Water Gas Shift Catalysis

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Water Gas Shift CatalysisChandra Ratnasamy a & Jon P. Wagner aa Sud Chemie, Louisville, KY, USAVersion of record first published: 05 Aug 2009.

To cite this article: Chandra Ratnasamy & Jon P. Wagner (2009): Water Gas Shift Catalysis, CatalysisReviews: Science and Engineering, 51:3, 325-440

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Water Gas Shift Catalysis

Chandra Ratnasamy and Jon P. Wagner

Sud Chemie, Louisville, KY, USA

Developments in water gas shift (WGS) catalysis, especially during the last decade, arereviewed. Recent developments include the development of (1) chromium-free catalyststhat can operate at lower steam to gas ratios and (2) more active catalysts that canoperate at gas hourly space velocities above 40,000 h21. A current challenge is todevelop catalysts for use in fuel cell applications. Precious metal catalysts supported onpartially reducible oxide supports (Pt-ceria, Pt-titania, Au-ceria, etc.) are the currentfront runners. A critical review of the mechanism of the WGS reaction is alsopresented.

Keywords Water gas shift, Hydrogen production, CO conversions, Fuel processor,Fuel cell, Iron oxide catalysts, Copper-Zinc oxide catalysts, Pt catalysts,Redox mechanism, Formate mechanism, High temperature shift, Lowtemperature shift, Sour gas shift, Chromium-free catalysts

1. INTRODUCTION

‘‘Water gas’’ is a mixture of hydrogen and carbon monoxide. It is used

extensively in the industry for the manufacture of ammonia, methanol,

hydrogen (for hydrotreating, hydrocracking of petroleum fractions and other

hydrogenations in the petroleum refining and petrochemical industry),

hydrocarbons (by the Fischer-Tropsch process) and metals (by the reduction

of the oxide ore). It is manufactured by the reaction of a carbonaceous material

(coal, coke, natural gas, naphtha, etc.) with steam [Eqs. (1, 2)], oxygen [Eq. (3)]

or carbon dioxide [Eq. (4)]:

CzH2O<COzH2 H2=CO~1; DH~131:2 kJ=molð Þ, ð1Þ

CH4zH2O<COz3 H2 H2=CO~3; DH~206:3 kJ=molð Þ, ð2Þ

Received 24 August 2008; accepted 17 February 2009.Address correspondence to Chandra Ratnasamy, Sud Chemie, Louisville, KY 40210,USA. E-mail: [email protected]

Catalysis Reviews, 51:325–440, 2009

Copyright # Taylor & Francis Group, LLC

ISSN: 0161-4940 print 1520-5703 online

DOI: 10.1080/01614940903048661

325

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CH4z0:5 O2<COz2H2 H2=CO~2; DH~{35:6 kJ=molð Þ, and ð3Þ

CH4zCO2<2COz2H2 H2=CO~1; DH~247:4 kJ=molð Þ: ð4Þ

Reactions 1, 2, and 4 are endothermic while reaction 3 is exothermic. It may be

noted that the molar ratio of H2 to CO varies depending on the source of

carbon/oxygen.

Steam reforming [Eq. (2)] is the most popular mode of generating water

gas, especially if the ultimate objective is the generation of pure hydrogen

since it provides the highest molar ratio of H2/CO of 3. The exothermic partial

oxidation [Eq. (3)] providing a H2/CO molar ratio of 2 is used in the

manufacture of water gas when (a) a lower H2/CO ratio (( 2, for example)

is needed (e.g., dimethyl ether or Fischer Tropsch synthesis), or (b) due to

difficulties in external heat supply, internal heat generation (autothermal

reforming) is needed as in the case of fuel processors for fuel cell applications.

‘‘Dry reforming’’ or ‘‘CO2 reforming’’ (Eq. 4) is an additional source of water gas

with a very low H2/CO molar ratio of one. This process is used in the

manufacture of water gas from natural gas for the reduction of iron ore

wherein CO has been found to be as good a reductant (if not better) as H2.

The water gas shift reaction [Eq. (5)] was first reported in 1888 (1), but it

came into popular usage later, as a source of hydrogen for the Haber process

for the manufacture of ammonia:

CO gð ÞzH2O gð Þ<CO2 gð ÞzH2 gð Þ DH~{41:1 kJ=molð Þ: ð5Þ

In the initial stages of the Haber ammonia process, the hydrogen needed for

the process was obtained from the water gas generated by Eq. (1). In coal or

coke gasification [Eq. (1)], when steam is contacted with incandescent coke (at

about 1000uC), CO2 is an additional product [Eq. (6)] especially at lower

temperatures:

Cz2H2O<CO2z2H2 DH~90 kJ=molð Þ: ð6Þ

While CO2 could be easily removed from the products of the reaction (by

absorption in water), CO had to be removed by liquefaction or copper liquor

scrubbing. A catalytic process to remove the CO from (CO + H2) mixtures was

needed. In 1914, Bosch and Wild (2) discovered that the oxides of iron and

chromium could convert a mixture of steam and CO into CO2 at 400–500uC,

according to Eq. 5, and, in the process, generate additional hydrogen for the

Haber process. Thenceforth, the water gas produced from the carbonaceous

source by steam reforming was passed over the iron-chromium catalyst to shift

326 Ratnasamy and Wagner

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the CO to CO2 by the water gas shift reaction. Iron-based catalysts are still

used today industrially. There are four general types of water gas shift

catalysts. One of them is the promoted iron oxide catalyst. Catalysts of this

type promote the shift reaction at moderately high temperatures (350–450uC)

and are therefore called high temperature shift (HTS) catalysts. The second

type is copper-zinc oxide catalyst and is called the low temperature shift (LTS)

catalyst because it is used at relatively low temperatures (190–250uC). The

third type employs cobalt and molybdenum sulfides as the active ingredients.

Catalysts of this type are sulfur-tolerant and can be used in sulfur-containing

‘‘sour gas’’ streams and are therefore called sour gas shift catalysts. There was

interest (in the past) in a fourth type of catalyst, medium temperature shift or

MTS catalyst that operates at temperatures between the HTS and LTS

catalysts. Normally, these are copper-zinc catalysts that are actually LTS

catalysts modified (usually with iron oxide) to operate at slightly higher

temperatures (275–350uC) than a standard LTS catalyst. In addition to the

above four, precious metal- based catalysts (mainly platinum and gold) have

been under intensive investigation during the last decade for use in fuel cell

applications. Promoters, like Cu and Al2O3 are added to the conventional iron

oxide - chromium oxide HTS catalyst compositions in some modern versions.

At lower temperatures, the iron-based catalysts are less active.

Equilibrium concentrations of CO are lower only at low temperatures (section

II), however. Hence, to achieve higher conversions of CO at lower

temperatures (190–250uC), a second, more active catalyst, based on Cu-ZnO,

was developed in the early 1960s and is used in the industry extensively. It

may be noted that Cu- based catalysts had been patented as early as 1931 (3).

Today, the industrial WGS process takes place in a series of adiabatic

converters where the effluent from the reformer system is converted in two

WGS reactors (HTS and LTS converters, respectively), with the second WGS

reactor at a significantly lower temperature in order to shift the equilibrium

towards the favored hydrogen product (Fig. 1). The modern, two stage WGS

converter systems reduce the CO concentrations to about 0.3%(wt) from the

high levels (10–50%) in the outlet from the reformers. During the last couple of

decades, fuel cells, generating electricity from the reaction of hydrogen with

oxygen, for stationary and mobile applications, have become popular. A crucial

prerequisite for the techno-economic success of fuel cells, especially those that

operate either at low temperatures (like the Polymer Electrolyte Membrane

Fuel Cells, PEMFC) or in mobile applications (as in automobiles), is the

discovery of improved reforming and WGS catalysts for the generation of

hydrogen which are much more active than those used in chemical plants. The

requirements of WGS catalysts for fuel cell applications are quite different

from those of the traditional Fe2O3-Cr2O3 or Cu-ZnO based catalysts. The

catalyst bed must have a reduced volume and weight to be economical and

Water Gas Shift Catalysis 327

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have sufficient durability to withstand rapid start-up and shut-down

conditions. In addition, the catalyst must not require controlled and elaborate,

pre-reduction procedures (as is the case with the Cu- based LTS catalysts),

must be non-pyrophoric and oxidation- tolerant on exposure to air. The two

HTS and LTS catalysts are extremely pyrophoric when activated (reduced)

and, therefore, safety from runaway heat generation and fires cannot be

ensured upon air exposure. In response to these needs, noble metal- based

reforming and WGS catalysts are under intense development worldwide for

fuel cell applications.

Newsome had provided an excellent review of the WGS literature up to

1980 (4). Lloyd et al. have reviewed the industrial developments in this area

up to 1996 (5). A comprehensive review by Kochlofl in 1997 (6, 7) covers both

the fundamental and applied aspects of the field. More recently, in 2003,

Ladebeck and Wagner have given a brief survey of the WGS catalyst

developments, especially for fuel cell applications (8, 9).

This review, after highlighting the significant features of the conventional

HTS and LTS catalysts and processes based on the Fe2O3-Cr2O3 and Cu-ZnO-

Al2O3 catalysts, respectively, critically examines the extensive results (both in

the journal and patent literature) from the study of noble metal based WGS

catalysts in the last decade or so, with particular emphasis on catalyst surface

Figure 1: Syngas generation and water gas shift reactors for NH3 Synthesis (8–9).


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structures, active sites, reaction intermediates and mechanisms. This subject

has developed significantly, mainly, in the last decade due to its relevance to

the fuel cell industry.

2. THERMODYNAMICS

The feed composition to the HTS reactor can vary depending on the end-

application of the outlet of the WGS stage. Table 1 gives the composition at the

inlet to the WGS stage for some typical applications. The high N2 content in

the feed, for eventual application in ammonia synthesis, comes from the

addition of air in the secondary reformer to provide the N2 reactant for NH3.

The water gas shift reaction is moderately exothermic (Eq. 5) and conversions

are equilibrium- controlled. The equilibrium constant decreases with increas-

ing temperature (Fig. 2) and, in the temperature range 315–480uC, is given (8)

by Eq. 7:

Kp~exp 4577:8=Tð Þ{4:33½ �, ð7Þ

where T is in K. Accordingly, high conversions are favored at low

temperatures and are not affected, significantly, by changes in total pressure.

The reaction is reversible and the forward rate is strongly inhibited by the

reaction products, H2 and CO2. When operated under adiabatic conditions

(typical in industry), the exothermic rise in the catalyst bed temperature can

Table 1: Feed compositions and process conditions at the inlet to the WGS stagefor some typical applications.

Application Code 1 A B C

Feed Composition (mole %, dry)CO 12.8 10.3 46CO2 7.8 11.4 6.9H2 56.4 74.5 47N2 22.4 0.1 —CH4 0.3 3.7 0.1Ar 0.3 — —

Inlet steam/gas, molar ratio 0.6 0.9–1.0 1.0–2.2Pressure, bar 25–30 20–30 12–30Inlet Temperature, uC 343–399 343–399 343–399Outlet Temperature, uC 399–466 399–454 371–454Space Velocity, h21 2500 1500–2000 500–1400Number of beds 1 1 3Outlet CO, mole %, dry 2.0–3.5 2.0–3.0 1.5–3.5

(1) Application Code. A: NH3 Plant based on steam-hydrocarbon reforming; B: H2 Plant basedon steam-hydrocarbon reforming; and C: H2 plant based on partial oxidation of oil feed.


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inhibit, due to thermodynamic reasons, CO conversions. This limitation can,

however, be mitigated by using two or more beds with heat removal between

them. CO levels at the exit of the HTS reactor are around 3–5 wt% while

values around 0.3 wt% can be achieved at an exit temperature of 200uC in the

LTS reactor. The lower limit of the operating temperature in the LTS reactor

is the dew point of water at the operating pressure (190–200uC at 30 bar).

Condensed steam affects, adversely, the catalytic activity of the Cu-based,

LTS catalysts. Even in the case of HTS, exposure to liquid water originating

from condensation should be avoided, since leaching of the water-soluble Cr 6+

ion is equally unwelcome. Similarly, the lower limit on the pressure is the

operating pressure of the downstream units (10–60 bars). The water content

has a strong influence on CO conversion. The water entering the WGS reactor

can be varied by controlling the amount added upstream at the reforming

stage or by injecting water before or between the stages of the WGS reaction.

In contrast, the CO, CO2, and H2 concentrations at the inlet to the HTS

reactor are more dependent on the reformer operation, which, in turn,

Figure 2: Variation of equilibrium constant (Kp) for the water-gas shift reaction withtemperature (5).


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determines the thermodynamic equilibrium conditions. The effect of water

concentrations at various temperatures on the equilibrium CO concentration

is shown in Figs. 3 and 4 for typical HTS and LTS operations, respectively (8,

9). The gas composition used for these calculations are shown in Table 2 and

represents a syngas generated from autothermal reforming (ATR) and

excludes any residual hydrocarbons which may also be present. By increasing

the molar steam to dry gas (CO+CO2) ratio from 0.25 (20% H2O) to 0.75 (42.9%

H2O), the equilibrium temperature (for 1%CO) increases by 100uC. By

operating at 100uC higher temperature, a significant reduction of the reactor

size can be achieved by utilizing the more favorable kinetics at the higher

temperature. The CO concentration at the inlet to the HTS reactor can vary

widely in the range 12–40% (dry basis) depending on the raw material

(natural gas or coal) and the reforming process (steam or autothermal

reforming) utilized to generate the CO. The HTS exit (and LTS inlet)

concentrations are in the range of 3–5% (dry basis) and depend on the

operating temperature of the HTS catalyst bed. Too low of values of the steam/

dry gas ratio can lead to catalyst deactivation (due to coke laydown) while

values much higher than that stoichiometrically needed by Eq. (5) increase the

energy costs and adversely affect the process economy.

The method of producing the syngas will also affect the WGS equilibrium

compositions (8, 9). Autothermal reforming produces a syngas with lower H2

Figure 3: Equilibrium CO concentrations in HTS gas from an autothermal reformer at varioussteam/gas ratios (8–9).


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concentration (due to the dilution with nitrogen) compared to steam

reforming. The lower H2 concentration increases the equilibrium CO

conversion whereas the high H2 concentrations expected with steam

reforming lower the equilibrium CO conversions. Figures 5 and 6 (8, 9) show

the equilibrium CO composition as a function of H2 content at constant CO

and CO2 concentrations for HTS and LTS gases, respectively. To achieve 1%

CO at the reactor outlet, the temperature must be decreased by nearly 40uCwhen the H2 is increased from 35 to 74%. In other words, the outlet CO

concentration would be 1.66% for steam reforming at the same temperature

required to achieve 1% CO for a feedstock from autothermal reforming. The

effect of H2 concentration is not as significant as the steam/dry gas ratio, but,

it is not trivial and must be considered when trying to maximize efficiency and

Figure 4: Equilibrium CO concentrations in LTS gas from an autothermal reformer as afunction of steam/gas ratios (8–9).

Table 2: Representative, methane-free, inlet gas compositions from autothermalreforming of methane (vol%) (7).

HTS (%) LTS (%)

CO 9 3CO2 7 13H2 24 30N2 28 28H2O 32 26


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Figure 5: Equilibrium CO in HTS as a function of H2 concentration (8–9).

Figure 6: Equilibrium CO in LTS as a function of H2 concentration (8–9).


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minimize the volume of the WGS reactor, especially in fuel processors for fuel

cell applications.

3. HIGH TEMPERATURE SHIFT CATALYSTS

3.1. Iron Oxide– Chromium Oxide CatalystsThe high temperature shift reaction using Fe2O3-Cr2O3 catalysts has been

in commercial use for more than 60 years. Many excellent reviews are

available (4–11). The important structural and textural roles of Cr2O3 in the

catalyst formulation has also been investigated in detail (11). Two stage CO

conversion systems employing WGS using Fe2O3-Cr2O3 catalysts and

methanation using nickel-based catalysts for CO removal was the common

and economical design in ammonia synthesis up to the late 1950s. Most of

those plants employed the Fe2O3-Cr2O3 HTS catalyst in the first, high

temperature reactor as well as in the second stage converter at temperatures

as low as 320uC. The conventional Fe2O3-Cr2O3 catalysts worked extremely

well for these high temperature applications but their relatively poor

performance in the lower temperature, second bed of these reactors motivated

further investigations. The early development of unsupported metallic copper

catalysts or copper supported on Al2O3, SiO2, MgO, pumice or Cr2O3 were

characterized by relatively short life and low space velocity operations (400–

1000 h21). Important progress was made by the addition of ZnO or ZnO-Al2O3.

These Cu-ZnO-Al2O3 catalysts exhibited not only a considerable increase in

lifetime, but also an increase in the turnover number by an order of

magnitude. Today, in industrial adiabatic converters, the syngas effluent

from the reformer system is converted in two steps, with the second step at a

significantly lower temperature in order to shift the equilibrium towards the

favored hydrogen product. In the first step, catalysts based on the Fe2O3 –

Cr2O3 oxides are applied at a reactor inlet temperature of 300–360uC and a

total pressure between 10 and 60 bars. Under normal operating conditions,

the temperature rises, progressively, through the reactor bed and can increase

up to 500uC. At exit gas temperatures of 400 to 500uC, the CO content can be

reduced in an industrial HTS converter to 5 vol% or lower. In this section, we

review the main features of the HTS catalyst/process and highlight the

developments during the last decade.

/Conventional Fe2O3-Cr2O3 catalysts contain about 80–90%(wt) of Fe2O3,

8–10% Cr2O3, the balance being promoters and stabilizers like copper oxide,

Al2O3, alkali, MgO, ZnO, etc. The BET surface areas of these catalysts vary

between 30–100 m2/g depending on the Cr2O3 and Al2O3 contents and

calcination temperatures. One of the major functions of Cr2O3 and Al2O3 is

to prevent the sintering, and, consequent loss of surface area of the iron oxide


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crystallites during the start – up and further operation (11). Pure Fe2O3, when

used as a HTS catalyst, deactivates fast due to sintering of the iron oxide

crystallites. In addition to being a textural promoter preventing the sintering

of iron oxide crystallites, Cr2O3 also functions as a structural promoter to

enhance the intrinsic catalytic activity of Fe2O3. As supplied, the Fe2O3-Cr2O3

catalyst is a solid solution of a - Fe2O3 and Cr2O3, wherein the Cr3+ ion

substitutes, isomorphously and partially, the Fe3+ ions in the a - Fe2O3 lattice

framework. Even though most of the chromium ions in the fresh catalyst are

present in the Cr3+ state, a small fraction, especially on the surface, is present

in the hexavalent state, as CrO3. During start-up in the industrial reactor,

Fe2O3 is reduced to Fe3O4 in syngas at 300–450uC (5) (Eqs. 8–9):

3Fe2O3zH2<2Fe3O4zH2O DH~{16:3 kJ=molð Þ, and ð8Þ

3Fe2O3zCO<2Fe3O4zCO2 DH~z24:8 kJ=molð Þ: ð9Þ

The reduction has to be done carefully and the reaction heat removed, to avoid

further reduction of Fe3O4 (Eqs. 10–14).

