22].
2.3. Effects of Support
The data of Table 1 show that the inhibition of CH4 oxidation by H2O on Pd catalysts is
dependent upon the support. Pd/Al2O3 shows significantly more inhibition with 10% H2O added to

191
the feed than either the Pd/SnO2 or Pd/ZrO2 catalysts. More detailed data from Kikuchi et al. [16]
comparing CH4 light-off curves for a 1.1 wt.% Pd/Al2O3 catalyst and a 1.1 wt.% Pd/SnO2 catalyst
with H2O added to the feed over a range of concentrations (1–20 vol.%), are shown in Figures 5
and 6. By increasing the H2O concentration, the CH4 light-off curves for both catalysts shift to
higher temperatures. However, the temperature shift is larger over the Pd/Al2O3 catalyst than
the Pd/SnO2. The authors completed a simplified kinetic analysis of the CH4 oxidation rate
data to show that the enthalpy of adsorption of H2O is strongest on the Pd/Al2O3 catalyst
(¨Had ~ í49 kJ/mol), from which they concluded that the significant loss in activity of the
Pd/Al2O3 in the presence of H2O is due to a high coverage of the active sites by H2O [16]. These
results could also be interpreted according to the more recent proposals by Schwartz et al. [44,57],
that hydroxyl accumulation on the support hinders oxygen migration and exchange, and hence
CH4 oxidation. The strong adsorption of H2O determined by kinetic analysis on the Pd/Al2O3
catalyst [16] is consistent with a large hydroxyl accumulation on the catalyst surface that could
inhibit the O exchange.

Figure 5. Catalytic combustion of CH4 over 1.1 wt.% Pd/SnO2 with different amounts of
water added (vol.%). Reaction conditions: CH4, 1 vol.%; O2, 20 vol.%; H2O, 0–20 vol.%;
N2, balance; GHSV 48,000 hí1. Reprinted with permission from [16]. Copyright©
2002 Elsevier.
The rate of deactivation during CH4 oxidation in the presence of H2O has been shown to be
reduced by using a support with high oxygen surface mobility. At temperatures below 450 °C,
Ciuparu et al. [30] reported the inhibition effect of H2O to be dependent upon the oxygen mobility
of the support. Comparing PdO supported on oxides with increasing surface oxygen mobility:
Al2O3 < ZrO2 < Ce0.1Zr0.9O2, they show that the resistance to H2O inhibition during CH4 oxidation
increases in the same order. The deactivation rate of PdO was also compared over Al2O3, MgO,
and TiO2 supports by Schwartz et al. [44,57] at temperatures <450 °C. Deactivation is shown to be
a consequence of reduced oxygen mobility due to hydroxyl adsorption. They also reported that
PdO/MgO catalyst has a slower deactivation rate compared with Al2O3 and TiO2 supports because
of the higher oxygen surface mobility on the MgO [44,57]. However, Pd catalysts dispersed on other

192
supports such as MCM-41, which have high surface area (1113 m2/g) and lower oxygen mobility than
MgO and Al2O3, did not deactivate either, suggesting that other factors also play a role, depending on
the catalyst and the support.

Figure 6. Catalytic combustion of CH4 over 1.1 wt.% Pd/Al2O3 with different amounts of
water added (vol.%). Reaction conditions: CH4, 1 vol.%; O2, 20 vol.%; H2O, 0–20 vol.%;
N2, balance; GHSV 48,000 hí1. Reprinted with permission from [16]. Copyright ©
2002 Elsevier.
Another study compared the stability of Pd/SiO2 and Pd/ZrO2 during CH4 oxidation using a dry
feed gas [53]. The data (Figure 7) show that the Pd/ZrO2 is stable after 40 h time-on-stream, while
the CH4 conversion over the Pd/SiO2 catalyst increases from 13% to 32% in the first 3 h, and then
decreases to 22% after 96 h (see Figure 7). Although the Pd/ZrO2 catalyst is more stable than the
Pd/SiO2 catalyst, its conversion is lower than for the Pd/SiO2 catalyst. The lower deactivation rate
observed on the Pd/ZrO2 is consistent with the higher oxygen mobility of this catalyst compared to
Pd/SiO2, as noted above.
Metal-support interactions, support stability and the tendency of the support to encapsulate Pd,
may also play a role in the deactivation of Pd catalysts during CH4 oxidation. Gannouni et al. [58]
compared Pd catalysts supported on silica and mesoporous aluminosilicas and showed that,
according to the light-off curves measured with 1% CH4, 4% O2 in He, CH4 oxidation activity is
enhanced on the pure silica support, whereas on the aluminosilica, the beneficial effect of Al3+ on
metal dispersion and catalytic activity is counterbalanced by partial metal encapsulation. Above
500 °C in the presence of H2O, the structural collapse of the support, metal sintering, and metal
encapsulation by the support all occur [58]. Similar effects were reported with SiO2 supports by
Zhu et al. [59]. SiO2 desorbs chemisorbed H2O (silanol groups –Si-OH) at ~397 °C [46] and the
formation of hydroxides according to the reaction: ‹ଶሺୱሻ ൅ ʹଶ ሺ୥ሻ ՞ ‹ሺሻସሺ୥ሻ is feasible at
temperatures above 700 °C [60,61]. Hydroxyl mobility can change the extent of metal-support
interactions [45,46]. Zhu et al. [59] reported the encapsulation of PdO by SiO2 during CH4
oxidation at only 325 °C. The authors suggested that silica migration by (i) formation of a

193
palladium silicide during H2 reduction at 650 °C that is subsequently oxidized during CH4
oxidation and (ii) migration of SiO2 during CH4 oxidation caused by the water formed during
reaction, are important related factors facilitating the encapsulation of PdO by the SiO2. Migration
of SiO2 onto the metal crystallites in other catalyst systems containing H2O has also been reported
in the literature [46,62].

Figure 7. Methane conversion over time of Pd/ZrO2 and Pd/Aerosil130 catalysts.
Reaction conditions: 1.5% CH4; 6% O2; total flow = 90 cm3miní1, balanced in He;
temperature = 325 °C; catalyst mass = 0.2 g. Reprinted with permission from [53].
Copyright © 2005 Elsevier.
Yoshida et al. [63] also examined the effects of various metal oxide supports of Pd on the low
temperature oxidation of CH4 as summarized in Table 3. The catalytic activity varies with the
support, but the support oxides with moderate acid strength (Al2O3 and SiO2) give maximum CH4
conversion. For these catalysts higher activity corresponds to a higher oxidation state of Pd (bulk
PdO). The lower activity of Pd on basic supports is attributed to the formation of binary oxides
from PdO and the support (such as Pd/MgOx), in spite of a high Pd oxidation state.
The effect of metal oxides added to Pd/Al2O3 to improve the hydrothermal stability has been
reported by Liu et al. [36] who showed in particular, that the addition of NiO or MgO improved the
hydrothermal stability of Pd/Al2O3 through the formation of NiAl2O4 and MgAl3O4 spinel
structures. According to the authors, the spinel results in weakened support acidity that suppresses
the formation of Pd(OH)2 during hydrothermal aging.

194
Table 3. Effect of support on properties of 5 wt.% Pd catalysts and their CH4 oxidation
conversion. Data adapted from [63].
Support
MgO
ZrO2
Al2O3
SiO2
SiO2-ZrO2
SiO2-Al2O3
SO42í-ZrO2
a

Support Acid Strength
(Ho)
22.3
9.3
3.3
-5.6
-8.2
-11.9
-13.6

Pd Dispersion
Fresh
Used
0.21
0.20
0.41
0.12
0.35
0.20
0.09
0.11
0.16
0.13
0.12
0.06
0.02

CH4 conversion a, %
12
3
59
58
20
10
11

measured at 350 °C in 0.25% CH4/3%O2 in He at GHSV of 1,200,000 hí1.

A comparison of initial CH4 oxidation activity as a function of temperature for Pd-Pt catalysts
on Al2O3, ZrO2, LaMnAl11O19, Ce-ZrO2, and Y-ZrO2 was reported by Persson et al. [64]. Monolith
catalysts were tested in a tubular quartz flow reactor at atmospheric pressure in 1.5 vol.% CH4 in
dry air and at a space velocity of 250,000 hí1. In steady-state experiments, reaction temperature
was set initially at 470 °C and then increased to 720 °C stepwise in 50 °C increments, with 1-h
holds at each temperature. The Pd-Pt/Al2O3 catalyst had the highest activity at lower temperatures
(470–570 °C), while the Pd-Pt/Ce-ZrO2 catalyst had the highest activity between 620 °C and
800 °C [64]. The authors suggested that the higher surface area of the Al2O3 compared to the other
supports (e.g., 90 m2/g for Al2O3 versus 10 m2/g for Ce-ZrO2) accounts for the higher activity of
Pd-Pt/Al2O3 at lower temperatures, due to higher dispersion of Pd-Pt oxides, while at higher
reaction temperatures the Pd-Pt catalyst probably undergoes reduction to the metal. A combination
of lower activity for Pd metal and its propensity for rapid sintering probably explain the lower
activity. The authors also suggested that the Ce-ZrO2 likely enhances the stability of the PdO,
similar to the enhanced stability observed on CeO2 [30]. In addition, ZrO2 has high oxygen
mobility [30] and the ability to re-oxidize metallic Pd into PdO should be higher. Indeed,
Pd/alumina is re-oxidized very slowly, whereas Pd supported on ceria-stabilized ZrO2, is re-oxidized
more rapidly.
Since H2O adsorption on the Pd and/or the support is an important step in inhibiting CH4
oxidation over Pd, support hydrophobicity may be expected to impact the inhibition effect of H2O.
Araya et al. [53] studied this effect on the deactivation of Pd-based catalysts by preparing 1 wt.%
Pd on two different commercial silicas, Aerosil130 and Aerosil R972. The Aerosil R972 is
hydrophobic since the OH groups have been replaced by methyl groups. Both 1% Pd/A130 and 1%
Pd/R972 were tested at 325 °C in a gaseous mixture of 1.5% CH4 and 6% O2 in He at a total flow
rate of 90 cm3 miní1 with addition of 3% H2O after 2 h As shown in Figure 8, the effect of H2O
addition to the feed gas is approximately the same for the hydrophobic silica, Pd/R972, and the
hydrophilic Pd/A130. In both cases, a large decrease in CH4 conversion is observed with the
introduction of H2O to the reactor. The authors reported a reaction order with respect to H2O of í0.25
for both Pd/A130 and Pd/R972, emphasizing that the hydrophobicity of the support does not affect the
extent of H2O inhibition observed on either catalyst.

195

Figure 8. (A) Pd/Aerosil130 catalyst, (B) Pd/R972 catalyst. Reaction conditions: total
flow = 90 cm3 (STP) miní1, temperature = 325 °C; catalyst mass = 0.2 g. Open
symbols: dry feed 1.5% CH4; 6% O2;balance He; closed symbol: wet feed 1.5% CH4;
6% O2 with 3% H2O, balance He. Reproduced with permission from [53]. Copyright ©
2005 Elsevier.
2.4. H2O Inhibition and Hydroxyl Formation
Although Pd(OH)2 has been postulated as a cause for deactivation of PdO catalysts in the
presence of H2O [18,31,32,40], and while this mechanism is consistent with many of the
observations discussed above, recent evidence obtained from FTIR and isotopic labeling
experiments that monitor the formation and conversion of hydroxyls on the catalyst surface during
reaction, suggest an alternative mechanism of deactivation.
Using DRIFTS, Persson et al. [35] reported an increase in signal intensity from surface
hydroxyls weakly H-bonded to the support (3200–3800 cmí1) [65] after introducing 1.5% CH4 in
air to a PdO/Al2O3 catalyst at low temperature (200 °C; Figure 9). The peak at 3016 cmí1 in Figure
9a, assigned to gas phase CH4, increases with time-on-stream because of catalyst deactivation. The
hydroxyls have characteristic adsorptions at 3733, 3697, 3556 and 3500 cmí1, with the hydroxyls at
3697 and 3733 cmí1 assigned to bridged and terminal isolated hydroxyl species, respectively. Upon
CH4 removal from the feed (Figure 9b), the peaks associated with OH species remain, highly
consistent with a slow desorption of OH species produced during CH4 oxidation. Hence, Persson et

196
al. [35] suggested that catalyst deactivation on PdO/Al2O3 might be due to the formation and
accumulation of hydroxyls on the catalyst surface, bound either to the PdO, Al2O3 or the interface
between the two [30]. Gao et al. [32] reported similar hydroxyl bands at 3733, 3697, 3556 and
3500 cmí1 during lean-burn CH4 oxidation (0.4% CH4 in air) at 250 °C. The FTIR spectra from
reaction with 2 vol.% H2O added to the CH4-O2 feed also yield a broad band at 3445 cmí1 that is
associated with OH species on Al2O3 [32].

Figure 9. Cont.

Figure 9. FTIR spectra of 5 wt.% Pd/Al2O3 at 200 °C (a) during the CH4-O2 reaction;
(b) desorption when CH4 was removed. Reproduced with permission from [35].
Copyright© 2007 Elsevier.

197

Figure 10. FTIR spectra at highest surface coverage and 350 °C on (1) PdO/Al2O3
during CH4-O2 reaction, (2) PdO/Al2O3 and (3) Al2O3 when injecting H2O pulses.
Reproduced with permission from [30]. Copyright © 2004 Elsevier.
Ciuparu et al. [30] also identified three well-defined peaks at 3732 (OHI), 3699 (OHII), and
3549 (OHIII) cmí1 associated with surface hydroxyls generated during CH4 oxidation on a
PdO/Al2O3 catalyst (3.5 wt % Pd) at 350 °C using a feed gas of 0.128% CH4 and 17.3% O2 in
He/N2 (Figure 10). The spectrum was compared to that measured at the same temperature when
injecting pulses of ~3% H2O into an air flow over the PdO/Al2O3 catalyst and the Al2O3 support
(see Figure 10). Since Al2O3 has been shown to have a significantly lower hydroxyl coverage
compared to PdO/Al2O3 when injecting H2O pulses at 350 °C (the spectrum of Al2O3 is magnified
by a factor of 15 in Figure 10), they concluded that the three peaks are associated with the presence
of OH adsorbed on the PdO catalyst surface. The higher hydroxyl coverage during CH4 oxidation
compared to pulse injection of H2O onto the PdO/Al2O3 catalyst, indicates that (1) adsorbed H2O is
dissociated on the surface of PdO/Al2O3 and (2) hydroxyls formed from H2O pulses are less
strongly bound to the surface than hydroxyls produced by the CH4 oxidation reaction.
Since the frequencies of the OHI and OHII species are shifted to higher wave numbers for OH
species more weakly bound to Pd, Ciuparu et al. [30] suggested that the high frequency peaks
(OHI, OHII) can be assigned to terminal and bridged hydroxyl species, respectively, and the low
frequency peak at ~3549 cmí1 with broad maximum values can be associated with OH species
bound to different sites (multi-bound OHs; OHIII) (Figure 10). Transient temperature experiments
show that the hydroxyl binding energy increases in the order OHI < OHII < OHIII [30].
The peak areas of the terminal, bridged, and multi-bound hydroxyls were monitored with
time-on-stream at different temperatures during reaction, as illustrated by Figure 11 for reaction at
175 °C [30]. Upon removal of CH4 from the feed, the peak areas for the bridged and multi-bound
OH species continue to increase, whereas the area of the terminal OH species decreases (Figure 11).
This decrease is attributed to the conversion of terminal OH species to bridged or multi-bound OH
species. Based on the intensities of the various hydroxyl species at different temperatures, the

198
authors proposed the inter-conversion among the OH species as: ୍୍୍ ՞ ୍୍ ՞ ୍ ՜ ଶ ሺ୥ሻ
where only terminal OH species recombine and desorb as H2O and the transformation of bridged
OH species to terminal OH species is the rate determining step (RDS) for hydroxyl desorption and
hence low temperature CH4 oxidation [30]. Importantly the authors show that the surface coverage
by the hydroxyls (Figure 11) correlates with the activity loss at low temperature, meaning that the
activity loss and surface coverage have similar timescales, from which they conclude that the
former is likely an effect of the latter [30].

Figure 11. The normalized peak areas of different surface OH species generated during
lean-CH4-O2 reaction at 175 °C. Reproduced with permission from [30]. Copyright
© 2004 Elsevier.
FTIR spectra measured during CH4 oxidation at 325 °C with 0.1% CH4/4%O2 in He over a
series of 3 wt.% PdO catalysts supported on Al2O3, MgO, TiO2 and MCM-41 [44] show that the
hydroxyl coverage is dependent on the support. On Al2O3, well defined peaks similar to those
identified by Ciuparu et al. [30] are observed, but no common peak among all catalysts that would
provide evidence for Pd-OH bond formation, are present. Furthermore the large contribution from
OH bonding on the supports makes it impossible to directly identify the presence of Pd(OH)2 on
these supports [32,44]. However, by using 18O isotopic labeling and FTIR, the authors demonstrate
that peaks associated with the accumulation of hydroxyls on PdO are not present at 325 °C. Hence,
the more recent evidence suggests that deactivation by Pd(OH)2 formation is unlikely, in agreement
with the experimental observation that Pd(OH)2/C decomposes in N2 at about 250 °C [66]. In
addition, evidence from temperature-programmed desorption studies of H2O adsorbed on PdO(101)
thin films, suggests the formation of an OH-H2O complex at low temperature (<127 °C) and low
coverage (< ½ monolayer), whereas H2O preferentially chemisorbs in molecular form at higher
coverages [67].
Schwartz et al. [44] showed, however, that catalyst deactivation during CH4 oxidation correlates
with hydroxyl accumulation on the oxide support. The redox mechanism for CH4 combustion on

199
Pd/PdO generally assumes dissociation of a CH4 molecule to yield a methyl fragment and a hydroxyl
group (CH4 + Pd-O + Pd-*ĺPd-OH + Pd-CH3, where Pd-* represents a vacancy) [68,69]. H atoms
are abstracted sequentially from the methyl group by neighboring Pd-O to form surface hydroxyl
groups (Pd-OH). Recombination of surface hydroxyls yields water and a surface vacancy
(2Pd-OHĺH2O + Pd-O +Pd-*), that is regenerated by oxygen (2Pd-* + O2ĺ2Pd-O) [68,69].
Based on their experimental studies, Schwartz et al. [44,57], proposed that during lean-burn CH4
oxidation, O2 molecules dissociate on Pd-* sites and exchange with oxygen on the support
so that Pd active sites are re-oxidized with oxygen atoms from the support during the catalytic
reaction as follows:
Pd-O + S-* ļ Pd-* + S-O
Pd-* + S-Os ļ Pd-Os + S-*

(1)
(2)

and overall:
Pd-O + S-OsļPd-Os + S-O

(3)

where S represents the support, S-* is an O vacancy on the support and Os represents an O atom
associated with the solid oxide. This proposed mechanism suggests the possibility that a primary
cause for catalyst deactivation is hydroxyl accumulation on the support, which hinders oxygen
migration and exchange processes.
Evidence for O exchange with the support is provided by the isotopic labeling experiments
summarized in Figure 12, during which Pd18O/Al216O3 and Pd18O/Mg16O were exposed to 18O2/He
flow at 400 °C [57]. An increase in 16O18O signal intensity with time is proposed to arise from
oxygen exchange with the catalyst support [44]. The 16O18O signal (see lower, separate dashed line
in Figure 12) is reduced when H216O is injected to the feed and is recovered when H216O is
removed. Apparently, hydroxyl groups tend to migrate to the oxide support rather than desorb. By
increasing the concentration of hydroxyl groups, through addition and dissociation of H2O, oxygen
exchange of Pd-* active sites with the oxide support (S-Os) is interrupted. Thus, the number of PdO
sites participating in the CH4 oxidation reaction decreases with time, as H2O dissociates and OH
coverage of the support increases, with a consequent decrease in CH4 conversion [44]. This
proposed mechanism of catalyst deactivation is believed to occur at temperatures below 450 °C.
Finally, the authors note that the rate of deactivation on Pd/Al2O3 catalysts, with higher
concentrations of hydroxyl during reaction, is higher than on catalysts containing a support with
higher oxygen mobility (Pd/MgO) [44,57].
Ciuparu et al. [70] also reported on pulsed experiments with 18O2 over pure Pd and Pd/ZrO2
catalysts, oxidized before reaction, to clarify the effect of hydroxyls on the surface oxygen
exchange. They determined that due to the slow recombination of hydroxyls and hence H2O
desorption from the Pd catalyst surface during CH4 oxidation (2Pd-OHĺH2O + Pd-O +Pd-*), the
isotopic exchange of oxygen with the Pd sites (see Figure 13) occurs before H2O desorption from
the surface. The oxygen vacancies on the PdO surface resulting from H2O desorption are thus

200
rapidly filled by oxygen from the PdO bulk or oxide support (Pd-* + S-OsļPd-Os + S-*). In fact,
in this unsteady-state experiment, the labeled oxygen pulsed through the catalyst bed, is purged
from the reactor before H2O is desorbed [70]. These observations are in agreement with the studies
of Schwartz et al. [44,57] already discussed and confirm that the accumulation of hydroxyls on the
Pd catalyst surface impedes the oxygen exchange and limits Pd catalyst activity.

Figure 12. Oxygen exchange of (a) 3 wt.% Pd18O/Al216O3 (top) and (b) 3 wt.%
Pd18O/Mg16O (bottom) with catalyst supports in a flow of 18O2/He at 400 °C. H216O
was injected at some time to probe its effect on oxygen exchange. Reproduced with
permission from [44]. Copyright ©2012 American Chemical Society.

Figure 13. Schematic of oxygen exchange during CH4 oxidation using labeled pulsed
experiments. Reproduced with permission from [70]. Copyright © 2002 Elsevier.

201
3. The Use of Pd-Bimetallic Catalysts for CH4 Oxidation
Pd-bimetallic catalysts have been studied to improve stability of Pd catalysts for CH4
oxidation [19,51,71,72]. Pd-bimetallic catalysts are usually less active than Pd alone [64,73–75]
simply because they contain less Pd, the most active metal for CH4 oxidation [20,25]. The lower
activity of the bimetallic compared to Pd alone may also be due to the presence of smaller amounts
of PdO as a result of alloy formation between Pd and Pt [64], or the transformation of PdO to Pd
metal [76]. According to Ozawa et al. [77] the addition of Pt improves PdO/Al2O3 catalyst stability
by preventing the growth of PdO and Pd–Pt particles during CH4 oxidation at high temperature
(800 °C) [77].
Several studies have reported higher initial activity of Pd-bimetallic catalysts compared to Pd
alone [17,19,51,78]. These researchers suggest that the second metal added to Pd dissociates
O2 and the resulting O atoms are adsorbed by Pd, helping to maintain PdO active sites.
Ishihara et al. [78] reported a T50 of 533 °C for a 1 wt.% Pd/Al2O3 catalyst, whereas for
a Pd-Ni/Al2O3 catalyst (Pd:Ni = 9:1) T50 was 380 °C. In another study, it was reported that a higher
dispersion of PdO on PdO-Pt/Į-Al2O3 catalyst (27%) compared to PdO/Į-Al2O3 (14%) results in
higher initial activity and higher stability of the bimetallic catalyst [51] . After exposing the
PdO/Į-Al2O3 catalyst to the reaction feed stream for 6 h at 350 °C, an increase in average particle
size from 8 to 11 nm is observed, whereas the average particle size does not change significantly
for the PdO-Pt/Į-Al2O3 catalyst [51].
Persson et al. [73] examined a series of Pd-bimetallics supported on Al2O3 finding that the
metallic phase structure has a significant influence on the catalyst stability. For example, in several
bimetallic systems (PdAg, PdCu, PdRh, and PdIr) separate phases of each metal oxide are formed
after calcination (at 1000 °C for 1h followed by 1000 °C for 2h after loading the supported metal
oxide powders onto a cordierite monolith) and this enhances catalyst stability in the case of the
PdCu and PdAg (as measured stepwise at temperatures from 400–800 °C in 1.4% CH4 in dry air at
a space velocity of 250,000 hí1). Formation of a Co or Ni aluminate spinel in PdCo and PdNi
bimetallics, however, does not improve catalyst stability, whereas alloy formation in PdPt and
PdAu on Al2O3 increases hydrothermal stability in the presence of 15% H2O/air at 1000 °C for 10
h. In another study by Persson et al. [64], Pd-Pt bimetallic catalysts on various supports (alumina,
zirconia) were shown to have higher thermal stability than monometallic Pd during CH4 oxidation
in dry air (1.5% CH4 in air at a GHSV 250,000 hí1). The stability of the Pd-Pt catalysts improved at
lower temperatures (up to 620 °C). At temperatures of 520 °C and 570 °C CH4 conversion on Pd-Pt
catalysts increased with time-on-stream. Above 620 °C (especially at 670 °C and 720 °C)
conversion decreased with time-on-stream. Those catalysts with higher initial activity also had
higher deactivation rates. The deactivation cannot be attributed to PdO decomposition because the
initial activity test showed that PdO decomposition started at higher temperature (770 °C with
1.5 vol.% CH4 in air). According to XRD results, no PdO decomposition was observed at
temperatures below 800 °C for the Pd/Al2O3, although PdO decomposition at ~700 °C may have
yielded Pd that was not detectable by XRD (due to low concentration or high dispersion).

202
The amount of second metal added to the Pd can also affect the stability of the bimetallic
catalyst. Persson et al. [74] reported that Pd-Pt bimetallic catalysts with Pd:Pt ratios of 2:1 and 1:1
are stable. Time-on-stream CH4 oxidation experiments (in 1.5% CH4 in air at a space velocity of
250,000 hí1) for both a 5 wt.% Pd/Al2O3 and a 2:1 Pd:Pt/Al2O3 bimetallic with total metal loading of
5 wt.% were studied over a wide range of temperatures (470–720 °C) [64]. The temperature was
increased from 470 °C to 720 °C stepwise by 50 °C and held for 1 h at each temperature. CH4
conversion over the Pd/Al2O3 and Pd-Pt/Al2O3 catalyst decreased during the 1 h reaction time at
each temperature. However, the decrease in conversion was lower for the bimetallic catalyst
compared to the Pd catalyst. The decrease in activity was higher at higher temperatures (670 °C
and 720 °C), especially for the Pd catalyst. In situ XRD spectra of the Pd-Pt bimetallic catalysts are
shown in Figure 14. At room temperature, a sharp peak corresponding to Pd-Pt (111) and a small
peak corresponding to PdO (101) are observed for the PdPt-Al2O3 catalyst. By increasing the
temperature to 300 °C, the PdO peak disappears and then reappears at 500 °C. The Pd-Pt peak
intensity reaches a maximum at 700 °C while the PdO peak disappears at this temperature. The
formation of Pd-Pt instead of PdO is consistent with deactivation of the bimetallic catalyst at high
temperature (700 °C).

Figure 14. High-temperature in situ XRD profiles of PdPt-Al2O3 during heating.
Reprinted with permission from [64]. Copyright © 2006 Elsevier.
Steady-state experiments using a 18.7 wt.% Pd/Al2O3 catalyst with different loadings of Pt (1.6,
3.1 and 3.9 wt.%) (Figure 15) reported by Ozawa et al. [77], also provide some insight into the
improved stability of bimetallic catalysts as Pt content increases. In this study, reaction temperature
was held at 800 °C and CH4 combustion rate was measured over a 10 h period using a 1% CH4 in air
feed gas at a GHSV of 1,500,000 mL/(gcat-h). Deactivation rate is shown to decrease as the Pt loading
of the Pd-Pt bimetallics increases. For example, the combustion rate for the 18 wt.% Pd-3.9 wt.%
Pt/Al2O3 decreases from 710 ȝmol sí1 gí1 to 460 ȝmol sí1 gí1 after 10 h, whereas it decreases to
400 ȝmol sí1 gí1 for the 18.4 wt.% Pd-1.6 wt.% Pt/Al2O3 catalyst.

203
XRD analysis of the catalysts studied by Ozawa et al. [77] after 10 h reaction indicates PdO to
be present in the Pt-doped catalysts while no Pd0 is observed. However, Pd0 is present in the Pd
monometallic catalyst, likely because of the decomposition of PdO at the high temperature of the
reaction (800 °C). In addition, the crystallite size of the PdO (101) in the Pd catalyst is larger than
for the Pd-Pt catalysts. Table 4 compares changes in PdO particle size and BET surface area before
and after 10 h reaction for the same Pd and Pd-Pt catalysts. From these data it is clear that the
extent of sintering of the Pd catalyst is greater than for the Pd-Pt catalysts. The time-on-stream
conversion data reported by Ozawa et al. [77] (Figure 15) were fitted to a deactivation equation
with two terms, the first representing rapid transformation of PdO to Pd0 of the Pd-Pt alloy phase,
and the second associated with the slow growth of the PdO crystallite [77]. The deactivation is
affected more by the second term suggesting that particle growth of the PdO is the main cause of
catalyst deactivation at the chosen reaction conditions [77].

Figure 15. CH4 combustion rate at 800 °C with time on stream. Combustion
conditions: 1 vol.% CH4, 99 vol.% air, CH4/air flow = 450 L.hí1, catalyst weight = 0.3
g. Catalyst 1, 2, 3, and 4 represent 18.7 wt.% Pd, 18.4 wt.% Pd-1.6 wt.% Pt, 18.1 wt.%
Pd-3.4 wt.% Pt, and 18.0 wt.% Pd-3.9 wt.% Pt over Al2O3 catalysts. Reprinted with
permission from [77]. Copyright © 2004 Elsevier.
Table 4. Changes in Pd and Pt-Pd catalyst properties before and after aging. Adapted
from [77].
Catalyst, wt.% on Al2O3

18.7% Pd

18.4% Pd-1.6% Pt

18.1% Pd-3.1% Pt

18.0% Pd-3.9% Pt

BET area, m2/g

Fresh
Aged

56
46

51
46

51
46

46
52

PdO size, nm

Fresh
Aged

12.5
17.9

15.3
18.0

15.2
16.7

14.7
16.2

These results are in a good agreement with the results reported by Yamamoto et al. [72] in which a
Pd-Pt bimetallic catalyst was more active for CH4 conversion than Pd (as measured by the temperature

204
required for 50% CH4 conversion) and the conversion was maintained following 2500 h
time-on-stream at 385 °C. XRD analyses showed that the crystallite growth as a function of time for
both Pd (111) and PdO (101) was faster on the Pd (10 g/L)/Al2O3 catalyst than the Pd(10 g/L)-Pt
(10 g/L)/Al2O3 catalyst. Hence one concludes that the presence of Pt retards the sintering of PdO.
Effects of H2O on deactivation of Pt versus Pt-Pd catalysts have also been reported, at both
thermal and hydrothermal aging conditions [17,19,71]. Pieck et al. [17] reported that the T50 of a
0.4% Pt-0.8% Pd/Al2O3 catalyst after thermal treatment at 600 °C for 4 h in wet air (60 cm3 miní1
air flow with 0.356 cm3 hí1 water ), is ~50 °C lower than that obtained over a Pd catalyst.
Lapisardi et al. [19] reported that a fresh Pd0.93-Pt0.07/Al2O3 catalyst (total metal loading 2.12 wt.%
with Pd:Pt molar ratio of 0.93:0.07) is as active as a fresh Pd/Al2O3 catalyst in a dry feed [19].
Interestingly, the Pd0.93-Pt0.07/Al2O3 catalyst is less affected by addition of 10 vol.% steam to the
feed stream than the 2.2 wt.% Pd/Al2O3 catalyst. The T50 for the Pd-Pt bimetallic increases from
320 °C to 400 °C when 10 vol.% steam is added to the feed stream, whereas the corresponding
increase in T50 for the Pd/Al2O3 catalyst is from 320 °C to 425 °C. Thus, the Pd-Pt bimetallic,
containing only 0.26 wt % Pt is more active and stable than the Pd catalyst for CH4 oxidation in the
presence of steam.
The stabilities of Pt and Pt-Pd catalysts each loaded on a wash coated monolith have also been
reported [71]. A feed stream with 4067 ppmv CH4 in air was reacted over these catalysts as reaction
temperature increased from 300 to 700 °C stepwise in 50 °C increments. CH4 conversion was
monitored for a period of 1 h at each temperature. Subsequently the temperature was decreased to
300 °C also in 50 °C steps, again holding at each temperature for 1 h. The conversion of CH4 was
compared for both heating and cooling cycles. The results show that the Pt-Pd catalyst is more
active than the Pt catalyst. The comparison between the heating and cooling cycles was also done
for steam-aged catalysts, in which the catalysts were exposed to the feed stream at 650 °C with
5 vol.% water for 20 h. Table 5 lists the T50 for both fresh Pt and Pd-Pt catalysts, the steam-aged
catalysts during tests in a dry feed and the steam-aged catalysts tested in a wet feed, containing
5 wt % H2O. The data show that the fresh Pd-Pt catalyst is more active than the fresh Pt catalyst.
Higher activities are also observed for steam-aged Pd-Pt catalysts tested in dry or wet feed gas.
Table 5. T50 for fresh and steam aged Pd and Pt-Pd catalysts operated in dry and wet
feed. Combustion conditions: 4067 vol. ppm CH4; total flow rate of 234.5 cm3/min;
500 mg catalyst; 5 vol.% water in wet feed. Adapted from [71].
Temperature at 50% CH4 conversion (T50), °C
Catalyst

Fresh

Steam-aged

Steam-aged

Dry feed

Dry feed

Wet feed

Pt

540

610

610

4:1 Pt-Pd

400

470

535

205
4. Kinetic Consequences of H2O on CH4 Oxidation over Pd Catalysts
The rate of CH4 oxidation over Pd catalysts is influenced by temperature, reactant partial
pressures, the state of the Pd at reaction conditions (Pd0, PdO or a sub-oxide), possibly Pd
crystallite size (i.e., may be structure-sensitive), and inhibition by products H2O and CO2.
Consequently, kinetic parameters reported in the literature vary over wide ranges; this is especially
true of the apparent activation energy for CH4 oxidation [20]. As noted by Carstens et al. [79], rate
data must account for the inhibition effect of H2O when determining the activation barrier, but
Ciuparu et al. [43] has shown that the correction is complicated by the fact that the effect of H2O
inhibition is temperature dependent. For example, the apparent activation energy for CH4 oxidation
over a Pd/ZrO2 catalyst is estimated to be 180 kJ/mol from data measured at temperatures below
192 °C, whereas a value of 87 kJ/mol is obtained at temperatures above 192 °C [42]. The higher
value of the apparent activation energy at lower temperatures is attributed to the strong inhibiting
effect of H2O on the Pd catalyst.
Zhu et al. [80] reported kinetic parameters for CH4 oxidation over a series of model Pd and PdO
surfaces and foils, and compared the values to literature data on supported Pd catalysts (Table 6). From
Table 6 the reaction orders for CH4 and O2 are probably not sensitive to the structure of the Pd catalyst,
although on the supported catalysts the reaction orders for H2O vary from í0.25 to í1.3. Taking
account of the error in the Ea estimates (±20 kJ/mol), Zhu et al. [80] concluded that on the large
single-crystal model catalysts, the activation energies are similar and the combustion of CH4 over Pd or
PdO is not sensitive to the structure of the catalyst. Larger Ea values are reported for the
Pd/oxide-supports (150–185 kJ/mol) corrected for the effect of H2O (assuming an order of í1) [79],
whereas the much smaller Ea for the Pd/zeolite catalysts (72–77 kJ/mol) are possibly associated with
the high acidity and high OH surface concentration of zeolites, in obvious contrast to the observed
inhibition by OH groups for PdO supported on conventional supports. The negative orders of reaction
for H2O are indicative of the varying degrees of inhibition of CH4 oxidation by H2O on Pd and PdO
surfaces and catalysts.
The role of H2O in the inhibition of PdO catalysts during CH4 oxidation has been documented in
this review to relate to the adsorption and slow desorption of H2O on active sites during reaction.
Kikuchi et al. [16] proposed a kinetic model assuming competitive adsorption between H2O and
CH4 on PdO sites, where dissociative CH4 adsorption was assumed to be the rate determining step
(RDS) and the coverage by C-species was assumed to be negligible. The main elementary steps of
the reaction are postulated as follows:
ଶ ሺ୥ሻ ൅ • ՜ ଶ  െ •
ସሺ୥ሻ ൅ ʹ• ՜ ଷ െ • ൅  െ •

(4)
(5)

from which the following rate expression is derived [16]:
‫ ݎ‬ൌ ݇௥

ܲେୌర
ͳ ൅ ‫ܭ‬ୌమ୓ ܲୌమ୓

(6)

206
where r is the reaction rate,  ୰ is the rate constant for H abstraction,  ୌమ୓ is the H2O adsorption
equilibrium constant, and େୌర and ୌమ ୓ are the partial pressures of CH4 and H2O, respectively.
 ୌమ ୓ is exponentially dependent upon the H2O adsorption enthalpy (ο‫ܪ‬௔ௗ௦ ). To increase the
activity and durability of the Pd catalysts in the presence of H2O,  ୌమ୓ should be small according
to the above reaction model. Based on the measured ο‫ܪ‬௔ௗ௦ values for water on supported Pd
catalysts, water adsorbed on Pd/Al2O3 has the highest negative adsorption enthalpy (ο‫ܪ‬௔ௗ௦ ൌí49
kJ molí1) compared to Pd/SnO2 (í31 kJ molí1) and Pd/Al2O3-36NiO (í30 kJ molí1) (Table 7)
despite the lower activation energy calculated for Pd/Al2O3 (see Table 7) [16]. A higher µο‫ܪ‬௔ௗ௦ µ
implies stronger H2O adsorption on the surface and is evidence of a higher coverage of active sites
by H2O molecules on Pd/Al2O3 catalysts and consequently lower catalyst activity. However, the
larger negative enthalpy also predicts a more rapid decrease in KH2O with increasing temperature
for Pd/Al2O3.
Table 6. Kinetic parameters for CH4 oxidation over Pd catalysts.
Catalyst

Ea,
kJ/mol

Pd foil
Pd (111)
Pd (100)
PdO foil
PdO(111)
PdO(100)

125
140
130
125
140
125

Pd black
8.5% Pd/Al2O3
0.5% Pd/Al2O3a
10% Pd/ZrO2
5% Pd/ZrO2
1% Pd/ZrO2
1% Pd/SiO2
2.8% Pd/HMord.
2.5% Pd-H-beta
a

T Range
°C

Refs

296–360
296–360
296–360
296–360
296–360
296–360

[81]
[80]
[80]
[80]
[80]
[80]

135
150
60
185
185
172
-

Reaction order
O2
H2O
Model Catalysts
0.7
í0.1
0.05
0.7
í0.1
0.05
0.9
0.01
0.07
0.7
0.2
í0.9
0.8
í0.1
í0.9
0.8
0.1
í1.0
Supported Catalysts
0.7
0.1
í0.8
1
0
-1
0.90
0.08
í1.3 to í0.9
1
0
í1
1.1
0.1
í1.0
1
0
í1.0
1
0
í0.25

296–360
232–360
240–400
232–360
250–280
227–441
227–441

[81]
[80,82]
[83]
[80,82]
[68]
[53]
[53]

77

0.7

-0.1

í0.4

342–417

[33]

72

0.5

0.2

í0.5

342–417

[33]

CH4

Ea determined under dry reaction conditions, correction for H2O inhibition.

The larger negative value in the order of H2O for the 1% Pd/ZrO2 catalyst, compared to the
Pd/SiO2 catalyst, as reported by Araya et al. [53] (Table 6), reflects stronger H2O adsorption on
ZrO2 than on the SiO2 [53]. Hurtado et al. [83] observed a change in the power-law reaction order
of H2O from í1.3 to í0.9 as temperature increased from 300 °C to 350 °C using a H2O-CH4-O2
reactant mixture and a commercial 0.5 wt.% Pd/Ȗ-Al2O3 catalyst. Considering the equation
proposed by Kikuchi et al. [16], with  ୌమ ୓ ୌమ୓ ‫ͳ ب‬, the H2O reaction order will reduce to í1 but
if  ୌమ ୓ ୌమ୓ is small, the H2O reaction order reduces to a value approaching zero.

207
Table 7. Estimated kinetic parameters for CH4 oxidation using the rate equation
୔ిౄర
[16].
‫ ݎ‬ൌ ݇௥
ଵା୏ౄమ ో ୔ౄమ ో

Pd loading
(wt.%)

ࡱ‫܉‬

οࡴ‫ ܛ܌܉‬for H2O

kJ/mol

kJ/mol

Pd/Al2O3

1.1

81

í49

Pd/SnO2

1.1

111

í31

Pd/Al2O3-36NiO

1.1

90

í30

Catalyst

Hurtado et al. [83] also attributed the inhibition effect of H2O during reaction to the adsorption
of H2O on Pd catalysts. Based on this assumption the authors examined several Eley-Rideal,
Langmuir-Hinshelwood and Mars-van Krevelen kinetic models finding that by considering
competitive adsorption between H2O and CH4 on Pd oxide sites and slow desorption of products,
the following kinetic model could be derived:
‫ݎ‬ൌ

௞భ ௞మ ௉ిౄర ௉ోమ

(7)

௞భ ୔ోమ ൫ଵା୏ౄమ ో ୔ౄమ ో ൯ାଶ௞మ ௉ిౄర ାሺ௞భ ௞మ Τ௞య ሻ௉ోమ ௉ిౄర

where k1, k2, and k3 are the rate constants for (1) irreversible oxygen adsorption, (2) surface
reaction with CH4, and (3) product desorption steps in the mechanistic sequence, respectively. This
model provides the best fit of their measured rate data. The ο‫ܪ‬ୟୢୱ for water estimated from
equation (7) is í54.5 kJ/mol, in agreement with the data of Table 7. The inhibiting effects of H2O
are assumed to be a consequence of a competitive adsorption between CH4 and H2O on PdO sites.
Deactivation by H2O was previously thought to be due to formation of inactive Pd(OH)2 that does
not participate in the CH4 oxidation reaction and is reversible at temperatures above 250 °C [66].
Hurtado et al. [83] also note that the formation of Pd(OH)2 is thermodynamically favored from
PdO sites rather than from Pd0. However, the more recent mechanism involving H2O inhibition of
the O exchange between Pd sites and oxide supports, proposed by Schwartz et al. [44,57] (see
earlier discussion) appears to be supported by more definitive data.
5. Conclusions
Studies of the past decade provide new insights into the effects of H2O on Pd catalysts during
CH4 oxidation, especially at lower temperatures. The principal effects of H2O are:
(a) reaction inhibition by H2O adsorption
(b) deactivation due to formation of Pd(OH)2 and
(c) H2O-assisted sintering at high reaction temperatures (>500 °C)
Reaction inhibition by H2O increases with (a) decreasing reaction temperature at <500 °C and
(b) higher H2O concentrations, while this effect is generally negligible at >500 °C. O surface
mobility of supports apparently influences H2O inhibition, i.e., high O mobility (on CeO2 and
ZrO2) results in less inhibition by H2O than for Al2O3.

208
The main cause of partially reversible deactivation has been related to hydroxyl adsorption
on the support and PdO. Although earlier studies suggested that formation of inactive Pd(OH)2
could be the cause of deactivation, recent studies provide definitive evidence that adsorbed
hydroxyls suppress O exchange between the support and Pd active sites causing suppression of
catalyst activity.
H2O-assisted sintering of supported Pd catalysts is observed at >500 °C. Catalysts with stabilized
supports or core-shell structures have higher resistance to hydrothermal sintering. Several studies show
that Pd bimetallic catalysts also improve catalyst stability, although explanations for the role of the
second metal are not well-defined. Suppression of PdO sintering, enhanced oxygen mobility and
suppression of hydroxide formation are postulated to play a key role in higher stability of Pd
bimetallic catalysts.
Rate expressions from kinetic studies of CH4 oxidation at conditions relevant to natural gas
vehicles are based on the assumptions of (a) product inhibition by H2O is a consequence of a
competitive adsorption mechanism between CH4 and H2O on PdO sites; and (b) deactivation by
H2O is due to the formation of inactive Pd(OH)2 . None of the previous kinetic studies have linked
the observed kinetic effects of H2O to O mobility that recent studies show is critical during
CH4 oxidation.
Acknowledgements
The financial support of Westport Innovations Inc. and Natural Sciences and Engineering
Research Council of Canada (NSERC) is gratefully acknowledged.
References
1.

2.
3.
4.
5.
6.

7.

US Government. The World Factbook 2013–14. Washington, DC: Central Intelligence
Agency, 2013. Available online: https://www.cia.gov/library/publications/the-worldfactbook/geos/xx.html (accessed on 1 September 2013).
Natural Gas Vehicle Knowledge Base, Statistics. Available online: http://www.iangv.org/
category/stats/2014 (Accessed on 25 March 2015).
Bartholomew, C.H.; Farrauto, R.J. Fundamentals of Industrial Catalytic Processes, 2nd ed.;
J. Wiley and Sons Inc: Hoboken, NJ, USA, 2006.
Barbier, J., Jr.; Duprez, D. Steam Effects in Three-Way Catalysis. Appl. Catal. B 1994, 4,
105–140.
Lox, E.S.J.; Engler, B.H. In Handbook of Heterogeneous Catalysis; Ertl, G., Knozinger, H.,
Weitkamp, J., Eds.; Wiley-VCH: Weinheim, Germany, 1997; Volume 4.
Ciuparu, D.; Lyubovsky, M.R.; Altman, E.; Pfefferle, L.D.; Datye, A. Catalytic Combustion of
Methane over Palladium-Based Catalysts. Catal. Rev. Sci. Eng. 2002, 44, doi:10.1081/CR120015482.
Silver, R.G.; Summers, J.C. In Catalysis and Automotive Pollution Control III Studies in Surface
Science and Catalysis; Frennet, A., Bastin, J.M., Eds.; Elsevier Science: Amsterdam, The
Netherlands, 1995; pp. 871–884.

209
8. Klingstedt, F.; Neyestanaki, A.K.; Byggningsbacka, R.; Lindfors, L.; Lundén, M.; Petersson, M.;
Tengström, P.; Ollonqvist, T.; Väyrynen, J. Palladium Based Catalysts for Exhaust
Aftertreatment of Natural Gas Powered Vehicles and Biofuel Combustion. Appl. Catal. A
2001, 209, 301–316.
9. Farrauto, R.J.; Hobson, M.C.; Kennelly, T.; Waterman, E.M. Catalytic Chemistry of Supported
Palladium for Combustion of Methane. Appl. Catal. A 1992, 81, 227–237.
10. Lampert, J.K.; Kazi, M.S.; Farrauto, R.J. Palladium Catalyst Performance for Methane
Emissions Abatement from Lean Burn Natural Gas Vehicles. Appl. Catal. B 1997, 14, 211–223.
11. Bounechada, D.; Groppi, G.; Forzatti, P.; Kallinen, K.; Kinnunen, T. Effect of Periodic
lean/rich Switch on Methane Conversion over a Ce–Zr Promoted Pd-Rh/Al2O3 catalyst in the
Exhausts of Natural Gas Vehicles. Appl. Catal. B 2012, 119–120, 91–99.
12. Gelin, P.; Primet, M. Complete Oxidation of Methane at Low Temperature over Noble Metal
Based Catalysts: A Review. Appl. Catal. B 2002, 39, 1–37.
13. Choudhary, T.V.; Banerjee, S.; Choudhary, V.R. Catalysts for Combustion of Methane and
Lower Alkanes. Appl. Catal. A 2002, 234, 1–23.
14. Centi, G. Supported Palladium Catalysts in Environmental Catalytic Technologies for Gaseous
Emissions. J. Mol. Catal. A 2001, 173, 287–312.
15. Forzatti, P.; Groppi, G. Catalytic Combustion for the Production of Energy. Catal. Today
1999, 54, 165–180.
16. Kikuchi, R.; Maeda, S.; Sasaki, K.; Wennerström, S.; Eguchi, K. Low-Temperature
Methane Oxidation over Oxide-Supported Pd Catalysts: Inhibitory Effect of Water Vapor.
Appl. Catal. A 2002, 232, 23–28.
17. Pieck, C.L.; Vera, C.R.; Peirotti, E.M.; Yori, J.C. Effect of Water Vapor on the Activity of
Pt-Pd/Al2O3 Catalysts for Methane Combustion. Appl. Catal. A 2002, 226, 281–291.
18. Roth, D.; Gélin, P.; Primet, M.; Tena, E. Catalytic Behaviour of Cl-Free and Cl-Containing
Pd/Al2O3 Catalysts in the Total Oxidation of Methane at Low Temperature. Appl. Catal. A
2000, 203, 37–45.
19. Lapisardi, G.; Urfels, L.; Gélin, P.; Primet, M.; Kaddouri, A.; Garbowski, E.; Toppi, S.; Tena,
E. Superior Catalytic Behaviour of Pt-Doped Pd Catalysts in the Complete Oxidation of
Methane at Low Temperature. Catal. Today 2006, 117, 564–568.
20. Li, Z.; Hoflund, G.B. A Review on Complete Oxidation of Methane at Low Tempertaure.
J. Nat. Gas Chem. 2003, 12, 153–160.
21. Colussi, S.; Gayen, A.; Camellone, M.F.; Boaro, M.; Llorca, J.; Fabris, S.; Trovarelli, A.
Nanofaceted Pd-O Sites in Pd-Ce Surface Superstructures: Enhanced Activity in Catalytic
Combustion of Methane. Angew. Chem. Int. Ed. 2009, 48, 8633–8636.
22. Cargnello, M.; Jaén, J.J.D.; Garrido, J.C.H.; Bakhmutsky, K.; Montini, T.; Gámez, J.J.C.;
Gorte, R.J.; Fornasiero, P. Exceptional Activity for Methane Combustion over Modular
Pd@CeO2 Subunits on Functionalized Al2O3. Science 2012, 337, 713–717.
23. Hellman, A.; Resta, A.; Martin, N.M.; Gustafson, J.; Trinchero, A.; Carlsson, P.-A.; Balmes,
O.; Felici, R.; van Rijn, R.; Frenken, J.W.M.; Andersen, J.N.; Lundgren, E.; Gronbeck, H. The
Active Phase of Palladium during Methane Oxidation. J. Phys. Chem. Lett. 2012, 3, 678–682.

210
24. Briot, P.; Primet, M. Catalytic Oxidation of Methane over Palladium Supported on Alumina:
Effect of Aging Under Reactants. Appl. Catal. 1991, 68, 301–314.
25. Gélin, P.; Urfels, L.; Primet, M.; Tena, E. Complete Oxidation of Methane at Low
Temperature over Pt and Pd Catalysts for the Abatement of Lean-Burn Natural Gas Fuelled
Vehicles Emissions: Influence of Water and Sulphur Containing Compounds. Catal. Today
2003, 83, 45–57.
26. Bartholomew, C.H. Mechanisms of Catalyst Deactivation. Appl. Catal. A 2001, 212, 17–60.
27. Kang, S.; Han, S.; Nam, S.; Nam, I.; Cho, B.; Kim, C.; Oh, S. Effect of Aging Atmosphere on
Thermal Sintering of Modern Commercial TWCs. Top. Catal. 2013, 56, 298–305.
28. Shinjoh, H. Noble Metal Sintering Suppression Technology in Three-Way Catalyst:
Automotive Three-Way Catalysts with the Noble Metal Sintering Suppression Technology
Based on the Support Anchoring Effect. Catal. Surv. Asia 2009, 13, 184–190.
29. Fathali, A.; Olsson, L.; Ekstrom, F.; Laurell, M.; Andersson, B. Hydrothermal Aging-Induced
Changes in Washcoats of Commercial Three-Way Catalysts. Top. Catal. 2013, 56, 323–328.
30. Ciuparu, D.; Perkins, E.; Pfefferle, L. In Situ DR-FTIR Investigation of Surface Hydroxyls on
Ȗ-Al2O3 Supported PdO Catalysts during Methane Combustion. Appl. Catal. A 2004, 263,
145–153.
31. Burch, R.; Urbano, F.J.; Loader, P.K. Methane Combustion Over Palladium Catalysts: The
Effect of Carbon Dioxide and Water on Activity. Appl. Catal. A 1995, 123, 173–184.
32. Gao, D.; Wang, S.; Zhang, C.; Yuan, Z.; Wang, S. Methane Combustion over Pd/Al2O3
Catalyst: Effects of Chlorine Ions and Water on Catalytic Activity. Chin. J. Catal. 2008, 29,
1221–1225.
33. Park, J.; Kim, B.; Shin, C.; Seo, G.; Kim, S.; Hong, S. Methane Combustion over Pd Catalysts
Loaded on Medium and Large Pore Zeolites. Top. Catal. 2009, 52, 27–34.
34. Stasinska, B.; Machocki, A.; Antoniak, K.; Rotko, M.; Figueiredo, J.L.; Gonçalves, F.
Importance of Palladium Dispersion in Pd/Al2O3 Catalysts for Complete Oxidation of Humid
Low-methane–air Mixtures. Catal. Today 2008, 137, 329–334.
35. Persson, K.; Pfefferle, L.D.; Schwartz, W.; Ersson, A.; Järås, S.G. Stability of
Palladium-Based Catalysts during Catalytic Combustion of Methane: The Influence of Water.
Appl. Catal. B 2007, 74, 242–250.
36. Liu, Y.; Wang, S.; Gao, D.; Sun, T.; Zhang, C.; Wang, S. Influence of Metal Oxides on the
Performance of Pd/Al2O3 Catalysts for Methane Combustion Under Lean-Fuel Conditions.
Fuel Process. Technol. 2013, 111, 55–61.
37. Zhu, G.; Han, J.; Zemlyanov, D.Y.; Ribeiro, F.H. Temperature Dependence of the Kinetics for
the Complete Oxidation of Methane on Palladium and Palladium Oxide. J. Phys. Chem. B
2005, 109, 2331–2337.
38. Okumura, K.; Shinohara, E.; Niwa, M. Pd Loaded on High Silica Beta Support Active for the
Total Oxidation of Diluted Methane in the Presence of Water Vapor. Catal. Today 2006, 117,
577–583.

211
39. Nomura, K.; Noro, K.; Nakamura, Y.; Yoshida, H.; Satsuma, A.; Hattori, T. Combustion of a
Trace Amount of CH4 in the Presence of Water Vapor over ZrO2-Supported Pd Catalysts.
Catal. Lett. 1999, 58, 127–130.
40. Cullis, C.F.; Nevell, T.G.; Trimm, D.L. Role of the Catalyst Support in the Oxidation of
Methane over Palladium. Faraday Trans. 1972, 68, 1406–1412.
41. Eriksson, S.; Boutonnet, M.; Järås, S. Catalytic Combustion of Methane in Steam and Carbon
Dioxide-Diluted Reaction Mixtures. Appl. Catal. A 2006, 312, 95–101.
42. Ciuparu, D.; Katsikis, N.; Pfefferle, L. Temperature and Time Dependence of the Water
Inhibition Effect on Supported Palladium Catalyst for Methane Combustion. Appl. Catal. A
2001, 216, 209–215.
43. Ciuparu, D.; Pfefferle, L. Support and Water Effects on Palladium Based Methane Combustion
Catalysts. Appl. Catal. A 2001, 209, 415–428.
44. Schwartz, W.R.; Ciuparu, D.; Pfefferle, L.D. Combustion of Methane over Palladium-Based
Catalysts: Catalytic Deactivation and Role of the Support. J. Phys. Chem. C 2012, 116,
8587–8593.
45. Hansen, T.W.; DeLaRiva, A.T.; Challa, S.R.; Datye, A.K. Sintering of Catalytic Nanoparticles:
Particle Migration Or Ostwald Ripening? Acc. Chem. Res. 2013, 46, 1720–1730.
46. Lamber, R.; Jaeger, N.; Schulz-Ekloff, G. Metal-Support Interaction in the Pd/SiO2 System:
Influence of the Support Pretreatment. J. Catal. 1990, 123, 285–297.
47. Nagai, Y.; Hirabayashi, T.; Dohmae, K.; Takagi, N.; Minami, T.; Shinjoh, H.; Matsumoto, S.
Sintering Inhibition Mechanism of Platinum Supported on Ceria-Based Oxide and
Pt-oxide–support Interaction. J. Catal. 2006, 242, 103–109.
48. Xu, Q.; Kharas, K.C.; Croley, B.J.; Datye, A.K. The Sintering of Supported Pd Automotive
Catalysts. ChemCatChem 2011, 3, 1004–1014.
49. Escandón, L.S.; Niño, D.; Díaz, E.; Ordóñez, S.; Díez, F.V. Effect of Hydrothermal Ageing on
the Performance of Ce-Promoted PdO/ZrO2 for Methane Combustion. Catal. Commun. 2008,
9, 2291–2296.
50. Muto, K.; Katada, N.; Niwa, M. Complete Oxidation of Methane on Supported Palladium
Catalyst: Support Effect. Appl. Catal. A 1996, 134, 203–215.
51. Narui, K.; Yata, H.; Furuta, K.; Nishida, A.; Kohtoku, Y.; Matsuzaki, T. Effects of Addition of
Pt to PdO/Al2O3 Catalyst on Catalytic Activity for Methane Combustion and TEM
Observations of Supported Particles. Appl. Catal. A 1999, 179, 165–173.
52. Zhang, B.; Wang, X.; M'Ramadj, O.; Li, D.; Zhang, H.; Lu, G. Effect of Water on the
Performance of Pd-ZSM-5 Catalysts for the Combustion of Methane. J. Nat. Gas Chem. 2008,
17, 87–92.
53. Araya, P.; Guerrero, S.; Robertson, J.; Gracia, F.J. Methane Combustion over Pd/SiO2
Catalysts with Different Degrees of Hydrophobicity. Appl. Catal. A 2005, 283, 225–233.
54. Lu, J.; Fu, B.; Kung, M.C.; Xiao, G.; Elam, J.W.; Kung, H.H.; Stair, P.C. Coking- and
Sintering-Resistant Palladium Catalysts Achieved through Atomic Layer Deposition. Science
2012, 335, 1205–1208.

212
55. Forman, A.J.; Park, J.; Tang, W.; Hu, Y.; Stucky, G.D.; McFarland, E.W. Silica-Encapsulated
Pd Nanoparticles as a Regenerable and Sintering-Resistant Catalyst. ChemCatChem 2010, 2,
1318–1324.
56. Adijanto, L.; Bennett, D.A.; Chen, C.; Yu, A.S.; Cargnello, M.; Fornasiero, P.; Gorte, R.J.;
Vohs, J.M. Exceptional Thermal Stability of Pd@CeO2 Core-Shell Catalyst Nanostructures
Grafted Onto an Oxide Surface. Nano Lett. 2013, 13, 2252–2257.
57. Schwartz, W.R.; Pfefferle, L.D. Combustion of Methane over Palladium-Based Catalysts:
Support Interactions. J. Phys. Chem. C 2012, 116, 8571–8578.
58. Gannouni, A.; Albela, B.; Zina, M.S.; Bonneviot, L. Metal Dispersion, Accessibility and
Catalytic Activity in Methane Oxidation of Mesoporous Templated Aluminosilica Supported
Palladium. Appl. Catal. A 2013, 464–465, 116–127.
59. Zhu, G.; Fujimoto, K.; Zemlyanov, D.Y.; Datye, A.K.; Ribeiro, F.H. Coverage of Palladium by
Silicon Oxide during Reduction in H2 and Complete Oxidation of Methane. J. Catal. 2004,
225, 170–178.
60. Jacobson, N.S.; Opila, E.J.; Myers, D.L.; Copland, E.H. Thermodynamics of Gas Phase
Species in the Si–O–H System. J. Chem. Thermodyn. 2005, 37, 1130–1137.
61. Opila, E.; Jacobson, N.; Myers, D.; Copland, E. Predicting Oxide Stability in High-Temperature
Water Vapor. JOM 2006, 58, 22–28.
62. Lund, C.R.F.; Dumesic, J.A. Strong Oxide-Oxide Interactions in Silica-Supported Fe3O4: III.
Water-Induced Migration of Silica on Geometrically Designed Catalysts. J. Catal. 1981, 72,
21–30.
63. Yoshida, H.; Nakajima, T.; Yazawa, Y.; Hattori, T. Support Effect on Methane Combustion
over Palladium Catalysts. Appl. Catal. B 2007, 71, 70–79.
64. Persson, K.; Ersson, A.; Colussi, S.; Trovarelli, A.; Järås, S.G. Catalytic Combustion of
Methane over Bimetallic Pd–Pt Catalysts: The Influence of Support Materials. Appl. Catal. B
2006, 66, 175–185.
65. Guerrero, S.; Araya, P.; Wolf, E.E. Methane Oxidation on Pd Supported on High Area
Zirconia Catalysts. Appl. Catal. A 2006, 298, 243–253.
66. Card, R.J.; Schmitt, J.L.; Simpson, J.M. Palladium-Carbon Hydrogenolysis Catalysts: The
Effect of Preparation Variables on Catalytic Activity. J. Catal. 1983, 79, 13–20.
67. Kan, H.H.; Colmyer, R.J.; Asthagiri, A.; Weaver, J.F. Adsorption of Water on a PdO(101)
Thin Film: Evidence of an Adsorbed HO-H2O Complex. J. Phys. Chem. C 2009, 113,
1495–1506.
68. Fujimoto, K.; Ribeiro, F.H.; Avalos-Borja, M.; Iglesia, E. Structure and Reactivity of
PdOx/ZrO2 Catalysts for Methane Oxidation at Low Temperatures. J. Catal. 1998, 179,
431–442.
69. Chin, Y.; Iglesia, E. Elementary Steps the Role of Chemisorbed Oxygen, and the Effects of
Cluster Size in Catalytic CH4-O2 Reactions on Palladium. J. Phys. Chem. C 2011, 115,
17845–17855.
70. Ciuparu, D.; Pfefferle, L. Contributions of Lattice Oxygen to the overall Oxygen Balance
during Methane Combustion over PdO-Based Catalysts. Catal. Today 2002, 77, 167–179.

213
71. Abbasi, R.; Wu, L.; Wanke, S.E.; Hayes, R.E. Kinetics of Methane Combustion over Pt and
Pt–Pd Catalysts. Chem. Eng. Res. Des. 2012, 90, 1930–1942.
72. Yamamoto, H.; Uchida, H. Oxidation of Methane over Pt and Pd Supported on Alumina in
Lean-Burn Natural-Gas Engine Exhaust. Catal. Today 1998, 45, 147–151.
73. Persson, K.; Ersson, A.; Jansson, K.; Iverlund, N.; Järås, S. Influence of Co-Metals on
Bimetallic Palladium Catalysts for Methane Combustion. J. Catal. 2005, 231, 139–150.
74. Persson, K.; Ersson, A.; Jansson, K.; Fierro, J.L.G.; Järås, S.G. Influence of Molar Ratio on
Pd–Pt Catalysts for Methane Combustion. J. Catal. 2006, 243, 14–24.
75. Persson, K.; Jansson, K.; Järås, S.G. Characterisation and Microstructure of Pd and Bimetallic
Pd–Pt Catalysts during Methane Oxidation. J. Catal. 2007, 245, 401–414.
76. Kuper, W.J.; Blaauw, M.; Berg, F.V.; Graaf, G.H. Catalytic Combustion Concept for Gas
Turbines. Catal. Today 1999, 47, 377–389.
77. Ozawa, Y.; Tochihara, Y.; Watanabe, A.; Nagai, M.; Omi, S. Deactivation of Pt·PdO/Al2O3 in
Catalytic Combustion of Methane. Appl. Catal. A 2004, 259, 1–7.
78. Ishihara, T. Effects of Additives on the Activity of Palladium Catalysts for Methane
Combustion. Chem. Lett. 1993, 22, 407–410.
79. Carstens, J.N.; Su, S.C.; Bell, A.T. Factors Affecting the Catalytic Activity of Pd/ZrO2 for the
Combustion of Methane. J. Catal. 1998, 176, 136.
80. Zhu, G.; Han, J.; Zemlyanov, D.Y.; Ribeiro, F.H. The Turnover Rate for the Catalytic
Combustion of Methane over Palladium is Not Sensitive to the Structure of the Catalyst. J. Am.
Chem. Soc. 2004, 126, 9896–9897.
81. Monteiro, R.S.; Zemlyanov, D.; Storey, J.M.; Ribeiro, F.H. Turnover Rate and Reaction
Orders for the Complete Oxidation of Methane on a Palladium Foil in Excess Dioxygen.
J. Catal. 2001, 199, 291–301.
82. Ribeiro, F.H.; Chow, M.; Dallabetta, R.A. Kinetics of the Complete Oxidation of Methane
over Supported Palladium Catalysts. J. Catal. 1994, 146, 537–544.
83. Hurtado, P.; Ordóñez, S.; Sastre, H.; DÕғez, F.V. Development of a Kinetic Model for the
Oxidation of Methane over Pd/Al2O3 at Dry and Wet Conditions. Appl. Catal. B 2004, 51,
229–238.

214

Effect of Ce and Zr Addition to Ni/SiO2 Catalysts for
Hydrogen Production through Ethanol Steam Reforming
Jose Antonio Calles, Alicia Carrero, Arturo Javier Vizcaíno and Montaña Lindo
Abstract: A series of Ni/Cex Zr1−x O2 /SiO2 catalysts with different Zr/Ce mass ratios were
prepared by incipient wetness impregnation. Ni/SiO2 , Ni/CeO2 and Ni/ZrO2 were also prepared
as reference materials to compare. Catalysts’ performances were tested in ethanol steam reforming
for hydrogen production and characterized by XRD, H2 -temperature programmed reduction (TPR),
NH3 -temperature programmed desorption (TPD), TEM, ICP-AES and N2 -sorption measurements.
The Ni/SiO2 catalyst led to a higher hydrogen selectivity than Ni/CeO2 and Ni/ZrO2 , but it could not
maintain complete ethanol conversion due to deactivation. The incorporation of Ce or Zr prior to
Ni on the silica support resulted in catalysts with better performance for steam reforming, keeping
complete ethanol conversion over time. When both Zr and Ce were incorporated into the catalyst,
Cex Zr1−x O2 solid solution was formed, as confirmed by XRD analyses. TPR results revealed
stronger Ni-support interaction in the Cex Zr1−x O2 -modified catalysts than in Ni/SiO2 one, which
can be attributed to an increase of the dispersion of Ni species. All of the Ni/Cex Zr1−x O2 /SiO2
catalysts exhibited good catalytic activity and stability after 8 h of time on stream at 600 ◦ C . The
best catalytic performance in terms of hydrogen selectivity was achieved when the Zr/Ce mass ratio
was three.
Reprinted from Catalysts. Cite as: Calles, J.A.; Carrero, A.; Vizcaíno, A.J.; Lindo, M. Effect of Ce
and Zr Addition to Ni/SiO2 Catalysts for Hydrogen Production through Ethanol Steam Reforming.
Catalysts 2015, 5, 58–76.
1. Introduction
Energy sustainability and reduction of CO2 emissions will be joined with a decrease in fossil fuel
use and the development of green energies. In this sense, hydrogen could be the energy carrier of
the future due to its clean and non-polluting nature [1–4]. However, the current hydrogen production
routes imply the use of fossil fuel-derived products, like methane, as feedstock and, for this reason,
the search of new alternatives for hydrogen production based on renewable resources is essential [5].
In line with this, hydrogen production from ethanol steam reforming is an attractive option, since
ethanol can be obtained from a wide variety of biomass feedstocks, and therefore, it can minimize
CO2 emissions. In addition, ethanol has a high H/C atomic ratio, low toxicity and can be easily and
safely manipulated and transported [6,7].
The ethanol steam reforming process (ESR) can be represented by the following equation:
CH3 CH2 OH + 3H2 O  6H2 + 2CO2

(1)

which involves several steps. The main reaction mechanism comprises dehydrogenation or
dehydration routes, which implies the formation of intermediates, like acetylene or ethylene,

215
respectively. Coke formation reactions also participate in the reaction pathway leading to catalyst
deactivation [8–10].
Ethanol steam reforming has been investigated using different catalysts [10,11]. In fact, Ni/SiO2
has been widely used, because it provides good activity and, at the same time, high selectivity to H2
and COx at a relatively low cost [12–19]. This catalytic behavior cannot only be ascribed to the nickel
active phase, since the support also affects catalysts activity and product distribution [11,15–18].
In this sense, the use of silica in Ni-Cu-supported catalysts gave better results in comparison to
other carrier materials, like alumina and MCM-41 [20]. Despite the good catalytic performance, the
literature also describes that Ni/SiO2 can be deactivated along time on stream, by metal sintering and
coke deposition.
Trying to get over this drawback, a number of attempts have been made to prevent deactivation
by modifying catalysts through the addition of different elements [16–18]. In this sense, increasing
attention has been given to CeO2 or ZrO2 -CeO2 mixed oxide, because the oxygen mobility improves
catalytic stability through coke gasification [9,11,21–24]. However, CeO2 and ZrO2 -CeO2 oxides
have a very low surface area to be used as catalytic supports. To increase the textural properties, silica
has been incorporated to ZrO2 -CeO2 [25,26], and the literature also reports how CeO2 or ZrO2 -CeO2
were supported on silica to be used in catalytic applications [27,28]. Other researchers describe the
addition of ZrO2 to ceria and silica to increase their thermal resistance and prevent the growth of Ni
crystallites [29–31].
Based on the above premises, this work proposes the use of Ce- and Zr-oxides as promoters
in the synthesis of Ni/SiO2 catalyst to reduce catalyst deactivation, allowing nickel particles
to accommodate on a porous support in order to avoid or minimize the sinterization of the
metallic phase.
2. Results and Discussion
2.1. Ethanol Steam Reforming over Reference Catalysts
Commercial silica, ceria and zirconia with BET surface areas of 276, 64 and 20 m2 /g, respectively,
were used as supports of the reference catalysts. The physicochemical properties of the catalyst
obtained after the incorporation of nickel into these supports are summarized in Table 1.
Table 1. Physicochemical properties of the reference catalysts.

a

Catalyst

Ni a
wt%

SBET
m2 /g

Acidity b
meq-NH3 /g

DcN iO
nm

DdN i
nm

Ni/S
Ni/C
Ni/Z

6.9
6.7
6.8

263
53
15

0.275

16.4
-

17.6
-

ICP-AES; b NH3 -temperature programmed desorption (TPD) analysis; c calculated from the (200)
reflection of NiO in XRD; d calculated from the (111) reflection of Ni in XRD.