Fe3O4zH2<3FeOzH2O DH~{63:8 kJ=molð Þ, ð10Þ

Fe3O4zCO<3FeOzCO2 DH~{22:6 kJ=molð Þ, ð11Þ

FeOzH2<FezH2O DH~{24:5 kJ=molð Þ, ð12Þ

FeOzCO<FezCO2 DH~{12:6 kJ=molð Þ, and ð13Þ

Fe3O4z4H2<3Fez4H2O DH~{149:4 kJ=molð Þ ð14Þ

Importantly, neither pure hydrogen nor H2-N2 mixtures should be used to

reduce the HTS catalysts to avoid the occurrence of the strongly exothermic

reduction to metallic Fe (Eq. 14). It is Fe3O4 that is the active phase

responsible for the WGS reaction. The CrO3 phase present must also be

reduced to Cr2O3 during start-up (Eqs. 15 and 16):

2 CrO3z3H2<Cr2O3z3 H2O DH~{684:7 kJ=molð Þ, and ð15Þ

2CrO3z3 CO<Cr2O3z3 CO2 DH~{808:2 kJ=molð Þ ð16Þ


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The ratios H2O/H2 and CO2 /CO determine the relative stabilities of the Fe2O3

and Fe3O4 as well as those of Cr2O3/CrO3 phases. Under normal operating

conditions of HTS (H2O/H2 . 0.4 and CO2/CO . 1.2), the Fe3O4 and Cr2O3

phases are more stable; neither FeO nor metallic Fe are formed under these

conditions. Formation of metallic Fe (due to low H2O/H2 ratios) can trigger the

highly exothermic methanation and Boudouard reactions (Eqs. 17 and 18,

respectively) and lead to runaway conditions and catalyst deactivation:

COz3 H2<CH4zH2O DH~{206:2 kJ=molð Þ, and ð17Þ

2CO<CzCO2 DH~{172:5 kJ=molð Þ ð18Þ

While maintaining sufficiently high H2O/H2 ratios is important, passing

steam, in the absence of reductants like H2 and CO, over the reduced iron-

oxide – chromium oxide catalyst, can reoxidize the Fe3O4 to Fe2O3 (Eq. 19) and

thereby lower catalytic activity:

2Fe3O4zH2O<3Fe2O3zH2 ð19Þ

The Fe2O3-Cr2O3 catalysts are rugged and have a lifetime of 3–5 years

depending, mainly, on the temperature of operation. Unlike the Cu-ZnO (LTS)

catalyst, the Fe2O3-Cr2O3 catalyst is not extremely sensitive to the presence of

sulfur and can tolerate the presence of substantial amounts of sulfur due to

the facile reversibility of the sulfidation reaction (Eq. 20):

Fe3O4z3H2SzH2<3FeSz4H2O DH~{75:0 kJ=molð Þ ð20Þ

The value of the equilibrium constant (Kp 5 PH2P3H2S/P4

H2O) varies from

3 6 10210 to 45 6 10210 in the range of 300–450uC. The rate of the HTS

reaction is limited by pore diffusion and linearly dependent on the steam

partial pressures under industrial conditions (12). A power law – type rate

Eq.satisfactorily fits the experimental data (13). Apart from an increase in the

pressure-drop across the catalyst bed during use due to inadequate

mechanical crushing strength of the catalyst pellets, catalyst deactivation is

mainly due to loss of iron oxide surface area by thermal sintering. In addition

to the abovementioned risks of thermal sintering during start- up reduction of

the catalyst, the exothermic nature of the WGS reaction (Eq. 5) also generates

a large amount of heat during the operational phase of the process. For

example, it has been estimated (5) that the conversion of 1 vol% of CO results

in a temperature rise of about 7–10uC in the catalyst bed. The syngas

feedstocks to the HTS reactor may contain anywhere from 8 vol% (steam

reformer) to 45% (partial oxidation/ autothermal reforming of methane or coal/

coke) of CO. Hence, temperatures in the catalyst bed may rise by 500uC (to


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about 800–850uC) if heat removal is inadequate. Distributing the catalyst in

two or three beds and providing inter-bed coolers can restrict the exit

temperatures to about 450uC and outlet CO content to 3–5 wt%. The Fe2O3-

Cr2O3 catalysts can tolerate sulfur up to, about, even 1000 ppm. Their major

drawbacks are: (a) the toxicity of the water-soluble Cr6+ ions posing health

hazards during catalyst manufacture and handling, and (b) the low volumetric

catalytic activity (GHSV510,000 – 15,000h21), especially at low temperatures,

when CO conversion is favored thermodynamically, necessitating the use of

large catalyst bed volumes. The latter handicap is of crucial importance in fuel

cell applications.

3.1.1. Influence of Catalyst Composition and Preparation Methods

The preparation method of the Fe2O3-Cr2O3 catalyst has a strong

influence on their properties (14, 15). They are usually prepared by

coprecipitation of the hydroxides followed by drying and calcining them to

the corresponding oxides. The oxides are reduced in situ before use. The

precipitation method involves the conventional coprecipitation of the mixed

iron and chromium nitrates with ammonium hydroxides. In an alternate

impregnation method, iron hydroxide gel is first prepared and then

impregnated with chromium nitrate solution. Chromium retards the sintering

of the iron oxide crystallites both during activation and the WGS reaction (14,

15). X-ray photoelectron spectroscopy revealed that there was surface

enrichment of Cr ions in fresh samples prepared by both the coprecipitation

and impregnation routes. The surface concentration of chromium was higher

in the impregnated samples. However, after activation and running the WGS

reaction, it was observed that the relative surface concentration of chromium

had decreased significantly in both the samples, suggesting that, during the

activation and WGS reaction, some Cr ions had migrated from the surface into

the bulk. This is in agreement with the earlier findings of Edwards et al. (16)

who had shown that Cr 3+ (d3) goes into the magnetite (Fe3O4) spinel lattice

and occupies, preferentially, the octahedral sites because of its high crystal

field stabilization, in contrast to the Fe3+ ion (d5) which does not have any

preferred site. When the Cr ions occupy tetrahedral sites, they cause strain in

the magnetite lattice, thereby decreasing the average particle size of the iron

oxide. The lower particle size, in turn, increases the surface area of the iron

oxide and leads to enhanced catalytic activity. The presence of chromium also

caused (16) an increase in the surface Fe2+/Fe3+ ratio (from XPS data) in the

fresh samples and this effect was more pronounced for the impregnated

sample. After reaction, however, these ratios were smaller for both the

precipitated and impregnated samples than the fresh sample. The Mossbauer

spectra of the pure magnetite sample indicated that the Fe3+ ions are in


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tetrahedral sites (A sites) and Fe2+ ions are in octahedral sites (B sites) of Fe3O4.

Chromium addition to magnetite (both by the precipitation and impregnation

methods) produced a decrease in the hyperfine magnetic fields in both sites due

to the partial replacement of iron ions by chromium. This decrease was more

pronounced in B sites, indicating that Cr 3+ ions entered preferentially the

octahedral B sites. The sample prepared by the impregnation method was more

active in the WGS reaction. The role of chromium is twofold: (a) as a textural

spacer/stabilizer for iron oxide crystallites (stabilization of their smaller

crystallite size) and (b) as a structural promoter in increasing the intrinsic

catalytic activity of iron oxide crystallites due to (a) the increase in lattice strain

caused by the substitution of Fe3+ by Cr3+ ions in the magnetite lattice and (b) an

increase in the surface area of the iron oxide crystallites as mentioned above.

During the last two decades, due to the rising cost of hydrocarbon

feedstocks, plants have been forced to keep operating costs low by being as

energy efficient as possible. One method used in improving energy efficiency is

by reducing the overall steam to gas ratio of the plant (usually starting at the

inlet to the reformer). Depending on the level, these lower steam-to-gas ratios

can cause over-reduction of the Fe2O3 in HTS catalyst and result in the

formation of iron carbides. Iron carbides are very effective catalysts for the

formation of hydrocarbons by the Fischer-Tropsch reactions. Products from

the Fischer-Tropsch reactions would negatively impact both LTS catalyst

performance and plant efficiency. The over-reduction can, also, result in a

volume shrinkage within the catalyst pellet that weakens it and may lead to

an increased rate of mechanical breakdown. To minimize the low steam to gas

ratio effects on HTS catalysts, the Sud-Chemie Inc. group developed and

introduced, in the late eighties, a copper promoted iron oxide- chromium oxide

formulation that successfully suppressed the Fischer- Tropsch reactions in

commercial operation. Figures 7 and 8 compare the by-products formation

Figure 7: Comparison of CH4 formation over standard iron oxide - chromium oxide andcopper promoted, iron oxide-chromium oxide catalysts.


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(methane and C2+ hydrocarbons, respectively) across a conventional iron

oxide-chromium oxide catalyst and a copper - promoted, iron oxide- chromium

oxide catalyst. It may be seen that there is a significantly lower by-product

formation in the latter. It is speculated that the presence of copper suppresses

C-O cleavage (in CO), prevents the formation of iron carbides and thereby

avoids the hydrogenation of the adsorbed carbon (to hydrocarbons) and

facilitates its desorption as CO or CO2.

3.2. Influence of Process Variables on Reaction RatesKey process variables affecting the performance of the HTS converter

involve the temperature and inlet steam/dry gas ratio since these influence

both the equilibrium CO content and reaction kinetics. Other factors to be

considered are pressure and catalyst activity.

Temperature: Since the reaction is exothermic, higher CO conversions can

be obtained by reducing the temperature at which the gas leaves the reactor.

However, this principle applies only to a catalyst that is equilibrium- and not

kinetically- limited. A reactor that is operating with the exit gas CO

concentration above equilibrium (kinetically limited) may benefit from higher

gas/ bed temperatures. The exit gas temperature determines both the catalyst

reaction rate in the bottom of the bed and the CO equilibrium value of the

outlet gas. For a reactor loaded with a highly active catalyst, the exit

temperature is determined primarily by the inlet temperature, CO concentra-

tion, and steam/ gas ratio. Exit CO equilibrium is usually achieved in these

cases. The higher the inlet CO concentration and lower the inlet steam/gas

ratio, the larger will be the overall temperature rise through the bed. The

temperature rise is also somewhat dependent on the other gas components

and their composition because of heat capacity effects. Temperature rises of

Figure 8: C2+ hydrocarbon production over iron oxide - chromium oxide and copper -

promoted, iron oxide - chromium oxide catalysts.


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30–75uC are common across commercial reactors. For some of the recent

copper-promoted iron oxide-chromium oxide catalysts, the maximum operat-

ing temperature is around 510uC.

Steam to gas ratio: Both laboratory and commercial data indicate that

higher steam/dry gas ratios in commercial ranges also increase the water gas

shift reaction rate. As a result of the steam/ dry gas ratio effect on both the

thermodynamic and kinetic properties of the process, higher values give

higher CO conversions and a lower exit CO content in the gas. In most plant

configurations, the inlet steam/gas ratio cannot be independently controlled in

the HTS reactor. Other considerations, such as downstream gas purity

requirements and the overall site energy balance determine the inlet reformer

steam/ gas ratio and, as a result, fix the value at the inlet of the HTS

converter. Commercial operating conditions are such that the equilibrium CO

concentrations at the exit of the HTS reactor are usually about 2.0 to 5.0%.

Figures 9–12 show the relative impacts of both inlet steam/ gas ratio and exit

temperature on the equilibrium CO concentration at the exit of the HTS

reactor. They also illustrate the differences in achievable CO levels as a

function of the end-product (ammonia or hydrogen; Figs, 9 and 10,

respectively) as well as the type of reformer feedstock (partial oxidation of

natural gas or fuel oil; Figures 11 and 12, respectively) used to generate the

HTS feed gas. In addition to CO conversions, the steam to gas ratio can also

affect the production of hydrocarbons (mainly methane) by the Fischer-

Tropsch reaction. To minimize such undesirable reactions, a minimum steam

to gas ratio of 0.4 and a maximum CO/CO2 ratio of 1.6 is ensured in the HTS

reactor.

Figure 9: Influence of HTS reactor inlet steam/gas ratio and exit temperature on equilibriumCO concentrations.


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Pressure: The equilibrium CO concentration is virtually unaffected by

system pressure. The pressure will, however, have an impact on the system

kinetics due to pore diffusion limitations and partial pressure effects of the

reactants. Higher pressures will improve overall CO conversion in kinetically-

limited applications.

Figure 10: CO equilibrium vs Inlet S/Gas Ratio and Exit Temperature – Hydrogen Plant.

Figure 11: CO equilibrium Vs Inlet Ratio and Exit Temperature – Partial Oxidation of naturalgas feed.


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3.3. Catalyst DeactivationThe primary deactivation mechanism of the HTS catalyst is due to

thermal sintering of the iron crystallites. The degree of thermal sintering is a

function of time and operating temperature and is irreversible. Thermal

sintering occurs more rapidly at higher temperatures. Hence, the maximum

bed/exit gas temperatures are usually limited to less than 510uC. Some

thermal sintering is unavoidable in the start-up and normal operation of the

catalyst. Commercial data suggest that there is approximately a 50% loss in

total surface area over the first few months of operation and, then, a further

25% loss throughout the remaining life of the catalyst. As a result of these

changes in surface area, the deactivation rate for the catalyst is faster during

the first few months of operation and then stabilizes with very gradual aging

after the first year. Unlike LTS catalysts which can show distinct zones of

deactivation (completely inactive, partially deactivated and essentially fresh,

non-deactivated), the HTS catalyst undergoes a more gradual deactivation

that is spread throughout the bed. The more typical symptom of activity loss is

a gradual spreading out of the reactor temperature profile and an increase in

the CO leakage. Figure 13 shows a typical change in the temperature profile

with time- on- stream in commercial reactors. For a relatively new catalyst,

the temperature increases more sharply in the top 50–60% of the bed. After

some time-on-stream, the catalyst in the top is less active and the temperature

profile changes. Throughout the catalyst aging period, however, there

continues to be some catalyst activity through all of the beds. As the profile

spreads more throughout the bed, it may become necessary to increase the

inlet gas temperature in order to maintain acceptable exit CO levels.

Apart from thermal sintering, activity loss for HTS catalyst is most

commonly due to the presence of poisons in the feedstock and the deposition of

solids on the catalyst. The latter (mainly entrained carbon, boiler water solids,

Figure 12: CO equilibrium vs Inlet S/G and Exit Temperature—Partial Oxidation of Fuel Oil.


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and/or silica) will coat the outer surface of the catalyst and block available

pores. In most modern plants, the HTS catalyst is downstream of equipment

operating at higher temperatures such as reformers and heat exchangers. As a

result, there may be a gradual depositing of steam- volatile compounds into

the HTS bed. In time, these deposits can plug the pores in the catalyst and/ or

the void space between catalyst tablets. As a consequence, the activity declines

and the pressure drop may increase. Proper design of the pore size distribution

and geometric shape of the catalyst pellets can minimize such effects. The

presence of sulfur in the feed gas will affect the size of the converter, as

allowances must be made for the adverse effect of sulfur on catalytic activity.

The presence of oxygen (from the secondary reformer or the partial oxidation

reactor) may also influence the design since the oxygen will be converted to

water through an exothermic reaction. Thus, when the shift feed contains

appreciable oxygen, an allowance may be necessary for the accompanying

temperature rise due to this reaction. Any saturates, e.g. methane, ethane,

propane or unsaturates (ethylene, propylene) in the process gas will

essentially pass through the shift converter unchanged. There is no conclusive

evidence to indicate that the saturates will crack, or that the unsaturates will

be hydrogenated to any significant degree. Even if the unsaturates do

hydrogenate, this side reaction apparently does not affect the catalytic

activity. Acetylene, on the other hand, can be troublesome, because it does

hydrogenate and impair the catalytic activity. When both diolefins and nitric

oxide are present together in the feed stream, polymeric gums are usually

formed and the shift catalyst could be subjected to a serious fouling problem.

3.4. HTS Catalytic Reactor Design ConsiderationsIt must be noted that, unlike the LTS catalyst, the HTS reactor system is

not designed to achieve equilibrium CO leakages for the major part of the

catalyst life. Although equilibrium CO leakages are often experienced at the

Figure 13: HTS Bed Temperature Profile at Start, Middle, and End of Run.


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start of run and for some period of time, the reactor size and catalyst volume to

achieve and maintain equilibrium throughout the total charge life would be

substantially higher and cost prohibitive. As a result of these cost considera-

tions, the reactor is usually designed to be kinetically, rather than

equilibrium, limited. This means that any factor influencing the overall

reaction kinetics will have a much more important impact on the required

catalyst volume for a given application. Typical design lives for a HTS catalyst

are 3–5 years before there is a need for catalyst replacement.

Since HTS catalysts/ reactors are usually designed to be kinetically

limited, the inlet gas temperature will have a significant impact on the

required catalyst volume. Lower inlet gas temperatures will require

increased catalyst volumes to achieve similar levels of performance.

Figure 14 shows the impact of temperature on reaction rates across a typical

Fe2O3-Cr2O3 HTS catalyst. At the higher operating temperatures for HTS

reactors, the WGS reaction is much more pore diffusion limited compared to

the LTS reaction. An increase in reaction rates can be achieved by

incorporating a catalyst with high geometric surface area per unit loaded

volume of the reactor and/or increasing the size of the pores. Although

pressure has no impact on the WGS equilibrium CO levels, there is a

significant influence on the reaction rate because of pore diffusion

considerations. Figure 15 shows the relative influence of system pressure

on reaction rates and the corresponding required catalyst volumes. The rate

increases with reactor pressure up to about 21 bar. Required catalyst

volumes would correspondingly decrease with increasing pressure. Operating

pressure for a HTS plant is usually set more by an examination of overall

economics (related to feedstock supply pressure and equipment costs) rather

than catalyst reaction rate effects.

Figure 14: Effect of Temperature on relative HTS reaction rates over a commercial Fe2O3 -Cr2O3 catalyst.


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3.5. Reaction Mechanisms over Iron Oxide-Chromium OxideCatalystsThe kinetics and reaction mechanism of the HTS WGS reaction has been

studied extensively and various mechanisms proposed (3–5). Temkin et al.

proposed, more than 50 years ago, that the WGS reaction proceeds by an Eley-

Rideal type mechanism, via alternate reduction and oxidation of the surface of

iron oxide (17–19):

COz Oð Þ<CO2zðÞ, and ð21Þ

H2OzðÞ<H2z Oð Þ, ð22Þ

where (O) is an oxygen atom on the oxide surface; and ( ) is a vacant site (an

oxygen anion vacancy) on the surface caused by the removal of an oxygen

atom. The surface is reduced by CO (Eq. 21), and subsequently, oxidized by

H2O [Eq. (22)}. This mechanism is referred to, in subsequent literature, as the

‘‘redox mechanism’’. The term redox mechanism denotes that the catalyst

itself undergoes changes in oxidation state during the course of the

mechanism. It does not refer to the oxidation state changes of the reactants

or products or associated intermediates. It should be pointed out, here, that

the Eley-Rideal mechanisms refer to an adsorbed species reacting with a gas-

phase species. The redox mechanism implied in Eqs. (21) and (22) is

essentially the same as the Mars - Van Krevelen mechanism, with the

difference that the oxygen used to oxidize the catalyst [Eq. (22)] comes from

the water rather than the gas-phase oxygen. Whether this oxygen atom

(extracted from gas-phase H2O) is better referred to as an ‘‘adsorbed’’ species

or, rather as an occupant of the surface lattice position is a moot point.

Figure 15: Effect of pressure on relative LTS reaction rates over a commercial Cu-ZnOcatalyst.


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A multistep Langmuir- Hinshelwood type mechanism (Eqs. 23–27) was

proposed by Oki et al.(20–22) in 1973. From simultaneous exchange rate

measurements, they concluded that while the evolution of gaseous H2 from

adsorbed H atoms [Eq. (27)] is the rate determining step at low CO

conversions, adsorption of CO [Eq. (23)} controls the overall reaction rate at

steady state, near-equilibrium, conditions prevalent in industrial reactors:

CO gð Þza<CO að Þ, ð23Þ

H2O gð Þz3a<2H að ÞzO að Þ, ð24Þ

CO að ÞzO að Þ<CO2 að Þza, ð25Þ

CO2 að Þ<CO2 gð Þza:, and ð26Þ

2H að Þ<H2 gð Þz2a ð27Þ

In the above Eqs., ‘‘a’ refers to an adsorption site. It can be located either on

the support or the metal oxide. The HTS reaction on Fe2O3-Cr2O3 catalysts

probably proceeds by an oxidation- reduction mechanism (See Section 9).