216
The actual nickel content measured by ICP-AES was near the nominal loading. Compared to
the supports, the BET surface area of the catalysts decreased, although the trend is maintained,
Ni/S >Ni/C >Ni/Z. Since the catalyst acidity can influence the products’ distribution in the steam
reforming reaction, it was measured by NH3 -TPD, verifying that only the Ni/Z catalyst exhibited
measurable acidity with a desorption peak around 220 ◦ C , corresponding to 0.275 meq-NH3 /g.
Figure 1 shows the XRD diffractograms of the calcined and reduced catalysts, where peaks
corresponding to the carrier materials and those associated with the Ni phase can be observed.
For calcined samples, reflections corresponding to the planes (111), (200) and (220) of cubic
NiO can be observed at 2θ = 37.3◦ , 43.3◦ and 62.9◦ (JCPDS 78-0643), while peaks attributed
to the planes (111) and (200) of metallic Ni appear at 2θ = 44.4◦ and 51.8◦ (JCPDS 70-1849)
for the reduced catalysts. The intensity of these reflections is hardly noticeable in the case
of the Ni/C sample and overlap with ZrO2 peaks in the Ni/Z sample. Apart from the
diffraction of the Ni phase, the pattern of Ni/S shows only one diffuse reflection around 2θ
values of 22.5◦ , typical of certain amorphous materials, while diffraction peaks corresponding
to cubic cerium (IV) oxide (JCPDS 89-8436; 2θ = 28.5◦ , 33.0◦ , 47.5◦ , 56.3◦ , 59.1◦ and 69.3◦ )
can be clearly observed for Ni/C, and those corresponding to a mixture of monoclinic and
tetragonal zirconium (IV) oxide (JCPDS 74-1200 and 80-0784, respectively; main peaks at
2θ = 28.3◦ , 31.5◦ and 50.2◦ ) are seen for the Ni/Z catalyst. These analyses were used to calculate
the Ni phase crystallites size of the Ni/S sample from the NiO(200) reflection at 2θ = 43.3◦ and
the Ni(111) reflection at 2θ = 44.4◦ (Table 1). However, accurate crystallite diameters could not be
calculated, neither for the Ni/C sample, due to the low intensity of the diffraction lines, nor for the
Ni/Z sample, where Ni species peaks overlapped with those of ZrO2 .
The TPR analysis of calcined reference catalysts is shown in Figure 2. The profiles show a main
peak between 270 and 350 ◦ C , attributed to the reduction of NiO particles to Ni. In the case of the
Ni/C and Ni/Z samples, the maximum of the reduction peak is placed at lower temperatures, and
also, an additional shoulder can be observed, which is attributed to a fraction of the NiO particles
strongly interacting with the support. Finally, in the case of the Ni/C sample, another peak arises
around 800 ◦ C , due to the partial reduction of Ce4+ to Ce3+ [21]. TPR analysis of the supports (not
shown) demonstrated that, under our experimental conditions, only ceria may give rise to hydrogen
consumption with maxima at 800 ◦ C , corresponding to the diffusion-limited partial reduction of bulk
CeO2 particles. The low reduction temperature of Ni/C and Ni/Z catalysts suggests also a weaker
metal support interaction in these two catalysts.
Table 2 summarizes the results obtained for the ethanol steam reforming reaction after 8 h of
time on stream (TOS). While the Ni/C sample kept ethanol conversion above 99%, it decreased to
95.0% and 97.1% after 8 h TOS for the Ni/Z and Ni/S samples, respectively. Regarding intermediate
products, acetaldehyde can be detected when using Ni/S and Ni/C, while ethylene is found with the
Ni/Z sample due to the presence of acid sites, favoring the ethanol dehydration reaction. Methane
selectivity decreases as hydrogen selectivity increases, indicating that methane steam reforming is
one of the main steps in the ethanol steam reforming reaction mechanism. Thus, a slight hydrogen
selectivity is reached by the Ni/S catalyst in comparison with the rest of samples.

217

Figure 1. XRD patterns of the calcined (a) and reduced (b) reference catalysts.

Figure 2.
Temperature programmed reduction (TPR) profiles of the calcined
reference catalysts.

218
Table 2. Results of ethanol steam reforming obtained with the reference catalysts
(atmospheric pressure; T: 600 ◦ C ; weight hourly space velocity; WHSV: 12.7 h−1 ;
RH2O/EtOH : 3.7 mol/mol; time on stream (TOS): 8 h).
Catalyst
Ni/S
Ni/C
Ni/Z

XEtOH

Selectivity (mol%)

mol%

H2

CO2

CO

CH4

C2 H4

CH3 CHO

97.1
99.1
95.0

84.2
82.5
80.7

49.4
48.0
45.3

31.5
31.6
31.8

17.4
19.6
20.4

0
0
2.4

1.7
0.7
0

Based on the above results and taking into account that silica is a porous support that could
accommodate Ce an Zr oxides, the next study was the modification of Ni/SiO2 catalysts by Ce or
Zr addition.
2.2. Ni/SiO2 Catalysts Modified by Ce or Zr
2.2.1. Characterization of Ce- or Zr-Modified Silica
For this study, four silica-supported Ni catalysts were prepared and modified by Ce or Zr addition,
both before and after the incorporation of Ni. Thus, amorphous silica was used as the support of the
Ce and Zr phases, and the XRD patterns of the resulting materials are shown in Figure 3, together
with that of the Ni/S sample (described in Section 2.1), where Ce and Zr were later incorporated.
In the case of the CeS sample, the main diffraction peaks characteristic of the planes (111), (200),
(220) and (311) of cubic CeO2 (JCPDS 89-8436) are observed at 2θ = 28.5◦ , 33.0◦ , 47.5◦ and
56.3◦ . On the other hand, in the diffraction pattern corresponding to the ZrS sample, no clear
peaks of Zr species can be seen, although a small broad slope appears around 2θ = 30.2◦ , which
may be attributed to the plane (101) of highly dispersed ZrO2 in the tetragonal phase (JCPDS
80-0784). This suggests that while Ce trends to aggregate into CeO2 particles [16], Zr trends to
form ZrO2 crystallites too small to be detected by XRD. This fact agrees with the results published
by Sánchez-Sánchez et al. [15] with Zr/Al2 O3 , where for Zr loadings up to 6.6 μmol Zr/m2 , Zr ions
exists in atomic dispersion. The BET surface area of CeS and ZrS was 267 and 227 m2 /g, respectively,
slightly lower in comparison to the silica used as the support, due to the metals’ incorporation.
Finally, TPR analysis of CeS and ZrS materials (not shown) evidenced that no detectable hydrogen
consumption occurs at the metal loading used in this study (around 10 wt%), in spite of the partial
reduction observed for the bulk CeO2 used as a support in Section 2.1.

219

Figure 3. XRD patterns of the Ce- and Zr-modified silica supports.
2.2.2. Characterization of Ce- or Zr-Modified Ni/SiO2 Catalysts
The physicochemical properties of the prepared Ni catalysts are summarized in Table 3. ICP-AES
analysis revealed that silica support was loaded with near the desired amounts of Ni, Ce and
Zr. Although the incorporation of a second metal into the previously described materials slightly
decreased the value of their textural properties, all of the catalysts exhibited a high BET surface
area. The acidity of the materials was measured by NH3 -TPD and is summarized in Table 3. Ce/NiS
and Ni/CeS catalysts show negligible acidity, while Zr incorporation into silica support results in the
presence of acid sites, which is more pronounced when Zr is incorporated after Ni, probably due to
easier accessibility to these acid sites (not covered by Ni).
The XRD spectra of the calcined and reduced catalysts are shown in Figure 4. All of the calcined
samples exhibit the peaks characteristic of cubic NiO (JCPDS 78-0643) at 2θ = 37.3◦ , 43.3◦ and
62.9◦ , while the reduced samples show the reflections corresponding to metallic Ni (JCPDS 70-1849)
at 2θ = 44.4◦ and 51.8◦ . As observed for the corresponding modified support, the main diffraction
peaks characteristic of cubic CeO2 appear in the case of the Ni/CeS and Ce/NiS catalysts, and the
broad slope attributed to highly-dispersed ZrO2 can be seen for the Zr/NiS and Ni/ZrS catalysts. The
mean size of the CeO2 particles, calculated by means of the Scherrer equation, resulted in being
higher for the Ce/NiS sample, which indicates that CeO2 particles are better dispersed on silica than
on Ni/S. Regarding the Ni phase, the mean crystallite sizes are summarized in Table 3. Concerning
the calcined samples, both Ce- and Zr-modified catalysts have slightly larger NiO crystallites than
the unmodified Ni/S sample (Table 1), mainly when Ni is added to catalyst after Ce or Zr. A similar

220
trend is found for the reduced samples. However, while the Ni phase crystallite size increases after
the reduction process for the Ce-modified samples, attributed to the thermal effect, the addition of Zr
leads even to a slight reduction of the crystallite size, which may be attributed to higher thermal
stability together with lattice contraction from NiO to Ni due to the different molar volumes of
these phases.
Table 3. Physicochemical properties of the Ni/SiO2 catalysts modified by Ce or Zr.
Catalyst

Ni a
wt%

Ce a
wt%

Zr a
wt%

SBET
m2 /g

Acidity b
meq-NH3 /g

DcCeO2
nm

DdN iO
nm

DeN i
nm

Ce/NiS
Zr/NiS
Ni/CeS
Ni/ZrS

6.7
6.7
6.8
6.9

9.7
9.2
-

9.4
9.9

228
236
247
213

0.144
0.133

8.1
7.0
-

19.8
20.0
21.8
21.6

19.9
19.9
23.2
19.7

a

ICP-AES measurements; b NH3 -TPD analysis; c calculated from the (220) reflection of CeO2 in
XRD; d calculated from the (200) reflection of NiO in XRD; e calculated from the (111) reflection
of Ni in XRD.

Figure 4. XRD patterns of the Ce- and Zr-modified Ni/SiO2 catalysts: (a) calcined;
(b) reduced.
The reducibility of these catalysts was determined by H2 -TPR experiments. The profiles shown
in Figure 5 exhibit more than one peak attributed to the direct reduction of Ni2+ species to Ni,
which means the presence of different reducible nickel species depending on the interaction with the
support. Quantitative analysis of the TPR profiles revealed total reduction of the Ni2+ species for
all of the samples. As in the case of the Ni/S sample (Figure 2), the peak at the lower temperature

221
is attributed to the typical reduction of NiO to Ni, which may correspond to relatively large NiO
particles, while the shoulder at the higher temperature is ascribed to nickel species in close contact
with the support [22,23]. The Ce/NiS sample presents a lower reduction temperature than Ni/S
catalyst, indicating the promotion of Ni reducibility by CeO2 incorporation, an effect typically
found for lanthanide elements [16]. On the contrary, reduction of the Zr/NiS sample takes place
at higher temperatures, which may indicate an intimate contact between the Ni phase and the
promoter due to partial covering of the NiO particles by the highly-dispersed Zr phase added to
the Ni/S sample. On the other hand, when Ni was impregnated after Ce or Zr, reduction occurs in
a broad temperature range, similarly to the reduction of Ni/S (Figure 2). A shoulder arises in the
profile of the Ni/CeS catalyst at slightly lower temperatures than the beginning of the reduction
of the Ni/S sample, probably due to the presence of oxygen vacancies in CeO2 , promoting the
reduction of NiO [24,32], an effect facilitated by the small dimensions of the CeO2 crystallites
(Table 3). The profile corresponding to the Ni/ZrS sample reaches higher temperatures, probably due
to the easier contact between NiO and ZrO2 , which is well dispersed over the support, unlike the CeO2
particles, as the XRD patterns evidenced. The shift of the Zr-modified catalysts profiles to higher
reduction temperatures in comparison to the rest of the samples indicates higher thermal stability
and accounts for the decrease in the Ni phase crystallite size after the reduction of the catalyst, as
previously observed from the XRD results.

Figure 5. Temperature programmed reduction (TPR) profiles of the calcined Ce- or
Zr-modified Ni/SiO2 catalysts.
2.2.3. Catalytic Results Using Ni/SiO2 Catalysts Modified by Ce or Zr
The catalytic results obtained by these catalysts after 8 h of TOS on ethanol steam reforming at
600 ◦ C are shown in Table 4. Complete ethanol conversion was maintained over all of the catalysts,
except in the case of the Ce/NiS sample, whereas it slightly decreased up to 98.5 after 8 h. On the

222
other hand, hydrogen selectivity is lower for the samples where incorporation of Ni was carried out
before the addition of the Ce or Zr, which may be due to some Ni particles covered by the promoters.
This effect is more pronounced in the case of the Zr/NiS sample, which reached the lowest hydrogen
selectivity, due to the high dispersion of the Zr phase, probably in the form of an overlayer over
the Ni particles. However, total ethanol conversion is achieved with this sample, because the acidic
nature of this catalyst favors ethanol dehydration to ethylene, as observed from the relatively high
C2 H4 amount obtained among the products. Ethylene selectivity is smaller for the Ni/ZrS sample,
as expected from its lower acidity (Table 3). Therefore, regarding the impregnation order of the
different elements, the incorporation of Ce or Zr before Ni on SiO2 leads to lower selectivity to
intermediate products and, consequently, higher selectivity towards the main products (H2 and CO2 ).
This may be ascribed to the easier accessibility of reactants to the Ni sites. Finally, the CO2 /CO ratio
among the gaseous products is increased by the presence of CeO2 , since it favors the water-gas shift
reaction [33,34].
Table 4. Results on ethanol steam reforming obtained with the Ni/SiO2 catalyst modified
by Ce or Zr (atmospheric pressure; T: 600 ◦ C ; weight hourly space velocity; WHSV:
12.7 h−1 ; RH2O/EtOH ; 3.7 mol/mol; TOS: 8 h).
Catalyst
NiS
Ce/NiS
Zr/NiS
Ni/CeS
Ni/ZrS

Selectivity (mol%)

XEtOH

Coke

mol%

H2

CO2

CO

CH4

C2 H4

CH3 CHO

g/gcat h

97.1
98.5
100
100
100

84.2
82.0
79.2
84.3
88.3

49.4
49.6
43.5
56.2
56.5

31.5
29.9
29.1
25.1
28.0

17.4
19.5
21.2
18.7
12.9

0
0
6.2
0.0
2.6

1.7
1.0
0
0
0

0.369
0.230
0.475
0.370
0.518

The results shown in Table 4 indicate that coke formation occurs at a much higher rate on
Zr-modified catalysts, because, as explained above, the acid sites joined with the presence of Zr
promote ethanol dehydration to ethylene, which acts as a hard precursor for carbon deposition [35].
Although the highest coke amount was found on the Ni/ZrS sample, no loss of activity was observed
for this sample, since carbon was deposited in the form of nanofibers (Figure 6). This type of
coke does not embed metal particles [36], keeping the metal surface accessible for reactants, which,
together with the high surface area of the silica support to accommodate large amounts of coke, has
been described to have a minor effect on catalyst deactivation [37,38]. However, high coke formation
is not desirable, because it would probably result in reactor plugging. On the other hand, when
Ce was added over the Ni/S sample (Ce/NiS), coke formation diminished as a consequence of the
enhancement of carbon gasification by CeO2 [16,22]. However, in the case of the Ni/CeS sample,
this reduction of the coke amount is not observed, due to the higher Ni particles of the catalyst that
favors the growth of carbon nanofibers [39,40].

223

Figure 6. TEM image of coke formed on catalysts after being used for ethanol steam
reforming for 8 h: (a) Ni/S; (b) Ce/NiS; (c) Ni/ZrS.
Consequently, taking into account that the highest hydrogen selectivity was reached with the
Ni/ZrS catalyst at complete ethanol conversion and Ce/NiS led to a lesser amount of carbon deposited

224
on the catalyst, the next study consisted of the modification of Ni/SiO2 catalysts by both Ce and Zr
incorporation before the Ni phase.
2.3. Ni/Cex Zr1−x O2 /SiO2 Catalysts: The Effect of the Zr/Ce Ratio
2.3.1. Characterization of the Ni/Cex Zr1−x O2 /SiO2 Catalysts
For this study, three Ni catalysts supported on Ce- and Zr-modified silica were prepared with
Zr/Ce mass ratios of 1/3, 1 and 3. Their physicochemical properties are shown in Table 5, where it
can be observed that there is a good agreement between the metal contents determined by ICP-AES
and the theoretical values. All of the catalysts show a high surface area in comparison to the values
corresponding to the Ni/C and Ni/Z catalysts (Table 1), as expected from the use of silica as a support.
However, they exhibit lower BET surface areas than the Ni/S sample, with a slight decreasing trend
as the Zr content increases (Table 3). As expected, the acidity of the materials measured by NH3 -TPD
increased with the Zr loading, the Ni/ZrS catalyst having the highest number of acid sites.
Table 5. Physicochemical properties of the Ni/Cex Zr1−x O2 /SiO2 catalysts.
Catalyst

Ni a
wt%

Ce a
wt%

Zr a
wt%

SBET
m2 /g

Acidity b
meqN H3 /g

DcCeO2
nm

DdN iO
nm

DeN i
nm

Ni/CeZrS-1/3
Ni/CeZrS-1
Ni/CeZrS-3

6.8
7.0
6.6

6.0
4.2
2.4

2.3
4.5
6.6

243
242
238

0.062
0.077
0.088

6.3
5.0
3.4

19.1
18.1
17.6

19.7
16.6
15.7

a

ICP-AES measurements; b NH3 -TPD analysis; c calculated from the (220) reflection of CeO2 in
XRD; d calculated from the (200) reflection of NiO in XRD; e calculated from the (111) reflection
of Ni in XRD.

The XRD spectra of the calcined and reduced catalysts are shown in Figure 7, where samples
containing both Ce and Zr have a pattern similar to that of Ni/CeS sample (Figure 4), but CeO2
peaks become less intense and shift to higher 2θ values with increasing Zr loading. This fact
has been ascribed to the formation of a Cex Zr1−x O2 solid solution, where the unit cell parameter
changes with the Zr/Ce ratio [21,23,26]. The size of the CeO2 crystallites on calcined catalysts were
calculated from the width of the (220) characteristic peak. The values displayed in Table 5 show that
crystallites found on these catalysts are smaller than those reported for unsupported Cex Zr1−x O2 [21],
which evidences that the relatively high surface area of SiO2 enhances their dispersion, and their
sizes become smaller as the Zr/Ce ratio increases. Only a slight increment of these particles sizes
(≈ 0.1–0.3 nm) was detected after reduction treatment.
Additionally, the peaks characteristic of cubic NiO (JCPDS 78-0643) and those corresponding
to metallic Ni (JCPDS 70-1849) can be observed for the calcined and reduced samples, respectively.
Concerning the calcined samples, the mean NiO crystallites sizes, calculated by applying the Scherrer
equation to the broadening of the (200), are summarized in Table 5. The NiO crystallite size decreases

225
as the Zr content is increased, which has been previously reported for unsupported Cex Zr1−x O2
catalysts [15,20]. In the case of the reduced samples, the Ni crystallite sizes follow a similar trend.
It is noticeable that Ni crystallites of the reduced Ni/CeZrS-1 and Ni/CeZrS-3 catalysts are smaller
than the Ni particles of the corresponding calcined samples and even smaller than Ni crystallites of
Ni/S catalyst, Ni/CeZrS-3 being the catalyst with the smallest Ni crystallites.

Figure 7. XRD patterns of the Ni/Cex Zr1−x O2 /SiO2 catalysts: (a) calcined; (b) reduced.
The reducibility of these catalysts was determined by H2 -TPR experiments, and the
corresponding profiles shown in Figure 8 evidenced a similarity with those of the Ni/CeS and Ni/ZrS
samples. These profiles have a peak at a lower temperature due to the reduction of NiO particles and
a shoulder at a higher temperature ascribed to the reduction of nickel species in close contact with
the support. Quantitative analysis of the TPR profiles revealed total reduction of the Ni2+ species
for all of the samples. By increasing the ZrO2 content, the reduction profiles slightly shift towards
higher temperatures, and this peak area increases while the area under the low temperature zone
decreases. This fact indicates a stronger interaction of Ni2+ species with the support as ZrO2 loading
increases [15,18,19], which accounts for the hindered sintering of the Ni phase when Zr loading
increases, as observed from the difference between the mean size of the nickel phase crystallites
before and after reduction (see Table 5). The decrease of the Cex Zr1−x O2 particle size as the Zr
loading increases would make it possible for a higher contact between this phase and the NiO, which
results in higher metal-support interaction [16].

226

Figure 8. TPR profiles of the calcined Ni/Cex Zr1−x O2 /SiO2 catalysts.
2.3.2. Catalytic Results Using the Ni/Cex Zr1−x O2 /SiO2 Catalysts
These catalysts were tested at 600 ◦ C in the ethanol steam reforming, and the obtained results
are shown in Table 6. All of the catalysts showed complete ethanol conversion after 8 h of TOS.
On the other hand, it can be observed how product selectivity depends on the Zr/Ce ratio. The only
H-containing intermediate compound found among the reaction products was methane, while neither
ethylene nor acetaldehyde could be detected. Besides, as the Zr/Ce ratio is increased, selectivity
towards CH4 decreases, and thus, hydrogen selectivity increases. This may be attributed to both a
decrease of Ni crystallite size and an increase of thermal stability when the Zr content increases,
as determined by XRD and TPR (Table 5 and Figure 8). Regarding the CO2 /CO ratio, it is higher
for Ni/Cex Zr1−x O2 /SiO2 catalysts than for the Ni/S sample and increases with the Zr/Ce ratio. An
enhanced oxygen storage capacity due to the formation of the Cex Zr1−x O2 solid solution has been
reported to favor the water-gas shift reaction [23,26].
Table 6. Results on ethanol steam reforming obtained with the Ni/Cex Zr1−x O2 /SiO2
catalysts (atmospheric pressure; T: 600 ◦ C ; weight hourly space velocity (WHSV):
12.7 h−1 ; RH2O/EtOH : 3.7 mol/mol; TOS: 8 h).
Catalyst
Ni/CeZrS-1/3
Ni/CeZrS-1
Ni/CeZrS-3

XEtOH

Selectivity (mol%)

Coke

mol%

H2

CO2

CO

CH4

C2 H4

CH3 CHO

g/gcat h

100
100
100

84.6
85.6
85.8

52.4
54.0
54.3

29.6
29.0
28.9

18.0
17.0
16.8

0
0
0

0
0
0

0.500
0.443
0.308

Finally, Table 6 compares catalysts in terms of coke formation, which decreases as the
Zr/Ce mass ratio in the catalyst increases. However, since ceria and zirconia improve oxygen

227
mobility [32,41,42], favoring carbon gasification, coke deposition on promoted catalysts decreased
by increasing Zr content in the Cex Zr1−x O2 solution, which agrees with the smaller size of NiO and
Ni crystallites found in samples with a higher Zr content.
3. Experimental Section
3.1. Catalysts Preparation
A series of catalysts were prepared to study the effect of the addition of Ce and Zr to Ni/SiO2
catalyst (Ni: 7 wt%). The modified catalysts contain a total loading of Ce and Zr of 10 wt% and
Ce/Zr mass ratios of 0, 1/3, 1, 3 and ∞. Moreover, the order of the Ce or Zr incorporation with
SiO2 , before or after Ni addition, was also studied. The addition of metals was accomplished by
incipient wetness impregnation of amorphous silica (Ineos, Warrington, UK) with aqueous solutions
of the metal precursors: Ni(NO3 )2 ·6H2 O (Scharlab, Sentmenat, Barcelona, Spain), Ce(NO3 )3 ·6H2 O
(Scharlab, Sentmenat, Barcelona, Spain) and ZrO(NO3 )2 ·6H2 O (Sigma–Aldrich, St. Louis, MO,
USA). After the incorporation of each metal (Ni, Zr or Ce), the solid sample was air-dried overnight
and further calcined at 500 ◦ C for 5 h with a heating rate of 1.8 ◦ C /min. Afterwards, the addition
of the other metal, if needed, was carried out by following the same procedure. Ni/SiO2 , Ni/CeO2
and Ni/ZrO2 were prepared as reference materials by impregnation of Ni over commercial silica
(Ineos, Warrington, UK), ceria (Riedel-de-Haën, Hanover, Germany) and zirconia (Sigma–Aldrich,
St. Louis, MO, USA). Catalyst are denoted as A/BY-x, where A/B are the metals added to the support
(A, secondly, and B, firstly added), Y is the support (S: silica, C: ceria and Z: zirconia), and x is the
Zr/Ce mass ratio in the final sample.
3.2. Catalysts Characterization
X-ray powder diffraction (XRD) analysis was used to find out the crystalline phases and calculate
mean crystallite size in the calcined and reduced catalysts. Measurements were carried out on
a Phillips (Eindhoven, The Netherlands) X’Pert PRO diffractometer using Cu Kα radiation. The
patterns were recorded with a 2θ increment step of 0.020◦ and a collection time of 2 s. Mean metallic
crystallite diameters were calculated by applying the Scherrer equation.
The inductively-coupled plasma atomic emission spectroscopy (ICP-AES) technique was used
to determine the total Ce, Zr and Ni content in the catalysts using a Varian (Palo Alto, CA,
USA) VISTA-PRO AX CCD-Simultaneous ICP-AES spectrophotometer. Samples were previously
dissolved with an acidic solution (HF and H2 SO4 ).
Textural properties of the materials were calculated by nitrogen adsorption-desorption at 77 K in
a Micromeritics (Norcross, GA, USA) TRISTAR 2050 sorptometer. Samples were first outgassed
under vacuum at 200 ◦ C for 4 h. Surface areas were determined using the BET method.
Acid properties of the catalysts and supports were obtained by ammonia temperature programmed
desorption analysis (NH3 -TPD) in a Micromeritics (Norcross, GA, USA) Autochem 2910 equipment.
Samples were previously outgassed under a He flow (50 NmL/min) at 560 ◦ C for 30 min. After

228
saturation of the sample with ammonia and removal of the physisorbed fraction by flowing He at
180 ◦ C , NH3 was desorbed from the sample by increasing the temperature up to 550 ◦ C with a
heating rate of 15 ◦ C /min, keeping this temperature constant for 30 min.
The reducibility of supported metals was determined by hydrogen temperature programmed
reduction analysis (H2 -TPR). Catalysts were analyzed with the same apparatus described for
NH3 -TPD. Catalysts were previously degasified in flowing argon (35 mL/min) for 30 min at 110 ◦ C
with a heating rate of 15 ◦ C /min. Afterwards, the H2 -TPR profile was obtained by flowing 10% H2
in Ar (35 NmL/min) from 25 ◦ C to 800 ◦ C with a heating rate of 5 ◦ C /min.
The morphology of catalyst particles and metal crystallites was analyzed by transmission electron
microscopy (TEM). Samples were prepared by dispersion of the powdered material in acetone and
following deposition on a copper grid with carbon support. Micrographs were attained on a Phillips
(Eindhoven, The Netherlands) TECNAI 20 equipped with a LaB6 filament and an accelerating
voltage of 200 kV.
Carbon deposited on the used catalysts was evaluated by thermogravimetric analysis (TG).
TG measurements were carried out on a TA instruments (New Castle, DE, USA) SDT 2960
thermobalance using an air flow of 100 mL/min and a heating rate of 5 ◦ C /min up to 700 ◦ C .
3.3. Catalytic Test
Catalysts activity and selectivity were measured in the ethanol steam reforming reaction on a
MICRO ACTIVITY-PRO unit, as described elsewhere [19]. The catalyst was placed into the reactor
and in situ reduced under flowing pure hydrogen (30 NmL/min) for 4.5 h at 550 ◦ C with a heating
rate of 2 ◦ C /min. After activation, the catalytic tests were carried out at atmospheric pressure and
600 ◦ C . A liquid water/ethanol (3.7 molar ratio) stream (0.075 mL/min) was fed into the system,
vaporized at 150 ◦ C and further diluted by N2 (30 NmL/min). Weight hourly space velocity (WHSV),
defined as the water-ethanol mass flow rate related to the mass of the catalyst, was fixed at 12.7 h−1 .
The composition of the output gas stream was determined online by a gas chromatograph, Varian
(Palo Alto, CA, USA) CP-3380, equipped with Hayesep Q and Molecular Sieve 13X columns, a
thermal conductivity detector and using He as the carrier and reference gas. Condensable vapors
(ethanol, water and acetaldehyde) were retained in a condenser and analyzed by chromatography.
4. Conclusions
Ethanol has been successfully converted into hydrogen though steam reforming over Ni catalysts
supported on SiO2 , CeO2 , ZrO2 and Ce- and/or Zr-modified SiO2 with different Zr/Ce mass ratios.
The Ni/SiO2 catalyst led to higher hydrogen selectivity than Ni/CeO2 and Ni/ZrO2 , but it could
not maintain complete ethanol conversion. Incorporation of Zr to Ni/SiO2 increased the Ni-support
interaction, leading to complete ethanol conversion along time and higher hydrogen selectivity.
However, large amounts of coke were formed as a consequence of ethylene formation induced by
acid sites found on Zr-modified catalysts. On the contrary, lower coke formation was observed when
Ce was incorporated into Ni/SiO2 , which can be ascribed to the presence of oxygen vacancies in

229
CeO2 promoting coke gasification. Concerning the impregnation order, the incorporation of Ce and
Zr prior to nickel on the silica support led to steam reforming catalysts with better performance.
When both Zr and Ce were incorporated into Ni/SiO2 , a Cex Zr1−x O2 solid solution was formed,
increasing the Ni-support interaction and producing smaller Ni crystallites as the Zr/Ce mass ratio
increased. As a result, complete ethanol conversion and hydrogen selectivity above 84 mol% were
reached with the Ni/Cex Zr1−x O2 /SiO2 catalysts. Concretely, using a Zr/Ce mass ratio of three, a
trade-off between high selectivity to the main products, relatively low coke deposition and high
thermal stability was achieved.
Acknowledgments
The authors acknowledge the financial support from the Ministerio de Economía y
Competitividad through the project “Hydrogen and liquid fuels production from microalgae by
thermochemical proceses” (Ref: CTQ2013-44447-R), and from the Comunidad de Madrid through
the project “Residuals to Energy part 2” project (Ref: P2013/MAE-2882).
Author Contributions
This work was conceived by Jose Antonio Calles and Alicia Carrero. Montaña Lindo did the
experimental work, assisted by Arturo Javier Vizcaíno who also draft the first version of manuscript.
Furthermore, Arturo Javier Vizcaíno, Jose Antonio Calles and Alicia Carrero wrote and revised the
final version of paper.
Conflicts of Interest
The authors declare no conflict of interest.
References
1. Züttel, A.; Remhof, A.; Borgschulte, A.; Friedrichs, O. Hydrogen: the future energy carrier.
Phil. Trans. R. Soc. A 2012, 368, 3329–3342.
2. Guo, L.J.; Zhao, L.; Jing, D.W.; Lu, Y.J.; Yang, Y.J.; Bai, B.F.; Zhang, X.M.; Ma, L.J.;
Wu, X.M. Solar hydrogen production and its development in China. Energy 2009, 34,
1073–1090.
3. Olateju, B.; Kumar, A. Hydrogen production from wind energy in Western Canada for
upgrading bitumen from oil sands. Energy 2011, 36, 6326–6339.
4. Zahedi, A. Hydrogen as storage option for intermittent renewable technologies such as solar
and wind. In Proceedings of the 2006 Australasian Universities Power Engineering Conference
(AUPEC’06), Melbourne, Australia, 10–13 December 2006.
5. Goltsova, V.A.; Veziroglu, T.N.; Goltsova, L.F. Hydrogen civilization of the future—A new
conception of the IAHE. Int. J. Hydrogen Energy 2006, 31, 153–159.

230
6. Grzegorczyk, W.; Denis, A.; Gac, W.; Ioannides, T.; Machocki, A. Hydrogen Formation
via Steam Reforming of Ethanol Over Cu/ZnO Catalyst Modified with Nickel, Cobalt and
Manganese. Catal. Lett. 2009, 128, 443–448.
7. Vizcaíno, A.J.; Carrero, A.; Calles, J.A. Hydrogen Production from Bioethanol. In Hydrogen
Production: Prospects and Processes; Nova Science Publishers, Inc.: New York, NY, USA,
2012; pp. 247–294
8. Papadopoulou, E.; Delimaris, D.; Denis, A.; Machocki, A.; Ioannides, T. Alcohol reforming
on cobalt-based catalysts prepared from organic salt precursors. Int. J. Hydrogen Energy 2007,
37, 16375–16381.
9. Navarro, R.M.; Peña, M.A.; Fierro, J.L.G. Hydrogen production reactions from crbon
feedstocks: Fossil fuels and biomass. Chem. Rev. 2007, 107, 3952–3991.
10. Ni, M.; Leung, D.Y.C.; Leung, M.K.H. A review on reforming bio-ethanol for hydrogen
production. Int. J. Hydrogen Energy 2007, 32, 3238–3247.
11. Auprête, F.; Descorme, C.; Duprez, D. Bio-ethanol catalytic steam reforming over supported
metal catalysts. Catal. Commun. 2002, 3, 263–267.
12. Comas, J.; Mariño, F.; Laborde, M.; Amadeo, N. Bio-ethanol steam reforming on Ni/Al2 O3
catalyst. Chem. Eng. J. 2004, 98, 61–68.
13. Fatsikostas, A.N.; Verykios, X.E. Reaction network of steam reforming of ethanol over
Ni–based catalysts. J. Catal. 2004, 225, 439–452.
14. Frusteri, F.; Freni, S.; Spadaro, L.; Chiodo, V.; Bonura, G.; Donato, S.; Cavallaro, S. H2
production for MC fuel cell by steam reforming of ethanol over MgO supported Pd, Rh, Ni and
Co catalysts. Catal. Commun. 2004, 5, 611–615.
15. Sánchez-Sánchez, M.C.; Navarro, R.M.; Fierro, J.L.G. Ethanol steam reforming over
Ni/M x Oy -Al2 O3 (M = Ce, La, Zr and Mg) catalysts: Influence of support on the hydrogen
production. Int. J. Hydrogen Energy 2007, 32, 1462–1471.
16. Calles, J.A.; Carrero, A.; Vizcaíno, A.J. Ce and La modification of mesoporous Cu-Ni/SBA-15
catalysts for hydrogen production through ethanol steam reforming. Micropor. Mesopor. Mater.
2009, 119, 200–207.
17. Vizcaíno, A.J.; Carrero, A.; Calles, J.A. Ethanol steam reforming on Mg- and Ca-modified
Cu-Ni/SBA-15 catalysts. Catal. Today 2009, 146, 63–70.
18. Lindo, M.; Vizcaíno, A.J.; Calles, J.A.; Carrero, A. Ethanol steam reforming on Ni/Al-SBA-15
catalysts: Effect of the aluminium content. Int. J. Hydrogen Energy 2010, 35, 5895–5901.
19. Vizcaíno, A.J.; Lindo, M.; Carrero, A.; Calles, J.A. Hydrogen production by steam reforming
of ethanol using Ni catalysts based on ternary mixed oxides prepared by coprecipitation. Int. J.
Hydrogen Energy 2012, 37, 1985–1992.
20. Vizcaíno, A.J.; Carrero, A.; Calles, J.A. Hydrogen production by ethanol steam reforming over
Cu-Ni supported catalysts. Int. J. Hydrogen Energy 2007, 32, 1450-1461.
21. Biswas, P.; Kunzru, D. Steam reforming of ethanol for production of hydrogen over
Ni/CeO2 -ZrO2 catalysts: Effect of support and metal loading. Int. J. Hydrogen Energy 2007,
32, 969–980.

231
22. Biswas P.; Kunzru, D. Oxidative steam reforming of ethanol over Ni/CeO2 -ZrO2 catalyst.
Chem. Eng. J. 2008, 136, 41–49.
23. Youn, M.H.; Seo, J.G.; Cho, K.M.; Park, S.; Park, D.R.; Jung, J.C.; Song, I.K.
Hydrogen production by auto-thermal reforming of ethanol over nickel catalysts supported on
Ce-modified mesoporous zirconia: Effect of Ce/Zr molar ratio. Int. J. Hydrogen Energy 2008,
33, 5052–5059.
24. Wan, H.; Li, X.; Ji, S.; Huang, B.; Wang, K.; Li, C. Effect of Ni loading and Cex Zr1−x O2
promoter on Ni-based SBA-15 catalysts for steam reforming of methane. J. Nat. Gas Chem.
2007, 16, 139–147.
25. Rocchini, E.; Vicario, M.; Llorca, J.; Leitenburg, C.; Dolcetti, G.; Trovarelli, A. Reduction and
oxygen storage behavior of noble metals supported on silica-doped ceria. J. Catal. 2002, 211,
407–421.
26. Raju, V.; Jaenicke, S.; Chuah, G.K. Effect of hydrothermal treatment and silica on thermal
stability and oxygen storage capacity of ceria-zirconia. Appl. Catal. B 2009, 191, 92–100.
27. Reddy, B.M.; Thrimurthulu, G.; Saikia, P.; Bharali, P. Silica supported ceria and ceria-zirconia
nanocomposite oxides for selective dehydration of 4-methylpentan-2-ol. J. Mol. Catal. A 2007,
275, 167–173.
28. Reddy, B.M.; Saikia, P.; Bharali, P.; Katta, L.; Thrimurthulu, G. Highly dispersed ceria and
ceria-zirconia nanocomposites over silica surface for catalytic applications. Catal. Today 2009,
141, 109–114.
29. Takahashi, R.; Sato, S.; Sodesawa, T.; Yoshida, M.; Tomiyama, S. Addition of zirconia in
Ni/SiO2 catalyst for improvement of steam resistance. Appl. Catal. A 2004, 273, 211–215.
30. Seo, J.G.; Youn, M.H.; Song, I.K. Effect of SiO2 -ZrO2 supports prepared by a grafting
method on hydrogen production by steam reforming of liquefied natural gas over Ni/SiO2 -ZrO2
catalysts. J. Power Sources 2007, 168, 251–257.
31. Roh, H.S.; Potdar, H.S.; Jun, K.W. Carbon dioxide reforming of methane over co-precipitated
Ni-CeO2 , Ni-ZrO2 and Ni-Ce-ZrO2 catalysts. Catal. Today 2004, 93–95, 39–44.
32. Trovarelli, A. Catalysis by Ceria and Related Materials; Imperial College Press: London,
UK, 2002.
33. Djinovic, P.; Batista, J.; Pintar, A. WGS reaction over nanostructures CuO-CeO2 catalysts
prepared by hard template method: Characterization, activity and deactivation. Catal. Today
2009, 147, 191–197.
34. Andreeva, D.; Ivanov, I.; Ilieva, L.; Sobczak, J.W.; Avdeev, G.; Tabakova, T. Nanosized gold
catalysts supported on ceria and ceria-alumina for WGS reaction: Influence of preparation
method. Appl. Catal. A 2007, 333, 153–160.
35. Trim, D.L. Coke formation and minimisation during steam reforming reactions. Catal. Today
1997, 37, 233–238.
36. Carrero, A.; Calles, J.A.; Vizcaíno, A.J. Hydrogen production by ethanol steam reforming
over Cu-Ni/SBA-15 supported catalysts prepared by direct synthesis and impregnation.
Appl. Catal. A 2007, 327, 82–94.

232
37. Wang, S.; Lu, G.Q. Reforming of methane with carbon dioxide over Ni/Al2 O2 catalysts: Effect
of nickel precursor. Appl. Catal. 1998, 69, 271–280.
38. Juan-Juan, J.; Román-Martínez, M.C.; Illán-Gómez, M.J. Effect of potassium content in the
activity of K-promoted Ni/Al2 O2 catalysts for the dry reforming of methane. Appl. Catal.
2006, 301, 9–15.
39. Carrero, A.; Calles, J.A.; Vizcaíno, A.J. Effect of Mg and Ca addition on coke deposition over
Cu-Ni/SiO2 catalysts for ethanol steam reforming. Chem. Eng. J. 2010, 163, 395–402.
40. Helveg, S.; Sehested, J.; Rostrup-Nielsen, J.R. Whisker carbon in perspective. Catal. Today
2011, 178, 42–46.
41. Natesakhawat, S.; Watson, R.B.; Wang, X.; Ozkan, U.S. Deactivation characteristics of
lanthanide-promoted sol-gel Ni/Al2 O2 catalysts in propane steam reforming. J. Catal. 2005,
234, 496–508.
42. Bellido, J.D.A.; Assaf, E.M. Nickel catalysts supported on ZrO2 , Y2 O3 -stabilized ZrO2 and
CaO-stabilized ZrO2 for the steam reforming of ethanol: Effect of the support and nickel load.
J. Power Sources 2008, 177, 24–32.

233

Deactivation Pattern of a “Model” Ni/MgO Catalyst in the
Pre-Reforming of n-Hexane
Giuseppe Trunfio and Francesco Arena
Abstract: The deactivation pattern of a “model” Ni/MgO catalyst in the pre-reforming of
n-hexane with steam (T, 450 °C; P, 5–15 bar) is reviewed. The influence of the steam-to-carbon ratio
(S/C, 1.5–3.5) on the rate of catalyst fouling by coking is ascertained. Catalyst fouling leads to an
exponential decay in activity, denoting 1st-order dependence of the coking process on active sites
availability. Hydrogen hinders the coking process, though slight activity decay is due to sintering of
the active Ni phase. Deactivation by thiophene causes a sharp, almost linear, drop to nearly zero
activity within only 6 h; this deactivation is likely due to dissociative adsorption of thiophene with
subsequent strong, irreversible chemical adsorption of S-atoms on active Ni sites, i.e., irreversible
poisoning. Modeling of activity decay curves (Į, at/a0) by proper kinetic equations allows assessing
the effects of temperature, pressure, S/C, H2 and thiophene feed on the deactivation pattern of the
model Ni/MgO catalyst by coking, sintering, and poisoning phenomena.
Reprinted from Catalysts. Cite as: Trunfio, G.; Arena, F. Deactivation Pattern of a “Model”
Ni/MgO Catalyst in the Pre-Reforming of n-Hexane. Catalysts 2014, 4, 196-214.
1. Introduction
There is a strong commitment worldwide to enhance the practical and economical feasibility
of the steam reforming (SR) technology through efficiency improvement and feedstock versatility,
which could result in significant improvements in fuel processing economy [1–7]. In this context,
pre-reforming constitutes an established technology with economic and operational benefits on the
overall syngas production, representing an important tool, especially for the revamping of older SR
plants [2–4,8–10]. A pre-reforming unit consists in a tubular adiabatic reformer, allowing whatever
hydrocarbon feed (NG, LPG, VN) be converted to CH4 and COx at low temperature, typically in
the 450–550 °C range, with many practical advantages consisting of [2–4]:
i.
ii.
iii.
iv.

an increased production capacity with smaller reformer furnaces;
a higher feedstock flexibility;
enhanced SR tube and catalyst lifetime;
design of innovative process configuration for low energy consumption and investment costs.

Furthermore, the pre-reforming unit drastically limits the risk of carbon formation inside tubular
reformers, ensuring a total conversion of higher hydrocarbons, and allowing, then, SR operations at
lower steam-to-carbon (S/C) ratio, with a higher preheat temperature and heat flux and capacity [9–11].
Moreover, the pre-reforming catalyst ensures the total elimination of S-containing impurities,
contributing to enhance SR catalysts’ lifetime [2–5,8,11–14].
Considering the importance for the improvement of syngas producing plants, little previous
literature deals with the basic aspects of the steam reforming of higher hydrocarbons reaction under

234
low temperature and high pressure conditions [12,14–17]. In fact, only after the pioneering work of
Rostrup-Nielsen on the steam reforming of C1-C7 alkanes on magnesia-supported metal catalysts,
which led to the development of the SPARG process [2,10,14], has light been shed on the reaction
mechanism and kinetics, leading to further improvements in catalysts’ performance [11,14,16,17].
Although noble-metal-based systems feature a high reactivity in the reforming of superior
hydrocarbons to methane and carbon oxides, typical formulations of pre-reforming catalysts
include nickel as the active phase [6–8,16,17]. Due to different operating conditions, pre-reforming
catalysts are generally richer in Ni loading (50–70 wt%), with the addition of rare earth elements
(e.g., La, Sm, Ce, Y, Zr, etc.) as promoters, and are characterized by larger total (SA) and metal
surface area (MSA) than SR counterparts [6,16]. Loading of the active phase and type of
promoters, however, depend upon process characteristics and targets [6,16].
Generally, coking processes on Ni-based catalysts in steam reforming and methanation reactions
proceed through common carbon intermediates leading to the formation of several types of deposits
depending on operating conditions [2,8,12–15,18–24]. Moreover, effects of metal sintering on the
active surface area of typical pre-reforming catalysts have been assessed [25]. Catalyst poisoning is
generally caused by strong chemisorption of sulfur impurities present in the feed stream as
inorganic and/or organic sulfides in most naturally occurring feedstock; sulfur poisoning depends
on many variables including the kind of sulfur compound, nature of the catalyst and operating
conditions [7,26,27].
In earlier papers, we documented that Ni/MgO catalysts obtained by an original non-aqueous
preparation route exhibit a remarkable performance in both “steam” and “dry” reforming of
methane [28–30]. Hindering the formation of the NiMgOx solid solution, indeed, such a preparation
procedure allows tuning both physico-chemical and catalytic properties through a proper selection
of Ni loading, calcination and reduction treatments [28,29,31].
This review provides an overview of reaction and deactivation networks occurring early (i.e.,
during the first 17 h) on a “model” Ni/MgO catalyst in the steam reforming of n-hexane under
pre-reforming conditions (T, 450 °C; P, 5–15 atm). Rates, selectivities, rate constants, and
equilibrium constants of important reaction paths as a function of time and rate constants for
deactivation by sintering, coking, and poisoning from the previous literature (with an emphasis on
the authors' previous works) are presented, and the effects of catalyst deactivation on activity,
selectivity, and stability are analyzed and discussed.
2. Results and Discussion
2.1. Effects of the Experimental Conditions on the Activity-Selectivity-Stability Pattern
Activity data of the “model” Ni/MgO catalyst in the pre-reforming of n-hexane at standard
reaction conditions (T, 450 °C; P, 10 bar; PC6H14, 0.18 bar; PH2O, 3.0 bar; S/C, 2.8) are shown in
Figure 1 in terms of n-hexane conversion (Ȥ) and product selectivity vs. reaction time (t.o.s.). The
only detected products are CH4, CO2, and CO [11–17] with a relative distribution which depends
on the conversion level. After 30 min, taken as the initial reaction time (t0, a0), the hexane
conversion is §70% with CH4, CO2 and CO selectivities of ca. 55, 43 and 2%, respectively. A drop

235
in conversion of about 50% over a period of only 17 h involves a SCH4 decrease from 55 to 30%,
counterbalanced by a rise in SCO2 from 45 to 70%, although no significant changes in SCO (2–3%)
are observed.
Figure 1. n-hexane conversion (Ȥ) and product selectivity (SȤ) vs. t.o.s. at standard
reaction conditions (adapted from ref. [11] with the kind permission of Elsevier,
copyright 2004).
0.8

χ
SCO2

χ - SX

0.6

SCH4

0.4

0.2
SCO
0.0
0

2

4

6

8

10

12

14

16

18

t.o.s. (h)

Activity data at 450 °C with varying space velocity (GHSV), feed stream composition (PH2O and
PH2) and pressure are summarized in Table 1 in terms of initial and final hexane conversion (Ȥ, %),
reactor output (molC6H14·hí1·gcatí1), product selectivity (SȤ, %) and weight gain of the “used”
catalyst samples (WC, g·gcatí1) due to carbon deposition and/or coke formation, determined by
TGA-DSC measurements. In addition, the influence of experimental conditions on the “relative
activity” (Į = Ȥ/Ȥ0, the ratio between conversion at the time “t” and “t0” respectively) is shown in
Figure 2.
Table 1. Initial and final (17 h) activity data and amount of carbon on the “used”
catalyst samples under different reaction conditions (T, 450 °C; pC6H14/p0, 0.018).
Run

P
(atm)

S/C

GHSV a
(hí1)

ȤC6H14
(%)

Reactor output
(molC6H14 hí1 gcatí1)

SCH4
(%)

SCO2
(%)

SCO
(%)

SC2–C5
(%)

WC b
(g gcatí1)

a
b
c
d
ec
f
g
hd

10
10
10
10
10
15
5
10

1.5
2.8
3.5
2.8
2.8
2.8
2.8
2.8

12,000
12,000
12,000
24,000
24,000
12,000
12,000
24,000

53–08
71–25
76–34
38–11
77–73
69–37
67–21
65–00

0.26 – 0.04
0.35–0.12
0.38–0.17
0.38–0.11
0.77–0.73
0.51–0.27
0.17–0.05
0.65–0.00

55–9
56–28
56–42
39–5
98–97
60–50
49–25
97–97

43–88
42–69
43–57
59–93
1–2
39–49
49–73
2–2

2–3
2–3
1–1
2–2
0–0
1–1
1–2
1–1

1–1
-

1.32
0.96
0.89
1.23
0.03
0.75
1.02
0.21

a

FC6H14/Vcat;

d

b

hydrogen (H2/C, 1) and thiophene (50 ppm on C6H14) in the feed.

weight gain due to carbon deposition and/or coke formation;

c

hydrogen (H2/C, 1) in the feed;

236
Data in Table 1 indicate that feed composition and operating conditions affect both hexane
conversion and CH4, CO2 and CO selectivities, while trace amounts of C2-C5 hydrocarbons
(SC2-C5 § 1%) are detected only in presence of hydrogen (run e).
Figure 2. Relative activity vs. t.o.s. at different experimental conditions (see Table 1).
1.0
0.8
0.6

Legend

α

run a
run b
run c
run d
run e
run f
run g
run h

0.4
0.2
0.0
0

2

4

6

8
10
t.o.s. (h)

12

14

16

18

At a constant pressure of 10 atm, the most obvious effects of the S/C ratio increase from 1.5 to
3.5 (runs a–c) consist in a progressive increase of hexane conversion (from 53 to 76%), indicating a
positive effect of water pressure on the reactor output, and a better catalyst stability (Figure 2)
coupled to a relative influence on product selectivity during reaction time (Table 1). The pressure
increase from 5 to 15 atm (S/C, 2.8) has no significant effect on the conversion level (runs b, f, g),
as an increase from 0.17 to 0.51 mmolC6H14āhí1āgcatí1 (Table 1) mirrors a fairly constant pressurenormalized reactor output (0.035 molC6H14āhí1āgcatí1/atm).
Figure 3. Overall conversion-selectivity (SȤ) data and H/Oout ratio at different
conditions (see Table 1) and reaction times (adapted from ref. [15] with the kind
permission of Wiley, copyright 2006).
1.0

200
hydrogen co-feed

0.8

S CO2

100
4

H/Oo ut

SX

3

0.4

S CH4

2

0.2

S CO

1

0.0

0
0.0

0.2

0.4

χ C6H14

0.6

0.8

H/Oout

0.6

237
In addition, an enhancement of SCH4 (49–60%) and, even more so, of catalyst stability are
evident, especially at 15 atm (run f).
Furthermore, a twofold increase of GHSV from 12,000 to 24,000 hí1 (runs b, d) causes an
almost proportional decrease of the initial conversion (38%) according to an unchanging reactor
output (0.35–0.38 molC6H14āhí1āgcatí1), a drop in SCH4 from 55 to 39% with a corresponding growth
of SCO2 (Table 1), and a considerably stronger activity decay (Figure 2).
In all the cases, a decreasing trend of the relative activity curves denotes noticeable deactivation
phenomena, the extents of which depend on S/C ratio, P, and GHSV (Figure 2). The observed loss
of activity can be largely accounted for by the weight loss recorded by TGA-DSC analysis of
“used” catalyst samples (Table 1) during combustion of the carbonaceous deposits. This points to
carbon formation and coking as the main causes of catalyst deactivation [11,15]. However, such
experiments did not allow to differentiate between carbon and coke nor quantitatively measure CO2
and H2O.
Under simulated pre-reforming conditions (run e) with H2 in the feed [3–6,16–18], a markedly
higher initial conversion (77%) and a two-fold higher reactor output are observed relative to the
standard tests (runs b, d). Moreover, the addition of hydrogen (H2/C = 1) to the feed increases SCH4
to 97–98% and greatly improves catalyst stability, with only a minor (ca. 5%) activity loss during
17 h of t.o.s. (Figure 2, run e).
All the conversion-selectivity data at different conditions (see Table 1) and reaction times
provide some general relationships showing that the selectivity pattern depends only on the activity
level (Figure 3). Thus, SCH4 increases steadily with increasing conversion, reaching an asymptotic
value of 55% in the conversion range of 60–80%, while SCO2 decreases from 90 to 45% and SCO
from 2 to 0.5% over the full range of measured conversion (8–75%). Changes in CO/CO2 are consistent
with the equilibrium composition predicted from the WGS reaction [11,15,16,18,19,32–37]. The
overall result of these changes in product selectivities is a trend of increasing output
hydrogen/oxygen atomic ratio (H/Oout) from zero at 8% conversion to a maximum value of 2.4 at
67% conversion (Figure 3), which is essentially the same as the H/C ratio of the hexane molecule
of 2.33. The apparent increase in the “hydrogenation” functionality with increasing conversion
level is consistent with a reaction network involving (1) cracking of hexane; (2) water dissociation;
(3) primary formation of CO, which is further transformed to CO2 via the water-gas-shift
(WGS); and (4) CH4 via methanation (MET) paths respectively [11,15], according to the following
surface reaction network, which includes two irreversible (1 and 8) and six reversible (2–7)
steps [11,13–16,19,21,33,35].
The above reaction scheme explains the inhibition of the methanation path (5) at low PCO and
PH2, that is the case at low conversion, while the excess of steam favors the formation of CO2 via
the WGS path (6). Thus, hydrogen “availability”, determined by the extent of hexane conversion, is the
key-parameter controlling the CH4/CO2 distribution [38]. Therefore, enabling the gasification of
coke-precursor species [11,15–17,21,22,35,36], the methanation step is the “competitor” process of
coking, explaining the high catalyst stability and the negligible weight gain of the used catalysts under
H2 co-feeding (Table 1). In fact, the same carbon-intermediate (CĮ) can undergo methanation at high H2

238
concentrations or, at low H2 concentrations, generate the following less reactive carbon species,
precursors of C-deposit buildup [19–22,35,36], by ageing
•
•
•
•
CRK

amorphous carbon (Cȕ),
vermicular carbon (CV),
bulk Ni carbide (CȖ), and
crystalline, graphitic carbon (CC).
C6H14 + n *

ĺ
KWS

WS

H2O + p *

'

HC

H2 + 2 *

'

GAS

CĮ(s) + O(s)

MET

CĮ(s) + 4 H(s)

WGS

CO(s) + O(s)

DES
COK

CO2(s)

KH2

KGAS

'

KMET

'

KWGS

'

KCO2

'

(6 í m) CĮ(s) + m CHx(s) + [(14 í x)/2] H2(g)

(1)

H2(g) + O(s)

(2)

2 H(s)

(3)

CO(s) + (n/6 + m - q) *

(4)

CH4(g) + (n/6 + 4 í t) *

(5)

CO2(s) + (m + q í s) *

(6)

CO2(g) + s *

(7)

kCOK

z CĮ  Cȕ (“encapsulating” + “wiskers”) – w *

(8)

2.2. Modeling of Deactivation Phenomena
From a mechanistic point of view, the coking process can be viewed as a consequence of the
surface accumulation of carbon moieties due to a lower gasification rate relative to the
accumulation rates of Calpha, polymeric carbon and coke, by the CO decomposition (scheme 1a) or
hexane cracking (scheme 1b).
Scheme 1. Carbon formation by CO decomposition (a) or hexane cracking (b).

C6H14

k0

k4

k

5

coke
k1

C

k-1

CO

k2
k-2

CO2

k3

k-3

CH4

a

b

TEM pictures of the used catalyst samples in Figure 4 confirm that catalyst deactivation, feeding
or not hydrogen, proceeds by different paths to form different inactive carbon or coke species.
According to previous TGA results, under standard reaction conditions (run b) the catalyst grains
(Figure 4B) appear embedded into an array of large carbon fibers (dfiber = 20–50 nm), while no

239
carbon deposits are apparent on the used catalyst in presence of H2 (Figure 4C) [11,15,17,35],
although halos surrounding several crystallites may be due to films of coke precursors.
Figure 4. TEM views (A–C) and PSD (D) of the “fresh” and “used” catalyst samples:
(A) “fresh”; (B) “used” (run b); (C) “used” (run e); (D) CSD’s of samples reproduced
in panels A–C.

(A)

(B)
35

percentage

30

(sample a) dAV=10.2 nm
(sa mple b) dAV=10.9 nm
(sample c) dAV=11.2 nm

25
20
15
10
5
0
1

(C)

2

3

4

5

d (n m)

6

7

8

(D)

It should be emphasized that the principal cause of deactivation in prereforming or reforming of
C2+ hydrocarbons at relatively low reaction temperatures is the formation of a hydrocarbonaceous
film on the nickel surface [19] which poisons nickel sites. This work establishes that deactivation is
minimized at high H2 concentrations.
In both Runs b and e, an observed “smoothing” of the cubic habit of the magnesia carrier
(Figure 4B,C) is consistent with moderate restructuring of the support, probably due to reaction
with steam [15,16]; this is likely the origin of very modest metal sintering, apparent from minor
changes in Ni crystallite size distributions (Figure 4D), accounting for ca. a 10% decrease in metal
dispersion in agreement with H2 uptake data [11].
As might be expected, thiophene co-feeding causes the most rapid decrease in reaction rate with
a linear drop during the first 6 h, (Figure 2, run h) [7]. The amount of C5H5S fed during this time
(0.05 g/kg hexane) corresponds to ca. 3 ȝmol (0.51 ȝmol/h), while the H2 uptake of the fresh

240
catalyst (120 μmol/gcat) corresponds to a surface Ni atom concentration for a 0.025 g sample of 6.0
μmol (i.e., 240 μmol/gcat of surface Ni sites assuming a H/Ni adsorption stoichiometry of 1 [7,15]).
Thus, during pre-reforming with the above specified feed of thiophene (run h), the maximum S/Ni
surface atomic ratio (i.e., average chemisorption stoichiometry) would be 0.5, assuming complete
and irreversible adsorption of sulfur atoms, which is likely under these conditions; in other words,
each thiophene molecule deactivates two Ni atoms. However, due to the rapid and irreversible
adsorption of S atoms on surface Ni atoms, the rate of poisoning could be influenced by pore
diffusional resistances; thus the amount of adsorbed S in the catalyst sample would drop off sharply
at the front of the bed and the sharp interface of sulfur saturated nickel would travel slowly, similar
to a chromatographic wave, through the catalyst bed. Hence, the equations modeling thiophene
adsorption are very complex and by necessity would include expressions for diffusion into the pores,
adsorption, dissociation, and the slowly moving axial concentration gradient through the bed, as
illustrated by the sophisticated two-dimensional model of H2S adsorption on a Ni steam reforming
catalyst described by Rostrup-Nielsen [39].
From the previous discussion, the deactivation pattern of the Ni/MgO catalyst can be assessed
taking into account the overall effects of coking, sintering, and poisoning. However, other than Run
h, the behavior of deactivation versus time observed in Figures 2 and 5 is due primarily to the
formation of carbon and coke, since negligible sulfur was present and effects of sintering were
negligibly small. With the exception of Run h, in the presence of thiophene, deactivation was
log-linear (see Figure 5), consistent with an exponential activity decay
α = exp[-kdeactāt]

(9)

Accordingly, the first order deactivation rate equation is
−

da
= k deact ⋅ a or - lnα = k deact ⋅ t
dt

(10)

These can be also expressed in terms of active sites availability (Csite), assuming that activity is
directly proportional to the concentration of surface sites [20],
(11)
C
α = site
0
C site
and, thus

-

dC site
= k deact ⋅ C site
dt

or

§C
- ln ¨¨ site
0
© C site

·
¸¸ = k deact ⋅ t
¹

(12)

where Csite and C°site are the instantaneous and initial concentrations of active sites, respectively.
Hence, the exponential decay of activity by coking is described by the equation

α = e ( cok )
− k

⋅t

(13)

which accounts for the general exponential-decay of catalytic activity for each of the runs in
Figure 2. From the slopes of the linear fits of the deactivation rate data in Figure 5, the deactivation
rate constants (kdeact) listed in Column 3 of Table 2 are obtained. It is further assumed that the rate

241
of sintering (which depends mainly on temperature) is negligibly small for all runs given the very
small change in crystallite diameter during reaction.
Figure 5. Plot of the relative activity at different experimental conditions (see Table 1)
vs. t.o.s..
0

1.0

-2

Legend

0.5

run a
run b
run c
run d
run e
run f
run g
run h

-3

-4
0

2

α

ln α

-1

0.0
4

6

8
10
t.o.s. (h)

12

14

16

18

Table 2. Deactivation kinetic constants for coking.
Run
a
b
c
d
e
f
g

r2
1.00
1.00
0.99
0.97
0.96
0.98
0.97

kcok a (hí1)
1.1 × 10í1 ± 2 × 10í3
5.5 × 10í2 ± 6 × 10í4
7.0 × 10í2 ± 1 × 10í3
7.7 × 10í2 ± 5 × 10í3
3.0 × 10í2 ± 1 × 10í3
6.3 × 10í2 ± 3 × 10í3
a kcok= kdeact.

It is emphasized that the calculation of rate constants in Table 2 (Run h involving sulfur
poisoning excepted) relies on the following assumptions:
1

2
3

The rates of deactivation by coking and sintering are small relative to the rate of the main
reaction. Hence the pseudo steady-state approximation can be made. This is a reasonable
assumption in view of 2–3 orders of magnitude differences between reaction and coking
rates at any condition and t.o.s. [15];
The rate of coking is slow enough that it is not influenced by a pore diffusional resistance
for the deactivation process;
Pore diffusion resistance for the main reaction is neglected which, at high reaction rates and
conversions, would lower the concentration of hydrocarbon coke precursors into the pores,
then lowering coking rates along the pores; similarly concentration gradients through the
bed due to high conversions should cause a decrease in adsorbed carbon concentration
through the bed. This is probably a poor assumption in view of the high conversions of this
study. Thus, the rate constants could be significantly in error due to differences in pore

242
diffusional resistance between runs and represent at best rough averages across pellets
and bed.
Figure 6. Log-plot of the kinetic constant of coking vs. water partial pressure (A); and
Relationship between the weight gain of the used samples (Table 1) and the kinetic
constant of coking (Table 2) at different reaction conditions (B).
1.4

-2.0

WC (g⋅gcath-1)

ln kdeact /(h-1)

slope, -1.10 ± 0.05
-2.5

-3.0

Legend
Effect of S/C ratio
Effect of Pressure
Effect of GHSV × 2

1.2

1.0

0.8

0.6

-3.5
-3.0

-2.8

-2.6

-2.4

ln PH2O/(atm-1)

(A)

-2.2

-2.0

0.03

0.06

0.09

0.12

kcok (h-1 )

(B)

Values of kcok in Table 2 are not constant as predicted by the simple first order deactivation
model; indeed, they vary by a factor of 20. In fact, the log-plot of kcok (Figure 6A) discloses an
inverse dependence (í1.10 ± 0.05) on PH2O, confirming that water concentration is a key-factor in
controlling coking rate at low H2 concentrations (typical of steady-state), a logical result given that
water is also a gasifying agent for carbon. The straight-line relationship between kcok and weight
gain in the used samples in Figure 6B supports a link between deactivation at low H2
concentrations to fouling by C-deposits for a range of typical prereforming conditions.
2.3. Effect of Deactivation on Product Selectivity and Surface Functionalities
Thermodynamic and kinetic evaluations of the outlet stream composition [2,9,15,21–23] shed
light on driving force for carbon deposition and, in turn, how carbon deposits influence selectivity.
In particular, assuming hexane cracking to be irreversible [9,15,22], a rough estimate of the
outlet stream composition can be made from a thermodynamic analysis of gasification (GAS),
water-gas-shift (WGS), and methanation (MET) reactions (Table 3) [9–11,15]. However, also the
Boudouard reaction has been considered to evaluate the incidence of CO dismutation on carbon
formation and coking [11]. Specifically, Px values obtained from conversion-selectivity data were
substituted in the equilibrium expressions to calculate experimental Kexp values of MET, GAS,
WGS and DISP reactions as a function of t.o.s. [11].

243
Table 3. Thermodynamic constants of the various reactions of the steam reforming
network (adapted from ref. [11] with the kind permission of Elsevier, copyright 2004).
Reaction
CRK
GAS
WGS
MET
DISP

C6H14
C + H2 O
CO + H2O
C + 2H2
2 CO

Stoichiometry
Œ
'
'
'
'

6 C + 7 H2
CO + H2
CO2 + H2
CH4
C + CO2

KP a
5.09 × 10í4
7.47 × 100
5.41 × 101
1.47 × 104

Equilibrium pressure constant values at 450 °C with Px values in MPa.

The ratio (ȕ) of Kexp to the corresponding calculated value of the equilibrium constant Keq at
450 °C (from Table 3) provides a measure of the distance from equilibrium for each reaction,
allowing a reliable assessment of the effects of the S/C ratio (Figure 7A–C) and H2 feeding
(Figure 7D) on activity versus time patterns. Figure 7 shows that irrespective of t.o.s. and S/C ratio,
the WGS reaction is always near thermodynamic equilibrium conditions, given that ȕWGS values
vary between 0.6 and 1.0. In contrast, gasification (ȕGAS ≈ 0.5) and, even more so, methanation
(ȕMET, 0.2–0.05) reactions operate far from equilibrium and are thus kinetically controlled. It
follows from the previous discussion that lower rates of GAS and MET relative to hexane cracking
constitutes a major driving force to carbon deposits build-up [17,19,20,35]. Furthermore, the
observation that ȕDISP is always larger than one implies a [PCO2/(PCO)2] ratio greater than that
allowed by thermodynamics, indicating that under prereforming conditions on Ni/MgO, the
forward reaction, which also contributes to carbon formation, is more kinetically favorable than the
reverse gasification of carbon by CO2. However, this deviation could also be an artifact of the
approximate reaction model.
It is interesting that while ȕGAS and ȕMET decrease with trends similar to that of the relative
activity (Į), ȕDISP increases progressively with t.o.s. (almost exponentially under the conditions of
Figure 7A), suggesting that carbon and coke reactivity decrease accordingly by ageing (i.e., by
conversion of coke precursors to highly polymerized coke and by conversion of CĮCȕ, CV, CȖ,
CC) [16,17,19,21,36]. In fact, the changes in product selectivity with the activity level (Figure 3)
mirror the different influences of coke and carbon on different functionalities of Ni catalysts,
consistent with the 14CO and 14CO2 isotopic labeling experiments of Jackson et al. who found that
Ni/Al2O3 catalyst retains high activity towards the scrambling of C-atoms and surface carbon
gasification, despite a considerable decrease in the availability of metal surface area due to growth
of gum-like carbonaceous deposits [35].
In the present paper, the role of the methanation reaction in modeling of prereforming has been
emphasized. Bartholomew emphasized that methanation and steam reforming functionalities of Ni
catalysts are closely related on the basis of the activation energy values [22]. Furthermore, RostrupNielsen stressed that hydrogenolysis and steam reforming reactions share many surface steps, requiring
also analogous ensembles (n, 2.5–2.7) of active sites [17]. Hence, effects of deactivation on
prereforming catalyst functionalities can be assessed by inspecting the trends of the relative rates of
MET, GAS, and WGS reactions (Įx = Į⋅SȤ), as shown in Figure 8. The semi-log plot of the relative
rates of CH4 (ĮCH4), CO (ĮCO), and CO2 (ĮCO2) formation provides satisfactory linear trends

244
(Figure 8), the slopes of which (kMET, kGAS, kWGS) represent the decay constants for the above
functionalities, listed in Table 4.
Figure 7. Pre-reforming of the n-hexane (T, 450 °C; P, 10 bar). Effect of S/C ratio and
H2 feeding on α and ȕ values of GAS, WGS, MET and DISP reactions vs. t.o.s.:
(A) S/C, 1.5; H2/C, 0; (B) S/C, 2.8; H2/C, 0; (C) S/C, 3.5; H2/C, 0; (D) S/C, 2.8; H2/C,
2. Legend: (z) α; (z) βWGS; (c) βDISP; (z) βGAS; (z) βMET (reproduced from ref. [11]
with the kind permission of Elsevier, copyright 2004).
23

23

22

22

21

21

20

β

2-1
2-2

α, β

α, β

20

2- 1
2- 2
2- 3

2-3

2- 4

2-4

2- 5

2-5

2- 6

2-6
0.25

0.5

1

2

4

8

16

0.25

32

0.5

1

(A)

4

8

16

32

8

16

32

(B)

23

23

22

22

21

21
20

2- 1

2- 1

β

20

2- 2

α, β

α, β

2

t.o.s. (h)

t.o.s. (h)

2- 2

2- 3

2- 3

2- 4

2- 4

2- 5

2- 5
2- 6

2- 6
0.25

0.5

1

2

4

8

16

t.o.s. (h)

(C)

32

0.25

0.5

1

2

4

t.o.s. (h)

(D)

These figures reveal the following trends:
1. MET rate (ĮCH4) decreases much more steeply than relative activity;
2. GAS rate (ĮCO) decreases similar to relative activity;
3. WGS rate (ĮCO2) decreases less than relative activity.