3.6. Metal- promoted Iron Oxide –Chromium Oxide HTSCatalystsThe possibility of increasing the activity of Fe2O3-Cr2O3 catalysts by

promotion has been studied by Trimm and coworkers (23–25). Small amounts

of precious metals were found to increase the rate of the forward reaction CO +H2O R CO2 + H2 and to increase the rate (Fig. 16). Platinum was found to

increase the reactivity of all the oxides with the promotional effect being most

pronounced with Cr2O3, U3O8 and CeO2-ZrO2 supports. Comparisons were

also made with Pt-U3O8 which was as efficient (on a weight basis) as Pt-

Fe2O3-Cr2O3. In fact, the specific activity, on an area basis, of Pt-U3O8 (BET

area 5 2.3 m2/g) was more than 25 times that of Pt-Fe2O3-Cr2O3 (BET area 5

63 m2/g) (Table 3). However, catalytic activity over this catalyst dropped

quickly as temperature was reduced. Trimm (19, 21) has also compared the

kinetics of the WGS reaction for the promoted and unpromoted Fe2O3-Cr2O3

catalysts. The general power rate law expression remained unchanged in the

absence and presence of the noble metal promoter indicating that it is the

number of active sites that has increased by promotion (by noble metals).


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Figure 16: Apparent activation energy plots for promoted iron-chromia catalysts (23).

Table 3: Rates and apparent activation energies for water gas shift over 1%Pt/oxide catalysts (21).

Catalyst BET area (m2/g21)Rate at 450uC

(mmol(CO) gcatalyst21s21) Ea(kJ/mol)

Pt/Cr2O3 22 0.174 41 ¡ 2Cr2O3 0.022 78 ¡ 1Pt/Cr2O3-Fe3O4

a63 0.149 50 ¡ 3

Cr2O3-Fe3O4a 0.124 70 ¡ 2

Pt/U3O8 2.3 0.142 59 ¡ 3U3O8 0.01 24 ¡ 2Pt/CeO2-ZrO2

a67 0.079 28 ¡ 1

CeO2-ZrO2a 0.008 55 ¡ 1

Pt/CeO2-Fe3O4

an.m.b 0.07 50 ¡ 1

Pt/CeO2 122 0.055 52 ¡ 1Pt/MgO 77 0.034 41 ¡ 1Pt/V2O5 6 0.032 52 ¡ 3Pt/ZrO2 n.m.b 0.026 24 ¡ 1Pt/Fe3O4 29 0.022 55 ¡ 3Fe3O4 0.023 48 ¡ 2Pt/MoO3 1.6 0.02 49 ¡ 3Pt/Bi2MoO6 2.1 0.018c 62 ¡ 4Pt/MnO2 17 0.016c 53 ¡ 2Pt/Al2O3 272 0.014 47 ¡ 1

aThe composition of the mixed oxides were as follows: 8 wt% Cr2O3-Fe3O4, 8 wt% CeO2 – Fe3O4,50 wt% CeO2-ZrO2. bNot measured. cNo measurement at 450uC; calculated from Arrheniusparameters.


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Apparent activation energies were found to be similar (, 50 kJ/mole) for Pt

supported on CeO2, Fe3O4-Cr2O3, CeO2-Fe3O4, Fe3O4, V2O5, MgO, MnO, and

Al2O3, despite up to 15-fold differences in rates of reactions (Table 3;

Figure 16). Since it is unlikely that surface diffusion of oxygen will have

similar activation energies for such a variety of solids, the authors suggested

that diffusion of oxygen on the surface or across the surface of the support to

react with CO adsorbed on the metal cannot be rate controlling in WGS

reactions on Fe3O4-Cr2O3 promoted with noble metals. Rhodium was found to

be the most active promoter for Fe2O3-Cr2O3 oxides. More recently, this group

had probed (24) the origin of rhodium promotion of Fe3O4-Cr2O3 catalysts for

the HTS reaction using various kinetic techniques and concluded that, of the

two steps that may restrict the rate of the WGS reaction over iron- chromium

oxide catalysts (reduction by CO and H2 generation through reoxidation by

water), rhodium acts primarily by accelerating the latter.

Although the promoted catalysts are more efficient than the unpromoted

Fe2O3-Cr2O3 catalysts above 300uC, they are still less active than the copper-

based catalysts at temperatures below 300uC. As a result, the CO concentra-

tion is reduced but, at about 3–4%, it is still too high for many applications

(e.g., fuel cells, ammonia synthesis). Low temperature WGS is required to

reduce CO concentrations still further.

During the last decade, attempts to develop improved HTS catalysts have

been along two main lines:(A) replacing, at least partially, Fe by more active

elements (like noble metals), and (B) replacing Cr, partially or completely, by

non-toxic elements like Cu, Ca, Ce, Zr, La etc. (26–34). Promotion of the

Fe2O3-Cr2O3 catalysts by 2 wt% Ag, Cu, Ba, Pb and Hg was explored by

Rhodes et al. (31). The catalysts were prepared by coprecipitation. Boron was

found to poison the activity slightly whereas the others did increase the

activity between 350–440uC, with the relative order being Hg.Ag, Ba

.Cu.Pb. unpromoted Fe2O3-Cr2O3 . B. From their results, the barium or

silver- promoted Fe2O3-Cr2O3 catalysts appear promising: a 10–15% increase

in CO conversion was observed (when Ag or Ba was incorporated in the

conventional Fe2O3-Cr2O3 catalysts) at their reaction conditions (400uC, 27

bar, GHSV 5 1.2 6 10 6 h21; the volume of steam was 75 volume% of the dry

gas). The promoters decreased the activation energy of the reaction (Table 4).

Based on the compensation effect (Fig. 17) seen when the activation energies

were plotted against the corresponding pre-exponential factors (in the

Arrhenius Eq.), the authors concluded that CO adsorption is an important

factor controlling the relative catalytic activities of the various samples in the

WGS reaction. Andreev et al.(29) studied the effect of the addition of CuO,

CoO, and ZnO (5 wt%) on the activity of Fe2O3-Cr2O3 catalysts. The Cu –

promoted sample was found to be the most active at 380uC. Kappen et al. (30)

investigated the state of their Cu promoter (0.17–1.5 wt%) in Fe2O3-Cr2O3


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catalysts and found that Cu was in the metallic state under the WGS reaction

conditions. However, it was reoxidized easily when exposed to the atmosphere.

Most of the current generation, industrial, HTS catalysts contain oxides of Fe,

Cr and Cu.

3.7. Chromium-free HTS CatalystsWhen chromium oxide is used as a component of a catalyst, especially in

hexavalent form which is soluble in water, expenditures must be incurred to

guarantee worker safety both during production and later handling of the

Table 4: Effect of additives on the performance of Fe3O4/Cr2O3 water gas shiftcatalysts (27).

Additive CO conversiona (%) Activation energyb (kJ/mol)

None 18.8 112B 18.7 108Pb 25 90Cu 27.9 81Ag 32.9 74Ba 33.5 83Hg 37.4 82

a CO conversion at 400uC, 27 bar, GHSV51.2 6 106 h21. b ¡ 4kJ/mol.

Figure 17: ‘‘Compensation effect’’ plot for the modified Fe2O3/Cr2O3 catalysts (31).


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catalyst, and health hazards cannot be fully ruled out despite considerable

effort. In addition, the spent catalyst ultimately poses a hazard to man and the

environment and must be disposed of in accordance with the laws for the

disposal of toxic waste. HTS catalysts completely free from Cr and containing

Ca, Ce or Zr were first claimed by Chinchen (34). Their catalytic activities,

however, were low. Similar low- active catalysts, based on Mg and Zn ferrites,

were also reported by Rethwisch and Dumesic (35). More active catalysts

based on alkali- promoted Co-, Cu-and Fe- manganese oxide systems were

reported by Hutchings and coworkers (36, 37). The relative first order rate

constants in the WGS reaction for Fe-Cr, Fe-Mn, Cu-Mn and Co-Mn catalysts

were found to be (36, 37) 1.0, 0.06, 0.75 and 1.75, respectively. Their Co-Mn

catalyst, however exhibited significant methanation activity and the Cu-Mn

catalysts were more sensitive to sulfur than the Fe-Cr formulations. Ladebeck

and Kochloefl (38) had found that chromia-free, iron oxide catalysts containing

about 5 wt% of Al2O3, 2 wt% of Cu and 2.5 wt% of CeO2 were very active for the

HTS reaction. The incorporation of ZrO2, La2O3 or MnO instead of CeO2

resulted in catalysts with a high initial activity but with a poorer stability.

Araujo and Rangel (28) investigated the catalytic performance of Al-doped, Fe-

based catalysts with small amounts of copper (3 wt%), prepared by the

coprecipitation (for Al and Fe) - impregnation (for Cu) method, in the HTS

reaction. The aluminium and copper-doped iron catalyst was studied at 370uCand showed similar activity compared to the commercial Fe-Cr-Cu catalyst.

Costa et al. subsequently examined (33) the use of thorium, instead of

chromium, in Fe- Cu – based catalysts for the HTS reaction. These Fe-Th-Cu

catalysts were more active than the commercial Fe-Cr-Cu catalyst at H2O/CO

5 0.6 and 370uC. Its high activity was attributed to an increase in surface

area due to the presence of thorium. From a detailed study of chromium-free,

iron-based HTS WGS catalysts, Natesakhawat et al. (26) concluded that a

combination of copper and aluminum is a potential replacement for Cr in HTS

catalysts. Further improvements in HTS activity of Fe-Al catalysts could be

achieved by the addition of small amounts of copper or cobalt. The CO

conversions (at 400uC, CO/H2O/N2 5 1/1/18 (vol) and feed GHSV of 6000 h21)

were 43, 46, 27, 12 and 16% for Fe-Cr, Fe-Cu-Al, Fe-Al, Fe-Ga and Fe-Mn,

respectively. As a textural promoter, aluminum oxide (like chromium oxide)

prevented the sintering of the iron oxide crystallites and stabilized the active

phase, magnetite (Fe3O4) by retarding its further reduction to FeO and

metallic Fe. The promotional effect of Cu was found to be strongly dependent

on the preparation method. Fe-Cu-Al catalysts prepared by a one-step method

(simultaneous coprecipitation of all the three component hydroxides) had

higher CO conversions than those prepared by a two-step, coprecipitation –

impregnation method (coprecipitation of the Fe and Al components followed by

impregnation of Cu on the precipitate obtained by coprecipitation) (Fig. 18).


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Although the activities were similar at 250uC, the Fe-Cu-Al catalyst pre-

pared by the one-step method was more active at higher temperatures. The

better stabilization against sintering, at higher temperatures, of the copper

crystallites in the coprecipitated samples is probably the reason for its

superior performance. A significant difference in the temperature-pro-

grammed-reduction profile was also observed between the two samples

(Fig. 19). The low temperature reduction profile has contributions from

three different reduction sites. The reduction of hematite to magnetite

appears to have shifted from 300–290uC. The major peak from reduction of

Cu species appears at 260uC with a very weak shoulder around 220uC. This

shoulder is possibly due to reduction of the Cu species which are on the

external surface of the catalyst and which can be reduced easily. The rest of

the copper species appear to be more difficult to reduce as seen by a shift in

reduction temperature from 220–260uC, possibly due to stronger interaction

with the hematite matrix. While the peak resulting from the reduction of

hematite to magnetite shifts to lower temperatures (from 300–290uC), the

peak from a further reduction of magnetite changes very little compared to

the non-promoted, Fe-Al sample. These results suggest that the preparation

method makes a significant difference in the way Cu promoter is

incorporated into the catalyst structure. During the last decade, more

active and chromium-free, noble metal-based HTS catalysts are under

development for use in fuel cell applications. These will be described later,

in Section 6.

Figure 18: Effect of preparation method (one vs. two steps) of Fe2O3-Al2O3-CuO catalysts onCO conversion (26).


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4. LOW TEMPERATURE WATER GAS SHIFT CATALYSTS

4.1. Cu-ZnO-Al2O3 CatalystsAn excellent perspective of the historical background for the evolution of

the low temperature water gas shift catalysts has been provided by Twigg et

al. (5). The development of highly efficient sulfur removal hydrodesulfurisa-

tion technologies using Co(Ni)- MoO3- Al2O3 catalysts in the 1960s provided

ammonia manufacturers with syngas streams containing less than 1.0 ppm

sulfur. This, in turn, enabled the use of the otherwise sulfur-sensitive Cu-ZnO

catalysts at sufficiently low temperatures (190–200uC) when the equilibrium

CO concentrations, at the exit of the LTS converters, can be below 0.3%. It

may be recalled that the Fe-Cr catalysts are not active enough below 350uC, at

which temperature, the equilibrium CO concentrations are around 3–5%. As a

direct consequence of having such low levels of CO (below 0.3%wt) from the

Cu-ZnO catalysts, it was economic to incorporate a methanation stage in the

process in place of the more complicated copper liquor scrubbing system that

Figure 19: Effect of preparation method (one vs. two steps) on the temperatureprogrammed reduction profiles, in H2, of Fe2O3-Al2O3-CuO catalysts (26).


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was formerly used to remove residual CO, thereby enhancing the technoeco-

nomic viability of large ammonia plants. The activity of metallic copper in the

WGS reaction has, of course, been known for a long time (4–7). The problem

was the easy copper sintering and the subsequent loss of copper surface area

during the activation (reduction of the precursor copper oxide) and use of

copper catalysts. Various stabilizers, like SiO2, Cr2O3, Mn - Cr2O3 etc., were

evaluated for their ability to stabilize the copper surface area (39). The

introduction of the Cu-ZnO catalyst in the early 1960s (5) and, later, the Cu-

ZnO-Cr2O3, the Cu-Zn-Mn-Cr2O3 and, especially Cu-ZnO-Al2O3 formulations

enabled the production of catalysts with high and stable copper surface areas

and established the LTS process as a standard operation in any scheme of

hydrogen production from carbonaceous raw material. Currently, Cu-ZnO-

Al2O3 based catalysts are used almost exclusively for industrial LTS

operations.

The feed gas to the LTS reactor is that exiting the HTS unit cooled, either

by direct or indirect heat exchange, to approximately 200uC. LTS converters

are employed more frequently in hydrogen and ammonia - producing plants

than in methanol or hydrocarbon (Fischer-Tropsch) plants. Hydrogen plants

normally begin with primary steam/hydrocarbon reforming of natural gas to

syngas which is then water gas shifted over HTS and LTS catalysts to

maximize the hydrogen mole fraction of the effluent. Following CO2 scrubbing

and methanation to remove unreacted CO, the product hydrogen is then

utilized for hydrocracking, hydrogenation or other service. In ammonia plants,

there is a secondary reformer between the primary reformer and the HTS

units for the introduction of the requisite nitrogen. Process conditions in the

WGS section are more severe than in hydrogen plants. This is because the

downstream ammonia process is considerably more sensitive to the purity of

the hydrogen produced. Not only does a lower level of hydrogen reduce

ammonia production, but also the corresponding higher level of inerts (like

CH4 and CH3OH) increases the purge rate from the synthesis loop. The

principal deactivation mechanism for LTS catalysts is poisoning by sulfur and

chlorides contained in the process gas. If only a single bed of LTS catalyst is

employed, this deactivation process begins as soon as the catalyst is placed on

stream and, normally, within 6–12 months, a rise in CO leakage will be

detected. In an ammonia plant, a rise of 0.1% CO in the LTS converter effluent

is roughly equivalent to a production loss of 30 T/day in a 3000 T/day of

ammonia plant. To minimize this production loss and maintain a low CO

leakage for a long period of time, many plants have installed guard beds of

LTS catalysts immediately ahead of the main LTS unit. These beds are

usually about 1/4 the size of the main LTS bed and serve to sacrificially screen

poisons from the main bed and to promote additional water gas shift. In

general an ammonia process will tolerate up to 0.4% CO in the LTS effluent


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before process economics dictate a catalyst change. Hydrogen plants may

tolerate a slightly higher CO leakage.

The preparative chemistry of the Cu-ZnO, with or without Al2O3 or Cr2O3,

has been studied extensively and is still a subject of interest since the nature

of the precursor mixture and its evolution during the preparation steps seem

to influence the catalytic properties. In an early publication, Uchida et al.(40)

tested several catalyst combinations, Cu/Zn, Cu/Al, Cu/Al/Zn, Cu/Fe and Cu/

Cr, prepared by the coprecipitation of their mixed hydroxides/carbonates/

hydroxycarbonates, and compared their catalytic activity and stability in the

WGS reaction. Addition of zinc to copper increased the catalytic activity which

reached a maximum around a Cu/Zn ratio of 0.4. They observed that the

method of preparation of the Cu-ZnO catalyst is extremely important in

determining the catalytic activity. They also established, using x-ray

diffraction, that the major constituents of a Cu-ZnO catalyst after use were

copper metal and zinc oxide (41). It was speculated, even at that early stage,

that copper metal can be the active ingredient (42, 43). Highly active catalysts

are prepared by coprecipitation from the corresponding aqueous solutions of

metal nitrates with sodium carbonate and having Cu/Zn atomic ratios

between 0.4 and 2.0 (42–45). To avoid extensive washing of the filter cake

for a reduction of its Na content, ammonium carbonate or hydroxide as a

precipitating agent was recommended by Sengupta et al. (46). The thermal

decomposition of the resultant, aqueous (Cu,Zn)(NH3)4(HCO3)2 complexes by

steam provides alkali-free Cu-Zn hydroxycarbonates. Petrini et al. (44, 45) had

also noted that highly active catalysts can be prepared if Al(OH)3 is added

during the Cu-Zn precipitation.