245
Figure 8. Relative rates of hexane conversion (Į, z), CH4 (ĮCH4, S), CO2 (ĮCO2, ¡)
and CO (ĮCO, T) formation vs. t.o.s. at different reaction conditions (see Table 1)
(adapted from ref. [15] with the kind permission of Wiley, copyright 2006).
(run a)

(run b)

0

0

-1
-2

α
aCH4
aCO2
aCO

ln α

ln α

-2
-3
-4

-4
α
αCH4
αCO2
αCO

-5
-6
0

2

-6

4

6

8

10

12

14

16

18

0

2

4

6

t.o.s. (h)

8

10

12

14

16

18

12

14

16

18

12

14

16

18

t.o.s. (h)

(run c)

(run d)
0

0

-2

α
αCH4
αCO2
αCO

-2
-3

α
αCH4
αCO2
αCO

ln α

ln α

-1

-4

-4
-6

-5
0

2

4

6

8

10

12

14

16

0

18

2

4

6

(run e)

10

(run f)

1

0
-1

-1

-2

α
αCH4
aCO2
aCO

-3

ln α

ln α

8

t.o.s. (h)

t.o.s. (h)

α
αCH4
aCO2
aCO

-3
-4

-5

-5
-7

-6
0

2

4

6

8

10

t.o.s. (h)

12

14

16

18

0

2

4

6

8

10

t.o.s. (h)

Indeed, direct relationships between kMET, kGAS, and kWGS decay and that of activity loss (kdeact)
result in straight-line correlations (Figure 9) with slopes equal to 2.7 ± 0.3, 1.1 ± 0.2 and 0.6 ± 0.1,
respectively. From this and similar figures it can be shown that MET, GAS, and WGS
functionalities depend on different ensembles of active sites [11,15,17,19–21,36]; thus, a slope of 1
for kGAS/kdeact suggests that hexane conversion involves single sites leading to CO, while MET and
WGS require larger and smaller ensemble of active sites for methane and carbon dioxide
formation, respectively.

246
Table 4. Deactivation constants of the various functionalities (T, 450 °C): hexane
conversion (kdeact), methanation (kMET), gasification (kGAS), water-gas-shift (kWGS).
Run
a
b
c
d
e
f

kdeact a
(hí1)
1.1 × 10í1
5.9 × 10í2
4.4 × 10í2
8.1 × 10í2
3.5 × 10í3
3.4 × 10í2

kcok a
(hí1)
1.1 × 10í1
5.6 × 10í2
4.0 × 10í2
7.7 × 10í2
0.000
3.0 × 10í2

kMET
(hí1)
2.2 × 10í1
9.7 × 10í2
6.0 × 10í2
2.2 × 10í1
3.5 × 10í3
4.4 × 10í2

kGAS
(hí1)
1.1 × 10í1
2.7 × 10í2
2.8 × 10í2
7.1 × 10í2
1.2 × 10í2
2.8 × 10í2

kWGS
(hí1)
6.6 × 10í2
3.2 × 10í2
2.7 × 10í2
6.1 × 10í2
1.9 × 10í2
2.2 × 10í2

a data taken from Table 2.

In fact, according to Beeckman et al. who adopted a stochastic approach to describe deactivation
by coking under different mechanisms [34], MET functionality is affected to the maximum extent
by the fouling of active sites because of its high number of metal atoms required for each site (e.g.,
formal site molecularity) in comparison to GAS and WGS ones [11,15,17,19,40,41].
Therefore, in agreement with literature data, kinetic and thermodynamic findings confirm that
surface carbon (i.e., CĮ) hydrogenation is the crucial step controlling activity, selectivity, and
stability against coking in the pre-reforming process. Activity decay is more pronounced at low
conversion, mostly under the occurrence of GAS and WGS reactions, as CO formation-decomposition
is mainly responsible for carbon build-up via an adsorbed CĮ intermediate [19,35].
Figure 9. Relationships between kMET (z), kGAS (T), kWGS (¡) and kdeact (see Table 4)
(adapted from ref. [15] with the kind permission of Wiley, copyright 2006).
0.25

Legend

kMET
kWGS
kGAS

kX (h-1)

0.20

kMET/kdeact, 2.7 ± 0.3

0.15

kGAS/kdeact, 1.1 ± 0.2
0.10

0.05

kWGS/kdeact, 0.6 ± 0.1
0.00
0.02

0.04

0.06

0.08

kdeact (h-1)

0.10

0.12

However, with addition of high concentrations of H2, similar values of kdeact and kMET
(Table 4), and a much steeper decay of gasification (kGAS/kdeact § 3.5) and WGS (kWGS/kdeact § 5)
functionalities [11,17,19] evidence that different reaction and deactivation paths are favored.

247
3. Experimental Section
3.1. Catalyst Preparation
The 19.1 wt% Ni/MgO catalyst was prepared by incipient wetness impregnation of “smoke” MgO
powder (UBE Ltd., Yamaguchi, Japan SABET, 30 m2āgí1) with an ethanolic solution of the Ni(NO3)2
salt [28–31]. The catalyst was dried at 80 °C, calcined at 400 °C (16 h) and pre-reduced for 6 h at
650 °C under H2 flow.
3.2. Catalyst Testing
Catalytic tests in the pre-reforming of n-hexane were performed at 450 °C and a pressure of
5–15 bar, using an isothermal “fixed bed” stainless steel microreactor (i.d., 6 mm), loaded with
25 mg of catalyst (40–70 mesh), diluted with same-sized SiC in a 1/30 weight ratio [11,15]. Before
tests, the catalyst was heated at 450 °C and further reduced in situ for 1 h in H2 flow (100 STP mL
miní1; P, 1 bar). The reaction mixture including n-C6H14 (1.8%), N2 (8.0%), H2O (16.2–37.8%) and He
(rest) for a variation of the S/C ratio between 1.5 and 3.5 respectively, was fed at the rate of
283 (GHSV, 12,000 hí1) or 566 (GHSV, 24,000 hí1) STP mL miní1. The effects of H2 (11.0%;
H2/C, 1.0) and thiophene (50 ppm in n-hexane) were probed at the highest GHSV and a S/C ratio of
2.8. Both water and n-hexane were fed as liquid by HPLC pumps. The reaction temperature was
controlled by a thermocouple in contact with the catalytic bed, while the reactor stream was
analyzed by a GC equipped with a three-columns analytical system connected to both TCD and
FID detectors for permanent gases and hydrocarbons analyses, respectively [11,15].
3.3. Catalyst Characterization
TEM analyses were performed using a PHILIPS CM12 Transmission Electron Microscope
(point-to-point resolution, 0.3 nm) on “fresh” and “spent” catalyst samples, dispersed ultrasonically
in ethanol and deposited over a thin carbon film supported on a standard copper grid. The Ni
particle size distribution (PSD) was obtained from an average of 150–200 particles, and the average
volume-area particle size calculated by the conventional statistical formula [29–31]:
d AV

¦n d
=
¦n d
i

3
i

i

2
i

i

(1)

i

The H2 uptake of “fresh” and “used” catalyst was evaluated by TPD measurements in the range
í80–620 °C using Ar as carrier gas (30 stp cm3⋅miní1) after in situ reduction of the samples in H2
flow at 450 °C for 1 h and subsequent saturation at r.t. (room temperature) and further at í80 °C.
The fraction of reduced Ni was evaluated by titration at 450 °C with O2 pulses (Ni + ½ O2  NiO)
after TPD analysis [31].
TGA-DSC measurements of the “used” catalyst samples (5 mg) in the range of 20–400 °C were
performed using a Simultaneous Thermal Analyser (STA 409C, Netzsch), operating in air with a
heating rate of 10 °C/min and an accuracy of 0.01 mg.

248
4. Conclusions
This review summarizes highlights of previous studies (mainly in our laboratory) of reaction and
deactivation paths on a “model” Ni/MgO catalyst in the pre-reforming of n-C6H14 with steam at
450 °C. A kinetic and thermodynamic analysis of the data from these studies provides (1) new
deactivation rate parameters for carbon deposition and coke formation and (2) insights into the
surface reaction network of the pre-reforming process, including hexane cracking, water-gas-shift,
and methanation reactions.
Important conclusions from this work can be summarized as follows:
•

•

•

•

Deactivation of Ni/MgO by carbon deposition and coking follows a pseudo 1st-order
dependence on active site concentration, consistent with an exponential decay with t.o.s.
That deactivation by coke and carbon formation is not simple first order in site
concentration, is evident from observed decreases in deactivation rate constants with
increasing S/C ratio, P, and contact time.
Comparison of thermodynamic predictions with experimental values of outlet stream
composition leads to the conclusion that the kinetics of hexane cracking/reforming,
carbon gasification and methanation reactions control the rate of coking; methanation is
the critical, rate limiting step responsible for coke deposition in the near absence of H2.
Carbon formation via the Boudouard reaction is apparently thermodynamically and
kinetically favored.
Under simulated industrial pre-reforming conditions at a high H2 concentration, hydrogen
substantially increases reaction throughput by speeding up hydrogenolysis, CO formation
via the water-gas-shift, and methanation reactions. H2 at high concentrations gasifies coke
and carbon precursors, thereby largely preventing formation of coke and inactive carbons.
S-poisoning by thiophene co-feeding causes a sharp drop in catalytic activity due to
irreversible poisoning; however, strong pore diffusional resistances in catalyst pellets and
bed prevent an accurate analysis of the poisoning mechanism and kinetics.

Author Contributions
This work was conceived by Francesco Arena, who arranged the manuscript, and Giuseppe
Trunfio, who carried out the experiments and also conceived and revised the manuscript.
Conflict of Interest
The authors declare no conflict of interest.
References
1.
2.

Piel, W.J. Transportation fuels of the future? Fuel Process. Tech. 2001, 71, 167–179.
Nagase, S.; Takami, S.; Hirayama, A.; Hirai, Y. Development of a high efficiency substitute
natural gas production process. Catal. Today 1998, 45, 393–397.

249
3.
4.
5.
6.
7.
8.
9.

10.

11.

12.

13.
14.
15.
16.
17.
18.
19.
20.
21.
22.

Cromarty, B.J.; Hooper, C.W. Increasing the throughput of an existing hydrogen plant. Int. J.
Hydrogen Energy 1997, 22, 17–22.
Aitani, A.M. Processes to enhance refinery-hydrogen production. Int. J. Hydrogen Energy
1996, 21, 267–271.
Fierro, J.L.G.; Peña, M.A.; Gómez, J.P. New catalytic routes for syngas and hydrogen
production. Appl. Catal. A 1996, 144, 7–57.
Rostrup-Nielsen, J.; Christiansen, L.J. Concept in Syngas Manufacture. Catalytic Science
Series; Imperial College Press: London, UK, 2011; Volume 10.
Bartholomew, C.H.; Farrauto, R.G. Fundamental of Industrial Catalytic Processes, 2nd ed.;
J. Wiley & sons, Inc.: Hobken, NJ, USA, 2006; pp. 339–371.
Wang, X.; Gorte, R.J. A study of steam reforming of hydrocarbon fuels on Pd/ceria.
Appl. Catal. A 2002, 224, 209–218.
Zou, X.; Wang, X.; Li, L.; Shen, K.; Lu, X.; Ding, W. Development of highly effective
supported nickel catalysts for pre-reforming of liquefied petroleum gas under low steam to
carbon molar ratios. Int. J. Hydrogen Energy 2010, 35, 12191–12200.
Shen, K.; Wang, X.; Zou, X.; Wang, X.; Lu, X.; Ding, W. Pre-reforming of liquefied
petroleum gas over nickel catalysts supported on magnesium aluminum mixed oxides. Int. J.
Hydrogen Energy 2011, 36, 4908–4916.
Arena, F.; Trunfio, G.; Alongi, E.; Branca, D.; Parmaliana, A. Modeling the activity-stability
pattern of Ni/MgO catalysts in the pre-reforming of n-hexane. Appl. Catal. A 2004, 266,
155–162.
Ayabe, A.; Omoto, H.; Utaka, T.; Kikuchi, R.; Sasaki, K.; Teraoke, Y.; Educhi, K. Catalytic
autothermal reforming of methane and propane over supported metal catalysts. Appl. Catal. A
2003, 241, 261–269.
Rostrup-Nielsen, J. Activity of nickel catalysts for steam reforming of hydrocarbons. J. Catal.
1973, 31, 173–181.
Rostrup-Nielsen, J. Catalysis, Science and Technology; Springer: Berlin, Germany, 1983.
Arena, F. Basic Relationships in the Pre-Reforming of n-Hexane on Ni/MgO Catalyst. AIChE
J. 2006, 52, 2823–2831.
Christensen, T.S. Adiabatic prereforming of hydrocarbons—An important step in syngas
production. Appl. Catal. A 1996, 138, 285–309.
Rostrup-Nielsen, J.; Alstrup, I. Innovation and science in the process industry: Steam
reforming and hydrogenolysis. Catal. Today 1999, 53, 311–316.
Sperle, T.; Chen, D.; Lødeng, R.; Holmen, A. Pre-reforming of natural gas on a Ni catalyst.
Criteria for carbon free operation. Appl. Catal. A 2005, 282, 195–204.
Bartholomew, C.H. Carbon deposition in methanation and steam reforming. Catal. Rev. Sci.
Eng. 1982, 24, 67–112.
Butt, J.B.; Petersen, E.E. Activation, Deactivation and Poisoning of Catalysts; Academic
Press: Inc.: San Diego, CA, USA, 1998.
Forzatti, P.; Lietti, L. Catalyst deactivation. Catal. Today 1999, 52, 165–181.
Bartholomew, C.H. Mechanism of catalyst deactivation. Appl. Catal. A 2001; 212, 17–60.

250
23. Kroll, V.C.H.; Swaan, H.M.; Mirodatos, C. Methane Reforming Reaction with Carbon
Dioxide Over Ni/SiO2 Catalyst. Deactivation Studies. J. Catal. 1996, 161, 409–422.
24. Tailleur, R.G.; Davila, Y. Optimal Hydrogen Production through Revamping
a Naphtha-Reforming Unit: Catalyst Deactivation. Energy Fuels 2008, 22, 2892–2901.
25. Sehested, J.; Carlsson, A.; Janssens, T.V.W.; Hansen, P.L.; Datye, A.K. Sintering of nickel
steam-reforming catalysts on MgAl2O4 spinel supports. J. Catal. 2001, 197, 200–209.
26. Ashrafi, M.; Pfeifer, C.; Pröll, T.; Hofbauer, H. Experimental Study of Model Biogas Catalytic
Steam Reforming: 2. Impact of Sulfur on the Deactivation and Regeneration of Ni-Based
Catalysts. Energy Fuels 2008, 22, 4190–4195.
27. Gallego, J.; Batiot-Dupeyrat, C.; Barrault, J.; Mondragón, F. Severe Deactivation of a LaNiO3
Perovskite-Type Catalyst Precursor with H2S during Methane Dry Reforming. Energy Fuels
2009, 23, 4883–4886.
28. Arena, F.; Parmaliana, A. Strategies of design of Ni/MgO catalysts for the reforming of
hydrocarbon to hydrogen/syngas. Curr. Top. Catal. 2006, 5, 69–88.
29. Parmaliana, A.; Arena, F.; Frusteri, F.; Coluccia, S.; Marchese, L.; Martra, G.; Chuvilin, A.
Magnesia Supported Nickel Catalysts. Part II: Surface properties and reactivity in CH4 steam
reforming. J. Catal. 1993, 141, 34–47.
30. Frusteri, F.; Spadaro, L.; Arena, F.; Chuvilin, A. TEM evidence for factors affecting the
genesis of carbon species on bare and K-doped Ni/MgO catalysts during the dry-reforming of
methane. Carbon 2002, 40, 1063–1070.
31. Arena, F.; Horrell, B.A.; Cocke, D.L.; Parmaliana, A.; Giordano, N. Magnesia-Supported
Nickel Catalysts. I: Structure and morphological properties. J. Catal. 1991, 132, 58–67.
32. Suzuki, T.; Iwanami, H.; Yoshinari, T. Steam reforming of kerosene on Ru/Al2O3 catalyst to
yield hydrogen. Int. J. Hydrogen Energy 2000, 25, 119–126.
33. Ming, Q.; Healey, T.; Allen, L.; Irving, P. Steam reforming of hydrocarbon fuels. Catal.
Today 2002, 77, 51–64.
34. Beeckman, J.W.; Nam, I.-S.; Froment, G.F. Stochastic modeling of catalyst deactivation by
site coverage. Stud. Surf. Sci. Catal. 1987, 34, 365–379.
35. Jackson, S.D.; Thomson, S.J.; Webb, G. Carbonaceous deposition associated with the catalytic
steam-reforming of hydrocarbons over nickel alumina catalysts. J. Catal. 1981, 70, 249–263.
36. Snoeck, J.-W.; Froment, G.F.; Fowles, M.J. Filamentous carbon formation and gasification:
thermodynamics, driving force, nucleation, and steady-state growth. J. Catal. 1997, 169,
240–249.
37. Chen, F.; Zha, S.; Dg, J.; Liu, M. Pre-reforming of propane for low-temperature SOFCs. Solid
State Ionics 2004, 166, 269–273.
38. Sughrue, E.L.; Bartholomew, C.H. Kinetics of carbon monoxide methanation on nickel
monolithic catalysts. Appl. Catal. 1982, 2, 239–256.
39. Rostrup-Nielsen, J. Sulfur Poisoning. In Progress in Catalyst Deactivation; Figuerido, J.L.,
Ed.; Martinus Nijhoff Publisher: The Hague, The Netherlands, 1982; pp. 209–227.

251
40. Sidjabat, O.; Trimm, D.L. Nickel-magnesia catalysts for the steam reforming of light
hydrocarbons. Top. Catal. 2000, 11/12, 279–282.
41. Fujita, S.; Terenuma, H.; Nakamura, M.; Takezawa, N. Mechanisms of methanation of CO
and CO2 over Ni. Ind. Eng. Chem. Res. 1991, 30, 1146–1151.

252

Inhibition of a Gold-Based Catalyst in Benzyl Alcohol
Oxidation: Understanding and Remediation
Emmanuel Skupien, Rob J. Berger, Vera P. Santos, Jorge Gascon, Michiel Makkee,
Michiel T. Kreutzer, Patricia J. Kooyman, Jacob A. Moulijn and Freek Kapteijn
Abstract: Benzyl alcohol oxidation was carried out in toluene as solvent, in the presence of the
potentially inhibiting oxidation products benzaldehyde and benzoic acid. Benzoic acid, or a product
of benzoic acid, is identified to be the inhibiting species. The presence of a basic potassium salt
(K2CO3 or KF) suppresses this inhibition, but promotes the formation of benzyl benzoate from the
alcohol and aldehyde. When a small amount of water is added together with the potassium salt, an
even greater beneficial effect is observed, due to a synergistic effect with the base. A kinetic model,
based on the three main reactions and four major reaction components, is presented to describe the
concentration-time profiles and inhibition. The inhibition, as well as the effect of the base, was
captured in the kinetic model, by combining strong benzoic acid adsorption and competitive
adsorption with benzyl alcohol. The effect of the potassium salt is accounted for in terms of
neutralization of benzoic acid.
Reprinted from Catalysts. Cite as: Skupien, E.; Berger, R.J.; Santos, V.P.; Gascon, J.; Makkee, M.;
Kreutzer, M.T.; Kooyman, P.J.; Moulijn, J.A.; Kapteijn, F. Inhibition of a Gold-Based Catalyst in
Benzyl Alcohol Oxidation: Understanding and Remediation. Catalysts 2014, 4, 89-115.
1. Introduction
As far as the chemical industry is concerned, biomass is forecast to be one of the major
successors of oil as a source of carbon for the production of organic molecules [1,2]. However, the
chemistry of biomass-derived molecules differs significantly from oil-derived molecules [1,3].
Their higher oxygen content renders them more sensitive to oxidation, requiring milder conditions
in selective oxidation processes. Furthermore, the solubility in water is enhanced when organics are
oxygenated, which can be either problematic or advantageous. It is now widely recognized that the
chemical industry will need to adapt to the new situation [1–3]. Once the catalytic abilities of gold
had been discovered [4], its high activity for oxidation reactions at mild conditions, down to room
temperature, was quickly noticed. This exceptionally low process temperature allows for a much
better control over the selectivity, in particular in selective oxidations. These reactions play an
important role in organic synthesis [1,5–11] and, as a consequence, much effort was put into
studying and benchmarking gold-based catalysts for selective oxidation of alcohols, ketones, and
carboxylic acids.
Benzyl alcohol (BnOH) selective oxidation to benzaldehyde (BnO) is one of these benchmark
reactions used extensively to assess the catalytic activity of gold catalysts [1,5–9,11–13].
Alcohol selective oxidation is often carried out in the presence of a base as “promoter” or
“co-catalyst” [1–3,7,8,11,12,14,15]. The base enhances the deprotonation of the alcohol, thereby
ensuring that the rate-limiting step is the catalytic oxidation step [1,15]. Furthermore, alkaline

253
conditions have also been reported to enhance the selectivity towards benzoic acid (BnOOH) [1,15].
However, the focus is usually on the initial turnover frequency (TOF) [5,7,12], whereas only a few
reports mention issues of deactivation and re-usability [1,6,9,11,14].
Deactivation can arise either from catalyst degradation (e.g., sintering) or from catalyst
poisoning or fouling [16]. Poisoning and fouling can sometimes be reversed by catalyst regeneration,
mainly under oxidative conditions [9]. Sulfur-containing impurities are often responsible for
poisoning of noble metal catalysts [16], and desulfurization catalysis has matured for decades to
answer this problem, for instance by a combination of hydrodesulphurization (HDS) and guard
beds. In a more general perspective, poisoning impurities in the feed can be eliminated by
dedicated treatments. In selective oxidation over noble metal catalysts, deactivation can also occur
due to the over-oxidation of Pt [14] and Pd [14,17] catalysts when an excess of (molecular) oxygen
is present. This over-oxidation means that too much atomic oxygen (a reaction intermediate) is
present on the catalytic sites, thereby blocking their accessibility for hydrocarbon adsorption.
Because of this, catalysts tested in the oxygen mass-transfer limited regime can exhibit higher
activity [1,11,14] than might be predicted based on data in the kinetic regime. It is generally
accepted that gold-based catalysts are resistant to over-oxidation [1,3,11], making them promising
candidates for oxidation reactions over extended periods of time. However, a reaction intermediate
or the product itself can be an inhibiting entity. It should be noted that inhibition is reversible
because the inhibition is remediated when the concentration of the inhibitor in the reaction medium
is sufficiently reduced, whereas poisoning is irreversible at the reaction conditions [16]. This
so-called product inhibition phenomenon is an even greater challenge, as the catalyst creates its
own poison while performing the desired reaction. This has been frequently observed both in
oxidation and dehydrogenation reactions. For instance, Dimitratos et al. [18] attributed deactivation
of Au-Pd and Au-Pt catalysts in octanol oxidation to inhibition by the carboxylate formed. They
also reported the alleviation of this inhibition when NaOH was present. Zope and Davis [3]
reported similar effects for the selective oxidation of glycerol to glyceric acid. They performed the
reaction in the presence of 19 different compounds: either products or intermediates in the glycerol
oxidation reaction, or species that might be formed from condensation of intermediates and/or
products. Ketones, condensation products of ketones or secondary alcohols (forming ketones upon
oxidation) were found to be inhibiting compounds. In contrast, simple carboxylic acids such as
acetic acid and propionic acid did not show appreciable inhibition, nor did diacids such as malonic
or succinic acid, or primary alcohols such as methanol. To the best of our knowledge, and despite
reviews mentioning the occurrence of product inhibition during the oxidation of alcohols on gold
and platinum group metal catalysts in general [11,14], no detailed study of this phenomenon for
BnOH oxidation over gold-based catalysts has been reported.
The current study concerns the Au-catalyzed partial oxidation of BnOH to BnO and benzyl
benzoate (BnOOBn) in toluene over the commercial AUROlite™ Au/Al2O3 catalyst. BnOH, which
is a primary alcohol, is oxidized to BnO and subsequently to BnOOH, while also BnOOBn can be
formed (Figure 1). This system suffers from deactivation, which is particularly observed when the
catalyst is re-used in batch-wise operation. This deactivation can be suppressed by the addition of
an inorganic base. The aim of the current study is to analyze this deactivation process, identify the

254
possible inhibiting species, elucidate the deactivation mechanism, and evaluate the beneficial
effect of the inorganic base. To accomplish this, the reaction was carried out under various
reaction conditions, including experiments in the presence of reaction products, bases, and
water. Additionally, kinetic modeling was performed in order to confirm the reaction and
deactivation mechanisms.
Figure 1. Reaction network: (a) oxidation of benzyl alcohol to benzaldehyde; (b) oxidation
of benzaldehyde to benzoic acid; (c) esterification of benzyl alcohol and benzoic acid, and
(d) esterification of benzyl alcohol and benzaldehyde under oxidative conditions to benzyl
benzoate [19].

2. Results and Discussion
2.1. Catalyst Deactivation in Base-Free Conditions
Figure 2 displays the concentration profiles for two subsequent experiments performed under
the same conditions and in the absence of a base. It should be noted that for all experimental data
points, the mass balance of the 4 main components (BnOH, BnO, BnOOH and BnOOBn) closes to
100% within measurement errors, with rare exceptions where discrepancies up to 4% are present.
Therefore, disproportionation and dehydration reactions, as reported by Alhumaimess et al. [5], can
be neglected. In the first experiment, the initial TOF is 0.7 sí1 and conversion levels off at around
55–60% after about 120 min. The selectivity to BnO is 93%, as reported in Table 1. In view of the
large discrepancies in TOF reported in literature [1], comparison of our data to previously
published ones is not straightforward. It is satisfactory that the values of 0.6 to 0.8 sí1 reported in
the present study are in the order of magnitude of those of 0.04 to 0.22 sí1 reported in the recent
review by Davis et al. [1] for nanoparticulate gold and of 2.8 to 4.4 sí1 for a gold foil reported in
the same review [1]. As indicated by Davis et al., the values span 2 orders of magnitude. The
reason for this is most likely the wide range of reaction conditions used in different studies.

255
Figure 2. Catalyst performance in base-free conditions. Concentration of (Ƈ) benzyl
alcohol, (Ŷ) benzaldehyde, (Ɣ) benzoic acid and (Ÿ) benzyl benzoate vs. reaction time
for (a) first run using fresh AUROlite™ and (b) second run using spent AUROlite™.
Reaction conditions: T = 80 °C, 0.8 g AUROlite™, CBnOH, t=0 = 3.0 × 10í4 mol·gí1 in 80
mL of toluene, 200 mL·miní1 air flow. Concentrations are expressed in moles per unit
mass of liquid in the reactor (mol·gí1). The symbols with error bars are the
experimental results and the lines represent the kinetic model.
(a)
0.3

30

(b)

0.3

0
0

50

100

t / min

150

200

250

-1

20

-3

0.1

10

-1

-1

0.0

C / 10 mol g

-3

C / 10 mol g

10

0.2

-6

0.1

-6

20

C / 10 mol g

0.2

C / 10 mol g

-1

30

0.0

0
0

50

100

150

200

250

t / min

In the second experiment, which is identical to the first one but with re-use of the same catalyst
sample after washing with toluene, virtually no conversion is observed, indicating that the catalyst
was completely deactivated. Since no other reactants than toluene and benzyl alcohol were present
in the reaction mixture, the deactivation must be caused either by catalyst deterioration or by an
inhibitor formed during the first experiment.
The potential presence of sulfur-containing impurities was investigated by analyzing the toluene
and benzyl alcohol by gas chromatography (GC). No sulfur-containing compounds could be
detected, implying that their concentration was below 50 ppb (the pulsed flame photometric
detector (PFPD) detection limit). Accounting for the quantities of these chemicals introduced in the
reaction mixture, this corresponds to 0.4 nmol of sulfur components at maximum in the reactor.
That is five orders of magnitude lower than the total amount of gold present. Since similar
concentration profiles as shown in Figure 2 were obtained when using reagents that were pretreated with activated carbon to remove any strongly adsorbing impurities, we conclude that feed
contaminants, including sulfur compounds, are not responsible for the observed deactivation. As a
consequence, activated carbon cleaning was deemed not to be necessary and was omitted for the
experiments reported here.

256
Table 1. Benzyl alcohol conversion rate, turn over frequency (TOF), conversion at 240 min
and selectivity to benzaldehyde at 60% conversion for AUROlite™ catalyst in different
test conditions.
Catalytic system

Base-free
Base-free re-test
Base-free–BnO
Base-free–BnOOH
K2CO3
K2CO3 re-test
KF
KF re-test
K2CO3–BnOOH
K2CO3–H2O
Base-free re-test–H2O

Figure

BnOH conversion
rate a

2a
2b
3
4
6a
6b
7a
7b
10
11
12

/mmol·miní1
0.37 ± 0.02
0.06 ± 0.04
0.09 ± 0.04
0
0.5 ± 0.1
0.20 ± 0.02
0.3 ± 0.1
0.5 ± 0.2
0.10 ± 0.03
0.43 ± 0.03
0.03 ± 0.05

TOF b
í1

/s

0.7 ± 0.1
0.1 ± 0.1
0.2 ± 0.1
0
0.9 ± 0.4
0.4 ± 0.1
0.6 ± 0.3
1.0 ± 0.5
0.2 ± 0.1
0.8 ± 0.2
0.1 ± 0.1

X at 240 min c
/%

Sel to BnO d
/%

63 ± 3
1±5
20 ± 10
2±7
78 ± 1
72 ± 1
81 ± 1
95 ± 1
61 ± 7
97 ± 1
1±4

93 ± 1
/
/
/
91 ± 1
85 ± 1
91 ± 1
83 ± 1
92 ± 1
86 ± 1
/

Concentrations are expressed in moles per unit mass of liquid in the reactor (mol·gí1).
ΔCBnOH
ΔCBnOH
a BnOH conversion rate =
× wliq , where
is the conversion rate of benzyl alcohol
Δt
Δt
í1
í1
(mmol·g ·min ), calculated by linear regression of the concentration values between 0 and 15 min and
wliq is the total mass of liquid (g); b TOF =

M Au
ΔCBnOH 0.001
w × 0.01
×
× wliq ×
, where cat
is
60
wcat × 0.01× 0.22
Δt
M Au

the total amount of gold in the reactor (mol) and 0.22 is the amount of edge + corner atoms per amount of
gold for 2.5 nm gold nanoparticles [20] (mol·molí1);

SelBnO =

CBnO, X =60%
CBnO, X =60% + CBnOOH , X =60% + 2 × CBnOOBn, X =60%

c

X=

CBnOH ,t =0 − CBnOH ,t =240min
CBnOH ,t =0

×100% ;

d

×100% In parallel and series reaction networks,

selectivities have to be compared at the same conversion levels. 60% is chosen here, as it is the level in
the base free experiment after 240 min.

These results thus suggest that products or intermediates formed cause the observed
deactivation. In order to investigate whether product inhibition is indeed taking place, an
experiment was performed with fresh catalyst, where BnO was added to the reactor 20 min prior to
the actual start of the catalytic reaction under the same reaction conditions, i.e., before the
introduction of BnOH. The 20 min exposure time was selected based on the results of Figure 1a
where it can be seen that BnOOBn was already formed at that reaction time, thereby ensuring that
all potentially inhibiting products were present. The results in Figure 3 show a very low conversion
of 20%, confirming strong product inhibition. It is striking that despite the appreciable amount of
BnO introduced at t < 0, neither BnOOBn nor BnOOH could be detected until the BnOH was
introduced at t = 0. Clearly the sites for the sequential reactions of BnO (Figure 1a,b) are fully
blocked without the presence of alcohol. It should be noted that for t < 0 the BnO concentration
shows a slight decrease, indicating that a small amount of BnO is consumed without producing a
detectable amount of BnOOH in the liquid phase. We conclude that some product is formed which
remains on the catalyst and inhibits further turnovers. Carboxylic acid moieties are well known to

257
interact strongly with gold nanoparticle surfaces, even allowing the stabilization of small gold
clusters in colloidal systems [21]. BnOOH is therefore suspected to be the inhibitor, since it is the
logical product of BnOH oxidation, although its concentration in the solution stayed below the GC
detection limit of 0.20 μmol·gí1.
Figure 3. Catalyst performance after pre-addition of benzaldehyde. Concentration of
(Ƈ) benzyl alcohol, (Ŷ) benzaldehyde, (Ɣ) benzoic acid and (Ÿ) benzyl benzoate vs.
reaction time for catalytic reaction over fresh AUROlite™. Benzaldehyde was
introduced 20 min before benzyl alcohol was added. Reaction conditions: T = 80 °C,
0.8 g AUROlite™, CBnOH, t=0 = 3.0 × 10í4 mol·gí1, CBnO, t<0 = 1.7 × 10í4 mol·gí1, in
80 mL of toluene, 200 mL·miní1 air flow. Concentrations are expressed in moles per
unit mass of liquid in the reactor (mol·gí1). The symbols with error bars are the
experimental results and the lines represent the kinetic model.
8

0.3

-3

2

-1

C / 10 mol g

0.1

-6

4

C / 10 mol g

-1

6
0.2

0

0.0
0

50

100

150

200

250

t / min

In order to evaluate the effect of BnOOH, a similar experiment was performed in which BnOOH
was added 20 min prior to the actual start of the test. The concentration profiles are shown in
Figure 4. In this experiment, approximately 40 times less acid was introduced compared to the
amount of aldehyde introduced for the experiment in Figure 3. Nevertheless, an even more
dramatic inhibiting effect was observed: hardly any conversion of BnOH occurred. The
concentration of BnOOH measured by GC at t < 0 is 1 μmol·gí1, which is about 25% of what was
added, suggesting a strong interaction of BnOOH with the catalyst. The slight increase in
concentration of BnOOH upon addition of BnOH at t = 0 is attributed to competitive adsorption
between the alcohol and the acid. No ester was present at the beginning of the reaction, nor was it
detected during the course of the reaction. This suggests that the inhibiting product is either
BnOOH or a compound formed from BnOOH. Therefore, product inhibition particularly occurs on
catalysts on which BnO can react further to BnOOH. This interpretation is confirmed by a
nanostructured gold-based catalyst synthesized in our lab exhibiting 100% selectivity to BnO not
showing deactivation for 4 consecutive runs [22].

258
Figure 4. Catalyst performance after pre-addition of benzoic acid. Concentration of (Ƈ)
benzyl alcohol, (Ŷ) benzaldehyde, (Ɣ) benzoic acid and (Ÿ) benzyl benzoate vs.
reaction time for catalytic reaction over fresh AUROlite™. Benzoic acid was introduced
20 min prior to the beginning of the reaction. Reaction conditions: T = 80 °C, 0.8 g
AUROlite™, CBnOH, t=0 = 3.0 × 10í4 mol·gí1, CBnOOH, t<0 = 4.1 × 10í6 mol·gí1, in 80 mL
of toluene, 200 mL·miní1 air flow. Concentrations are expressed in moles per unit mass
of liquid in the reactor (mol·gí1). The symbols with error bars are the experimental
results and the lines represent the kinetic model.
25

0.3

-3

-6

10

0.1

5

-1

C / 10 mol g

15

C / 10 mol g

-1

20
0.2

0

0.0
0

50

100

150

200

250

t / min

2.2. Influence of a Base on Catalyst Deactivation
Figure 5 shows the reaction mechanism, in basic conditions, of the oxidation of alcohol to
aldehyde and the sequential oxidation to carboxylic acid [1,15,23–25]. Au* indicates an active site
on the gold surface. A dashed line represents a chemical interaction of a species with an adsorption
site on the gold surface. Electron transfers corresponding to bond cleavage or formation are
indicated by curved arrows. As indicated in Figure 5, reaction step a, the role of the base is to
deprotonate the alcohol, thereby ensuring that the rate limiting step is the reaction on the gold
surface, presumably the ȕ-hydride elimination forming the aldehyde (step b yielding R-CH=O) or
the carboxylic acid (step d yielding R-COOH) [1,15,23–25]. However, basic conditions have also
been reported to enhance the selectivity towards carboxylic acid, by favoring the conversion of
aldehyde to the corresponding geminal diol [1,15] as depicted by reaction step c (The diol is shown
as R-CH(OH)Oí, partially deprotonated and adsorbed on Au*). In parallel, the gold site that carries
the hydride (Au*-Hí) is regenerated by adsorbed molecular oxygen (Au*||||O2) via a peroxyl
intermediate (step e yielding Au*-O-OH) [1,15], or via dissociated oxygen from the catalyst
support [25]. This also regenerates an OHí (steps f and g) [15] and closes the catalytic cycle as
shown on the right-hand side of Figure 5. For stoichiometric reasons, it is clear that two hydrides
species must react per O2 molecule. However, it is unclear whether the second hydride reacts after
the O-O bond dissociation (step f), or if the O-O bond dissociation is assisted by the second hydride
(in the latter case, step f and g would be simultaneous). The question arises if the presence of a base
has an influence on the catalyst stability.

259
Figure 5. Reaction mechanism of benzyl alcohol oxidation to benzaldehyde and
sequential oxidation to benzoic acid over a gold catalyst, co-catalyzed by the base
HOí [1,15,23–25].

The catalytic reaction was carried out in the presence of two different bases: potassium
carbonate (pKb1 = 3.68) and potassium fluoride (pKb = 10.8). Although the pKb values (defined in
aqueous environment) are not directly transferable to the aprotic solvent (toluene), and these bases
hardly dissolve in toluene, the pKb values still indicate the relative basicity of both bases used. The
spent catalysts were also re-used. A significant loss of KF was observed during the recovery of the
catalyst after the first test, which was compensated for by the addition of another 2.1 g of KF to the
spent catalyst, resulting in an estimated total amount of KF present during the second experiment
with the spent catalyst of about 3 g. The concentration profiles of these experiments are presented
in Figures 6 and 7, respectively.
The results of the first run are similar to those for the base-free experiments, although a slight
increase in the initial TOF, from 0.7 to 0.9 sí1, is observed (Table 1). Increased activity upon
addition of a base has already been reported [1] and is probably due to the enhanced deprotonation
of BnOH. The conversion at 240 min is also higher: 78% compared with 63% in base-free
conditions. The selectivity to BnO does not change significantly (93% when no base is present vs.
91% in the presence of potassium carbonate). However, it should be noted that more ester is
formed in the presence of K2CO3 (at 240 min: 16 μmol·gí1 compared to 8 μmol·gí1 in base free
conditions), and the amount of BnOOH remains below detection limit. The most striking
difference, when compared to base-free conditions, is the largely maintained activity when re-using
the spent catalyst. The initial TOF is lower than that of the fresh catalyst (0.4 sí1 compared with
0.9 sí1), but the conversion at 240 min is comparable (72% compared with 78%). This shows that
the addition of a base to the reaction medium largely remediates the strong product inhibition
observed under base-free conditions.

260
Figure 6. Catalyst performance in the presence of K2CO3. Concentration of (Ƈ) benzyl
alcohol, (Ŷ) benzaldehyde, (Ɣ) benzoic acid and (Ÿ) benzyl benzoate vs. reaction time
for catalytic reaction over fresh AUROlite™ in the presence of K2CO3. (a) First run
using fresh AUROlite™ (b) second run using spent AUROlite™. Reaction conditions:
T = 80 °C, 0.8 g AUROlite™, 3.04 g K2CO3, CBnOH, t=0 = 3.0 × 10í4 mol·gí1, in 80 mL
of toluene, 200 mL·miní1 air flow. Concentrations are expressed in moles per unit mass
of liquid in the reactor (mol·gí1). The symbols with error bars are the experimental
results and the lines represent the kinetic model.
(b)

(a)
0.3

30

0.2

20

0.1

10

50

0.3

-1
-3

C / 10 mol g

C / 10 mol g

-3

50

100

150

200

20
0.1

-1

0

30

-6

-1

0

0.0

0.2

C / 10 mol g

-6

C / 10 mol g

-1

40

10
0.0

250

0
0

50

t / min

100

150

200

250

t / min

Figure 7. Catalyst performance in the presence of KF. Concentration of (Ƈ) benzyl
alcohol, (Ŷ) benzaldehyde, (Ɣ) benzoic acid and (Ÿ) benzyl benzoate vs. reaction time
for catalytic reaction over fresh AUROlite™ in the presence of KF. (a) First run using
fresh AUROlite™ and 2.1 g KF (b) second run using spent AUROlite™ and ~3 g KF.
Reaction conditions: T = 80 °C, 0.8 g AUROlite™, CBnOH, t=0 = 3.0 × 10í4 mol·gí1, in
80 mL of toluene, 200 mL·miní1 air flow. Concentrations are expressed in moles per
unit mass of liquid in the reactor (mol·gí1). The symbols with error bars are the
experimental results and the lines represent the kinetic model.
(a)
0.3

100

t / min

150

200

250

0.1

20

-1

40

-3

C / 10 mol g

-1
-3

50

0.2

-1

0
0

60

-6

0.0

-1

10

-6

20
0.1

0.3

C / 10 mol g

30

0.2

C / 10 mol g

C / 10 mol g

(b)
40

0

0.0
0

50

100

150

200

250

t / min

When K2CO3 is replaced by an equimolar amount of KF, the results are similar. TOF is slightly
lower (0.6 sí1) although the difference is within the measurement error. The conversion is 81%
after 240 min when KF is used, and 78% when K2CO3 is used (Table 1). Apparently, the strength
of the base does not affect the activity of the fresh catalyst. Upon re-using the spent catalyst
however, a higher initial reaction rate is observed (1.0 sí1) and the final conversion even reaches

261
95%. This increase is attributed to the additional potassium fluoride added. The preserved catalytic
activity proves that no other deactivation mechanism than product inhibition (e.g., sintering or
poisoning) is taking place. It is clear that the bases largely neutralize the acid responsible for
product inhibition. As this alleviation of the inhibition is always accompanied by an increase of
selectivity towards BnOOBn, apparently ester formation is faster in the presence of a base.
However, direct esterification of the carboxylic acid with the alcohol catalyzed by a base is
mechanistically not likely. After deprotonation of the BnOOH by a base, nucleophilic attack of the
alcohol (R-CH2-OH) on the functional carbon of carboxylate anion (R-COOí) is highly unlikely as
this would imply that a nucleophile (also seen as a Lewis base) would have to react with an
electron rich species. This is illustrated in Figure 8 (steps d and e). The classic acid-catalyzed
esterification between carboxylic acid and alcohol is depicted on the left hand-side of Figure 8 for
comparison. It shows that under acidic conditions, the functional carbon becomes positively
charged (Figure 8 step a, yielding R-C(OH)2+) and thus more prone to nucleophilic attack. Another
pathway involving base-catalyzed ester formation from an alcohol and an aldehyde has been
suggested by Rodríguez-Reyes et al. [19]. This pathway is illustrated in Figure 9, and can explain
why in our system a higher ester production is observed under basic conditions.
Figure 8. Reaction mechanism of acid catalyzed esterification of carboxylic acid and
alcohol. (a) protonation of carbonyl oxygen yielding an electrophilic carbocation;
(b) nucleophilic attack of the alcohol and (c) dehydration yielding the corresponding
ester. (d) deprotonation of the carboxylic acid by a base yielding a carboxylate anion;
(e) the nucleophilic attack of the alcohol is then greatly disfavored.

262
Figure 9. Reaction mechanism of base-catalyzed ester formation from alcohol and
aldehyde, adapted from [19]. (a) alcohol deprotonation, (b) nucleophilic attack of
alcoholate on carbonyl, followed by (c) ȕ-hydride elimination yields the corresponding
ester. The rest of the catalytic cycle consists of the oxidation of the hydride left on the
gold surface by molecular oxygen, which also regenerates the base as depicted in
Figure 5 (steps f and g).

To further elucidate the role of the base on product inhibition, an experiment was performed
with potassium carbonate and pre-addition of BnOOH. The concentration profiles are shown in
Figure 10 and are directly comparable with the ones shown in Figure 6a where K2CO3 was
present but no inhibitor was pre-added, and with Figure 4 where no base was present but BnOOH
was pre-added.
In comparison with Figure 6a where no BnOOH was added, the initial TOF is lower: 0.2 sí1
(Table 1). The conversion after 240 min is 61% and the selectivity to BnO is 92%. When compared
with the results in Figure 4 (no carbonate added), we can conclude that even though inhibition is
still observed, the presence of potassium carbonate greatly reduces it.
Based on our interpretation, the presence of a base results in (partial) neutralization of the
BnOOH, thereby alleviating inhibition. In parallel, basic conditions also enhance the subsequent
ester formation depicted in Figure 9, leading to more ester production and decreasing aldehyde
selectivity. An effect of the amount of potassium on the inhibition is also suspected based on
Figure 7a,b, but this has not been further quantified.

263
Figure 10. Catalyst performance in the presence of K2CO3 after pre-addition of benzoic
acid. Concentration of (Ƈ) benzyl alcohol, (Ŷ) benzaldehyde, (Ɣ) benzoic acid and (Ÿ)
benzyl benzoate vs. reaction time for catalytic reaction over fresh AUROlite™ where
benzoic acid was introduced prior to the beginning of the reaction. Reaction conditions:
T = 80 °C, 0.8 g AUROlite™, 3.04 g K2CO3, CBnOH, t=0 = 3.0 × 10í4 mol·gí1,
CBnOOH, t<0 = 4.1 × 10í6 mol·gí1, in 80 mL of toluene, 200 mL·miní1 air flow.
Concentrations are expressed in moles per unit mass of liquid in the reactor (mol·gí1).
The symbols with error bars are the experimental results and the lines represent the
kinetic model.

40

-3

20
10

0.0

-1

0.1

-6

C / 10 mol g

30
0.2

C / 10 mol g

-1

0.3

0
0

50

100

150

200

250

t / min

2.3. Influence of Water on Catalyst Deactivation
Water was already recognized to be crucial for Au-catalyzed gas phase CO oxidation by
Daté et al. [26], who demonstrated its great beneficial role. The mechanism was later proposed by
Daniells et al. [27]. In many of the reaction mechanisms discussed above, water plays a role. It is
produced in an amount equimolar to BnO (see Figures 1a and 5a). The mechanism of the oxidation
of aldehyde to carboxylic acid involves a base-catalyzed hydration of aldehyde to geminal diol.
Water is also a byproduct of the equilibrium-limited esterification. Yang et al. [28] studied the
influence of different water contents on the kinetics of oxidation and determined that water has a
promoting effect. They observed higher conversions of BnOH and higher selectivities to BnO when
an optimal amount of water was used. More recently, Chang and coworkers [29] conducted a
computational study to better understand this effect on methanol dehydrogenation/oxidation. They
concluded that the promoting effect originates from a facilitated peroxyl formation from O2 by
transfer of hydrogen from the water itself or from the alcohol via the water, where hydrogen bonds
are reported to play a key role in this mechanism. Therefore, water could play a role in the
deactivation of the catalyst in the present study. With this in mind, the reaction was carried out in
the presence of the small amount of water that adheres to the catalyst after immersion in water and
filtration. The mass difference before and after this treatment indicates about 0.5 g of water per
gram of catalyst, which corresponds to around 550 mol of water per mol of gold. The results of this
experiment are presented in Figure 11.

264
Figure 11. Catalyst performance in the presence of K2CO3 and a small amount of water.
Concentration profiles of (Ƈ) benzyl alcohol, (Ŷ) benzaldehyde, (Ɣ) benzoic acid and (Ÿ)
benzyl benzoate vs. reaction time for catalytic reaction over fresh AUROlite™ in the
presence of water. Reaction conditions: T = 80 °C, 0.8 g AUROlite™, 3.04 g K2CO3, ~0.4 g
H2O, CBnOH, t=0 = 3.0 × 10í4 mol·gí1, in 80 mL of toluene, 200 mL·miní1 air flow.
Concentrations are expressed in moles per unit mass of liquid in the reactor (mol·gí1).
The symbols with error bars are the experimental results and the lines represent the
kinetic model.

0.2

40

0.1

20

0.0

0

-3

-6

C / 10 mol g

-1

60

C / 10 mol g

0.3

-1

0

50

100

150

200

250

t / min

Compared with the dry conditions of Figure 6a the initial TOFs are equal within the
experimental error, being 0.8 sí1 in moist conditions and 0.9 sí1 in dry conditions. We do not
observe the kinetic effect of water found by Yang et al., but in the presence of water our catalyst
does not deactivate and the final conversion is close to 100%. The water thus positively influences
the catalyst stability. A possible explanation could be that water interacts with the inhibiting
product, thereby diminishing its interaction with the catalyst. However, since potassium carbonate
was also present in the reactor, a synergistic effect of base and water cannot be excluded. To
address this question, the deactivated catalyst obtained after the experiment of Figure 2b (where no
base was present) was tested again in the presence of water and without the addition of any base.
Concentration profiles of this experiment are displayed in Figure 12.
Clearly, catalytic activity could not be recovered by this treatment. Thus, water alone does not
remove the species responsible for deactivation, and the beneficial influence of water observed in
Figure 11 is due to a synergistic effect with the potassium carbonate, e.g., by an enhanced
dissolution of the inorganic base. It should also be noted that adding more water than used in these
experiments provoked phase separation in which the catalyst agglomerated in the water phase,
thereby eliminating the dispersion of the catalyst powder in the organic phase and causing the
reaction to proceed in a mass-transport limited regime.

265
Figure 12. Catalyst performance of the water-washed spent catalyst in base-free
conditions from Figure 2 in the presence of water. Concentration of (Ƈ) benzyl alcohol,
(Ŷ) benzaldehyde, (Ɣ) benzoic acid and (Ÿ) benzyl benzoate vs. reaction time for
catalytic reaction spent AUROlite™ in the presence of water. Reaction conditions:
T = 80 °C, 0.8 g AUROlite™, ~1 g H2O, CBnOH, t=0 = 3.0 × 10í4 mol·gí1, in 80 mL of
toluene, 200 mL·miní1 air flow. Concentrations are expressed in moles per unit mass of
liquid in the reactor (mol·gí1). The symbols with error bars are the experimental results
and the lines represent the kinetic model.

50
0.3

30

-3

-6

C / 10 mol g

0.2

C / 10 mol g

-1

40

20

-1

0.1
10
0.0

0
0

50

100

150

200

250

t / min

In order to further assess which hypothesis holds, the spent catalyst was analyzed by DRIFTS
before and after washing in boiling water for 12 h (Figure 13). The fresh catalyst barely shows any
features. In contrast, both the spent catalyst and the spent catalyst after boiling in water show clear
absorption features. The catalyst tested in the presence of KF, which did not deactivate after two
runs, exhibited different features. Identification of the species on the surface of the catalysts was
based on reference spectra recorded of the compounds present in the reactor adsorbed on alumina.
The band at 1200 cm-1 is attributed to Carom–CHO or Carom–CH2OH stretching vibrations [30] and
originates only from BnOH and BnO. Interestingly, this band is only present for the catalyst tested
with KF. A small absorption band at 1390 cm-1 is seen in practically all cases and corresponds to
the bending vibration of O–H bonds; thus, it cannot be used to differentiate the components of
interest. The sharp absorption at 1450 cm-1 is due to in-plane bending vibrations of protons at a
primary alcohol carbon and is specific to BnOH. Unfortunately, this band is often masked by a
broader one due to other O–H vibrations at the same wavenumber. Absorptions between 1500 and
1600 cm-1 are attributed to Csp2=Csp2 stretching vibrations present in all the compounds of interest.
The broad band at 1710 cm-1 is specific to benzaldehydes and is due to ʌ-conjugation of bonds
throughout the entire molecule. Again, this characteristic feature is only present for the catalyst
tested with KF. A smaller feature is also observed for BnOOH for the same reasons, but the
carboxyl moiety seems to alter it drastically. The bending mode of water, which should appear at
around 1600–1800 cmí1, is hardly visible due to the pre-treatment of the samples at 473 K. Some
features are also observed at higher wavenumbers. The small bands centered at around 2740 cm-1
and 2820 cm-1 arise from wagging and stretching vibrations respectively of the BnO carbonyl proton,
but these bands are not detected on any of the spent catalysts. The symmetric and anti-symmetric

266
stretching vibrations of primary alcohol methylene group protons give rise to absorptions at 2870 and
2930 cm-1, respectively. The three bands between 3030 and 3090 cm-1 are assigned to Csp2–H
vibrations, which stem from any aromatic compound in the reaction medium.
Figure 13. DRIFT spectra of (a) fresh AUROlite™ catalyst; (b) used with K2CO3;
(c) used with K2CO3 and washed in boiling water; (d) used with KF, (e)–(i) reference
compounds adsorbed on alumina. Catalyst samples were pretreated under He at 473 K
before recording.

(i) Toluene/Al2O3
(h) Benzyl benzoate/Al2O3

(g) Benzoic acid/Al2O3

(f) Benzaldehyde/Al2O3
(e) Benzyl alcohol/Al2O3
(d) AUROliteTM spent KF
(c) AUROliteTM spent K2CO3 washed
(b) AUROliteTM spent K2CO3
(a) AUROliteTM fresh

3200

3000

2800

2000

1800

-1

σ / cm

1600

1400

1200

1000

Figure 13 clearly demonstrates that water does not visibly wash off the species involved in the
reaction from the catalyst surface. Even after 12 h in boiling water, the intensities of the absorption
bands corresponding to aromatic species and oxygenated aromatic species do not show any sign of
decrease. Therefore, combined with the results of Figure 12, the hypothesis that water remediates
the product inhibition by enhancing desorption of the inhibiting product is refuted. The absorption
features on the catalyst tested in the presence of K2CO3 are very similar to those of BnOOH and
BnOOBn, although the observed bands are not specific to these compounds. In contrast, the
catalyst tested in the presence of KF shows absorption features similar to those of BnO, in
particular the characteristic band at 1710 cmí1. Since this catalyst did not deactivate whereas the
catalyst tested with K2CO3 showed some deactivation, the attribution of the inhibitor being BnOOH
or one of its products is supported. However, in view of the low degree of deactivation shown by
the catalyst tested in the presence of K2CO3, a signal corresponding to BnO would be expected. It
remains unclear why BnO seems to be absent from this catalyst surface despite the fact that KF and
K2CO3 have similar beneficial effects.

267
2.4. Kinetic Modeling
Since our experiments show that the BnOOH concentration in solution is very low at all times and
that the ester formation increases in the presence of a base, it is assumed that the esterification
runs entirely through the reaction of BnO with BnOH by H abstraction. The reactions involved in the
model are:
C6 H 5CH 2OH + 12 O2 ⎯⎯
→ C6 H 5CHO + H 2O

(1)

C6 H 5CHO + 12 O2 ⎯⎯
→ C6 H 5COOH

(2)

C6 H5CHO + C6 H5CH 2OH + 12 O2 ⎯⎯
→ C6 H5COOCH 2C6 H 5 + H 2O

(3)

The reaction model is accordingly assumed to consist of the set of surface reactions shown in
the following set of equations, in which ∗ stands for a catalytic oxidation site:
BnOH adsorption

ZZX C6 H 5CH 2OH ∗
C6 H 5CH 2OH + ∗ YZZ

(4)

BnO adsorption

ZZX C6 H 5CHO ∗
C6 H 5CHO + ∗ YZZ

(5)

BnOH oxidation

C6 H 5CH 2OH ∗ + ∗ ⎯⎯
→ C6 H 5CHOH ∗ + H ∗

(6)

1

O2

1
4
H ∗ ⎯⎯⎯
fast → ∗ + 2 H 2O

O2 ∗
C6 H 5CHOH ∗ ⎯⎯⎯
fast → C6 H 5CHO ∗ + HO2 ∗
3
1
HO2 ∗ ⎯⎯⎯
fast → 2 H 2O + 4 O2

BnO hydration

H 2O
ZZZZ
X
C6 H 5CHO ∗ YZZZ
Z C6 H 5CH ( OH )2 ∗
− H 2O

(7)

BnO oxidation

C6 H 5CH ( OH )2 * + ∗ ⎯⎯
→ C6 H 5CHO ( OH ) ∗ + H ∗

(8)

1

O2

1
4
H ∗ ⎯⎯⎯
fast → ∗ + 2 H 2O

O2 ∗
C6 H 5CHO ( OH ) ∗ ⎯⎯⎯
fast → C6 H 5COOH ∗ + HO2 ∗
3
1
HO2 ∗ ⎯⎯⎯
fast → 2 H 2O + 4 O2

BnOOH adsorption

ZZX C6 H 5COOH ∗
C6 H 5COOH + ∗ YZZ

(9)

BnOOBn adsorption

ZZX C6 H 5COOCH 2C6 H 5 ∗
C6 H 5COOCH 2C6 H 5 + ∗ YZZ

(10)

BnOOBn formation

C6 H5CHO∗ + C6 H5CH2OH ∗ ⎯⎯
→ C6 H5COOCH2C6 H5 ∗ + 2 H ∗

(11)

1

O2

2
2 H ∗ ⎯⎯⎯
fast → ∗ + H 2O

268
It is assumed that all three surface oxidation reactions are irreversible. These reactions require
two adjacent sites, one for the adsorbed species originating from the BnOH or BnO and one empty
site that was regenerated by the oxygen or, for the esterification, adjacently adsorbed BnOH or
BnO. It is also assumed that the hydride species, H*, and the peroxy species. HO2*, are quickly
removed in the excess of oxygen present, leading to the assumption that the occupancy of both
species is always very low. Adsorption of water and oxygen is assumed not to inhibit the
reaction. The surface reactions are described by the surface species reaction rates shown in the
following equations:
BnOH adsorption

rBnOH = kBnOH CBnOH NTθ∗ − k− BnOH NTθ BnOH ∗

(12)

BnO adsorption

rBnO = kBnOCBnO NTθ∗ − k− BnO NTθ BnO∗

(13)

BnOOH adsorption

rBnOOH = kBnOOH CBnOOH NTθ∗ − k− BnOOH NTθ BnOOH ∗

(14)

BnOOBn adsorption

rBnOOBn = kBnOOBnCBnOOBn NTθ∗ − k− BnOOBn NTθ BnOOBn∗

(15)

BnOH oxidation

r1 = k1' N T θ BnOH ∗θ ∗ s

(16)

BnO oxidation

r2 = k2' NT θ BnO∗θ∗ s

(17)

BnOOBn formation

r3 = k3' NT θ BnO∗θ BnOH ∗ s

(18)

Where:

(mol·gcatí1·sí1)

ri = reaction rate
ki = reaction rate constant (unit according to equations)
Ci = concentration (mol·gí1)
și = occupancy of the surface sites (-)
NT = total number of surface oxidation sites per unit catalyst mass (gcatí1)
s = number of adjacent sites per site (-)

The hydration of BnO followed by the sequential oxidation of the diol species has been lumped
here into a single step (17), as the equilibrium constant of the reversible hydration will be highly
correlated with the rate constant of the oxidation step.
In general, adsorption reactions are at quasi-equilibrium, while the surface reactions are
rate-limiting steps. This allows expression of the surface coverages as functions of the
concentrations and the adsorption constants:
BnOH adsorption

θ BnOH ∗ = K BnOH CBnOHθ∗ with K BnOH = kBnOH k− BnOH

(19)

BnO adsorption

θ BnO∗ = K BnOCBnOθ∗ with K BnO = kBnO k− BnO

(20)

BnOOH adsorption

θ BnOOH ∗ = K BnOOH CBnOOHθ∗ with K BnOOH = kBnOOH k− BnOOH

(21)

BnOOBn adsorption

θ BnOOBn∗ = K BnOOBnCBnOOBnθ∗ with KBnOOBn = kBnOOBn k− BnOOBn

(22)

Where: Ki = adsorption constant of compound i on a gold active site (g·molí1)
After substitution of all the surface coverage expressions in the site balance:

269

θ BnOH ∗ + θ BnO∗ + θ BnOOH ∗ + θ BnOOBn∗ + θ∗ = 1
θ∗ =

1
+
+
+
K
C
K
C
K
1
(
BnOH BnOH
BnO BnO
BnOOH C BnOOH + K BnOOBn CBnOOBn )

(23)
(24)

The kinetic parameter estimation based on the full model showed that not all adsorption
constants could be estimated properly due to a too weak sensitivity. It appeared that only the
adsorption term of the most polar species present, BnOOH, was significant allowing the estimation
of its adsorption constant and neglecting the adsorption terms of the BnOH, the BnO and the
BnOOBn. Since it was experimentally observed that the presence of a base, i.e., K2CO3 or KF,
results in significant ester formation, it was assumed that this reaction is base-catalyzed. Since
hardly any ester is formed in absence of base, it is concluded that the condensation reaction of
BnOOH and BnOH does not occur to any significant extent.
By elimination of the coverages of the surface species, the following reaction rate expressions
are obtained:
r1 =
r2 =

r3 =
Where:

k1

(1 + K BnOOH CBnOOH )

CBnOH

With: k1 = k1' K BnOH NT s

2

CBnO

With: k2 = k2' K BnO NT s

2

CBnO CBnOH

With: k3 = k3' K BnO K BnOH NT s

2

k2

(1 + K BnOOH CBnOOH )
Sbase k3

(1 + K BnOOH CBnOOH )

(25)

(26)

(27)

ri = reaction rate (mol·kgcatí1·sí1)
ki = lumped rate constant (unit according to equations)
Ci = concentration (mol·gí1)
Sbase = base strength (K2CO3 or KF) (-)

In the two catalytic experiments in which BnOOH was added prior to the BnOH (Figures 4
and 10), it was observed that the catalyst support Al2O3 adsorbs BnOOH more strongly than
BnOH. In the first experiment (Figure 4), it was observed that the molar quantity of BnOOH in the
liquid was much lower (16 μmol) than that originally added (780 μmol). The amount of BnOOH in
solution increased only slightly (to 74 μmol) after subsequent addition of the 22 mmol of BnOH,
demonstrating that BnOOH is adsorbed more strongly than BnOH. In the second experiment
(Figure 10), it was observed that despite the addition of a similar amount of BnOOH (712 μmol
yielding a theoretical concentration of 10.3 μmol·gí1), the concentration in solution remained
below detection limit (0.20 μmol·gí1) for the entire reaction. Since the total amount of gold present
in the reactor is in all cases 41 μmol, we conclude that adsorption on the catalyst support is
responsible for the missing quantities of BnOOH. The adsorption on the support is captured by site
coverages assuming to follow Langmuir behavior:

270

θ BnOH , Al O =
2 3

ads
K BnOH
, Al2 O3 C BnOH

θ BnOOH , Al O =
2 3

Where:

(28)

ads
ads
1 + K BnOH
, Al2 O3 C BnOH + K BnOOH , Al2 O3 C BnOOH
ads
K BnOOH
, Al2 O3 C BnOH

(29)

ads
ads
1 + K BnOH
, Al2 O3 C BnOH + K BnOOH , Al2 O3 C BnOOH

θ i , Al O = occupancy of the surface sites (-)
2

3

í1
Kiads
, Al2O3 = adsorption constant of compound i on the support (g·mol )
ads
ads
From these experimental data the adsorption constants K BnOH
, Al2O3 and K BnOOH , Al2O3 were

estimated to be 18.1 g·molí1 and 2.47 × 104 g·molí1, respectively, and the total adsorption site
density was found to be 1.12 mmol·gcatí1. With a specific surface area of 230 m2·gí1 and using
Avogadro’s number, the corresponding site density on the support is equivalent to 2.92 sites per
nm2, which is in good agreement with Matulewicz et al. [31] who report a value of 2 sites/nm2 for
their Ȗ-alumina.
This yields a relation between the overall acid concentration in the vessel, CBnOOH, tot, and the
actual acid concentration in the liquid phase, CBnOOH:
C BnOOH = C BnOOH , tot − σ OH , Al2O3

Where:

ads
K BnOOH
wcat
, Al2 O3 C BnOOH
ads
ads
wliq 1 + K BnOH , Al2O3 C BnOH + K BnOOH
, Al2 O3 C BnOOH

(30)

CBnOOH, tot = BnOOH concentration if no adsorption would take place (mol·gí1)
ıOH, Al2O3 = adsorption site concentration on the catalyst (mol·gcatí1)
wcat = amount of catalyst in the reactor (gcat)
wliq = amount of liquid in the reactor (g)

Since the amount of BnOH in the liquid phase is about three orders of magnitude larger than the
amount of surface adsorption sites, the influence of BnOH adsorption on the concentration in the
liquid phase is neglected. The actual acid concentration in the liquid phase can therefore be
calculated directly from this quadratic equation with respect to CBnOOH:
CBnOOH =

Where:

T = CBnOOH , tot −

(

ads
§
4CBnOOH , tot 1 + K BnOH
1¨
, Al2O3 C BnOH
T + T2 +
ads
2 ¨¨
K BnOOH , Al2O3
©

(1 + K

ads
BnOH , Al2O3

K

CBnOH

ads
BnOOH , Al2 O3

)−σ

OH , Al2O3

) ·¸
¸¸
¹

(31)

wcat
wliq

The strong adsorption of BnOOH could be the main cause of the very low activity of catalysts
reused for the same experiment after a test in absence of a base, since it was observed that this acid
remains on the catalyst during rinsing with toluene and also during boiling in water (Figure 13).
The amount of BnOOH that remains on the catalyst in these ‘second runs’ (Figures 2b, 6b, 7b
and 12) is not known but an estimate can be made. Assuming that the BnOOH concentration in the
liquid equals its detection limit of 0.20 μmol·gí1, a coverage of 0.88 is found from our simulations.
Thus, it is tentatively concluded that although the amounts of acid formed in the previous runs is

271
very low (often below the detection limit), the coverage is close to 1. Since the inhibition with a
reused catalyst was very small or negligible in the presence of K2CO3 or KF, it was assumed for
simplicity that in these cases all acid was removed, which seems acceptable in view of the time
(approximately half an hour, the time needed for heating plus temperature stabilization) that the
catalyst particles were in the close vicinity of the K2CO3 or KF crystals at reaction conditions
(80 °C and well mixed) before the BnOH was added.
The model also accounts for the time allowed between adding the BnO or the BnOOH to the
reactor and adding the main reactant BnOH, in all cases about 20 min at reaction conditions
(additionally from the heating time). In the case of BnO this causes the formation of significant
amounts of BnOOH, inhibiting the reaction (Figure 3).
Besides the three rate constants k1, k2, k3, and the three adsorption equilibrium constants,
ads
ads
KBnOOH, K BnOH
, Al2O3 and K BnOOH , Al2O3 , there are several other unknown parameters, which are related
to the effect of the base present. The experimental results show that the main effects of either
K2CO3 or KF are (i) a decrease of the inhibition and (ii) an increase in formation of the
BnOOBn ester.
In order to account for effect (i), it is assumed that the bases react to potassium benzoate with the
BnOOH formed. Although acid-base reactions are typically instantaneous reactions, a finite rate is
assumed to account for the transport from the catalyst pores to the insoluble crystals of K2CO3 or
KF. The reaction rate is assumed to be first-order with respect to the acid and independent of the
catalyst concentration:

rAcBase = k AcBaseCBnOOH

(32)

Effect (ii) is accounted for by defining parameters to describe the strength of either K2CO3 or
KF in the catalysis of the esterification of BnO with BnOH:
No base: Sbase = 1

(33)

K 2 CO3 : Sbase = 1 + S K 2CO3

(34)

KF : Sbase = 1 + S K 2CO3 S KF / K 2CO3

(35)

SK2CO3 is defined as the base strength of K2CO3 and SKF/K2CO3 as the relative base strength of KF
compared to K2CO3.
Since the experiments were performed in batch operation in a vessel that is assumed to be
ideally stirred, the reactor model used to describe the process is the batch reactor model:
dCi , L
dt

Where:

= Ccat ¦ν i , j rj
j

Ci, L = concentration of component i in the liquid phase (mmol·gí1)
t = time (s)
Ccat = catalyst concentration (mg·gí1)
Ȟi, j = stoichiometric coefficient of component i in reaction j (-)
rj = reaction rate of reaction j (mmol·mgcatí1·sí1)

(36)

272
The complete experimental dataset used for the parameter estimation contains 140 experimental
data points obtained in 14 batch experiments at various conditions. The parameter estimation was
carried out using the software package Athena Visual Studio [32], applying Bayesian estimation for
multiresponse experiments using the full covariance matrix [33]. The concentrations (expressed in
mol·gí1) of the four measured liquid components were used as the input for the objective function
to be minimized. Since the concentrations of the BnOOH and the ester were typically up to two
orders of magnitude smaller than those of the BnOH and BnO, the weight of BnOOBn was set at
10 and that of BnOOH at 100. For experiments with a very low conversion rate of BnOH, the
weight of BnO was increased to 10 as well.
With the exception of the experiment of Figure 4, BnOOH concentrations were mostly below
the detection limit of 0.20 ȝmol·gí1. Since the model predictions are very sensitive to the
acid concentration, it was necessary to estimate acid concentrations for experiments where these
were not detectable. In those cases, acid concentrations were arbitrarily assumed to be half of the
detection limit.
These data lead to fits of the seven kinetic parameters to the experimental dataset and the results
are shown in Table 2.
Table 2. Optimal estimates of the kinetic parameters using all experimental data.
Parameter

Unit

k1

/g·mgcatí1·sí1

k2

/g·mgcatí1·sí1
/g2·mmolí1·mgcatí1·sí1
/sí1
/g·mmolí1

k3
kAcBase
KBnOOH

K
K

ads
BnOH , Al2O3

í1

/g·mmol

ads
BnOOH , Al2O3

í1

/g·mmol

SK2CO3

/-

SKF/K2CO3
SSR*
(*) SSR =

v

n

¦w ¦( C

i , k , exp

i

i =1

Estimate

k =1

2.69 × 10

í3

2.37 × 10í4
í4

6.6 × 10

95% confidence range
Value

Relative/%
í4

± 19

± 5.9 × 10í5

± 25

± 5.0 × 10
± 3.5 × 10

í4

± 54

0.71

± 0.15

± 21

1.23 × 104

± 1.5 × 103

± 12

18.1

fixed

2.47 × 10

4

fixed

1.39

± 0.81

± 58

/-

1.37

± 0.39

± 29

/-

0.414

-

-

− Ci, k , mod ) (sum of the squared residuals) (wi = weight factor for response i,

v = number of responses, Ci,

2

k, exp

= experimental response of component i in experiment

k, Ci, k, mod = model response of component i in experiment k, n = number of experiments (samples).