Gines et al. (47) reported a detailed study of the influence of preparation

methods on the activity and structure sensitivity of the Cu-ZnO-Al2O3 mixed

oxide catalysts. Samples were prepared by coprecipitation from aqueous

solutions of the nitrates of Cu, Zn and Al with sodium carbonate at 60uC and a

constant pH around 7 in a stirred batch reactor. The precipitates were filtered,

washed with distilled water at 60uC until no sodium ions were detected

and dried at 90–100uC overnight. Finally, the samples were decomposed in

air for 8 h at temperatures between 400–700uC. Depending on the ratio of

Cu, Zn, and Al cations, different hydroxycarbonate phases were formed:

malachite[Cu2(OH)2CO3], which is capable of substituting Cu by zinc and

is called zincian-malachite or rosasite [(Cu,Zn)2 (OH)2CO3],hydrotalcite

[(Cu,Zn)6Al2CO3(OH)16. 4H2O], aurichalcite [Cu,Zn)5(CO3)2(OH)6] and hydro-

zincite [Zn5(CO3)2(OH)6]. The rosasite phase can transform to the aurichalcite

phase for Zn concentrations greater than 40 mol%. No trace of copper

hydroxynitrates like gherardite was observed. An important observation

was that hydrotalcite was selectively obtained as a single phase only in

preparations using a (Cu+Zn)/Al atomic ratio of 3, the stoichiometric, M2+/M 3+


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metal cation ratio in hydrotalcite. The BET surface areas increased with Al

content. On thermal decomposition, the mixed oxide, CuO-ZnO-Al2O3, was

obtained. X-ray diffraction data revealed the presence of crystalline CuO and

ZnO. Additionally, amorphous alumina was also present. Crystalline, spinel-

like, ZnAl2O4 was detected in trace amounts only in samples containing .13%

Al2O3. The concentration of free CuO and ZnO crystallite sizes were related to

the hydrotalcite content in the hydroxycarbonate precursor: higher the

amount of hydrotalcite in the precursor, the lower the CuO and ZnO

crystallite sizes in the resulting mixed, ternary oxide. The influence of the

hydroxycarbonate precursor structure was preserved throughout the calcina-

tion step and manifested itself in a different reducibility of the CuO/ZnO

precursors. After activating and reducing the samples, they were tested in the

WGS reaction. Cu (I) oxide is a probable intermediate in the reduction of CuO

to Cu metal. Completely reduced Cu clusters on ZnO constitute the active bulk

phase for the WGS reaction. A gas mixture consisting of 10%CO/30% N2/

30%H2/ 30% H2O was fed to the fixed bed reactor at a volumetric flow of 750 ml

STP/ min (catalyst weight 5 0.5 g). The reaction was carried out at 230uC and

1 bar. It should be pointed out that CO2, one of the products of the reversible

WGS reaction, was not included in the inlet gas mixture. A remarkable feature

of their catalytic result is that the turnover frequency (number of CO2

molecules produced per surface copper atom per second) was essentially

constant, not only when the copper metal surface area was varied between 3–

35 m2/g Cu, but also when the CuO loading was varied between 30 and 50 wt%,

the Al/Zn atomic ratios between 0 and 2.5, the copper dispersion between 0.5

and 5.0%, and the calcination temperature between 400–700uC, clearly

suggesting that the specific reaction rate is proportional to the copper metal

surface area. Based on these results, the authors concluded that (a) the WGS

reaction is a structure insensitive reaction and linearly proportional to the

surface area of metallic copper; and that (b) both the metallic copper

dispersion and catalytic activity were related to the amount of hydrotalcite

contained in the precursor precipitate; the higher the content of the

hydrotalcite in the precursor, the higher the catalytic activity of the resulting

catalyst.

Contrary results, namely, that the turnover frequency does vary, by an

order of magnitude, when the copper metal surface area was changed from 10

to 40 m2 /g, for the Cu-ZnO-Al2O3 system had been reported earlier, by

Chinchen and Spencer (48). These authors had carried out the WGS reaction

at 30 bar and their reaction mixture had included CO2 under conditions closer

to those in practice in the industry. Even though it is well established (4, 5)

that the catalytic activity of Cu-ZnO-Al2O3 catalysts in WGS reactions

increases with the surface area of metallic copper, there are no reports that,

under industrial conditions, the rate of the reaction correlates linearly with


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the metallic Cu area over the entire Cu-Zn composition range. While a high Cu

surface area is a necessary prerequisite for catalytic activity, additional

factors like the ‘‘microstrain’’ in the copper nanocrystallites due to the

presence of Zn ions probably affect catalytic activity. The hydroxycarbonate

precursors mentioned earlier probably influence the residual concentrations of

Zn in the Cu metal crystallites in Cu-ZnO. Similarly, oxygen vacancies in ZnO

formed, for example, during the reduction/activation of the catalyst or during

the WGS reaction will also influence the catalytic activity indirectly by

influencing the wetting behavior at the Cu/ZnO interface and, thereby, the

‘‘microstrain’’ in the Cu crystallites. Hence, bulk structural changes in the

ZnO or Cu metal crystallites resulting from the preparation procedures cannot

be ignored.

4.2. Promoted Cu-based LTS CatalystsAttempts have been made during the past decade to prepare alternate

base metal catalysts which are superior to the conventional Cu-ZnO-Al2O3

catalysts. Tanaka et al. (49–51) have explored the performance of Cu-Mn

spinel oxides in the LTS reaction. They had originally found (50, 51) that Cu-

Mn spinel catalysts which were prepared by coprecipitation with NH3, showed

a WGS activity comparable to that of Cu-ZnO-Al2O3 catalysts in spite of their

low surface area. Since Cu and Mn ions may not have coprecipitated

homogeneously due to formation of the copper amine complex, [Cu(NH3)4]2+

by the NH3 coprecipitation method, they later prepared (49) their catalysts by

citric acid complex, urea homogeneous coprecipitation or the Pechini method.

The last method involves the polymerization accompanied with esterification

of ethylene glycol and citric acid during the precipitation of the hydroxides of

Cu and Mn. Higher CO conversions were obtained for samples prepared by the

citric acid method. CO conversion was enhanced with a rise in the calcination

temperature of the Cu-Mn spinel prepared by the citric acid method. Partial

substitution of Fe or Al for Mn in the spinel lattice enhanced their CO

conversion activity to levels higher than that of conventional Cu-ZnO-Al2O3

catalysts when the temperature was increased to 300uC Fig. 20.

4.3. KineticsThe kinetics of the water gas shift reaction has been studied extensively

(52–59). An accurate description of the measured reaction rates from a data

set can be obtained from an expression where all kinetic parameters are fitted,

for example, to a power law. Such empirical kinetic expressions are essential

in reactor design calculations where it is necessary to have a very accurate

description of the reaction rate. Different mechanisms, however, can lead to


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the same overall kinetic expression. Hence, it is difficult to determine the

mechanism from an empirical kinetic expression alone. Microkinetic models

are useful here as they are based on the knowledge about elementary steps

and their energetics. They enable us to estimate surface coverages, reaction

orders, and activation enthalpy during reaction conditions. Ovesen et al. (58)

had analysed the microkinetics of the WGS reaction under industrial

conditions based on a model developed by them earlier (59). The reaction

was studied over three different Cu- based catalysts, Cu-ZnO-Al2O3, Cu-Al2O3

and Cu-SiO2. The Cu-ZnO-Al2O3 catalysts contained about 40% Cu, 22% Zn

and 5% Al. Ovesen et al.’s model (58, 59) is based on the surface redox

mechanism:

1. H2O (g) + * u H2O*

2. H2O* + * u OH* + H*

3. 2OH* u H2O* + O*

4. OH* + * u O* + H*

5. 2H* u H2 (g) + 2*

6. CO (g) + * u CO*

7. CO* + O* u CO2* + *

8. CO2* u CO2 (g) + *

Figure 20: CO conversion over Cu-Mn catalysts. Reaction conditions: H2 37.5%; CO, 5.0%;H2O, 25.0%; CO2, 12.5%; space velocity, 6400 h21 (49).


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where the asterisk signifies a free surface site and X* is an adsorbed species,

X. The expressions for the rate and equilibrium Eqs. that constitute the model

are shown in Table 5. The model describes the coverage of surface species in

addition to the overall rate. When this model was tested against measure-

ments for an industrial (Cu-ZnO-Fe2O3) catalyst at 1 bar by Van Herwijnen

and de Jong (52), a good agreement was found (58). From parallel

physicochemical measurements, it was deduced that the catalyst exposed

nanocrystallites of Cu (111) facets almost exclusively. The rate-determining

step was dependent, critically, on the composition of the feed gas mixture. It

was found that reaction step 2 above is rate limiting in a gas with a low ratio of

water to carbon monoxide whereas reaction step 7 is rate limiting in a gas with

a high ratio of water to CO. Reaction 4 was significant in a CO2 + H2 mixture.

However, when this model was tested against the high pressure data,

deviation between the calculated and experimental rates was found (58). To

describe the kinetics of the water gas shift reaction at industrial conditions it

was necessary to include the synthesis and hydrogenation of formate (reaction

steps 9–11 below):

9. CO2* + H* u HCOO* + *

10. HCOO* + H* u H2COO* + *

11. H2COO* + 4H* u CH3OH (g) + H2O(g) + 5*

The reaction step 9 was in equilibrium under the industrial conditions (high

pressure). The coverage of HCOO* was always low. Ovesen et al.’s reaction

Table 5: Rate and Equilibrium Eqs. for Kinetic Model (51).

K1PH2O

P0

~HH2O

r2~k2HH2OH{k2

K2

HOHHH

K2H2OH~HH2OHO

r4~k4HOHHO{k4

K4HOHH

K5H2H~

PH2

P0H2

K6PCO

P0H~HCO

r7~k7HCOHO{k7

K7HCO2

H

K8HCO2~

PCO2

P0H

Note: ki is Forward Rate Constant, Ki Equilibrium Constant, and hi Surface Coverage of Species i).


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sequence for the industrial LTS reaction (58) consists of the steps 1 through

11 with steps 2, 4, 7, and 10 as possible slow steps. The formate may be

present on the surface but, it is not a species in the catalytic cycle for CO

conversion to CO2. These conclusions from kinetic studies are similar to those

of a combined kinetic and DRIFTS study of Pt- and Au-based catalysts by

Meunier et al. (60–62). It should be pointed out, here, that Ovesen et al.’s rate

equations do not consider that co-adsorbed water molecules may influence the

rate of decomposition of the formate intermediate. It will be interesting to

explore the changes if this issue is taken into consideration (see sections 9 and

10). The satisfactory agreement between the calculated exit mole fraction of

CO from the microkinetic model and the experimental exit mole fraction of

CO for Cu-ZnO-Al2O3 is shown in Fig. (21). This model was refined further in a

later publication by Schumacher et al. (54). It was established that the

adsorption energies for CO and oxygen (the latter arising from H2O) can

describe, to a large extent, changes in the remaining activation and adsorption

energies through linear correlations. The model predicted well the order of

catalytic activities for transition metals although it failed to describe the

experimental data quantitatively. The discrepancy was due to the neglect of

adsorbate-adsorbate interactions which play an important role at high

coverages. The model also predicted that the activity of copper can be

improved by increasing the strength with which CO and oxygen are bound to

the surface, thus suggesting possible directions for improving the LTS

catalyst.

Figure 21: Calculated (from the microkinetic model) and experimental exit mole fraction (inwet gas) of CO for Cu/ZnO/Al2O3 (51).


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4.4. Deactivation of LTS Catalysts

4.4.1. Thermal Sintering

When formulated properly and operated under standard LTS condi-

tions, the Cu-ZnO-Al2O3 catalyst is quite rugged and lasts a few years.

The major sources of catalyst deactivation are thermal sintering of the

copper crystallites and poisoning by sulfur and chlorine compounds. Twigg

and Spencer have reviewed the deactivation of copper-based catalysts in

the WGS reaction (64). Due to the low melting point of copper metal

(1083uC), copper has low Tammam and Huttig temperatures. Cu-ZnO-

Al2O3 catalysts sinter and lose copper surface area, and, hence, catalytic

activity, when heated above 300uC. Indeed, one of the major roles of Al2O3

is to retard such growth of copper crystallites and function as a textural

promoter. Details of the mechanism of the thermal sintering of Cu

catalysts under hydrogen at elevated temperatures were studied by Tohji

et al. (65) using EXAFS techniques. As the temperature was increased in

hydrogen, a quasi- two-dimensional layer of copper metal epitaxially

developed over the ZnO support below 127uC. Between 127–230uC, small

copper metal clusters dispersed over ZnO start to appear. Above 250–

300uC, the small clusters fuse to give larger copper metal crystals

agglomerated on the support. Since these catalysts begin to lose copper

surface area and catalytic activity also above 250uC, it is reasonable to

assume that the active sites for the WGS reaction are associated with the

small copper clusters and their concentrations are diminished when the

small crystallites grow into larger ones. Thermal sintering leads to their

growth and consequent catalytic deactivation (66).

4.4.2. Sulfur Poisoning

The second major cause of deactivation of these catalysts is poisoning by

sulfur compounds present in the reaction gas stream. The LTS, Cu-ZnO-Al2O3

catalyst operates at 190–250uC, a temperature sufficiently low wherein

thermodynamics favors strong adsorption of poisons. Sulfur is a powerful

poison for Cu, as indicated by the change in enthalpy of the sulfidation

reaction [Eq. (28)] (64):

2 CuzH2S[Cu2SzH2 DH~{59:4 kJ=mol ð28Þ

Sulfiding of copper, hence, occurs. The corresponding equilibrium constant is

(64) about 1 6 10+5. Sulfur, accumulating on the surface, blocks the pores and

the active sites leading to catalytic deactivation. To retain the long term


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activity of copper catalysts, it is usual to maintain gas phase sulfur

concentrations below 0.1 ppm of S. In addition to keeping gas phase

concentrations of sulfur low, the ZnO component of the catalysts is also

engineered during catalyst manufacture to divert the sulfur away from the

small Cu crystallites and absorb it in ZnO as ZnS. It is essential to keep the

crystallite size of ZnO as small as possible to accomplish this absorption. The

reaction of H2S with ZnO [Eq. (29)] is quite exothermic (66) and proceeds

readily:

ZnO sð ÞzH2S gð Þ[ZnS sð ÞzH2O gð Þ ð29Þ

DHu 5 276.7 kJ /mol; DSu 5 23.0 J.mol21 K21.

There are two forms of zinc sulfide, wurtzite (a-ZnS), and sphalerite (b- ZnS)

and both forms are seen in discharged plant samples of zinc oxide absorbents.

Sphalerite is the more stable form and the above data refer to this form. The

equilibrium constant at 500 K is 7.4 6 10 7. The reaction is strongly favored

thermodynamically.

4.4.3. Chloride Poisoning

Chlorine compounds, like HCl, form low-melting cuprous chloride (m.p. 5

430uC) on reaction with copper in the Cu-ZnO catalyst. The ZnCl2 formed also

has a low melting point (283uC). Their formation is favored thermodynami-

cally (64) under the WGS reaction conditions [Eqs. (30, 31)]:

Cu sð ÞzHCl gð Þ[CuCl sð Þz0:5 H2 gð Þ ð30Þ

DHu 5 2 43.5 kJ /mol, and

ZnO sð Þz2HCl gð Þ[ZnCl2 sð ÞzH2O gð Þ ð31Þ

DHu 5 2 121.8 kJ /mol; DSu 5 117.2 J.mol21 K21.

These mobile chlorides facilitate the movement and sintering of copper as well

as the ZnO crystallites on the catalyst surface. The limits on HCl content to

avoid catalyst poisoning are more severe than for H2S poisoning, on the order

of 1 ppb. Unlike the case of sulfur poisoning, the ZnO cannot offer any

protection in the case of chloride poisoning.

In addition to the major poisons, sulfur and chloride, the Cu-ZnO catalysts

are also deactivated by the presence of As, trivalent phosphorous, silica, and

transition metals like Fe, Co and Ni, in the feed stream. Due to their low

temperature of operation, Cu-ZnO catalysts do not form significant amounts of

coke when operated with purified feedstock (65).


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5. SULFUR TOLERANT WGS CATALYSTS

The sulfur levels in natural gas or light petroleum naphtha are in the range of

5–50 ppm and conventional hydrodesulfurisation of the feedstock with Co (Ni)-

Mo- alumina catalysts is used before steam reforming them with nickel- based

catalysts. The latter are deactivated in the presence of sulfur. H2S can be

removed from natural gas as well as hydrodesulfuriser effluents by reaction

with ZnO at 370uC. Other sulfur compounds can be removed from natural gas

also by absorption at ambient temperatures on activated charcoal (loaded with

copper) or molecular sieves. The efficiency of these absorption systems

depends both on the type of sulfur compounds and on the amount of high

molecular weight hydrocarbons in the natural gas. Low boiling sulfur

compounds, like COS, are not strongly absorbed and condensable hydro-

carbons can rapidly saturate the absorbent. Catalytic hydrodesulfurisation

can remove COS. The removal of H2S by absorption in a hot ZnO bed is usually

not complete. Approximately 50 ppb H2S slips through and enters the

reformer upstream of the WGS reactor. After the volume expansion due to

the reforming reaction, the resulting H2S concentration in the gas entering

the WGS reactor is about 10 ppb. Due to the high temperatures in the reformer

and the low capacity of modern Fe2O3-Cr2O3-Cu HTS catalysts for sulfur

absorption, nearly all this residual sulfur exits the HTS stage and is removed

from the syngas by the Cu-ZnO-Al2O3 LTS catalyst located downstream.

Unlike Fe2O3-Cr2O3 catalysts, the Cu-ZnO-Al2O3 catalysts are adversely

affected by the presence of sulfur compounds in concentrations greater than

about 0.1 ppm. The deactivation is irreversible even when the sulfur is

removed from the feed gas stream. Normally the Cu-ZnO-Al2O3, LTS catalyst

reactors are designed for a space velocity of 1000–2500 h21 to take into account

poisoning by this sulfur. The actual catalyst volume in the reactor represents

approximately three times the volume needed by the kinetics. The Cu-ZnO-

Al2O3 catalysts, thus, serve also as a total sulfur absorber protecting

downstream processes (ammonia, methanol and Fischer-Tropsch syntheses,

hydrogenations, fuel cell electrodes, etc.) in industrial applications.

It is necessary to consider sulfur-tolerant WGS catalysts mainly when the

syngas is generated by the gasification and partial oxidation of heavy fuel oil,

tar sands, oil shale, coal, coke or biomass. Syngas from these raw materials

contain much larger concentrations of CO (up to 50%) (Table 1) and sulfur (up

to 3% wt) (5). In such cases, the Fe2O3-Cr2O3 catalyst had to be used as the

only WGS catalyst; conventional, Cu-based LTS catalysts cannot be used at

the high sulfur concentrations at the exit of HTS reactors in such operations.

The Fe2O3-Cr2O3 catalyst is sulfided during use and, in the sulfided state its

activity is much lower than in the oxide state. It is, hence, necessary to either

operate at higher H2O/CO ratios or remove the sulfur compounds from the

process gas over sulfided Co-Mo-alumina catalysts before it enters the HTS


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reactor. In view of the high energy costs of operating at high H2O/CO ratios,

the latter option is usually adopted (5). In addition to removal of sulfur

compounds, these Co-Mo-alumina catalysts also serve to remove CO (by the

WGS reaction) from the process gas and, thus, serve as sulfur-tolerant, sour

gas shift catalysts. In fact, these catalysts are active mainly in the sulfided

form. When such sour gas catalysts based on Co-Mo sulfides are used, the

preferred minimum inlet sulfur in the feed for acceptable perfomance is about

300 ppm. If these catalysts are adequately presulfided before use, then they

can operate satisfactorily even in feed streams that contain H2S at a level as

low as 35 ppm. Non-sulfided Co-Mo catalyst exhibits very little WGS activity.

Commercial sour gas converters with Co-Mo catalysts operate in the

temperature range of 250–350uC and at pressures from atmospheric to 40

bar. Typical process conditions in a Co-Mo- based sour gas shift catalytic

reactor in a H2 plant using Texaco partial oxidation process to generate syngas

from heavy oil are shown in Table 6. The syngas from the partial oxidation

reactor contains 0.25% sulfur. The sulfided Co-Mo catalyst is deployed in 3

beds. The CO content is reduced from 46% (vol) at the inlet to the first bed to

1% at the exit of the third bed. It should be noted that all the cobalt moly-based

sour gas shift catalysts convert H2S in the presence of CO into COS. Therefore,

the COS concentration at the outlet of the last sour gas shift reactor is at

equilibrium. At high operating pressures and relatively high steam/dry gas

ratio, the resulting COS concentration is usually well below 0.1 ppmv.

However, under certain circumstances, the COS concentration can be much

higher and downstream COS hydrolysis has to be considered. One of the

advantages of the sour gas shift reaction using sulfided, cobalt molybdenum

catalysts is that they operate at much lower temperatures (250–350uC) than

conventional HTS, iron oxide- chromium oxide catalysts (350–450uC).

Table 6: Typical process conditions for cobalt-molybdenum catalyst-based sourgas shift reactor.

Bed 1 Bed 2 Bed 3

Inlet Feed Composition (mole %)CO 46 16 3.1CO2 6.9 26 34.2H2 47 57.9 62.6CH4 0.1 0.1 0.1Sulfur 0.25 — —

Inlet steam/gas, molar ratio 0.96 0.7 0.61Pressure, bar 35 34 33Inlet Temperature, uC 266 288 278Outlet Temperature, uC 411 367 292Space Velocity, h21 2940 2220 1785Outlet CO, mole % 16 3.1 1


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Therefore, the water gas shift equilibrium is favored resulting in lower outlet

CO concentrations. The sour gas shift catalysts also need much less steam for

the same or even higher CO conversion since the possibility of metal formation

(and accompanying methanation and Fischer-Tropsch reactions) from these

sulfided Co-Mo catalysts is remote. However, these catalysts operate at lower

space velocities and, hence, need about 20% more catalyst than the

corresponding iron oxide- chromium oxide HTS catalysts. Additionally, they

also need sulfur in the syngas to be, and remain, in the active sulfided state.

They are used mainly for production of syngas from coal and heavy oil

gasification.

Addition of alkali to these sulfided catalysts promotes their WGS activity

(67, 68). It has also been reported (69, 70) that Co-Mo-based catalysts

promoted by Ti improve the WGS activity of the former in the presence of

sulfur compounds. The sulfided Co-Mo catalysts are not affected by poisons,

like NH3 or HCN when they are present in low concentrations (below about

0.5%). Phenol is a catalyst poison but the rate of deactivation is relatively low

at low concentrations of phenol. Phenol poisoning is reversible and the

catalyst can be regenerated with steam - air regeneration. A high benzene

concentration (above 10%) tends to decrease the catalyst’s activity. Chloride is

a major poison for these catalysts. Even at a 1–2 ppb level, chlorides have an

adverse effect on catalyst performance. The effect of chlorides is cumulative

and catalyst regeneration will not restore catalyst activity.

Mellor et al. (72) reported novel Co-MnO and CoCr2O4 catalysts tolerant to

sulfur up to levels of 220 ppm under WGS reaction conditions. However, in

coal-derived process gases containing between 0.25 and 0.3 mol% sulfur, and

at a reaction temperature of 400uC, the Co-MnO catalyst deactivated rapidly

and irreversibly with formation of bulk Co9S8 and a surface manganese sulfide

species. The CoCr2O4 catalyst deactivated only partially under similar

conditions. Bulk sulfiding of the CoCr2O4 catalyst to CoCr2S4 occurred at

550uC and this catalyst gave near equilibrium CO conversions in the WGS

reaction. A pre-sulfided cobalt chromium catalyst demonstrated typical sulfur

dependent mechanistic characteristics, with a maximum activity above 400uC(72). It may be noted that these sulfided Co-MnO and CoCr2O4 catalysts are

active in the WGS reaction only at temperatures considerably higher than the

sulfided cobalt moly catalysts. They are, hence, under a thermodynamic

handicap, vis-a-vis the latter regarding CO conversion.

6. Pt GROUP METAL-BASED WGS CATALYSTS

Even though the high WGS activity of the platinum group of metals was

known for many decades, their high price precluded their adoption in


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commercial practice. The need for compact catalyst beds in automobile

applications of fuel cells provided an impetus for intense research in this field

during the last decade. Based on the experience from earlier studies of

automotive exhaust catalysts, gold, platinum and other metal –onpartially

reducible metal oxide supports have been the frontrunners in this area (73–

115). It should, however, be noted that auto exhaust catalysis operates under

an oxidative atmosphere above 400uC while the WGS fuel processing

environment is a reducing atmosphere at temperatures between 180–450uCand wherein the partial pressures of H2 and CO2 are much greater. In

addition, noble metals exhibit lower activity in WGS reactions below about

250uC which limits the CO exit levels to about 0.5–1.0 wt%. Of the many

catalysts that have been studied, precious metals (mainly Pt, Rh, Ru, Au, and

Pd) deposited on partially reducible oxides (ceria, zirconia, titania, iron oxides,

and mixed oxides of ceria, like ceria- zirconia) have been the most

investigated. These catalysts are quite active in the 250–400uC range. Pre-

reduction of these catalysts is not required and they can be safely exposed to

air during cool down or start-up without significant loss of performance, a

crucial requirement of fuel processor catalysts. The reaction rate for these

catalysts is close to zero order for CO and, hence, advantageous in driving the

reaction to equilibrium with minimal volume as compared to conventional Cu-

ZnO, where the order for CO is close to one (101). A large number of different

formulations, combining precious metals with partially reducible oxides, have

been proposed as promising catalysts in the literature for the WGS reaction.

Some typical examples are : Au-Fe2O3(73, 74), Au-CeO2(74, 75), Au-TiO2(76),

Ru-ZrO2(77), Rh-CeO2(68) , Pt- CeO2 (74, 78–81), Pt-ZrO2(82)] ,Pt-TiO2(83)

,Pt-Fe2O3(85), and Pd-CeO2(86, 87). Some non-noble metal-based catalysts,

with partially reducible metal oxide supports, have also been reported [Cu-

CeO2 (88, 89), Ag-TiO2 (76), Cu-TiO2 (76), Cu-ZrO2 (90), Cu-Fe2O3 (91).

Grenoble et al. (53) and Panagiotopoulou and Kondaridis (92) had shown that

the precious metal-based catalysts are bifunctional; both the metal and

support have a significant influence on the overall performance. Ceria and

ceria- zirconia have been explored extensively as supports for LTS catalysts in

the last decade. The incorporation of Zr improves the thermal stability

(against sintering), oxygen storage capacity and WGS activity of the ceria

crystallites. The bulk structure of zirconia-doped ceria is well known. Several

different tetragonal phases of varying degrees of stability can be formed

depending upon the Zr doping level, preparation technique, crystallite size,

shape and thermal history. The metastable ‘‘t’’ phase is commonly formed

when Zr doping ranges from 15–28 wt%. It has a tetragonal oxygen ion

sublattice and a cubic Ce/Zr fluorite sublattice. Another metastable t’ phase,

that can form between 28–63 wt% Zr doping, is tetragonal (P42/nmc space

group) on both sublattices having a c/a lattice parameter ratio greater than


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unity. Fluorite structures, such as CeO2, commonly occur as crystallites that

maximize the most stable {111} surface face.

Apart from the noble metals and Au, other transition metals, such as

cobalt and nickel have also been investigated as WGS catalysts. They,

however, cause methanation of CO to CH4 under typical WGS reaction

conditions, especially, below 350uC [Eq. (17)].

A meaningful comparison and rating of all the reported catalysts is

difficult since the various authors had prepared their catalysts by different

methods, with different catalyst precursors, and had evaluated those using

different compositions of the feedstocks and at different reaction conditions.

In an attempt to bring some order in the picture, Thinon et al. (93) have

recently screened about 20 metal-on-oxide catalysts for the WGS reaction

under identical conditions using a model reformate as the reaction mixture.

They used commercial high-throughput equipment consisting of 16 parallel

reactors set-up to compare the activity and selectivity of these bifunctional

WGS catalysts. The catalysts were prepared by impregnation of the supports

with a solution of the corresponding metal precursors. The supports used were

commercial metal oxide powders with surface areas between 30–80 m2/g

except Fe2O3 (7 m2/g). The impregnated material were dried and calcined at

400uC. The feed stream to simulate a typical reformate consisted of 10% CO,

10% CO2, 20% H2O, 30% H2 and 30% Ar. Additional runs were also made with

a feed of 10% CO and 20% H2O diluted in Ar to investigate the forward

reaction in the WGS equilibrium. The catalytic activity was evaluated at 1 bar.

Catalysts based on Pt, Au, Cu, Rh, Pd, and Ru supported on ceria, alumina,

zirconia, Fe2O3 and TiO2 were evaluated. The Rh and Ru- based catalysts

were found to promote methanation reactions. The salient features of

their results are shown in Table 7. In this Table, catalytic activity has been

Table 7: Apparent activation energies and catalytic activities (at 300uC) of Pt, Auand Cu based catalysts ([81).

Catalyst Ea(kJ/mol) Activity (mmol/kg cats)

0.9%Pt/CeO2-Al2O3 70 271.5%Pt/ZrO2 58 202%Pt/CeO2 65 151.9%Pt/TiO2 23 391.5%Pt/Fe2O3 44 61.7% Pd/CeO2 43 85%Au/CeO2 9 271.5%Au/TiO2 29 125%Au/Fe2O3 21 121.5%Au/ZrO2 15 122.1%Cu/CeO2 43 168.9%Cu/CeO2 49 189.1%Au/Fe2O3 23 13


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defined as:

Activity mmol=kgcat:sð Þ~FCO| XCO=Wcatð Þ, ð32Þ

where FCO (mol/s) is the molar flow rate of CO, XCO is the fractional conversionand Wcat (kg) is the weight of the catalyst. Pt/TiO2 and Pt/CeO2-Al2O3 are themost active catalysts at 300uC. It must be pointed out that the inhibitingeffects of the products on the reaction rates are neglected in calculating thevalues in Table 7. Hydrogen and carbon dioxide have, generally, a negativeeffect on the activity and they can also be the reactants for the methanationreactions. The Pt-based catalysts show the highest values for the activationenergies, Cu- based catalysts intermediate values and Au low values. The lowactivation energy observed for the gold catalyst should make it attractive atlow temperatures, especially in combination with Pt/TiO2, provided poisoningby the products, H2O and CO2, is not significant. These conclusions, however,have to be validated by experiments (a) at higher pressures (10–40 bars) and(b) for longer periods of time before application in industry.

One of the drawbacks of TiO2 (vis-a-vis ceria) as a support for Pt in this

reaction is the higher temperatures needed to partially reduce the former. In

an attempt to address this problem, Gonzalez et al. (94) were able to improve

the low temperature activity of Pt –TiO2 catalysts by incorporating ceria in the

support. Pt supported on ceria - modified TiO2 catalyst showed better thermal

stability and lower temperature reducibility compared to TiO2 and a higher

WGS activity than titania or ceria supports (Fig. 22). The catalytic activity of

Figure 22: CO conversion for the WGS reaction on supported Pt catalysts: (m) Pt/TiO2, (&)Pt/Ce-TiO2 ($) Pt/CeO2 (reference). Reaction conditions: total pressure 1 atm, GHSV 5 21200Lh21 kgcat21, feed gas composition (mol%): H2 28%, CH4 0.1%, CO 4.4%, CO2 8.7%, N2 29.2%,H2O 29.6%. Dotted line shows thermodynamic equilibrium limit (94).


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Pt-TiO2 (as well as those of Pd- and Ir-TiO2) was also improved by the addition

of Re by Sato et al. (95) (Figure 23). Catalytic activities were evaluated by

them, in a closed gas circulation system, as an initial H2 formation rate in 10

Torr of CO and 10 Torr of H2O at 50–200uC. Among the Pt, Pd and Ir catalysts,

Pt-TiO2 was the most active catalyst lending further support to Thinon et al.’s

conclusions (81). Two important features were observed: (a) possible formation

of bimetallic surface clusters with Re in the case of Pt and Pd; and (b)

anchoring and ‘spacing’ of metal nanoparticles by highly dispersed Re over

TiO2 in the case of Ir. Panagiotopoulou et al. (96) studied the influence of the

source of the TiO2 support in Pt-TiO2 catalysts and found that the WGS

activity depended strongly also on the phase composition and particle size of

the TiO2 support; the activity increased with increasing reducibility of TiO2.

Both TPR and Raman spectroscopy data indicated that the titania could be

reduced by H2 or CO at temperatures as low as 150uC. Based on their results,

the authors suggested that the titania surface undergoes successive reduction

and oxidation by adsorbed CO and water, respectively, thereby cycling

between TiO22x and TiO2. Sato et al. (97), from CO adsorption, X-ray

photoelectron spectroscopy, in situ IR spectroscopy of adsorbed CO molecules

and catalytic studies of the Pt-Re-TiO2 system, observed that a bimetallic Pt-

Re alloy is formed under reaction conditions and that an additional surface

compound is formed between Pt and Re during the WGS reaction. It is not a

Figure 23: Influence of Re content on the H2 formation rates of WGS reaction over 2 wt% Pt-Re/TiO2 (100uC), 1 wt% Pd – Re/TiO2 (200uC) and 1 wt% Ir-Re/TiO2 (100uC) (95).


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mixture of Pt and ReOx. This surface compound, probably, accounts for the

greater activity of Pt-Re-TiO2. The binding energy of the Pt electrons is lower

in Pt-Re-TiO2 than in Pt-TiO2 suggesting that the Pt crystallites are slightly

negatively charged. CO is adsorbed more strongly on Pt in Pt-Re-TiO2 than on

Pt-TiO2. This is understandable since a more negative Pt will transfer

electrons more easily to the antibonding orbitals of CO, thereby stabilizing CO

in the adsorbed state. IR spectra of CO on Pt-TiO2 reveals only linearly

adsorbed CO. On Pt-Re-TiO2, bridged CO as well as formate ions are seen

additionally indicating that CO is more activated on Pt-Re-TiO2 than on Pt-

TiO2. A more activated CO is more likely to undergo further conversion to

CO2. Hence, the Pt-Re-TiO2 is more active.

When long-term catalytic runs were carried out over the promising Pt-

TiO2 catalysts, Azzam et al. (98) found that even though their Pt-TiO2 was a

very active and selective catalyst for the WGS reaction, they deactivated with

time on stream (Fig. 24). Catalyst deactivation during the WGS reaction was

also a problem with Pt- and Pd- ceria catalysts (86–87, 101). Wang et al. (86)

investigated the mechanism responsible for the irreversible deactivation of

ceria- supported precious metals for the WGS reaction through accelerated

aging tests. They showed that deactivation of Pd- ceria occurs more rapidly at

400uC than 250uC when operating with an integral reactor in 25 Torr each of

CO and H2O. By heating a fresh catalyst in H2, H2O, CO or CO2, it was

discovered that deactivation occurs due to the presence of CO. Similar

conclusions were also reached by Ruettinger et al. (86). Measurements of

Figure 24: WGS CO conversion for Pt/TiO2 with time on stream at 300uC. After 22 h tests, thecatalyst was subjected to the following treatment for 1 h each: (a) O2 at 450uC, (b) H2 at300uC, (c) N2 at 300uC, then tested in WGS. Testing conditions: PCO5 60mbar, PH2O5150mbar,P52 bar, and GHSV5410,000 h21, mcat 551 mg (98).


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metal dispersion by CO adsorption and by X-ray diffraction show (86) that

deactivation on Pd- ceria and Pt - ceria catalysts was due to loss of noble metal

surface area. Pd dispersion values, for example, decreased from 23% (fresh

catalyst) to 3% after a 10 hr treatment in CO at 400uC causing the CO

conversion to decrease from 25% (fresh catalyst) to 3% after treatment for only

2 hr in CO at 400uC. The corresponding dispersion values after similar

treatments in H2O and CO2 were 25 and 16, respectively. The CO conversions

were also not significantly decreased by similar treatments in H2, H2O or CO2

at 400uC. Finally, water gas shift rates on a series of Pd- Ceria catalysts with

ceria crystallite sizes ranging from 7.2 to 40 nm and Pd loadings of either 1 or

6 wt% demonstrated that the rates were strictly proportional to the CO

adsorption capacity and, hence, Pd surface area (Fig. 25). Later, using in-situ

FTIR spectroscopy in the OCO stretching region, Gorte et al. (106) observed

strongly – held, carbonate-like species on the surface, formed from CO. These

were postulated to be the major cause of catalyst deactivation. The authors,

however, do not show the C-H stretching region to confirm the presence/

absence of the formate. Since (a) high temperature treatments in H2 did not

reduce catalytic activity, and (b) high temperature oxidation also did not

restore the activity of deactivated catalysts, they ruled out the over-reduction

of ceria as a contributory factor to the deactivation of the catalyst. Just

because there is a coverage of a surface by a species does not mean that the

Figure 25: Differential water gas shift rates as a function of CO adsorption capacity for aseries of Pd/ceria catalysts in 25 TOrr each of CO and H2O at 250uC. X, 1 wt%Pd/ceria, withceria calcined at 600uC, &, 6 wt%Pd/ceria, with ceria calcined at 600uC, D, 1 wt% Pd/ceria,with ceria precipitated and calcined at 350uC; $, 1 wt% Pd/ceria, with ceria calcined at800uC; m, 1 wt% Pd/ceria, with ceria calcined at 950uC (86–87).


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species in question is causing deactivation. First, one has to see the increase of

that species as a function of time. Secondly, loss of metal support interaction

(e.g., growth of metal particle size) can cause the steady state coverage of the

intermediate (e.g., formate, carbonate) to increase because the metal may

assist in decomposing that intermediate and any loss in the metal’s interaction

with the support (e.g., ceria) would cause the inventory of the intermediate to

build up (since it would be reacting more slowly with loss of the metal’s

interaction). So, while the first point is necessary, the second point explains

why the first point is not sufficient. Zalc et al. (109) observed a strong

dependence of the deactivation rate on the presence of hydrogen in the feed

and suggested that irreversible over-reduction of ceria by hydrogen may,

under certain circumstances, be yet another cause of deactivation of the Pt-

ceria catalysts. It has also been proposed (110) that, yet another potential

cause of deactivation was the growth of ceria crystallites and occlusion of Pt

crystallites, and the consequent decrease of the BET surface area during the

reaction. While all the above-mentioned factors may potentially lead to

catalyst deactivation, it is difficult to extrapolate the validity and relevance of

these conclusions to the WGS reaction in industrial reactors in view of the

widely different methods of catalyst preparation, activation and reaction

conditions used by the various authors. Noble metal- based catalysts

containing a combination of Pt, CeO2 and TiO2 have, recently, been claimed

to be superior WGS catalysts (111). Baidya et al. (112) have compared the

structure, reducibility and catalytic activity of various solid solution oxides

containing cerium, titanium and platinum. Nanocrystalline Ce12xTixO2 (0, x

, 0.4) and Ce12x2yTixPtyO22d (x 5 0.15, y 5 0.01, 0.02) solid solutions,

crystallizing in the fluorite structure, were prepared by a novel, single step

solution combustion method. Their fluorite structure and solid solution

formation were confirmed by XRD Rietveld calculations. Temperature

programmed reduction and XPS study of Ce12x TixO2 (x 5 0.00–0.04) showed

complete reduction of Ti4+ to Ti3+ and reduction of , 20% of Ce4+ to Ce3+ state,

compared to 8% Ce4+ to Ce3+ reduction in the case of pure CeO2, below 675uC.

The insertion of both Pt and Ti ions in the ceria lattice enhanced the

reducibility of CeO2. Ce0.84Ti0.15 Pt0.01O22d crystallized with a fluorite

structure and Pt was ionically substituted with 2+ and 4+ oxidation states.

The amount of hydrogen adsorbed at 30uC over Ce0.84Ti0.15Pt0.01O22d was two

orders of magnitude larger than that over pure 8 nm Pt metal crystallites. CO

and hydrocarbon oxidation activities were also much higher over the Pt-Ti-

Ceria sample, Ce12x2yTixPtyO2 (x 5 0.15, y 5 0.01, 0.02), compared to the Pt-

Ceria sample, Ce12x PtxO2 (x 5 0.01, 0.02). Synergistic involvement of the

Pt2+/Pt0 and Ti4+/Ti3+ redox couples in addition to Ce4+/Ce3+ were held

responsible for the higher reducibility and catalytic activity in the oxidation of

CO (Figure 26). As may be seen from the figure, the Ce0.84Ti0.15Pt0.01O2-d has a


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much lower light-off temperature with T50 5 170uC compared to T50 5 260uCfor the Ce0.99Pt0.01O22d sample indicating that the incorporation of Ti in Pt-

ceria has enhanced the CO oxidation activity probably by increasing the

oxygen ion vacancy concentration and the consequent increase in the oxygen

storage capacity of the material. Parallel TPR and XPS measurements also

confirmed the greater reducibility of the Pt-titania-ceria samples. XPS data

also indicated the presence of Pt in the ionic (Pt2+) state. In view of the known

deactivation of both the Pt-TiO2 and Pt-CeO2 samples during prolonged WGS

reactions, it will be interesting to study the long term stability of the Pt-TiO2-

CeO2 solid solution catalyst in the WGS reaction.