While relatively good fits of predicted concentrations of BnOH, benzaldehyde and BnOOBn
with time are evident in Figures 2a and 6a for first batch runs with and without K2CO3, relatively
poor fits of benzaldehyde concentration are observed for second-time runs (Figures 2b and 6b) and
the first run with KF (Figure 7a). The poorer fit of the second runs in the presence of K2CO3 is
understandable since it was assumed in the model that all BnOOH was removed by the base in
between the experiments, which is probably not completely justified, indicating that some acid or

273
another inhibiting species remains on the catalyst. In all other experiments, approximate fits of
concentrations of one or both products or of all three species (alcohol, aldehyde and ester) are
observed (see Figures 3, 4, 7b, and 10–12). Thus, variations in how well the fit follows the data are
a logical consequence of attempting to simulate in a single model a wide range of concentrations
with and without base and in the absence and presence of strongly inhibiting aldehyde and acid
product species. Moreover, the model did not include effects of water and was limited to the three
most important reactions and four most important species. The 54% and 58% relative confidence
intervals for k3 and SK2CO3 originate from the strong correlation between these two parameters,
which is discussed later in this section. The approximate nature of the model can be attributed to:
(1) Assumptions that are only approximately valid, e.g., (a) arbitrary estimates of benzoic acid
concentration, and (b) the assumption that benzoic acid is completely removed from the catalyst by
interaction with insoluble K2CO3 or KF crystals.
(2) By practical necessity, the limited scope of the mechanistic scheme, e.g., (a) neglecting
effects of adsorption or inhibition of some species such as coverages of BnOH of BnOOBn and (b)
neglecting the positive effect of water.
Nevertheless, the model provides (1) accurate predictions of initial reaction rate for oxidation of
BnOH to BnO on a gold/alumina catalyst and (2) approximate predictions of the effects of BnOOH
inhibition and the neutralizing effect of potassium salts to alleviate this inhibition.
The value of the rate constant for the reaction of the BnOOH with the base to potassium
benzoate, kAcBase, represents the characteristic time for the transport of BnOOH from a catalytic site
to the K2CO3 or KF crystals. The order of magnitude can be compared with an estimate of the
characteristic time for diffusion [34] of BnOOH through the catalyst pores to the liquid bulk,
obtained from a typical diffusion distance (one third of the catalyst particle size 50 ȝm, estimated
from the sieve mesh size) and an effective diffusivity of 4.8 × 10í10 m2·sí1 (estimated using Wilke
and Chang’s relationship [35], using porosity-tortuosity ratio of 0.14, based on data of similar
aluminas and catalysts):
τ diff , BnOOH =

Where:

( 13 d p )

2

2 × Deff , BnOOH

(37)

IJdiff,BnOOH = diffusion time (s)
dp = catalyst particle diameter (m)
Deff, BnOOH = effective diffusivity (m2·sí1)

This yields a typical diffusion time of 0.3 s, which is about 4.8 times smaller than the typical
time (kAcBase)í1 = 1.41 s. The latter seems a plausible value in view of the additional transport
resistance that might be caused by the transfer from the external catalyst surface towards the
K2CO3 or KF crystals.
The correlation matrix for the estimated parameters, shown in Table 3, shows that the strongest
correlation occurs between k3 and SK2CO3 with a correlation coefficient of í0.97, which is in line
with our conclusion that the ester formation in our system is base-catalyzed. All correlations

274
between the parameters justify maintaining all parameters in the model since these do not exceed
the value of 0.99, which is accepted as the limit for a proper parameter estimation [36].
Table 3. Correlation matrix between all the parameters estimated using the optimized
kinetic model.
Parameter

k1

k2

k3

kAcBase

KBnOOH

SK2CO3

SKF/K2CO3

k1
k2
k3

1
0.52
0.06
0.12
0.51
0.02
í0.01

1
í0.13
í0.17
í0.30
0.12
í0.02

1
í0.02
0.09
í0.97
0.33

1
0.64
0.06
í0.05

1
í0.01
í0.01

1
í0.45

1

kAcBase
KBnOOH
SK2CO3
SKF/K2CO3

3. Experimental Section
Toluene (anhydrous, 99.8%), benzyl alcohol (>99%), potassium carbonate (>99.0%),
tetradecane (> 99%), potassium fluoride (>99.99%) and phosphorus pentoxide desiccant were
supplied by Sigma Aldrich and were used without further purification. AUROlite™ catalyst
(Au/Al2O3 1 wt.%, Au average particle size: 2–3 nm, specific surface area: 200–260 m2·gí1, from
supplier specifications) was supplied by Strem Chemicals in the form of extrudates. The extrudates
were crushed and sieved to a particle size <71 μm, thereby excluding diffusion limitations during
catalytic experiments as verified using different catalyst particle sizes. The resulting powder was
stored in a well-sealed container at 4 °C and in the dark. Catalytic testing under dry conditions was
performed with this powder used as such. For the catalytic tests involving water, the desired
amount of this powder was suspended in Milli-Q® water (18.2 Mȍ·cm) under sonication for
30 min, and then vacuum filtered (using a Büchner funnel). The resulting moist catalyst was
collected from the filter with a spatula and used as such for catalytic testing. The mass difference
before and after this step indicates that around 0.5 g of water is adsorbed per gram of catalyst.
Catalytic experiments were carried out in a 100 mL round-bottom vessel, the inner diameter of
which is 60 mm. The vessel was equipped with a reflux condenser and Teflon baffles, and
mechanically stirred at 1300 rpm with a 4-blade Teflon impeller. Upon varying the catalyst
quantity in preliminary tests, an initial reaction rate proportional to the catalyst quantity was
observed, indicating that mass transport limitations were absent. In a typical catalytic test, 3.04 g of
K2CO3 and 0.8 g of AUROlite™ are introduced in the vessel together with 80 mL of toluene. Two
complementary tests were performed using 2.07 g of KF instead of K2CO3. The vessel was heated
to 80 °C by means of an oil bath, and 200 mL·miní1 of air was bubbled through the reaction
mixture via a glass frit. When the temperature was stabilized, 2.4 g of BnOH was introduced using
a syringe, constituting the beginning of the test (t = 0 min). Small samples of 300 μL were taken at
recorded times and filtered from catalyst and K2CO3 powders with a 13 mm syringe Teflon filter of
0.2 μm pore size (diameter: 13 mm; pore size: 0.2 μm; PTFE membrane; VWR International) and

275
introduced in a GC sample vial together with 20 μL of tetradecane, the latter being used as
internal standard.
GC analyses were performed using a Varian CP-3380 equipped with a FID detector and
a CP-Sil 8 CB cat. no. 7453 column (length: 50 m; diameter: 0.25 mm; coating thickness:
0.25 μm). The initial temperature of the GC oven was 150 °C and was maintained for 4 min, then
increased with 100 °C·miní1 to 220 °C and then maintained at 220 °C for 6.3 min. After testing, the
catalyst was recovered by vacuum filtration, washed with 80 mL of toluene at room temperature
and stored over P2O5 in an evacuated desiccator. In the case of the water treated catalyst samples,
the catalyst was washed with 80 mL of toluene, followed by extensive washing with about 250 mL
of Milli-Q® water (18.2 Mȍ·cm) at room temperature.
Diffuse Reflectance Infra-Red Fourier Transform Spectroscopy (DRIFTS) spectra were
recorded on a Nicolet model 8700 spectrometer, equipped with a high-temperature DRIFTS cell,
and a DTGS-TEC detector. The spectra were recorded with 256 scans at 4 cmí1 resolution from
4000 to 500 cmí1 using potassium bromide (KBr) to perform background subtraction. The samples
were pre-treated at 473 K for 1 h in a helium flow of 20 mL·miní1.
4. Conclusions
Our study shows that benzoic acid or compounds formed from benzoic acid cause catalyst
inhibition in benzyl alcohol oxidation in toluene and in absence of a base. The introduction of a
potassium salt as a base prevents this inhibition by neutralizing the benzoic acid formed. Basic
conditions result in a decrease in selectivity to benzaldehyde and in an increase of ester production.
The enhanced ester formation probably occurs via condensation of alkoxy species (formed by
alcohol deprotonation by the base) with benzaldehyde under oxidative conditions, and is not the
result of an increased benzoic acid production followed by esterification, as might be expected
under acidic conditions. Water appears to have no influence on inhibition, but may enhance the
effect of the base described above by improved dissolution. Although effects of water were not
modeled, a kinetic effect for water cannot be excluded.
The concentration versus time data of the batch experiments in this study, which covered a wide
range and included effects of acid inhibition and base, were fitted to a comprehensive kinetic model
for (1) the primary reaction, oxidation of benzyl alcohol to benzaldehyde; (2) secondary oxidation
of benzaldehyde to benzoic acid; and (3) secondary esterification of benzyl alcohol and
benzaldehyde to benzyl benzoate. Effects of base (potassium salts) were also included in the model.
The resulting model predicts concentration-time trends approximately well, including inhibition by
benzoic acid and the neutralization of benzoic acid by potassium salts, forming potassium
benzoate. A precise fit of the model to experimental data was observed in two first batch runs, with
and without K2CO3. Variations in how well the fit follows the data are a logical consequence of
attempting to simulate in a single model a wide range of concentrations with and without base and
in the absence and presence of product species such as aldehyde and strongly inhibiting acid.
Moreover, the model did not include effects of water and was limited to the three most important
reactions and four most important species.

276
Nevertheless, the model provides (1) accurate predictions of initial reaction rate for oxidation of
benzyl alcohol to benzaldehyde on a gold/alumina catalyst and (2) approximate predictions of
the effects of benzoic acid inhibition and the neutralizing effect of a potassium base to alleviate
this inhibition.
Acknowledgments
The Dutch National Research School Combination Catalysis Controlled by Chemical Design
(NRSC-Catalysis) is gratefully acknowledged for financial support.
Conflict of Interest
The authors declare no conflict of interest.
References
1.

Davis, S.E.; Ide, M.S.; Davis, R.J. Selective oxidation of alcohols and aldehydes over
supported metal nanoparticles. Green Chem. 2013, 15, 17–45.
2. Lilga, M.A.; Hallen, R.T.; Gray, M. Production of oxidized derivatives of
5-Hydroxymethylfurfural (HMF). Topics Catal. 2010, 53, 1264–1269.
3. Zope, B.N.; Davis, R.J. Inhibition of gold and platinum catalysts by reactive intermediates
produced in the selective oxidation of alcohols in liquid water. Green Chem. 2011, 13,
3484–3491.
4. Haruta, M.; Yamada, N.; Kobayashi, T.; Iijima, S. Gold catalysts prepared by coprecipitation
for low-temperature oxidation of hydrogen and of carbon monoxide. J. Catal. 1989, 115,
301–309.
5. Alhumaimess, M.; Lin, Z.; Weng, W.; Dimitratos, N.; Dummer, N.F.; Taylor, S.H.; Bartley,
J.K.; Kiely, C.J.; Hutchings, G.J. Oxidation of benzyl alcohol by using gold nanoparticles
supported on ceria foam. ChemSusChem 2012, 5, 125–131.
6. Deplanche, K.; Mikheenko, I.P.; Bennett, J.A.; Merroun, M.; Mounzer, H.; Wood, J.; MacAskie, L.E.
Selective oxidation of benzyl-alcohol over biomass-supported Au/Pd bioinorganic catalysts.
Topics Catal. 2011, 54, 1110–1114.
7. Guo, H.; Kemell, M.; Al-Hunaiti, A.; Rautiainen, S.; Leskelä, M.; Repo, T. Gold-palladium
supported on porous steel fiber matrix: Structured catalyst for benzyl alcohol oxidation and
benzyl amine oxidation. Catal. Commun. 2011, 12, 1260–1264.
8. Guo, H.; Al-Hunaiti, A.; Kemell, M.; Rautiainen, S.; Leskelä, M.; Repo, T. Gold catalysis
outside nanoscale: Bulk gold catalyzes the aerobic oxidation of ʌ-activated alcohols.
ChemCatChem 2011, 3, 1872–1875.
9. Hao, Y.; Hao, G.P.; Guo, D.C.; Guo, C.Z.; Li, W.C.; Li, M.R.; Lu, A.H. Bimetallic Au-Pd
Nanoparticles Confined in Tubular Mesoporous Carbon as Highly Selective and Reusable
Benzyl Alcohol Oxidation Catalysts. ChemCatChem 2012, 4, 1595–1602.
10. Heeskens, D.; Aghaei, P.; Kaluza, S.; Strunk, J.; Muhler, M. Selective oxidation of ethanol in
the liquid phase over Au/TiO2. Phys. Status Solidi. A 2013, 250, 1107–1118.

277
11. Mallat, T.; Baiker, A. Oxidation of alcohols with molecular oxygen on solid catalysts.
Chem. Rev. 2004, 104, 3037–3058.
12. Quintanilla, A.; Butselaar-Orthlieb, V.C.L.; Kwakernaak, C.; Sloof, W.G.; Kreutzer, M.T.;
Kapteijn, F. Weakly bound capping agents on gold nanoparticles in catalysis: Surface poison?
J. Catal. 2010, 271, 104–114.
13. Sá, J.; Taylor, S.F.R.; Daly, H.; Goguet, A.; Tiruvalam, R.; He, Q.; Kiely, C.J.; Hutchings, G.J.;
Hardacre, C. Redispersion of gold supported on oxides. ACS Catal. 2012, 2, 552–560.
14. Mallat, T.; Baiker, A. Oxidation of alcohols with molecular oxygen on platinum metal
catalysts in aqueous solutions. Catal. Today 1994, 19, 247–283.
15. Zope, B.N.; Hibbitts, D.D.; Neurock, M.; Davis, R.J. Reactivity of the gold/water interface
during selective oxidation catalysis. Science 2010, 330, 74–78.
16. Moulijn, J.A.; van Diepen, A.E.; Kapteijn, F. Catalyst deactivation: Is it predictable?
What to do? Appl. Catal. 2001, 212, 3–16.
17. Keresszegi, C.; Bürgi, T.; Mallat, T.; Baiker, A. On the role of oxygen in the liquid-phase
aerobic oxidation of alcohols on palladium. J. Catal. 2002, 211, 244–251.
18. Dimitratos, N.; Villa, A.; Wang, D.; Porta, F.; Su, D.; Prati, L. Pd and Pt catalysts modified by
alloying with Au in the selective oxidation of alcohols. J. Catal. 2006, 244, 113–121.
19. Rodríguez-Reyes, J.C.F.; Friend, C.M.; Madix, R.J. Origin of the selectivity in the
gold-mediated oxidation of benzyl alcohol. Surf. Sci. 2012, 606, 1129–1134.
20. Fang, W.; Chen, J.; Zhang, Q.; Deng, W.; Wang, Y. Hydrotalcite-supported gold catalyst for
the oxidant-free dehydrogenation of benzyl alcohol: Studies on support and gold size effects.
Chem. A Eur. J. 2011, 17, 1247–1256.
21. Kimling, J.; Maier, M.; Okenve, B.; Kotaidis, V.; Ballot, H.; Plech, A. Turkevich method for
gold nanoparticle synthesis revisited. J. Phys. Chem. B 2006, 110, 15700–15707.
22. Juan-Alcañiz, J.; Ferrando-Soria, J.; Luz, I.; Serra-Crespo, P.; Skupien, E.; Santos, V.P.; Pardo, E.;
Llabrés i Xamena, F.X.; Kapteijn, F.; Gascon, J. The oxamate route, a versatile
post-functionalization for metal incorporation in MIL-101(Cr): Catalytic applications of Cu,
Pd, and Au. J. Catal. 2013, 307, 295–304.
23. Fristrup, P.; Johansen, L.B.; Christensen, C.H. Mechanistic investigation of the gold-catalyzed
aerobic oxidation of alcohols. Catal. Lett. 2008, 120, 184–190.
24. Abad, A.; Corma, A.; García, H. Catalyst parameters determining activity and selectivity of
supported gold nanoparticles for the aerobic oxidation of alcohols: The molecular reaction
mechanism. Chem. A Eur. J. 2008, 14, 212–222.
25. Pina, C.D.; Falletta, E.; Rossi, M. Update on selective oxidation using gold. Chem. Soc. Rev.
2012, 41, 350–369.
26. Daté, M.; Okumura, M.; Tsubota, S.; Haruta, M. Vital role of moisture in the catalytic activity
of supported gold nanoparticles. Angew. Chem. Int. Ed. 2004, 43, 2129–2132.
27. Daniells, S.T.; Overweg, A.R.; Makkee, M.; Moulijn, J.A. The mechanism of low-temperature
CO oxidation with Au/Fe2O3 catalysts: A combined Mössbauer, FT-IR, and TAP reactor
study. J. Catal. 2005, 230, 52–65.

278
28. Yang, X.; Wang, X.; Liang, C.; Su, W.; Wang, C.; Feng, Z.; Li, C.; Qiu, J. Aerobic oxidation
of alcohols over Au/TiO2: An insight on the promotion effect of water on the catalytic activity
of Au/TiO2. Catal. Commun. 2008, 9, 2278–2281.
29. Chang, C.R.; Yang, X.F.; Long, B.; Li, J. A water-promoted mechanism of alcohol oxidation
on a Au(111) surface: Understanding the catalytic behavior of bulk gold. ACS Catal. 2013, 3,
1693–1699.
30. Socrates, G. Infrared and Raman Characteristic Group Frequencies. Tables and Charts, 3rd
ed.; John Wiley & Sons Ltd: Heidelberg, Germany, 2004; pp. 50–167.
31. Matulewicz, E.R.A.; Kerkhof, F.P.J.M.; Moulijn, J.A.; Reitsma, H.J. Structure and activity of
fluorinated alumina. 1. Determination of the number of protonic sites by an infrared study of
adsorbed pyridines. J. Colloid Interface Sci. 1980, 77, 110–119.
32. Athena Visual Studio, Software for Modeling, Estimation and Optimization, Version 14.2.
Available online: www.AthenaVisual.com 1997–2009 (accessed 25 September 2013).
33. Stewart, W.E.; Caracotsios, M.; Sorensen, J.P. Parameter estimation from multiresponse data.
AIChE J. 1992, 38, 641–650.
34. Atkins, P.W. Physical Chemistry, 5th ed.; W.H. Freeman: New York, NY, USA, 1994.
35. Perry, R.H.; Green, D.W. Perry's Chemical Engineers' Handbook, 7th ed.; McGraw-Hill: New
York, NY, USA, 1997.
36. Stewart, W.E.; Caracotsios M. Computer-Aided Modelling of Reactive Systems; John Wiley &
Sons, Inc., Hoboken, NJ, USA, 2008.

279

Investigation of the Deactivation Phenomena Occurring in the
Cyclohexane Photocatalytic Oxidative Dehydrogenation on
MoOx/TiO2 through Gas Phase and in situ DRIFTS Analyses
Vincenzo Vaiano, Diana Sannino, Ana Rita Almeida, Guido Mul and Paolo Ciambelli
Abstract: In this work, the results of gas phase cyclohexane photocatalytic oxidative
dehydrogenation on MoOx/SO4/TiO2 catalysts with DRIFTS analysis are presented. Analysis of
products in the gas-phase discharge of a fixed bed photoreactor was coupled with in situ
monitoring of the photocatalyst surface during irradiation with an IR probe. An interaction between
cyclohexane and surface sulfates was found by DRIFTS analysis in the absence of UV irradiation,
showing evidence of the formation of an organo-sulfur compound. In particular, in the absence of
irradiation, sulfate species initiate a redox reaction through hydrogen abstraction of cyclohexane
and formation of sulfate (IV) species. In previous studies, it was concluded that reduction of the
sulfate (IV) species via hydrogen abstraction during UV irradiation may produce gas phase SO2
and thereby loss of surface sulfur species. Gas phase analysis showed that the presence of MoOx
species, at same sulfate loading, changes the selectivity of the photoreaction, promoting the
formation of benzene. The amount of surface sulfate influenced benzene yield, which decreases
when the sulfate coverage is lower. During irradiation, a strong deactivation was observed due to
the poisoning of the surface by carbon deposits strongly adsorbed on catalyst surface.
Reprinted from Catalysts. Cite as: Vaiano, V.; Sannino, D.; Almeida, A.R.; Mul, G.; Ciambelli, P.
Investigation of the Deactivation Phenomena Occurring in the Cyclohexane Photocatalytic
Oxidative Dehydrogenation on MoOx/TiO2 through Gas Phase and in situ DRIFTS Analyses.
Catalysts 2013, 3, 978-997.
1. Introduction
Photocatalytic oxidation reactions have been widely used in processes such as the
decontamination of water and air [1–8]. However, applications of heterogeneous photocatalysis for
the synthesis of compounds of commercial importance have been considered only in recent years.
The most studied photocatalytic reactions occur in slurry systems. Among them, partial oxidation of
cyclohexane is an important commercial reaction, as the resultant products, cyclohexanol and
cyclohexanone, are precursors in the syntheses synthesis of adipic acid, in turn intermediates in the
production of nylon [9]. In their review, Maldotti et al. [10] reported the main aspects concerning
the photocatalytic oxidation of cyclohexane. Li et al. [11] showed that the quantum size and
surface state are key factors governing the selectivity in photoxidation on TiO2 nanoparticles [12,13].
The photo-oxidation of cyclohexane on titanium dioxide was also investigated in neat cyclohexane
and in various solvents showing an influence of the solvent media on the cyclohexane oxidation
rate and on the selectivity to cyclohexanol and cyclohexanone [14,15].
From theoretical and practical points of view, the ideal solvent for the photo-oxidation of
cyclohexane is one that minimizes the strengths of adsorption of the desired products on titanium

280
dioxide and does not compete with cyclohexane and oxygen for adsorption sites. Otherwise, the
solvent could be strongly adsorbed but is non-reactive with itself upon forming a radical on the
illuminated titanium dioxide surface [16].
Supported transition-metal oxides can absorb light, and the transferred energy can be used to
activate C-H bonds in saturated hydrocarbons, a chemical step that is generally unselective [6]. In
this context, cyclohexanol, cyclohexanone, and polyoxygenates have been formed from
cyclohexane when polyvanadate or polyoxytungstate were supported on several oxides [17,18].
It was shown that attenuated total reflection Fourier transform infrared (ATR-FTIR) is a suitable
way to monitor in real time and in situ the light-induced heterogeneous oxidations [19], allowing deeper
knowledge of the complex phenomena occurring under irradiation to be obtained. ATR-FTIR
technique was also performed to study the photodegradation of organic pollutants in water on
TiO2 [20].
The generally low efficiency associated with liquid phase photocatalytic reactions, which
typically occur at low conversion, coupled with the difficulty of separating catalyst and products in
a liquid has motivated research of gas-solid systems, e.g., catalysts active in the gas-phase partial
oxidation of cyclohexane. However, selective photoxidation of cyclohexane yields different
reaction products in gas phase relative to liquid phase. In particular, cyclohexene or benzene are
selectively obtained through gas phase oxidative dehydrogenation of cyclohexane on MoOx/TiO2
photocatalysts, with UV illumination both in fixed [21] and fluidized bed reactors [22–24]. Higher
molybdenum loadings improved the benzene selectivity, whereas with only titania, total conversion
to carbon dioxide is obtained. The selective formation of benzene is attributed to the poisoning by
Mo-species of unselective sites on the titania surface which otherwise totally oxidize cyclohexane
to CO2 and water [24–26].
A mechanism for the catalytic photo-oxidative dehydrogenation of cyclohexane to benzene on
MoOx/TiO2 was recently proposed by analyzing the gas-phase coming from a photoreactor [21].
This mechanism involves dehydrogenation of cyclohexane to cyclohexene followed by
oxy-dehydrogenation to benzene on molybdenum oxide active sites via a detailed sequence of
elementary steps [21]. In the same paper, several hypotheses relating to the role of sulfate species in
promoting the selectivity to benzene formation were postulated: (1) sulfate present on TiO2 surface
may facilitate hydrogen abstraction from an adsorbed cyclohexane molecule, owing to its strong
acid properties, or (2) it may participate in the re-oxidation step of the octahedrally-coordinated
polymolybdate surface species [21]. With regard to the influence of sulfate concentration,
Ciambelli et al. [24,26,27] showed that during photooxidative dehydrogenation of cyclohexane, the
presence of sulfate species on the surface of titania favor a high benzene yield. Cyclohexene was
produced in low concentration and CO2 was not detected in gas-phase. Enhanced photooxidative
dehydrogenation activity of MoOx/TiO2 catalysts was attributed to the increase in surface acidity by
sulfate species, which enhances hydrocarbon adsorption coverage on the catalyst surface.
On the other hand, the addition of sulfate to MoOx/Ȗ-Al2O3 catalysts was found to promote the
selective mono-oxidative dehydrogenation of cyclohexane to cyclohexene [28,29]. An optimum in
MoO3 and SO4 loadings were found to be 8 and 2.4 wt%, respectively (corresponding to MoO3/SO4
molar ratio equal to 2.22), while a decrease in the catalytic activity at higher sulfate loading was

281
ascribed to MoOx decoration by sulfates. By studying the influence of the preparation method and
molybdenum loading on sulfated MoOx/Ȗ-Al2O3, it was shown that selectivity to benzene increases
with increasing catalyst acidity, as the latter favors cyclohexene adsorption and thus its conversion to
benzene. Close proximity of surface sulfates to octahedral polymolybdate appears to be a key
parameter promoting photoactivity of these catalysts [28].
While much has been learned from the previous work, an in situ study of photocatalyst surfaces
under UV irradiation is nevertheless needed to gain a better understanding of the surface
phenomena. Typical problems of molybdenum based catalysts that need to be addressed include the
role of sulfate and the deactivation occurring during gas-phase photocatalytic oxidative
dehydrogenation of cyclohexane. Thus, in this study we present the results of gas phase
cyclohexane photocatalytic oxidative dehydrogenation on MoOx/TiO2 through the analysis of
products present in gas-phase exit of the reactor, complemented by in situ DRIFTS analysis, to
monitor the photocatalyst surface in the absence and during UV irradiation.
2. Results and Discussion
2.1. Catalysts Characterization
The list of catalysts investigated and their characterization results are reported in Table 1.
Table 1. List of catalysts and their characterization results.
Catalyst
DT05
4.7MoDT0.5
DT2
2MoDT2
8MoDT2
a

MoO3 a
loading, wt%
0
4.6
0
1.8
7.6

SO4 b, wt%

SSA, m2/g

Mo=O stretching (Raman) c, cmí1

0.6
0.6
2.3
2.2
2.2

67
68
71
71
63

956
953
978

evaluated by ICP-MS; b evaluated by TG-MS analysis; c these results are obtained from previous studies [21,25,26].

Raman results of all the samples were reported in previous works [21,25,26]. In summary,
Raman bands related to Mo-species can be observed in the range 820–1000 cmí1. For all Mo-based
catalysts, bands at 956 cmí1 on 4.7Mo/DT05, at 953 cmí1 on 2MoDT2, at 978 cmí1 on 8MoDT2
and at 956 cmí1 on 8Mo2S are observed and assigned to terminal Mo=O stretching of octahedral
MoOx species [30]. The increasing of wavenumber was an indication of a higher polymerization
degree of Mo species at increasing loadings, corresponding to a Mo-nuclearity likely ranging
from 7 to 12 [21].
2.2. Adsorption Measurement in the Absence of UV Irradiation
2.2.1. DRIFTS Analysis
Cyclohexane was admitted in absence of UV irradiation on the catalysts 2MoDT2, 4.7MoDT05,
DT2 and DT05, and in Figure 1 the spectra after 1 h in presence of gas mix 2 are reported. The
cyclohexane peak at 1450 cmí1 appears to be present on the overall tested photocatalysts, indicating the

282
occurring of the adsorption of the hydrocarbon. This first finding was coherently with the occurring of
Langmuir–Hinshelwood mechanism [23]. Together with the cyclohexane adsorption, a carbonyl
vibration (~1681 cmí1) starts to grow while an organo-sulfur vibration, found at ~1354 cmí1,
decreases [31]. The decrease in the organo-sulfur vibration is more evident for the support with
higher sulfate amount (DT2). These results suggest that in the absence of irradiation sulfate species
initiate a redox reaction through hydrogen abstraction of cyclohexane and formation of sulfate (IV)
species. In previous studies, it was concluded that reduction of the sulfate (IV) species via
hydrogen abstraction may produce gas phase SO2 and thereby loss of surface sulfur species [27].
The extent of carbonyl vibration may be related to the content of either molybdenum or sulfate,
increasing in the combination of both sulfate and molybdenum. Simultaneously, an isosbestic point
on the bridging-OH TiO2 vibration, probably related to the displacement of water molecules by
cyclohexane molecules, can be observed in the range 3600–3750 cmí1. Surface sulfates could be
highly hydrated showing a weak absorption at 1635 cmí1 and a strong broad absorption ranging
from 3000 to 3500 cmí1 [32]. However, these signals are not visible because the main phenomenon
is the appearance of an organo-sulfur compound (found at ~1354 cmí1) resulting from a surface
reaction between sulfate and adsorbed cyclohexane that may have determined the loss of hydroxyls
bonded with sulfate.
Figure 2 shows in more detail the behavior of 2MoDT2 catalyst surface during dark adsorption
of gas mix 2. While the cyclohexane peak at 1450 cmí1 stabilizes in a few minutes, the signals at
1681 and 1354 cmí1 grow steadily in adsorption time, indicating a continuous surface modification.
Though the peaks evolve at the same time, their behavior is not completely parallel, which
indicates different surface phenomena. The 1354 cmí1 peak could originate from the oxidation of
an organic molecule coupled to the sulfate. Other sulfate related vibrations could not be followed.
Figure 3 shows the spectra after 1 h of gas mix 3 (with H2O in the feed) during dark adsorption
on 2MoDT2, 4.7MoDT05, DT2 and Hombikat. Once again, there is an increased carbonyl
vibration and a decreasing in organo-sulfur vibration. In the carbonyl area, only a single peak
occurred when H2O was absent in the gas feed (Figure 1), while when H2O is present (Figure 3) a
different product may be formed leading to 2 different vibrations in this area.
In addition, adsorbed water may be playing a role since it absorbs around 1630 cmí1 and an
isosbestic point may occur. To verify that the negative peak (~1354 cmí1) was related to sulfate, an
unsulfated anatase-TiO2 (Hombikat) was used as reference. The latter sample showed similar band
of adsorbed cyclohexane at 1450 cmí1, weak bands of carbonyl compounds with an intensity
similar to that one present on the sample with low sulfate coverage (4.7MoDT05), and
different hydroxyl bands. In particular, in the presence of adsorbed cyclohexane, the hydroxyl band
at 3680 cmí1 is negative, that is hindered by cyclohexane adsorption. The hydroxyls band at
3640 cmí1 is instead still present, while organo-sulfur (1350 cmí1) peak is not detectable. When
organo-sulfur compounds are present, the band of hydroxyls at 3640 cmí1 is negative for the
samples at high sulfate content. In the absence of sulfate, on the Hombikat catalyst, the band of
hydroxyls at 3680 cmí1, instead, is negative. For this reason, the isosbestic point on the
bridging-OH vibration is related to the displacement of water molecules by cyclohexane molecules,

283
preferably interacting with hydroxyls close to the sulfate, or with OH of titania in the absence
of sulfate.
Figure 1. DRIFT spectra after 1 h of dark adsorption for gas mix 2 on DT05, DT2,
2MoDT2 and 4.7MoDT05 catalysts. Cyclohexane spectrum is shown as reference.
0.3
0.2
0.1
cyclohexane

-0.0

Absorbance

-0.1

Isosbestic point
Bridging-OH in TiO2

-0.2

2MoDT2

-0.3

DT2

Organosulphur
vibration

-0.4
-0.5

4.7MoDT0.5

-0.6

DT05

-0.7
carbonyl

-0.8
3500

3000

2500
2000
Wavenumbers (cm-1)

1500

1000

Figure 2. DRIFT spectra during 1 h of dark adsorption for gas mix 2 on 2MoDT2 catalyst.