The long-term stability of ceria-based catalysts for WGS operation in fuel

cell applications was investigated by Zalc et al. (109) who prepared a variety of

Pt-ceria WGS catalysts and tested them in the range 250–450uC under feed

and reaction conditions typical of a reformer outlet. They observed first order

deactivation. Virtually identical deactivation rates were found for all the Pt-

ceria catalysts tested. Significantly lower deactivation rates were observed

when hydrogen was not present in the feed. Attempts to rejuvenate the

catalyst by heating under steam and under air were unsuccessful (109).

Catalyst deactivation is still a major obstacle in the commercialization of WGS

catalysts for fuel processing. The goal for most programs is 40,000 hours of

catalyst life. This is an ambitious goal of about 4.5 years of continuous

operation. There are indications, from studies in the industry that addition of

other rare earth elements like lanthanum, praseodymium, etc., to the ceria-

zirconia support can reduce, to some extent, the agglomeration of ceria

crystallites and, the consequent deactivation of Pt-Ceria-Zirconia catalysts.

Another potential cause of catalyst deactivation in the case of Pt group-based

Figure 26: (a)CO oxidation over Ce0.85Ti0.15O2, and Pt, Ti substituted oxides. CO52 vol%,O252 vol%, Flow rate 5 100sccm, GHSV543,000h21, W525mg (112).


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catalysts is the relatively larger formation of hydrocarbons, including

methane, over these catalysts compared to the conventional Cu-ZnO-Al2O3

catalysts. Once formed, these hydrocarbons can undergo further reactions,

like dehydrogenations/ hydrogenolysis/hydrogenation (over Pt), oligomerisa-

tion, carbon formation etc. This problem will be more significant at lower

temperatures? The known higher Fischer-Tropsch activity of the Pt group

metals (compared to copper) for the synthesis of hydrocarbons from CO and

H2, at 200–350uC, is a handicap in the WGS reaction. We may, perhaps, have

to reduce the Fischer-Tropsch activity of the noble metal component to the

level of copper without, however, sacrificing its greater catalytic activity in

WGS reactions. Incorporation of the Pt in the lattice sites of the partially

reducible cerium oxide (as in the work of Baidya et al. (99)) and preserving the

Pt in the ionic state under the WGS reaction conditions, may be one potential

solution since ionic Pt, while active in redox reactions (109), is not known to

possess Fischer –Tropsch activity.

7. Au-BASED WGS CATALYSTS

The low WGS activity of Pt-, Rh- and Pd-based catalysts below 250uC had led

to increased interest in more active catalysts to take advantage of the

favorable thermodynamics at these temperatures. In the past 10 years,

supported gold catalysts with remarkably high activity for the WGS reaction

have been discovered (116, 117). Gold catalysts can offer some advantages in

the range of 180–250uC where the Pt group metals are insufficiently active

(118–143). They are, also, not pyrophoric if exposed to air and require no

exceptional pre-treatment before use. Figure 27 illustrates the high activity of

gold when compared to the Pt and the Cu-Zn-based catalysts. First developed

as a low temperature catalyst for the preferential oxidation of carbon

monoxide (in a mixture of CO and H2) by Haruta et al. (118), it was soon

recognized that the catalytic activity was high only when the particle size of

gold was very small, of the order of 1–5 nanometers (119). Extension of the

studies to low temperature WGS over Au/ a-Fe2O3 (120) and Au-Fe2O3 - MOx

(121) showed that the catalysts were active at temperatures as low as 160uC.

Again, the activity was associated with highly dispersed gold (about 2 nm

particles) (120). The dissociative adsorption of water on the nano gold particles

followed by spillover of hydroxyl groups onto adjacent ferric oxide sites,

involving the redox couple Fe3+/ Fe2+, was postulated. Promotion of Au/Fe2O3

by Ru increased the WGS rate threefold at 120uC (122). Gold nano particles

supported on other supports including TiO2 (123–125), and ceria (126–128)

were also found to lead to catalysts active at temperatures below 200uC. Based

on their studies of Au-TiO2, Andreeva et al.(120) postulated the existence of


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gold in an ionic form at the interface between the Au and the TiO2 phases,

probably as Au ions inserted in the surface regions of the TiO2 lattice. In order

to estimate the relative contributions of the metallic and ‘‘ionic’’ gold to

catalytic activity, Fu et al. (137) measured the rates of the WGS reaction after

leaching out the metallic Au from Au-ceria with NaCN. The rates and

apparent activation energies were the same, before and after leaching with

NaCN, highlighting the importance of the fraction of the gold (presumably

ionic) that was not leached out by the NaCN treatment. The amount of such

‘‘ionic’’ gold inserted in ceria was found to increase with decreasing crystallite

size of ceria. Large crystallites of ceria did not retain any gold. Incorporation of

gold also increased the stability of the ceria microcrystallites. Catalytic

activity in the WGS reaction was also reasonably stable. When this idea was

extended and gold ions were stabilized in the framework of an ionic lattice, as

in Au2Sr5O8 or La2Au0.5O4 (128), not only was the sintering reduced and the

thermal stability of the catalyst increased but the catalytic activity was also

enhanced. It should be mentioned here that recent studies indicate that ionic

gold is unlikely to be present, in the steady state reducing conditions during

WGS reaction, especially, at higher pressures (129). In-situ time-resolved X-

ray diffraction and X-ray absorption spectroscopy were used by Rodriguez et

al. (129) to monitor the behavior of nanostructured Au-CeO2 catalysts under

the WGS reaction. Above 250uC, a complete AuOx ) Au transformation was

observed with high catalytic activity. Photoemission results for the oxidation

and reduction of Au nanoparticles supported on rough ceria films or a CeO2

(111) single crystal corroborated that cationic Aud+ species cannot be the key

Figure 27: CO conversion over supported Au and Pt catalysts (116).


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sites responsible for the WGS activity at high temperatures. They suggested

that the active sites in Au - ceria catalysts involve pure gold nanoparticles in

contact with O vacancies on the ceria. The role of cationic Au3+ and nonionic

Au0 species in the LTS reaction over Au-ceria catalysts was also studied by

Karpenko et al. (130–131) by comparing the reaction behavior of a cyanide-

leached catalyst with that of non-leached catalysts. Using rate measurements

as well as in situ spectroscopic and structure-sensitive techniques, they found

that, based on the Au mass balance, cyanide leaching removed all the Au

except for ionic Au3+ species, and that leaching resulted in pronounced decay

of the catalyst mass- normalized activity to 1–25% of that of a non-leached

catalyst. The extent of the activity loss strongly depended on the post- leaching

treatment of the leached catalyst. Both the catalyst pretreatment after the

leaching and, in particular, the WGS reaction resulted in considerable

reformation of Au0 aggregates and metallic Au0 nanoparticles as indicated

by Au(4f) signals at 85.8 ev(Au3+), 84.0 – 84.6 ev (up-shifted signal of small Au0

aggregates), and 84.0 ev (metallic Au0). Hence, they concluded (130) that (a)

Au0 species, including both small aggregates and metallic nanoparticles

contribute predominantly to the WGS activity, and (b) cationic gold has a

negligible contribution to the WGS activity in the steady state. Au ions are,

expectedly reduced to Au0 atoms in the reducing atmosphere during the WGS

reaction. Combining TEM, XRD, XPS, DRIFTS and activity studies, they

concluded (131), further, that for reaction up to 200uC, catalyst deactivation

was dominated by the formation of stable adsorbed monodentate carbonate

species. The influence of other effects, such as catalyst reduction/ oxidation

were less significant.

Au-CeO2 and Au-CeO2-Al2O3 catalysts were also investigated by Andreeva

et al. (132) who compared samples, prepared by a mechanochemical

activation, with those prepared by a conventional coprecipitation; the former

were more active. This was attributed to the smaller size of the Au and ceria

crystallites in the former. The main role of alumina was that of a textural

promoter in stabilizing the Au and ceria crystallites against agglomeration

during the WGS reaction and, thereby, maintaining a high catalytic activity in

the steady state. The addition of alumina to ceria results in smaller ceria

crystallites and, consequently, an increase in the number of oxygen vacancies

and oxygen storage capacity of ceria, as estimated from temperature-

programmed reduction experiments. A correlation was found between WGS

activity and the oxygen storage capacity of the samples (132).

In an attempt to gain insights into the reactivity of supported Au

nanoparticles, Janssens et al. (133) applied density – functional calculations,

adsorption studies of CO and oxygen on single crystal surfaces and WGS

activity measurements on well-characterised, supported gold particles. They

attributed the increasing activity of supported Au catalysts, with decreasing


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Au particle size, to the increasing number of low - coordinated Au atoms

present in such small particles. Their DFT calculations indicate that

adsorption of CO and oxygen on the densely–packed surfaces (which expose

Au atoms with high coordination numbers, like 8) is generally difficult or

thermodynamically not possible. On the other hand, adsorption was favored

on Au atoms with a lower coordination number. The effect of the Au

coordination number on the adsorption strength of CO and oxygen was found

to be larger than other electronic effects or strain and was, therefore, a crucial

parameter for the catalytic activity. The smaller particle size and support

effects influence the catalytic activity only indirectly through their influence

on exposing a larger number of low-coordinated Au atoms. Among such atoms,

the Au atoms located at the corners of Au crystallites (and, hence, with the

lowest coordination numbers) were the most reactive. During the synthesis of

the various supported Au catalysts, the properties of the support surface (i.e.,

quality and number of nucleation sites) influence the size, dispersion and

morphology of the Au nanoparticles, and, thereby, the concentration of active,

low coordinated sites. Moreover, during catalytic operation, the metal-support

interface energy, which is influenced largely by the support, has a significant

influence on the stability of the particles. A large interface (metal – support)

energy probably can retard the sintering of the Au nanoparticles. Figure 28

(133) shows that there is a clear relation between the adsorption energy of CO

(and oxygen) and the coordination number of the Au atoms to which these

molecules are attached. The lower the coordination number of Au, the stronger

Figure 28: Correlation between the binding energies for CO, O2, and O atoms on Au andthe coordination number of the Au atoms. The solid blue dots indicate experimentallydetermined values for CO adsorption energy on steps, edges and the Au (110)-(162) surface(133).


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the Au-CO bond. The coordination number effect on the adsorption energies

(Fig. 28) (133) can, in turn, be related to the changes in the surface electronic

structure. The low- coordinated Au atoms have high-lying metal d states,

which are in a better position to interact with the adsorbate valence state than

the low- lying states of the high coordination number Au sites of the close-

packed structure (133). This is one of the main reasons why the low-

coordinated transition metal atoms on surfaces are, generally, more active in

catalytic reactions. The trend that the CO adsorption strength on Au increases

with decreasing Au coordination number is also reflected in temperature-

programmed desorption spectra of CO on Au single crystal surfaces and well-

defined nanoparticles (Fig. 29). Janssens et al. (133) attributed the desorption

around 2103 to 283uC (in Fig. 29) to CO adsorbed on defect or corner sites, the

desorption around 2 123uC to CO adsorption on the (110)-(162) surface, and

Figure 29: Temperature programmed desorption of CO on various Au samples (133).


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the desorption around 2163 to 2143uC to CO adsorbed on the edge of the

nanoparticles, or step sites on the single crystal surfaces. For particles with a

given shape, the number of corner atoms per gold particle is independent of

the particle diameter. Hence, for a given amount of gold in the catalyst, the

larger the number of metal particles (i.e., larger the dispersion), the larger will

be the number of corner atoms and, hence, the catalytic activity. This

explanation (133) is fundamentally different from the quantum size effects

conventionally invoked and which ascribes the higher catalytic activity of

small Au particles to changes in electronic structure as the particle size

decreases. Though quantum size effects are important for very small particles

containing only a few atoms (134), they seem not to be necessary to explain the

catalytic effect for supported Au particles larger than about 1.5 nm. It may be

noted that Bond and Thompson had also pointed out earlier that significant

chemisorption of molecules like CO occurs only when an adequate number of

low-coordination surface Au atoms are present (135).

One of the major drawbacks of the gold catalyst is catalytic deactivation

during use. There are two potential causes of deactivation of Au catalysts

(135). The first is Au particle growth, giving larger, but less active particles,

and the second is the formation of unreactive species formed during the WGS

reaction and physical blocking of those sites at which participation of the

support is essential for high activity. Such species include carbonates,

bicarbonates, formates etc. In attempts to prolong the catalytic activity of

supported gold catalysts (for CO oxidation), Moreau and Bond (136) have

recently found that inclusion of Fe(OH)3 or lanthanum oxide during the

preparation of Au catalysts supported on ceria and zirconia, gave better

activity and much improved stability with time-on-stream. This effect was

linked to the ability of the FeOx phase to provide hydroxyl groups, stable at

the reaction temperatures, that are needed for the catalytic action and to form

anion vacancies (by replacement of a tetra- by a trivalent metal cation) at

which O2 or H2O molecules can chemisorb. The effect is similar to those seen

when La3+ or, Fe3+ are dispersed in the ceria lattice (136).

The relatively high activity of gold catalysts has been challenged. Jacobs

et al. (138) reported that a 5%Pt- ceria catalyst was much more active than a

5% Au-ceria catalyst. They attributed their distinctive results, essentially, to

(a) the higher content of Pt in their catalysts, (b) the complete reduction of

platinum oxide, and (c) the ‘‘careful activation’’ of their samples. Differences in

the details of catalyst preparation and determination of the dispersion of noble

metals on partially reducible oxides (see Section 10), adopted by the various

researchers, perhaps, explain such differences. In addition, differences in the

composition of the feedstocks used to evaluate the catalytic activities of the

Au- and Pt-based catalysts and the reaction conditions will also influence the

conclusions. The two metals, Pt and Au, respond differently (116) to changes


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in the concentrations of the reactants. The power law dependencies, on the

concentrations of the reactants and products, are different for Pt and Au (116).

Thus, under certain conditions, the order with respect to carbon monoxide is

negative for Pt but positive for Au. The rate for Au, for example, was given by:

R~k CO½ �0:7 H2O½ �0:6 CO2½ �{0:3H2½ �{0:9: ð33Þ

The positive order with respect to CO reflects the weak adsorption of CO on Au

in contrast to Pt where CO is much more strongly adsorbed, at least, up to

200–250uC, accounting for the negative order with respect to CO on Pt. In

addition to catalytic activity, the stability of the catalyst during prolonged use

under various process conditions is also of major importance. Here, Pt-based

catalysts are quite rugged and have a distinct advantage over Au-based

catalysts. The performance of the latter is more sensitive to conditions of

storage and operation. In view of the importance of resolving this issue (the

relative superiority of Pt- and Au- based catalysts) in the design of fuel

processors for fuel cells, a quantitative comparison of their kinetic behavior

using industrial feedstock and under identical, but, realistic conditions, is

desirable.

Can the performance of gold-based catalysts be improved? Two approaches

have been taken in the last few years: (a) Improving the metal function by

combining Au with another metal (like Pt) to form bimetallic catalysts, and (b)

incorporating promoters in the ceria support. Juan et al. (139) have reported

that when Pt, Pd, W, or Ni is added to Au – ceria, there is a synergistic effect

and the resultant bimetallic catalysts are more active than Au-Ceria or Pt-

Ceria. The catalysts were tested in the temperature range 150–500uC, with a

H2O/ CO ratio of 13.5 and GHSV of 52,000 h21 (Fig. 30). The Au-Pt -Ceria

clearly displayed a much higher activity compared to Au-Ceria at the same

Figure 30: CO conversion over Au-M bimetallic promoters on CeO2 (139).


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temperature. The WGS activities over these samples were ranked as: Au-Pt .

Pt . Au-Pd . Au-W . Pd . Au-Ni . Au. For a quantitative estimate of the

synergy between the two metallic functions (Au and Pt), the atomic loadings

as well as the dispersion of the two metals must be kept the same.

It has to be noted that the catalytic activities (in Fig. 30) of the bimetallic

gold and other metal- supported catalysts are expressed in terms of CO

fractional conversion without using their normalized specific activities

(activity per metal site) since the interactions between different metals over

these bimetallic gold catalysts is not, sufficiently, clear to differentiate

between CO adsorption on Au or Pt or both. The total number of metal sites

per gram of the catalyst (in mmoles per g), evaluated via CO chemisorption

with the assumption of 1:1 CO to metal site ratio, were 2.7(Au), 3.8(Pt) and 2.5

(Au-Pt), all of them supported on ceria (139). Among all the catalysts, a

3 wt%(Au-Pt)- CeO2 displayed the best catalytic activity in the WGS reaction.

Very interestingly, the WGS activity was strongly correlated with the surface

reducibility data from temperature-programmed reduction experiments

(Figure 31). The reducibility of the catalysts, in turn, depended on the

modified local electronic band structure of the promoted ceria. CeO2 shows two

distinct reduction peaks, one at 440uC (assigned to reduction of surface

oxygen) and another at 800uC (reduction of bulk ceria oxygen).The incorpora-

tion of gold and the metallic promoters in the ceria catalyst facilitated the

reduction of surface oxygen at lower temperatures while the reduction of bulk

oxygen remained unchanged. Among the different promoters studied by them,

the Au-Pt-Ceria combination was more effective than Pt, Pd, or Au alone on

ceria in giving higher WGS activity at lower temperatures. The temperature of

the surface oxygen reduction peaks, in Au-Pt, was 120uC. The corresponding

temperatures in the case of Pt and Pd supported on ceria were 130uC and

135uC, respectively. It may be noted that Fu et al. (140) and Andreeva et al.

(141) had also reported similar low temperatures (around 150uC) in theTPR

profile for reduction, in hydrogen, of their Au - ceria prepared by a deposition –

precipitation technique. There is also additional support from the literature

(140, 142) that Au facilitates reduction of surface oxygen at temperatures

lower than even noble metals. The ranked order (139) of the lowering in

temperature (Fig. 31) of the first reduction peak (from surface oxygen loss

from ceria) was Au-Pt . Pt . Au-Pd . Au-Ca . Au-W . Pd . Au-Ni . Au.

This order matched, closely, the relative order of the WGS catalytic activity of

these samples.

Ceria is a n-type semiconductor whose electronic band structure can be

modified by promoters. From the UV diffuse reflectance spectra of the

samples, a clear, blue shift of the absorption edge (O2p R Ce 4f) of the ceria

upon doping with Au or Au-Pd was observed. The degree of band gap widening

was found to relate to different bimetallic promoters (139). The order of


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bandgap widening was remarkably similar to both the orders of surface

oxygen reducibility and WGS activity suggesting an electron transfer

mechanism at the interface between ceria and the metallic components

facilitating the redox transformations occurring in ceria (139):

2Mz2 O½ �u2 Mzz2½ �zO2z2e, and ð34Þ

Figure 31: TPR profiles of CeO2 (a) monometallic and gold bimetallic doped CeO2 samples:Au-Pt/CeO2 (b); Au-Pd/CeO2 (c); Pd/CeO2 (d); Pt/CeO2 (e); Au-W/CeO2 (f); Au-Ca /CeO2

(g); Au-Ni/CeO2 (h) and Au/CeO2 (i) (139).


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2ez2Ce4zu2 Ce3z, ð35Þ

where M represents a metal, like Pt, Pd, Au etc., [ ] an oxygen anion vacancy

and [O] denotes an oxygen anion on the surface.

One of the problems in comparing, quantitatively, the specific WGS

reaction rate data over precious metal- reducible oxide catalysts from different

laboratories is the absence of a reference method for determining the

dispersion of the metal on the support. In the case of a metal dispersed over

a non-reducible oxide, like Pt-silica or Pt-alumina, the experimental

procedures (temperature of reduction, temperature of H2 or CO adsorption,

etc.) and the stoichiometry of the chemisorption for determination of

reproducible and accurate metal dispersion values are accepted and have

been standardized. An important difficulty originates when the support can

also adsorb (and, even react with) the probe molecules, H2 or CO in quantities

comparable or even more than the metal itself. Such is the case for redox

supports such as ceria, titania or ceria-zirconia on which hydrogen spillover

processes (from the metal to the oxide) occur easily in the presence of a

metallic phase, like Pt or Au. Perrichon et al. (144) determined the Pt

dispersion by chemisorption of H2 and CO in a series of Pt-ceria-zirconia

catalysts covering the full range of composition between ceria and zirconia

using volumetric techniques and FTIR spectroscopy. Using IR spectroscopy of

adsorbed CO to distinguish the CO adsorption on the Pt surface from that on

the ceria-zirconia support allowed them to validate a protocol of hydrogen

chemisorption for measuring the metal dispersion. A first method is based on

the CO adsorption isotherm analysis, using IR spectroscopy as the detection

tool. Apart from the quantitative analysis of the adsorbed/desorbed gas phase,

this method also gives information about the coordination mode of the CO

molecule on the metal particle, linear or bridged. The second hydrogen

chemisorption method is based on the use of a double isotherm of hydrogen

adsorption at 278uC , this low temperature being required to suppress the

hydrogen spillover from the metal to the ceria- zirconia support. The

irreversible adsorption of H2, measured either at saturation or by extrapola-

tion to zero pressure, leads to the most reliable metal dispersion values which

can be independently confirmed by FTIR spectroscopy of adsorbed CO. Pt

dispersion, measured by this method, was always higher on the mixed oxides,

ceria- zirconia, than on the pure ceria or zirconia supports (144).

8. MONOLITH-COATED WGS CATALYSTS FOR FUEL CELLS

Current fuel cells use hydrogen, produced by reforming (steam or auto-

thermal) and partial oxidation of natural gas or liquid fuel, to generate


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electricity. The WGS reaction is a critical step in reducing the CO

concentration in such H2 feed streams, especially in low temperature fuel

cells which are not tolerant to CO in concentrations more than 50– 100 ppm

(Fig. 32). The design of fuel processors for stationary fuel cells is less

constrained by the need for compaction and fast response as it is for

automotive applications. Compared to conventional industrial WGS plants for

the generation of hydrogen, however, a reduction in reformer and water gas

catalytic reactor sizes by over two orders of magnitude is necessary before fuel

cells can compete techno-economically with other modes of electricity

generation in automobile applications.

Among the various fuel cells, the Proton Exchange Membrane Fuel Cells

(PEMFC) offer great promise as an alternative to traditional fuel combustion

for generation of electricity for mobile and stationary applications. This H2

must have a CO concentration lower than 50 ppm. In a typical fuel processor

for a PEMFC, the hydrocarbon undergoes reforming by the steam reforming

(SR), autothermal reforming (ATR) or catalytic partial oxidation (CPO). The

reformate then undergoes a series of reactions with the goal of reducing the

concentration of CO and increasing the concentration of H2. The first is the

water gas shift reaction which reduces the concentration of CO in the

reformate from about 10% to less than 1% while increasing the hydrogen

concentration. Further CO clean-up methods, such as preferential CO

Figure 32: Water gas shift in a Fuel Processor for fuel cells (8–9).


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oxidation or selective CO methanation are needed to reduce CO levels to below

50 ppm.

Most of the papers on precious Pt- group or gold catalysts, described in the

earlier sections, report results with powder catalysts or those in the form of

extrudates or tablets. However, in many fuel cell applications, due to

requirements of very high space velocities (to reduce the volume of the

catalyst bed), low pressure drops and mechanical strength, the use of monolith

catalysts is almost mandatory. In response to this requirement, during the

last decade, many publications and patents (especially from industrial

research laboratories) have appeared that describe results with noble metal-

based WGS catalysts washcoated on ceramic or metallic monoliths (113–115).

The requirements of WGS catalysts for vehicular fuel cell applications are

quite different from those needed in NH3 or H2 plants (Table 8). The

development of robust WGS catalysts that can operate in such demanding

conditions is critical in the development of hydrogen generators for fuel cells.

Furthermore, much more active catalysts are sought to make the fuel

processor as compact as possible. The challenge is formidable to achieve such

high catalytic activity at low temperatures (111, 145–148). It is also desirable

to replace the two, HTS and LTS, reactors operating in the 350–450uC and

200–300uC , respectively, by a single, medium temperature, shift reactor in

the 200–350uC range. The activity of the Pt-group catalysts is inadequate

below 250uC. Since CO concentrations as high as1% are tolerated by recent

preferential oxidation (PROX) catalysts used in mobile fuel processors without

sacrificing too much efficiency, this enables one to run monolithic WGS

catalysts at temperatures as high as 300–350uC to reduce the residual CO

content to about 1% in the reformate gas. Even for stationary applications,

this concept of replacing the two HTS and LTS reactors by a single shift

reactor is appealing because of the immense volume/ weight savings and the

Table 8: WGS catalyst requirements for mobile and stationary applications (7).

WGS catalyst attribute Mobile application Stationary application

Volume reduction Critical, ,0.11kW-1 Not as constrainedWeight reduction Critical, ,0.11kgkW-1 Not as constrainedCost Critical, ,$1kW-1 Not as criticalRapid response Critical, , 15 secs Load followingNonpyrophoric Important Eliminate purgingAttrition resistance Critical No constraintSelectivity Critical ImportantNo reduction required Critical ImportantOxidation tolerant Critical ImportantCondensation tolerant Important ImportantPoison tolerant Desired DesiredPressure drop Important Important


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ruggedness of the monolith catalysts. While the efficiency of the monolith

catalysts is slightly lower (about 2–5%), the catalyst bed volume is about 90%

smaller than a comparable system with a particulate catalyst. At the present

stage of development of the Pt-group metal-based, single stage shift reactor, a

two-stage PROX or selective methanation catalyst may be necessary to reduce

the CO content of the reformate to , 10 ppm. The discovery of PEMFC anodes

that can tolerate higher amounts of CO in the H2 stream (Pt-Ru instead of Pt)

may improve the situation further. Additionally, replacement of the low

temperature (100uC) polyvinyl styrene – based electrolyte membrane by the

high temperature (200uC) polybenzimidazole - based membranes in the

PEMFC (enabling the fuel cell operation at 200uC) can also lead to a greater

tolerance of CO by the fuel cell Pt anodes since the poisoning by strongly-held

CO is less at higher temperatures.

8.1. Preparation of Monolith-Coated WGS CatalystsMonolith-coated WGS catalysts are comprised of essentially three

elements: (a) the honeycomb monolith made of cordierite or a metal, (b) the

active metal (the metals Pt, Pd, Rh, Au or their mixtures), and (c) the support

metal oxide (ceria, zirconia, lanthana, titania, alumina or their mixtures)

powder (the ‘washcoat’). The high geometric surface area of honeycomb

monoliths combined with their good mechanical strength and low pressure

drop make them particularly attractive for vehicular applications. The

performance and durability of the finished catalyst depends significantly on

the quality of the washcoat. It is extremely important that the washcoating

process produces a reproducible, uniform, layer of the washcoat. Apart from

the active metals and the support metal oxide, there are several other

materials which can act as additive, binder or adhesive to the washcoat slurry

which is deposited on the monolith prior to impregnation of the noble metal.

Acetic acid is added to most washcoating slurries as a peptizing and dispersing

agent to maintain an adequately low viscosity of the washcoating solution.

Major steps in slurry preparation are particle size reduction of the support

metal oxide powder and addition of appropriate acid/sol to adjust the pH,

viscosity and homogeneity of the support metal oxide slurry in water. Size

reduction of powder particles in water down to a few microns to achieve well

dispersed, homogeneous aqueous slurry can be accomplished by ball milling.

The particle size distribution of the washcoat affects the mechanical strength

of the finished washcoat and its adhesion to the monolith, as well as the

rheological properties of the slurry during the washcoating process. In the

next step, the materials are dispersed in an acidic medium in a tank with a

high-shear mixer. The solids content in the slurry is typically 30–50%wt. After

prolonged mixing, the slurry suspension becomes a stable colloidal system.


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The amount of washcoat that can be deposited on the monolith depends on the

properties of the monolith and the slurry. The catalyst samples are then

prepared by dipping the cordierite monolith in this slurry until the desired

loading is reached. The monolith is then air-knifed to remove excess slurry

followed by drying at 110uC and calcining at 550uC. The monolith is then

impregnated, with an aqueous solution containing the precious metal complex,

dried and calcined. The physicochemical properties of the washcoated monolith

catalyst, like chemical composition, BET surface area, pore volume, metal

dispersion and washcoat adhesion are then measured. When ceria-zirconia is

used as the support oxide for the platinum group metals, their content, in g/liter

volume of the final monolith catalyst, is between 200–500. The noble metal

content is between 2–10 g/liter of the monolith. The costs of the noble metal and

the rare earth oxides constitute a significant part of the cost of the fuel processor

in a fuel cell and efforts are in progress to reduce them further.

The ceramic honeycomb monolith, which is generally used, has some

disadvantages. It has a non-uniform flow distribution due to unidirectional

channels and a closed structure between channels and slow diffusion rate of

reactants to the catalyst surface due to low turbulence in the channels.

Further, when the catalyst particles are washcoated into the channels, the

catalyst particles are not uniformly deposited and are mainly deposited at

corners of square-shaped channels in the monolith, thus leading to lower

catalytic activity. In addition, the low thermal conductivity of the ceramic

material is also a disadvantage in dissipating away the heat generated in

exothermic reactions like the WGS reaction. Metallic monoliths have high

thermal conductivity, larger geometric surface area per unit volume, easy

fabrication and have uniformly deposited catalyst particles. The gas flux flows

in the channel direction as well as in the direction perpendicular to the

channels. Thus, a turbulent flow of the reactant gases results, leading to high

mass transfer rates. Consequently, the required reactor volumes are

decreased. The metallic monolith can be made of a refractory metal like

stainless steel or other iron-based corrosion resistant alloys (e.g., iron-

chromium alloy). They are typically fabricated from such materials by placing

a flat and corrugated metal sheet, one over the other, and rolling the stacked

sheets into a tubular configuration about an axis parallel to the configuration,

to provide a cylindrical –shaped body having a plurality of fine, parallel gas

flow passages, which can range from 200 to 1200 per square inch of face area

compared to about 400 for the cordierite monolith.

8.2. Catalytic PropertiesThe catalytic properties of the monolith catalyst are, broadly, similar to

those of the corresponding powder catalysts described in earlier sections. Only


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a few examples will be given here to illustrate the advantages of the monolith

catalyst. Figure 33 compares the catalytic activity of Pt- CeO2, Pt-ZrO2, Pt-Re-

CeO2, Pt-CeO2-ZrO2 and Pt-Re-CeO2-ZrO2. The Pt-Re-CeO2-ZrO2 catalyst

reaches equilibrium conversion levels around 275uC. The high space velocity

of operation (GHSV 5 20,000 h21) should especially be noted. The advantages

of the noble metal-washcoated, monolith catalysts are apparent from these

results. A conventional Cu-ZnO-Al2O3 LTS catalyst does not give equilibrium

conversions at such high space velocities. Particulate (extrudate or sphere)

catalysts will give rise to a very high pressure drop at these space velocities.

As anticipated, the Pt-Re combination is superior to the Pt-alone composition

and ceria-zirconia is a superior support to ceria or zirconia alone. The

influence of monolith geometry and external geometric surface area on CO

conversion is shown in Fig. 34. The 600 and 400 cpsi (cells per square inch)

monoliths have the same CO conversion while the catalyst with 225 cpsi has a

lower activity suggesting that mass transfer from the bulk fluid to the catalyst

surface, and, not the surface reaction, is controlling the rate of the reaction in

monoliths with lower geometric surface areas (like the 225 cpsi monolith). As

expected, catalytic activity decreases at high space velocities (Fig. 35). It may,

however, be noted that even at such high space velocities the catalyst has,

still, a fairly good catalytic activity. The amount of ceria-zirconia washcoated

per unit volume of the monolith (keeping the Pt content constant) varies,

usually, between 200–500 g/L of monolith. Ceria-zirconia amounts beyond

500 g/liter of the monolith do not increase the catalytic activity. The superior

Figure 33: CO conversion over supported Pt, Re and PtRe catalysts.


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Figure 35: Influence of gas velocities on CO conversion over monolith catalysts.

Figure 34: Influence of monolith geometry (CPSI) on CO conversion.


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activity of TiO2- based supports, illustrated earlier (Section 6), is also

reproduced in the case of TiO2 supports washcoated on monoliths (Fig. 36).

Pt-Re-TiO2 has a higher activity than Pt-Re-ceria-zirconia. It must be

mentioned that the catalytic activities of the monoliths shown in Figs. 33–36

above are initial activities. Like their powder and particulate analogs (Fig. 37),

Figure 36: CO conversion over Pt-Re-CeO2-ZrO2 and Pt-Re-TiO2 (Initial activity).

Figure 37: Deactivation of 1% Pt/CeO2($) and 1% Pt/MgCeO2(#) at 300uC with time;(6%CO, 16%H2, 1.6%CO2, 60%H2O, and 0.4% CH4)(v/v); (168).


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the monolith catalysts also undergo deactivation on prolonged use.

Deactivation of Pt group metal–ceria catalysts with time-on-stream in realistic

reformate gas streams have been reported (149, 150) and attributed to various

reasons, such as metal-induced over- reduction of ceria, precious metal

sintering after high temperature reaction aging of Pd and Pt- CeO2 catalysts

(86, 87) and carbonate formation (106). Deactivation due to carbonate

formation has also been reported for Pt-CeO2 (151) and Au-CeO2 (137).

Apart from their high cost, some of the technical drawbacks of noble metal

catalysts in WGS applications at low temperatures include (a) lower catalytic

activity (compared to Cu and Au) below 250uC, (b) formation of CH4 (up to 1%)

below 300uC (152), and (c) formation of strongly–held formates and carbonate

species on the surface, which, eventually cause catalyst deactivation. The

known Fischer-Tropsch activity of the noble metals in the 200 – 300uC range

(153) is, perhaps, relevant in accounting for their methanation activity under

WGS conditions. In 2007, researchers from the Honda Motor Company

reported (152) results of combinatorial catalysis for over 250,000 materials

and claimed that catalysts containing a combination of (a) one noble metal like

Pt or Rh, (b) one group 11 metal like Cu, Ag, Au, and (c) one partially reducible

oxide like ceria, zirconia, titania, lanthana,vanadia or their mixtures, form

superior WGS catalysts active and stable at low temperatures. While most of

the other elements of the above combination were well-known for their

importance in WGS catalysis, the inclusion of vanadia as a promoting support

for the WGS reaction is interesting and may open future possibilities.

Suppression of methanation activity of monolith catalysts during WGS

reactions at low temperatures by inclusion of basic oxides, like ZnO, MgO,

CaO, SrO, and BaO, in the catalyst support, has been claimed by Farrauto’s

group (154). Inclusion of any of these basic oxides in the catalyst formulation

has been claimed to reduce the methane content in the outlet of the WGS

reactor to less than 5 ppm.

A comparison of the catalytic activity, in the WGS reaction, of Pt-CeO2-

Al2O3 in the pellet form vis-a-vis the metal platelets-washcoated formulation

has been published (155). The authors prepared their Pt-CeO2-Al2O3 catalyst

by a sol-gel method, washcoated it in the micro channels of stainless steel

platelets and evaluated them for catalytic activity in the WGS reaction.

Microstructured metal platelets offer excellent temperature control due to

their good thermal conductivity and small dimensions. Moreover, the use of

thin washcoat layers of catalysts eliminates the intraparticle diffusion

limitations that occur for fast reactions. The superior catalytic activity of

the platelets compared to the pellet samples (Fig. 38) was attributed to the

diffusion limitations inside the pellet samples. The conversion over the pellet

samples above 290uC is lower than that of the platelet samples indicating that

there might be diffusion limitations inside the pellets, as the pellet size


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(250 mm) is substantially larger than the equivalent particle diameter of the

coated, catalyst layer (37 mm) inside the micro channel. To verify this

hypothesis, a simulation was performed that took into account explicitly the

diffusion of matter inside the pellets. The simulation was based on a value of

the average particle diameter of 250 mm, a tortuosity factor of 5 and the mean

pore diameter and pellet porosity as calculated from the N2 adsorption data.

Figure 38 compares the experimental data and the simulation. It can be

observed that at 260uC, the powder and the platelet indeed give similar initial

rates. Above this temperature, the rates (for the pellet) are lower due to

diffusion limitations inside pores, until the thermodynamic equilibrium is

reached, at which point, the rate for the reverse water gas shift is close to the

forward WGS rate and, hence, the overall rate is lower. They concluded that,

catalysts deposited on micro structured platelets lead to a better utilization of

the Pt metal.

9. SURFACE STRUCTURES, ACTIVE SITES AND REACTIONMECHANISMS

The mechanism of the WGS reaction has been studied thoroughly for many

decades. Even though there is some consensus on the redox mechanism

prevailing over the iron-chromia catalysts at high temperatures, there is

considerable uncertainty about the operative mechanism at low temperatures,

over the Cu-ZnO and precious metal- partially reducible metal oxide catalysts.

Figure 38: Comparison of CO conversion over pellets (&) and metallic platelet (m)substrates. symbols: experimental data, solid line: model calculations; (155).


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Newsome (4) and Kochloffl (6, 7) have reviewed the literature in this field upto

1975 and 1996, respectively. The mechanism over the Au based catalysts has

been reviewed by Burch (116). The WGS reaction involves the removal of an

oxygen atom from the H2O molecule to liberate H2 and addition of the oxygen

to CO to form CO2. The dissociation of H2O can occur on the metal, the support

or both. Similarly, the CO can react with the oxygen-containing species (H2O,

OH or O) either from the gas phase, the adsorbed state or the surface lattice.

All the mechanisms that have been proposed for the WGS reaction can be,

broadly, divided into two categories: (a) those that involve a rate –

determining step in which a molecule of H2O or CO, from the gas phase,

reacts with a surface species (oxygen vacancy or a surface oxygen atom) as

exemplified in Eqs. 21 and 22 or (b) those that involve a rate – determining

step in which the reaction takes place between two adsorbed species (the

Langmuir-Hinshelwood mechanism), illustrated, for example, in Eqs. 23–27.

The first is exemplified by the redox mechanism proposed long ago, in 1949, by

Temkin (17–19) and developed further during the past decades. The multistep

mechanism (Eqs. 23–27) proposed by Oki et al., in 1973 (20–22), as well as the

formate mechanism fall in the second category (the L-H mechanism). The

redox mechanism for the HTS reaction was supported by the results of

Boreskov et al., (156–157) who established that the rates of reduction and

oxidation of an iron oxide-based WGS catalyst were in good agreement with

reaction rates calculated from Eqs. 21 (surface oxidation of the catalyst) and

22 (surface reduction). Additional support for this mechanism was derived

from the results of Tinkle and Dumesic (158) who, from rates of adsorption/

desorption and interconversion of CO and CO2 (using isotope exchange

techniques) over iron oxide – chromium oxide catalysts, concluded that CO/

CO2 interconversion is fast compared with adsorption/desorption of CO and

CO2. Thus, Eq. 25 (the surface reaction between adsorbed CO and O moeities)

is fast and, not the rate determining step. The picture is more complex for the

mechanism of the LTS reaction. This issue is discussed below in more detail.