0.10
cyclohexane
1681

Absorbance

0.00

-0.10

-0.20
1354
-0.30

-0.40

1800

1700

1600

1500
Wavenumbers (cm-1)

1400

1300

1200

284

Figure 3. DRIFT spectra after 1 h of dark adsorption for gas mix 3 on DT2, 2MoDT2,
4.7MoDT05 and Hombikat catalysts. Cyclohexane spectrum is shown as reference.
0.3
0.2
0.1

Absorbance

0.0
-0.1

cyclohexane
Bridging-OH
on TiO2
4.7MoDT0.5

-0.2

Carbonyl

Organo-sulphur

2MoDT2

-0.3

DT2

-0.4
-0.5
-0.6

Hombikat

-0.7
3500

3000

2500
2000
Wavenumbers (cm-1)

1500

1000

2.2.2. Gas Phase Analysis in the Absence of UV Irradiation
At the run starting time, the gaseous feeding stream was passed through the reactor in the absence of
UV irradiation at room temperature. Adsorption of cyclohexane was observed by the decrease in its
concentration. Cyclohexane breakthrough time was about 10 min. Thereafter cyclohexane outlet
concentration slowly increased to reach the inlet value after about 50 min, indicating that adsorption
equilibrium of cyclohexane on the catalyst surface was attained. During this period, no products in
gas-phase were detected because they remained adsorbed on catalyst surface (as observed from
DRIFT analysis).
A support to the formation of organo-sulfur compounds in dark conditions observed by DRIFTS
could be found in an our paper [24] where the amount of cyclohexane adsorbed in the dark on
MoOx/TiO2 catalysts was linearly correlated with sulfate surface density. The linear increase in the
amount of the cyclohexane adsorbed in dark conditions has been determined to increase
cyclohexane reaction rate in presence of UV irradiation [24]. In the same paper, it was also
indicated that the adsorption of cyclohexane is correlated to the corresponding increase of catalyst
acidity. These results suggest that surface sulfate facilitates hydrogen abstraction from adsorbed
cyclohexane, increasing its storage on catalyst surface.

285

2.3. Photocatalytic Activity Tests
2.3.1. Gas Phase Analysis
Photocatalytic tests performed on titania based photocatalysts at low sulfate coverage are
reported in Figure 4. Cyclohexane conversion on DT05 and 4.7MoDT05 reached a maximum after
about 8 min of irradiation time and then decreased. The strong catalyst deactivation is particularly
evident after this time of irradiation.
On DT05, carbon dioxide was the only product detected in gas phase (Figure 5). It started to be
formed progressively; after that, the UV light was activated. In this case total carbon mass balance
was closed to about 100%. In presence of Mo species, on 4.7MoDT05, the main product was
benzene, whose maximum into the production was very delayed with respect both the maximum of
cyclohexane conversion, and the achieving of cyclohexane conversion steady state conditions. In
fact, the maximum outlet concentration of benzene was 7 ppm after 55 min. CO2 concentration
showed a steady state formation value of 10 ppm. Cyclohexene was formed before benzene in
agreement with the mechanism reported in [21], that considered consecutive steps of oxidative
dehydrogenation going through cyclohexene, as intermediate, to get finally benzene.
Figure 4. Cyclohexane conversion on DT05 and 4.7MoDT05 during UV irradiation.
cyclohexane conversion, %

40
35
30
25

DT05

20

4.7MoDT05

15
10
5
0
0

10

20

30

40

50

60

70

80

90

100

110

irradiation time, min

The comparison of cyclohexane conversion over DT2, 2MoDT2 and 8MoDT2 is shown in
Figure 6.
A maximum value was reached after about 5 min on all catalysts, then activity decreased
approaching a steady state conversion. On 2MoDT2, maximum cyclohexane conversion was about
45%, decreasing to about 10% in 90 min. With the same sulfate content, increasing Mo loading up
to 8 wt% MoO3, the initial maximum conversion was lower, about 15%, while steady state
conversion was 2.3% after 30 min. Therefore, the progressive coverage of the titania surface by
MoOx species resulted in a decreased initial and steady state cyclohexane conversions according to
our previous results [21]. On DT2 the initial maximum conversion was higher with respect to
8MoDT2 (about 25%), while steady state conversion was similar. Thus, with lower sulfate

286
coverage, steady state photocatalyst activity was smaller, so underlining the relevance of sulfate
presence on the photocatalysts performances.
Figure 5. (a) Benzene outlet concentration; (b) carbon dioxide outlet concentration and
(c) cyclohexene outlet concentration on DT05 and 4.7MoDT05 catalysts during
UV irradiation.
10

(a)

9

DT05

benzene, ppm

8

4.7MoDT05

7
6
5
4
3
2
1
0
400

(b)

350

CO2, ppm

300

DT05

250

4.7MoDT05

200
150
100
50

cyclohexene, ppm

0
2

(c)

1.6
1.2

DT05
4.7MoDT05

0.8
0.4
0
0

10

20

30

40

50

60

70

irradiation time, min

80

90

100 110

287
Figure 6. Cyclohexane conversion on DT2, 2MoDT2 and 8MoDT2 catalysts during
UV irradiation.
50

cyclohexane conversion, %

45
40
35

DT2

30

2MoDT2

25

8MoDT2

20
15
10
5
0
0

10

20

30

40

50

60

70

80

90

100

110

irradiation time, min

Figure 7. (a) Benzene outlet concentration, (b) carbon dioxide outlet concentration and
(c) cyclohexene outlet concentration on DT2, 2MoDT2 and 8MoDT2 catalysts during
UV irradiation.
25

benz ene, ppm

DT2

(a)

20

2MoDT2
8MoDT2

15
10
5

0
250

CO2, ppm

DT2

(b)

200

2MoDT2
8MoDT2

150
100
50

cyclohexene, ppm

0
2

(c)

1.6
1.2

DT2
2MoDT2

0.8

8MoDT2

0.4
0
0

10

20

30

40

50

60

70

irradia tion time , min

80

90

100 110

288
On all MoOx/TiO2 catalysts, reaction products were benzene, CO2 and few amount of
cyclohexene (Figure 7). Benzene concentration showed a maximum (17 ppm after 36 min on
8MoDT2 and 13 ppm after 45 min on 2MoDT2), while CO2 concentration was 100 ppm on
2MoDT2 and 6 ppm on 8MoDT2. As observed for DT05 catalysts, the only product observed on
DT2 was carbon dioxide and its concentration was the highest.
The higher cyclohexene concentration was recorded for 8MoDT2 whereas its formation was
negligible on 2MoDT2, reaching a value less than 1 ppm.
In summary, gas phase analysis of cyclohexane and reaction products evidenced that the presence of
MoOx species on the surface of titania, at same sulfate content, changes the selectivity of the catalyst
with increasing molybdenum loading. These results indicate that the interaction between titania and
supported molybdenum oxide plays an essential role in the catalyst selectivity. In addition also the
amount of surface sulfate influenced benzene yield, decreasing when the sulfate content is lower.
2.3.2. DRIFTS Analysis
The results of photocatalytic reaction are shown in Figure 8a,b for gas mix 2. The CH stretching
vibration of cyclohexane increases steadily during illumination probably because cyclohexane
adsorption was still occurring, possibly as the result of light induced dehydration of the surface,
making free new adsorption sites. So besides dehydration and the continued adsorption of
cyclohexane on TiO2 sites, also depletion of hydroxyl groups occurs during photoreaction, as
showed by the increasing negative absorption in the range 3200–3800 cmí1. This has been also
observed performing the reaction in liquid phase.
During UV irradiation, the peaks, likely to be ascribed to adsorbed cyclohexanone (1691 and 1671
cmí1) were found (Figure 4a), indicating the occurrence of an oxidation step in the reaction. The
formation of two different peaks could be related to the presence of different active sites at the surface,
but also to a different ketone. The peak at 1671 cmí1, in liquid phase cyclohexane photoxidation, was
attributed to a stronger adsorption site, while the 1691 cmí1 to a weaker cyclohexanone adsorption
site [19]. The position of these peaks is, however, different from the one observed during
cyclohexanone dark adsorption (1681cmí1), so during this step, the formation of different surface
products could be supposed. An examination of the literature was then performed with the aim to
support the latter hypothesis.
In a paper concerning the cyclohexene photo-oxidation over V/TiO2 catalysts [33],
cyclohexenone was formed during irradiation. One of the most intense band of this unsaturated
carbonyl compound is located at 1692 cmí1. So the signal at 1691 cmí1 could be ascribed also to
the presence of cyclohexenone, in turn formed from cyclohexene [33]. In addition, it should be
considered that C=C stretching frequency of alkenes lies in the range between 1680–1620 cmí1 [34],
therefore it overlaps to the signals due to adsorbed ketones. The formation of cyclohexene was also
found in cyclohexane oxidation at low temperatures using copper chloride in pyridine as
catalyst [35] and by liquid phase photocatalytic oxidation of cyclohexane on TiO2 in various
solvents [15]. In particular, by using dichloromethane as solvent, the presence of cyclohexene and
cyclohexenone was found. Thus it is not possible to exclude the formation of an unsaturated cyclic
hydrocarbon in the presence of molybdenum, taking into account also that the cyclohexene adsorbs

289
on the surface through the C=C double bond, and the contribution of the -HC=CH- stretching
would therefore be absent [36].
With regard to benzene detected in gas phase, it is not observed in DRIFTS analysis, in
agreement with the low affinity evinced in [21] for the titania surface. Both the two kind of
hydroxyls disappear during the photoreaction.
Figure 8. DRIFT spectra during 90 min of UV irradiation for gas mix 2 on 2MoDT2
catalyst in the range 1100–1925 cmí1 (a) and in the range 2750–3900 cmí1 (b).
cyclohexanone

(a)

0.070

1691
1671

0.060

Absorbance

0.050
0.040
0.030
0.020

water

0.010
0.000
-0.010
1900

1800

1700

1600

1500

1400

1300

1200

Wavenumbers (cm-1)
0.024

(b)

CH2
stretching

0.020
0.016

Absorbance

0.012

3040-3010 cm-1
C-H stretching of
-CH=CH-

0.008

Bridging-OH on TiO2

0.004
0.000
-0.004

90min
illumination

-0.008
-0.012
3800

3600

3400

3200

Wavenumbers (cm-1)

3000

2800

Figure 9 compares the photocatalytic activity of 8MoDT2, 4.7MoDT05 and 2MoDT2 with DT2
and DT05 supports. The DT05 and DT2 supports show similar carbonyl absorptions, with stronger
1671 cmí1 vibration. When molybdenum is added to the catalyst, a decrease in this peak is
observed and the 1691cmí1 peak becomes more prominent. Similar to 2MoDT2, 8MoDT2
photocatalyst shows two bands at 1671 and 1691 cmí1. The latter signal could be ascribed also in
this case to cyclohexenone. The contribution at about 1690 cmí1 is not visible for DT2 and DT05,
indicating that the formation of this compound is due probably correlated to the presence of

290
molybdenum. The band at 1690 cmí1 is also not present on 4.7MoDT05. This result could be
explained considering that gas phase results on 4.7MoDT05 (Figures 4 and 5) showed the lowest
activity and benzene production with together a strong deactivation during UV irradiation.
Therefore, in this case, the formation of carbonylic compounds is predominant. On the other hand, the
formation of oxygenated products (stronger 1671 cmí1 vibration) is observed for all catalysts, and a
decrease in the formation of these products occurs with increasing molybdenum loading confirmed by
the decrease of the band at 1671 cmí1 when molybdenum is added on TiO2 surface.
Figure 9. DRIFT spectra after 90 min of UV irradiation for gas mix 2 on 8MoDT2,
4.7MoDT05, 2MoDT2, DT2 and DT05 catalysts.
1691

0.04
-0.00

1671

8MoDT2

-0.04
4.7MoDT05
-0.08
-0.12 2MoDT2
-0.16
DT2
-0.20
-0.24 DT05
1800

1600
1400
Wavenumbers (cm-1)

1200

The other data to consider are the bands in the region from 1300 cmí1 to about 1450 cmí1,
ascribed to the carbon deposits on the catalyst surface during the photoreaction, which determine
the deactivation of catalyst observed in the literature [37]. These observations suggest that these
species are formed on bare titania, where total oxidation occurs preferentially, while Mo-species
hinder titania surface sites active for total oxidation. The photocatalytic tests performed on all
catalysts revealed that a rapid decays of activity occurred in the first minutes on stream (Figures 4
and 6). This initial decrease of activity is probably due to a poisoning of the surface by carbon
deposits, blocking of a part of the catalytic surface sites.
2.4. Further Test to Assess Intermediates
With the aim to confirm if carbonylic compounds are the responsible of catalyst deactivation, a
photocatalytic test on 8MoDT2 was carried out by feeding cyclohexanone with the same operating
conditions used for cyclohexane.

291
The obtained results are reported in Figure 10.
Figure 10. Cyclohexanone conversion and CO2 concentration during UV irradiation.
70

9

60

8

conversion

7

CO2 production

50

6

40

5
30

4
3

CO 2 , ppm

cyclohexanone conversion, %

10

20

2
10

1

0

0
0

5

10

15

20

25

30

irradiation time, min

Maximum cyclohexanone conversion was about 9%, and then activity decreased to 1% in
15 min. A steady state condition was obtained after about 21 min of irradiation with a conversion
of approximately 0.9%. Carbon dioxide was formed immediately after the UV sources were
switched on and reached a concentration of about 60 ppm after an irradiation time of 0.5 min. CO2
was the only observed product and no other reaction products were detected in gas phase. These
results evidenced that cyclohexanone (carbonyl compounds shown by DRIFTS spectra) is the
precursor for CO2 production. According to DRIFTS spectra, the accumulation on surface of
intermediates formed by cyclohexanone oxidation during irradiation is the responsible of
catalyst deactivation.
With regard to the influence of sulfate, a considerable decrease in the formation of carboxylates
is observed with higher molybdenum and sulfate content confirming that the simultaneous presence
of Mo-species and sulfate sharply increases photocatalytic activity and selectivity. The acidity
induced by the sulfate on TiO2 surface [38] furnishes hydrocarbon activation towards partial
oxidation and supplying the initial step in the absence of irradiation [24]. However, if the sulfate is
lost during dark adsorption, it can induce a deactivation of catalyst under illumination.
For gas mix 3, when H2O is also present, almost no reaction has been observed for the
molybdenum catalysts. The decrease in reactivity due to the hydration of metal oxides has been
discussed in literature [39].
An example is shown in Figure 11, for the 4.7MoDT05 catalysts, in which very small peaks
were observed under UV irradiation. The increase of CH stretching vibration indicates that
cyclohexane adsorption on the catalyst surface is still occurring evidencing a deactivation

292
phenomenon. The contribution of the reaction products to the spectra is very low, only visible by
the formation of small peaks between 1800–1000 cmí1 and the decrease of the TiO2 bridging-OH.
Figure 11. DRIFT spectra during 90 min of UV irradiation for gas mix 3 on
4.7MoDT05 catalyst.
0.060

0.060

0.050

0.050

Absorbance

0.040

0.030

3040-3010 cm-1
C-H stretching of
-CH=CH-

0.040

0.030

TiO2
Bridging-OH
Water

0.020

0.020

0.010

0.010

0.000

0.000

3800

3600

3400
3200
3000
Wavenumbers (cm-1)

2800

1800

1600
1400
1200
Wavenumbers (cm-1)

Figure 12. DRIFT spectra after 90 min of UV irradiation for gas mix 3 on DT2 and
Hombikat catalysts.
0.16

0.08
0.12

0.06
0.04

Absorbance

0.08
TiO2
Bridging-OH

0.04

3040-3010 cm-1
C-H stretching of
-CH=CH-

0.02
0.00

Hombikat

Hombikat

-0.02

0.00

-0.04
-0.06 DT2

-0.04
DT2

-0.08

-0.08

-0.10
3800

3600

3400
3200
Wavenumbers (cm-1)

3000

2000

1800

1600
1400
Wavenumbers (cm-1)

1200

Under gas mix 3 flow better results were observed when there was no molybdenum on the TiO2
surface, as can be seen in Figure 12. While on MoOx/TiO2 catalysts, the water presence in the gas feed
greatly decreased the product formation, in DT2 this effect was not observed. The reactivity of DT2
under gas mix 2 (without H2O) and 3 (with H2O) showed similar reaction products and spectral
intensity. As expected from previous work, no peaks related to unsaturated hydrocarbons formation
were observed on DT2 [21]. Hombikat is also shown as reference, which appears to catalyze the

293
formation of more surface products and higher bridging-OH deactivation than the ones observed for
sulfated TiO2. This last result is a further confirmation of the role of sulfate in the deactivation
of photocatalysts.
3. Experimental Section
3.1. Catalysts Preparation and Characterization
Two titanias were used as supports: two commercial titania samples (DT and DT51 by
Millenium Inorganic Chemicals) with different sulfate content (respectively 0.5 wt% and 2 wt%
expressed as SO3) contributed from the experimental procedure used for the synthesis of samples [40].
The samples are named, respectively, DT2 and DT05 with reference to the sulfate content.
MoOx-based catalysts were prepared by wet impregnation of titania with aqueous solution of
ammonium heptamolybdate (NH4)6Mo7O24 4H2O, followed by drying at 120 °C and calcination at
400 °C for 3 h. Unsulfated titania (Hombikat) was used as reference.
Thermogravimetric analysis (TG-DSC-MS) was carried out in air flow on powder samples with
a thermoanalyzer (Q600, TA) in the range 20–1000 °C with an heating rate of 10 °C/min. Sulfate
content has been evaluated from the weight loss in the range in which the release of SO2 occurred.
Specific surface area was obtained by N2 adsorption-desorption isotherm at í196 °C with a Costech
Sorptometer 1040. Powder samples were treated at 150 °C for 2 h in He flow (99.9990%) before
testing. Laser Raman spectra of powder samples were obtained with a Dispersive MicroRaman
(Invia, Renishaw), equipped with 785 nm diode-laser, in the range 100–2500 cmí1 Raman shift.
Chemical analysis of molybdenum loading was performed by inductive coupled plasma-mass
spectrometry (7500c ICP-MS, Agilent) after microwave digestion (Ethos Plus from Milestone) of
samples in HNO3/HCl and HF/HCl mixtures.
3.2. Gas phase Analysis
Photocatalytic tests were carried out feeding 1000 ppm cyclohexane, 1500 ppm oxygen and
1600 ppm water in N2 (total flow rate: 830 (stp) cm3/min) to a continuous gas-solid annular
photocatalytic fixed-bed reactor. The tests were carried out in presence of water in order to
minimize catalyst deactivation as reported in literature [37].
Oxygen and nitrogen were fed from cylinders, nitrogen being the carrier gas for cyclohexane
and water vaporized from two temperature controlled saturators containing pure cyclohexane and
water. The gas flow rates were measured and controlled by mass flow controllers (Brooks
Instrument). The annular section of the reactor [27] (reactor volume: 7 l) was realized with two
axially mounted 500 mm long quartz tubes of 140 and 40 mm diameter, respectively. The reactor
was equipped with seven 40 W UV fluorescent lamps providing photons wavelengths in the range
from 300 to 425 nm, with primary peak centered at 365 nm. One lamp (UVA Cleo Performance 40
W, Philips) was centered inside the inner tube while the others (R-UVA TLK 40 W/10R flood
lamp, Philips) were located symmetrically around the reactor.

294
Although light sources vary in their intensity, they have the same emission spectrum and in the
case of photocatalytic oxidative dehydrogenation of cyclohexane, the selectivity of the reaction was
not influenced by light intensity [23].
The overall system composed of UV lamps and photoreactor are covered by aluminum foils to
minimize the dispersion of the photons in the space surrounding the photoreactor. In order to avoid
temperature gradients in the reactor caused by irradiation, the temperature was controlled to
35 ± 2 °C by cooling fans. The catalytic reactor bed was prepared in situ, by coating quartz flakes
previously loaded in the annular section of a quartz continuous flow reactor with aqueous slurry of
catalysts powder. The coated flakes were dried at 120 °C for 24 h in order to remove the excess of
physisorbed water. This treatment resulted in uniform coating well adhering to the quartz flakes
surface. The gas composition was continuously determined by on line analyzers connected to a PC
for data acquisition. CO and CO2 concentration is measured by an on line non dispersive IR
analyzer (Uras 10, Hartmann & Braun), working on the basis of specific adsorption of IR radiation
(wavelength from 2 to 8 ȝm). Oxygen, cyclohexane and reaction products composition is
determined by an on line quadrupole mass detector (MD800, ThermoFinnigan) that can analyze the
outlet reactor gas, introduced into a heated silica capillary, up to m/z = 800.
3.3. DRIFTS Analysis
A layer of KBr powder was introduced in the DRIFTS (3 window cell) holder [41], over which each
of the different catalysts and supports were deposited and pressed. The catalysts were heated up to
120 °C for 30 min in He flow (app. 30mL/min) to remove weakly adsorbed water from the surface.
Different gas mixtures (100mL/min flow) were prepared for the catalytic tests:
Gas mix 1: 1000ppm Cyclohexane, 1500ppm O2
Gas mix 2: 1000ppm Cyclohexane, 5000ppm O2
Gas mix 3: 1000 ppm Cyclohexane, 250 ppm H2O, 5000ppm O2
Further, for DRIFTS analysis, nitrogen is the carrier gas for cyclohexane and water vaporized
from two temperature controlled saturators containing pure cyclohexane and water.
A background (128 scan averages) of the dried catalyst at room temperature was taken and used
for the adsorption step, which was followed with IR during 40–60 min, until surface stabilization
was achieved. A spectrum (128 scan averages) was taken after adsorption stabilization, which was
used as background for the reaction step. Reaction was continued for 90 min; the first minute of
collection was performed in the dark, and the rest was done under irradiation by a 150 W Xe light
with a light diffuser. All spectra, except for background spectra, consisted of 64 averaged scans and
water vapor correction has been applied to most of them.
Only the background spectrum was collected at room temperature, while during the irradiation
the temperature increased up to a value very similar to that one of the fixed bed reactor used for gas
phase analysis. In this way, the selectivity of the reaction was not different.

295
4. Conclusions
From coupling of gas phase and DRIFTS analysis, a deeper knowledge of phenomena occurring
during photocatalytic selective oxidation of cyclohexane on MoOx/TiO2 was obtained.
During adsorption measurements in the absence of UV irradiation, there is a clear indication of
organo-sulfur compound formation accompanied by H-abstraction by the acidic sulfate species. The
accompanying stepwise reduction of the sulfate by protons causes formation of SO2 and thereby loss
of surface sulfur. The decrease in surface sulfur during the dark adsorption experiment leads to a
subsequent initial decrease in photocatalytic activity during irradiation since the acidity induced by
the sulfate facilitates hydrocarbon activation in partial oxidation of cyclohexane after supplying the
initial step in the absence of irradiation. The further decrease of photoactivity during irradiation is due
to a poisoning of the surface by carbonaceous species derived from carbonylic compounds formed on
bare titania, blocking of a portion of the catalytic surface sites. A considerable decrease in the rate of
poisoning by carbonaceous species is observed at higher molybdenum and sulfate contents,
confirming that the simultaneous presence of Mo-species and sulfate enhances photocatalytic activity
and selectivity. Gas phase analysis of cyclohexane and reaction products evidenced that the
presence of MoOx species at same sulfate coverage increases the selectivity of the catalyst with
increasing molybdenum content indicating that the interaction between titania and supported
molybdenum oxide plays an essential role in changing the catalyst selectivity. In addition
increasing coverage of surface sulfate influenced benzene yield.
Conflicts of Interest
No author of the present manuscript has a direct financial relation with the commercial identities
mentioned in the paper that might lead to a conflict of interest.
References
1.

2.

3.
4.

5.

Sannino, D.; Vaiano, V.; Sacco, O.; Ciambelli, P. Mathematical modelling of photocatalytic
degradation of methylene blue under visible light irradiation. J. Environ. Chem. Eng. 2013, 1,
56–60.
Sacco, O.; Stoller, M.; Vaiano, V.; Ciambelli, P.; Chianese, A.; Sannino, D. Photocatalytic
Degradation of Organic Dyes under Visible Light on N-Doped TiO2 Photocatalysts. Int. J.
Photoenergy 2012, Article ID 626759:1–Article ID 626759:8.
Murcia, J.J.; Hidalgo, M.C.; Navío, J.A.; Vaiano, V.; Sannino, D.; Ciambelli, P. Cyclohexane
photocatalytic oxidation on Pt/TiO2 catalysts. Catal. Today 2013, 209, 164–169.
Stoller, M.; Movassaghi, K.; Chianese, A. Photocatalytic degradation of orange II in aqueous
solutions by immobilized nanostructured titanium dioxide. Chem. Eng. Trans. 2011, 24,
229–234.
Ibrahim, H.; de Lasa, H. Photo-catalytic conversion of air borne pollutants: Effect of catalyst
type and catalyst loading in a novel photo-CREC-air unit. Appl. Catal. B 2002, 38, 201–213.

296
6.

7.
8.
9.

10.
11.

12.
13.

14.

15.

16.

17.
18.
19.
20.

21.

Augugliaro, V.; Bellardita, M.; Loddo, V.; Palmisano, G.; Palmisano, L.; Yurdakal, S.
Overview on oxidation mechanisms of organic compounds by TiO2 in heterogeneous
photocatalysis. J. Photochem. Photobiol. C 2012, 13, 224–245.
Sannino, D.; Vaiano, V.; Isupova, L.A.; Ciambelli, P. Heterogeneous Photo-Fenton Oxidation
of Organic Pollutants on Structured Catalysts. J. Adv. Oxid. Technol. 2012, 15, 294–300.
Sannino, D.; Vaiano, V.; Ciambelli, P.; Isupova, L.A. Structured catalysts for photo-Fenton
oxidation of acetic acid. Catal. Today 2011, 161, 255–259.
Molinari, A.; Amadelli, R.; Mazzacani, A.; Sartori, G.; Maldotti, A. Tetralkylammonium and
sodium decatungstate heterogenized on silica: Effects of the nature of cations on the
photocatalytic oxidation of organic substrates. Langmuir 2002, 18, 5400–5405.
Maldotti, A.; Molinari, A.; Amadelli, R. Photocatalysis with organized systems for the
oxofunctionalization of hydrocarbons by O2. Chem. Rev. 2002, 102, 3811–3836.
Li, X.; Chen, G.; Po-Lock, Y.; Kutal, C. Photocatalytic oxidation of cyclohexane over TiO2
nanoparticles by molecular oxygen under mild conditions. J. Chem. Technol. Biotechnol.
2003, 78, 1246–1251.
Stoller, M.; Miranda, L.; Chianese, A. Optimal Feed location in a Spinning Disc Reactor for
the Production of TiO2 Nanoparticles. Chem. Eng. Trans. 2009; 17, 993–998.
de Caprariis, B.; Di Rita, M.; Stoller, M.; Verdone, N.; Chianese, A. Reaction-precipitation by
a spinning disc reactor: Influence of hydrodynamics on nanoparticles production. Chem. Eng.
Sci. 2012, 76, 73–78.
Boarini, P.; Carassiti, V.; Maldotti, A.; Amadelli, R. Photocatalytic oxygenation of cyclohexane on
titanium dioxide suspensions: Effect of the solvent and of oxygen. Langmuir 1998, 14,
2080–2085.
Almquist, C.B.; Biswas, P. The photo-oxidation of cyclohexane on titanium dioxide: an
investigation of competitive adsorption and its effects on product formation and selectivity.
Appl. Catal. A 2001, 214, 259–271.
Berg, O.; Hamdy, M.S.; Maschmeyer, T.; Moulijn, J.A.; Bonn, M.; Mul, G. On the
wavelength-dependent performance of Cr-doped silica in selective photo-oxidation. J. Phys.
Chem. C 2008, 112, 5471–5475.
Maldotti, A.; Amadelli, R.; Varani, G.; Tollari, S.; Porta, F. Photocatalytic processes with
polyoxotungstates: Oxidation of cyclohexylamine. Inorg. Chem. 1994, 33, 2968–2973.
Teramura, K.; Tanaka, T.; Yamamoto, T.; Funabiki, T. Photo-oxidation of cyclohexane over
alumina-supported vanadium oxide catalyst. J. Mol. Catal. A 2001, 165, 299–301.
Almeida, A.R.; Moulijn, J.A.; Mul, G. In situ ATR-FTIR study on the selective
photo-oxidation of cyclohexane over anatase TiO2. J. Phys. Chem. C 2008, 112, 1552–1561.
Xu, W.; Raftery, D.; Francisco, J.S. Effect of irradiation sources and oxygen concentration on
the photocatalytic oxidation of 2-propanol and acetone studied by in situ FTIR. J. Phys.
Chem. B 2003, 107, 4537–4544.
Ciambelli, P.; Sannino, D.; Palma, V.; Vaiano, V.; Bickley, R.I. Reaction mechanism
of cyclohexane selective photo-oxidation to benzene on molybdena/titania catalysts. Appl.
Catal. A 2008, 349, 140–147.

297
22. Ciambelli, P.; Sannino, D.; Palma, V.; Vaiano, V.; Mazzei, R.S. Improved Performances of a
Fluidized Bed Photoreactor by a Microscale Illumination System. Int. J. Photoenergy 2009,
Article ID 709365:1– Article ID 709365:7.
23. Palma, V.; Sannino, D.; Vaiano, V.; Ciambelli, P. Fluidized-Bed Reactor for the
Intensification of Gas-Phase Photocatalytic Oxidative Dehydrogenation of Cyclohexane.
Ind. Eng. Chem. Res. 2010, 49, 10279–10286.
24. Ciambelli, P.; Sannino, D.; Palma, V.; Vaiano, V. The effect of sulfate doping on nanosized
TiO2 and MoOx/TiO2 catalysts in cyclohexane photooxidative dehydrogenation. Int. J.
Photoenergy 2008, Article ID 258631:1–Article ID 258631:8.
25. Ciambelli, P.; Sannino, D.; Palma, V.; Vaiano, V. Photocatalysed selective oxidation of
cyclohexane to benzene on MoOx,/TiO2. Catal. Today 2005, 99, 143–149.
26. Ciambelli, P.; Sannino, D.; Palma, V.; Vaiano, V. Cyclohexane photocatalytic oxidative
dehydrogenation to benzene on sulfated titania supported MoOx. Stud. Surf. Sci. Catal. 2005,
155, 179–187.
27. Sannino, D.; Vaiano, V.; Ciambelli, P.; Eloy, P.; Gaigneaux, E.M. Avoiding the deactivation
of sulfated MoOx/TiO2 catalysts in the photocatalytic cyclohexane oxidative dehydrogenation
by a fluidized bed photoreactor. Appl. Catal.A 2011, 394, 71–78.
28. Ciambelli, P.; Sannino, D.; Palma, V.; Vaiano, V.; Eloy, P.; Dury, F.; Gaigneaux, E.M. Tuning
the selectivity of MoOx supported catalysts for cyclohexane photo oxidehydrogenation.
Catal. Today 2007, 128, 251–257.
29. Ciambelli, P.; Sannino, D.; Palma, V.; Vaiano, V.; Mazzei, R.S.; Eloy, P.; Gaigneaux, E.M.
Photocatalytic cyclohexane oxidehydrogenation on sulfated MoOx/gamma-Al2O3 catalysts.
Catal. Today 2009, 141, 367–373.
30. Cheng, C.P.; Schrader, G.L. Characterization of supported molybdate catalysts during
preparation using laser Raman spectroscopy. J.Catal. 1979, 60, 276–294.
31. Bellamy, L.J.; Mayo, D.W. Infrared frequency effects of lone pair interactions with
antibonding orbitals on adjacent atoms. J. Phys. Chem. 1976, 80, 1217–1220.
32. Bezrodna, T.; Puchkovska, G.; Shimanovska, V.; Chashechnikova, I.; Khalyavka, T.; Baran, J.
Pyridine-TiO2 surface interaction as a probe for surface active centers analysis. Appl. Surface Sci.
2003, 214, 222–231.
33. Mul, G.; Wasylenko, W.; Hamdy, M.S.; Frei, H. Cyclohexene photo-oxidation over vanadia
catalyst analyzed by time resolved ATR-FT-IR spectroscopy. Phys. Chem. Chem. Phys. 2008,
10, 3131–3137.
34. Coates, J. Vibrational spectroscopy: Instrumentation for infrared and Raman spectroscopy.
Appl. Spectrosc. Rev. 1998, 33, 267–425.
35. Schuchardt, U.; Pereira, R.; Rufo, M. Iron(III) and copper(II) catalysed cyclohexane oxidation by
molecular oxygen in the presence of tert-butyl hydroperoxide. J. Mol. Catal. A 1998, 135,
257–262.
36. Manner, W.L.; Girolami, G.S.; Nuzzo, R.G. Sequential dehydrogenation of unsaturated cyclic
C5 and C5 hydrocarbons on Pt(111). J. Phys. Chem. B 1998, 102, 10295–10306.

298
37. Einaga, H.; Futamura, S.; Ibusuki, T. Heterogeneous photocatalytic oxidation of benzene,
toluene, cyclohexene and cyclohexane in humidified air: Comparison of decomposition
behavior on photoirradiated TiO2 catalyst. Appl. Catal. B 2002, 38, 215–225.
38. Ciambelli, P.; Fortuna, M.E.; Sannino, D.; Baldacci, A. The influence of sulfate on the
catalytic properties of V2O5-TiO2 and WO3-TiO2 in the reduction of nitric oxide with
ammonia. Catal. Today 1996, 29, 161–164.
39. Amano, F.; Yamaguchi, T.; Tanaka, T. Photocatalytic oxidation of propylene with molecular
oxygen over highly dispersed titanium, vanadium, and chromium oxides on silica. J. Phys. Chem.
B 2006, 110, 281–288.
40. Taranto, J.; Frochot, D.; Pichat, P. Photocatalytic treatment of air: Comparison of various
tiO2, coating methods, and supports using methanol or n-Octane as tst pollutant. Ind. Eng.
Chem. Res. 2009, 48, 6229–6236.
41. Carneiro, J.T.; Moulijn, J.A.; Mul, G. Photocatalytic oxidation of cyclohexane by titanium
dioxide: Catalyst deactivation and regeneration. J. Catal. 2010, 273, 199–210.

MDPI AG
Klybeckstrasse 64
4057 Basel, Switzerland
Tel. +41 61 683 77 34
Fax +41 61 302 89 18
http://www.mdpi.com/
Catalysts Editorial Office
E-mail: catalysts@mdpi.com
http://www.mdpi.com/journal/catalysts

MDPI • Basel • Beijing • Wuhan • Barcelona
ISBN 978-3-03842-188-7
www.mdpi.com