9.1. High Temperature Shift CatalystsIron oxide can exist in three forms: hematite (Fe2O3), magnetite (Fe3O4)

and wustite (FeO). FeO is unstable below 570uC , when it decomposes to a- Fe

and Fe3O4. Below 570uC, the reduction of Fe2O3 to Fe metal proceeds in two

steps via an Fe3O4 intermediate. The reduction of Fe2O3 to Fe3O4 is

exothermic, whereas further reduction to the metal is endothermic.

Hematite crystallizes in the Al2O3 (corundum) structure with a closely packed

oxygen lattice, with Fe3+ cations occupying octahedral sites. Its structure can

be visualized (159) as being composed of Fe-O3-Fe units (triplets) of closely

packed oxygen atoms with Fe(III) on either side. The Fe(III) atoms in each of


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these Fe-O3-Fe units have opposite spins, being antiferromagnetically coupled

as a result of superexchange interaction through the triad of oxygen atoms.

Substitution of Fe cation sites by other metal (M) ion substituents in the

structure promotes formation of mixed or inverse A (12d)B d [A d[B(22 d)]O4

structures, where d is the degree of inversion. For a normal spinel, AB2O4, d 5

0, whereas for an inverse spinel, B[AB]O4, d 5 1. In magnetite- type

structures, octahedral sites are occupied by 2+ and 3+ ions, whereas

tetrahedral sites are occupied only by 3+ ions. Boreskov et al. (156–157)

had, earlier, demonstrated that the octahedral Fe2+ and Fe3+ ions located in

the magnetite-based structure function as a redox couple, and that magnetite-

based catalysts can be highly effective for the complete dissociation of H2O

into H2 and adsorbed oxygen under HTS reaction conditions. Water

dissociation causes the oxidation of Fe2+ to Fe3+ and liberates H2. The Fe3+

centers may, subsequently, be reduced to Fe2+ by CO or H2, thereby producing

CO2 (or H2O) to complete the reaction loop. Many of these substitutions

improve the thermal and textural stability of the structure while promoting

the reducibility of Fe2O3 to Fe3O4. Inverse and mixed spinel structures readily

undergo rapid electron exchange between the 2+ and 3+ states, thereby

catalyzing the WGS reaction. A detailed investigation of the structural

properties of the magnetite (Fe3O4) - based HTS catalyst system has been

published, recently, by Khan et al. (159). These authors prepared metal-doped,

iron oxide-based catalysts with nominal composition of Fe1.82 M0.18O3, (where

M5 Cr, Mn, Co, Ni, Cu, Zn and Ce) by the coprecipitation of the nitrates.

Dilute ammonia was used to precipitate the hydroxides at pH 5 8.5. The

resulting cake was dried at 80uC and further calcined at 500uC in an inert

environment. The structure of the materials was studied by various structural

techniques and evaluated for their catalytic activity in the WGS reaction at

350 – 550uC, different steam to CO ratios (1, 3.5, and 7) and a GHSV of

60,000 h 21. On activation, the hematite- like Fe1.82M0.18O3 phase transformed

into either an inverse or mixed Fe2.73M0.27O4 magnetite - like spinel phase.

The activity of Cr- and Ce- substituted Fe3O4 materials approached

equilibrium levels at high temperatures. At lower temperatures, the activity

of these magnetite- based catalysts was limited by the dissociation of steam.

Interestingly, they discovered that Ce- substituted Fe3O4 spinels are quite

promising HTS catalysts under steam-rich and high temperature conditions.

Khan et al. also carried out the temperature- programmed reduction in

hydrogen of their various promoted iron catalysts. Each promoter ion

influenced the reduction profile of iron oxide in a unique manner. In the case

of the Fe2O3-Cr2O3 catalyst, the first reduction peak at 225uC corresponded to

the reduction of Cr6+ to Cr3+. Further partial reduction of Cr3+ to Cr2+, which

would be expected at 490uC, was not observed. Reduction of Fe2O3 to Fe3O4

was observed at Tmax 350uC, whereas further reduction to FeO occurred at


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higher temperatures. Adding chromium to Fe2O3 did not improve the

reducibility of hematite to magnetite. Based on XRD, TPR and Mossbauer

studies of the Fe2O3-Cr2O3 spinel and earlier information, Khan et al.

proposed that iron – chromia forms an inverse spinel structure and that Cr3+

replaces equal amounts of Fe2+ and Fe3+ from the octahedral sites, with the

displaced Fe2+, consequently located in tetrahedral sites. In the Fe2O3-CuO

catalyst, the reduction of Cu2+ to Cu+ occurred at 143uC. An interesting

observation was that the addition of Cu to Fe2O3 decreased the reduction

temperature of hematite to magnetite considerably to 190uC, compared to

348uC for the pristine hematite sample. The addition of Cu to iron oxide thus

improved reducibility of Fe3+ to Fe2+ species. On doping with Cu, the mobility

of lattice oxygen and hydroxyl groups increased, due to the greater

electronegativity of Cu (1.9) compared to Fe (1.8), thereby perhaps improving

catalytic activity. Promoting iron oxide with cerium causes a shift in the

reduction temperature peak maxima of both hematite - to - magnetite and

magnetite – to - wustite to lower temperatures. In the Fe2O3-Cr2O3, the ceria

surface shell reduction occurs at 380uC, instead of 485uC as in pure ceria. The

temperature of bulk reduction of ceria, however, was not affected by the

presence of iron. These results are significant in the development of Cr-free

iron oxide-based, HTS catalysts.

The propensity of metallic iron and the lower oxides of iron to be oxidized

by steam and evolve hydrogen at high temperatures are well known.

Additionally, the adsorption and surface concentrations of CO at high

temperatures will also be low. The redox mechanism will thus be favored at

high temperatures and over those catalysts which can dissociate H2O into H2

and O2, a crucial requirement of the redox mechanism. Reviewing this area in

1996, Kochloffl (6, 7) concluded that the WGS reaction at high temperatures

over Fe2O3-Cr2O3 catalysts proceeds, most probably, by the redox mechanism.

9.2. Low Temperature Shift CatalystsFrom a mechanistic viewpoint, the LTS catalysts that have been studied

extensively may be divided into three groups; (a) Cu-ZnO-Al2O3, that is

currently used as the standard, industrial catalyst, (b) the Pt group metals

supported on partially reducible oxides, like ceria, titania, zirconia or their

mixtures, and (c) Au supported on the same, above-mentioned oxides. The

reaction mechanism on oxide-supported Au catalysts may be different from

that on supported platinum metal catalysts, because of (a) the lower

adsorption energy of CO on the Au nanoparticles; and (b) the inactivity of

Au (unlike the Pt- group of metals) for H2O dissociation. Some of the features

of the LTS reaction over Cu-ZnO catalysts that distinguish them from the HTS

reaction over iron oxide – chromia are: (1) the dissociation of H2O to H2 and O


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over copper metal or ZnO, at these low temperatures, is less documented

compared to that on the iron oxide, Fe3O4, at the higher temperatures (120);

(2) the amount of CO adsorbed on metallic copper, even at the low

temperatures, is less than that on the platinum group of metals at similar

temperatures and (3) the WGS rate is proportional to the CO partial pressure

to the first order over the Cu-ZnO compared to the zero order observed over

the Pt-based catalysts. The last feature implies that, reducing the CO

concentration from 1.0% (a typical value in the LTS reactor) to 0.5% requires

twice as much catalyst as reducing it from 2% to 1% thus leading to large

second stage LTS reactors. Hence, major efforts have been made in the last

two decades, in the field of fuel processors, to develop cost-effective catalysts

that are tolerant to oxygen exposure, have robust, high volumetric activities at

low temperatures and whose CO conversion rate is independent of the CO

concentration (zero order) in the range 3–0.3%. As mentioned earlier (Section

6), noble metal- based catalysts, like Pt- ceria, Pt-ceria-zirconia, Pt- titania

and their modified versions, like Pt-Na-ceria, Pt-Na-titania, Pt-Na-ceria-

titania appear promising. Though these catalysts have high initial activities,

they still undergo deactivation at temperatures below 250uC on prolonged use.

Currently, efforts are in progress to find the root cause of deactivation of these

catalysts so that they can be used successfully in fuel cells. It is in this context

that a better understanding of the basic mechanism of the LTS reaction over

these noble metal- based catalysts is of importance. A redox mechanism,

involving the reduction of the catalyst by CO and reoxidation (and, thereby the

regeneration of the catalyst by H2O), similar to the one described above for the

HTS reaction over the iron oxide- chromium oxide catalyst, has also been

proposed for the LTS reaction over the Pt-ceria catalysts. In this mechanism,

the CO abstracts an oxygen (forming CO2) from the ceria lattice at the Pt-ceria

interface. Two Ce4+ ions are reduced to the Ce3+ state in this process. The

resulting reduced ceria lattice is then reoxidised through the dissociation of

the incoming H2O. The O vacancy is, thereby, refilled, formally oxidizing two

Ce3+ to Ce4+ and releasing molecular H2 in the process. If the sites for the

adsorption of CO and H2O are different, there will be no competition between

CO and H2O for adsorption and the zero order rate dependence for CO may be

observed. If the sites for the adsorption of CO or H2O is occupied by the other

reactant (H2O and CO, respectively), or, by strongly adsorbed species (like the

formates, carbonates or carboxylates), the reaction order, may be different.

However, it should be noted that at the lower temperatures characteristic of

the LTS reactions over catalysts like Cu-ZnO, and Pt-ceria (190–250uC), the

ability of steam to reoxidise the partially reducible oxide supports (with or

without the presence of the noble metals) has not been demonstrated so far; In

addition, in the case of fuel cell conditions (for Pt-Ceria), this reoxidation of

Ce3+ to Ce4+ has to occur in the presence of a considerable amount of hydrogen.


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Techniques, like XANES, can, perhaps, be used to resolve this issue. In

addition to the redox mechanism, an associative mechanism involving a

surface formate intermediate has also been proposed.

9.2.1. The ‘Formate Surface Intermediate’ Hypothesis

From studies of the WGS reaction over copper chromite catalysts,

Armstrong and Hilditch, had, already in 1920, proposed (160) that the

reaction involved the adsorption of CO and H2O on the catalyst surface to form

a surface intermediate that subsequently decomposed to CO2 and H2. To

isolate and identify the ‘‘surface intermediate’’, they reacted it with NH3 and

obtained ammonium formate. The formation and decomposition of a formate-

or formic acid – type surface species was speculated to lead to the formation of

the products, H2 and CO2, over copper – based catalysts, like Cu-Cr2O3.

Similarly, using dimethyl sulphate as a methylating agent to trap the surface

formate, Deluzarche et al. (161) obtained dimethyl formate further supporting

the presence of formates on the surface during the WGS reaction. During the

last few decades, extensive research using a variety of techniques has,

conclusively, proven the existence of a formate species at low temperatures

over Cu-ZnO as well as precious metal- reducible metal oxide catalysts (162).

Boreskov and Davydov (163–164) had earlier carried out pioneering IR

spectroscopic studies supporting the associative mechanism over a wide

number of copper- based catalysts including the industrially important Cu-

ZnO. Additional early work supporting the formate associative mechanism for

Cu-ZnO include those of Herwijnen et al.(165–166) who observed the nearly

identical conversion rates for the water gas shift reaction and formate

decomposition. Rhodes et al.(167) have, however, raised some doubts as to

whether these surface formates constitute the only, or, even the main

intermediates in the reaction path of CO and H2O to H2 and CO2 or are merely

spectators formed by a parallel route from the reactants and/ or products

(167). Shido and Iwasawa (107, 108) investigated the WGS reaction over ZnO,

CeO2 and MgO using in-situ FTIR spectroscopy of the surface species. Their

results indicated that surface OH groups, formed by reaction of H2O with

oxygen vacancies on partially reduced CeO2, reacted with CO to form bridged

formates. The bridged formates were converted to bidentate formates above

170uC. This transformation occurred at room temperatures in the presence of

water. About 30% of these adsorbed bidentate formates were, in turn,

decomposed to the final products (H2 and CO2) and adsorbed, unidentate

carbonates. The rest decomposed back to the reactants, CO and H2O. These

transformations of the bidentate formates were also influenced by the

presence of H2O. Coadsorbed H2O also promoted the decomposition of

the unidentate carbonates to CO2. In addition to the unidentate carbonates,


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the presence of surface carboxylates and bidentate carbonates were also

identified by Shido and Iwasawa (107, 108). Binet et al. (168) and Fallah et al.

(169) observed IR bands at , 3650 cm21 on partially reduced (with H2) CeO2

samples and assigned it to Type II bridging OH groups. Jacobs et al. have,

indeed confirmed, using in-situ IR spectroscopy, (170–172) that the formates,

formed on Pt (1%) on ceria, decomposed to CO and OH in the absence of steam

in about 6 min at 300uC , while in the presence of steam, they decomposed

completely and much more rapidly in 8 min even at 140uC to produce H2 and

unidentate carbonates. Based on mechanistic studies, including kinetic

isotope effect (170) and isotope tracer studies (172), they, further, suggested

that the rate-limiting step in the LTS reaction is the cleavage of the C-H bond

of the surface formate. Bridging OH groups and surface formates have also

been identified, earlier, by IR spectroscopy over thoria and zirconia-based

catalysts (173–175). Jacobs et al. have also observed that Pt-thoria (176) and

Pt-zirconia (177) possess much higher WGS activity than the corresponding

oxides without the precious metals and attributed it to the more facile

formation of the Type II bridged OH groups, and the surface formates derived

from them, over the Pt-loaded catalysts. In addition, kinetic isotope effects

similar to those observed in the case of Pt-ceria were also observed for Pt

supported on thoria and zirconia suggesting that the rate-limiting step was

likely to be the C-H bond scission of the formate intermediate in these cases

also. To summarize, in the formate mechanism, a bidentate formate produced

from CO and surface OH groups acts as an intermediate. The bidentate

formate, then, decomposes to gaseous hydrogen and a surface unidentate

carbonate, which further decomposes to gaseous CO2. One of the major

contributions of Jacob et.al., is the discovery that co-adsorbed water plays a

crucial role in the selective decomposition of the formate intermediate to CO2

and H2. The main roles of Pt are (1) to catalyze the reduction of ceria, leading

to the formation of surface, terminal OH groups on ceria and (2) to catalyse the

decomposition of the formate to H2 and CO2. The rate determining step is

likely to be the decomposition of the unidentate formate to CO2 and H2.

If the decomposition of the formates to H2 and CO2 is the rate-limiting

step, then, factors that facilitate the cleavage of the formate C-H bond should

enhance the reaction rate. More specifically, addition of bases, like the alkali

ions, which are known to accelerate the decomposition of formates, should

enhance the WGS rates. Pigos et al. (178) have recently observed that the

incorporation of Na in Pt- zirconia catalysts does, indeed, enhance the WGS

reaction rates. Interestingly, their DRIFTS investigations suggest that

incorporation of Na in Pt- zirconia modifies the electronic properties of the

surface formate and weakens its C-H bond. Two significant features of their

DRIFTS results (178) are (a) The C-H band of the formate species was

shifted to lower wavenumbers from 2880 cm21 (Pt-ZrO2) to 2842 and


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2804 cm2 1 (Pt-Na-ZrO2), respectively, indicating a weakening of the C-H bond

on Na incorporation; and (b) The ratio of the intensity of the bridged to linear

Pt carbonyls increased from 1:5 to 4:5, thus, favoring bridge-bonded CO in the

Pt-Na-ZrO2 catalyst. The influence of catalyst basicity in increasing the

concentration of the bridged carbonyls had, already, been reported by Mojet et

al., who found (179) that, in Pt-SiO2 and Pt-K-L (LTL) zeolites, increasing the

K+ ion content also increased the concentration of the bridged carbonyls. The

faster decomposition of the surface formate over Pt-Na-ZrO2 is illustrated in

Fig. 39. 20% of the initial intensity of the IR band of formate is decreased in

5.2 min for the Pt-ZrO2 and 2.4 min over Pt-Na-ZrO2. CO2 was also seen in the

presence of steam. In addition to the formate, carbonates and carboxylates

were also seen after steaming. Pigos et al. (178) also compared the stability of

the formate species, on Pt-ZrO2 and PtNa-ZrO2, in the absence of steam, by

monitoring their thermal decomposition. It may be recalled that Shido and

Iwasawa (107, 108) had, earlier, observed that in the absence of steam,

formates on metal-ceria catalysts decompose, primarily, in the reverse

direction, back to CO and OH. To follow the thermal decomposition of the

formate in the absence of steam, C-H bond breaking was probed by Pigos et al.

(178) by flowing D2 and monitoring the exchange of the C-H and C-D formate

bands. The areas of the formate C-H bands (at 2880 and 2966 cm21) and their

corresponding C-D bands were quantified and plotted as a function of time

(Fig. 40). The formate H-D exchange rates were very close to the overall

formate decay rate. Moreover, the half-life for H-D exchange for Pt-Na-ZrO2

was approximately half that of the Pt-ZrO2 (Fig. 40) indicating that C-H bond

breaking in the formate is more facile for the Na-promoted catalyst. It may be

observed that in none of the aforementioned studies were the formate

Figure 39: Formate area response to steaming at 130uC for Pt/ZrO2 and PtNa/ZrO2 (178).


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decomposition rate constant and surface coverage determined simultaneously

at the steady state under WGS reaction conditions. Based on DRIFTS analyses

combined with the utilization of isotopic tracers, it had been shown (180) that

formates were less reactive than carbonyl and carbonate species under steady

state conditions whereas the reverse trend was observed during the non-

steady state, desorption experiments.

The reactivity of the species formed at the surface of a Au-Ce(La)O2

catalyst during the WGS reaction, in the steady state, was investigated by

Meunier et al.(60–62) using simultaneous DRIFTS and kinetic analysis. The

exchange of the product CO2 and formate and carbonate surface species were

followed during an isotope exchange of the reactant, CO, using a DRIFTS cell

as the reactor. In independent experiments, the DRIFTS cell yielded identical

reaction rates to that measured in a quartz, plug-flow reactor. The DRIFTS

signal was used to quantify the relative concentrations of the surface species

as well as that of CO2. The analysis of the formate exchange curves between

Figure 40: Formate area to D2:N2 at 225uC for (top) Pt/ZrO2 and (bottom) PtNa/ZrO2. The K-life of formate is indicated for formate decay and formate exchange from H to D. Fasterdecay and exchange rates are observed for PtNa/ZrO2, indicating a higher reactivity forformate C-H bond breaking. Reverse decomposition (178).


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155–220uC suggested the presence of two types of surface formates: ‘‘slow

formates’’ that display an exchange rate constant 10- to 20 fold slower than

that of CO2 and ‘‘fast formates’’ that exchanged on a time scale similar to that

of CO2. Figure 41 compares the molar rate of formate decomposition to that of

CO2 formation per unit mass of catalyst. The specific rate of CO2 formation

was determined from the CO2 concentration in the DRIFTS cell exhaust gas

(by gas chromatography), the sample weight in the DRIFTS crucible and the

flow rate of the reactants. The rate of formate decomposition was calculated

from the DRIFTS data as the sum of the decomposition of the ‘‘fast’’ and ‘‘slow’’

surface formates. The semilog plot shows that the rate of CO2 formation was

more than an order of magnitude (about 60-fold) higher than the rate of

decomposition of the (slow + fast) formates, indicating that the formates,

detected by DRIFTS, cannot be the only reaction intermediates in the

production of CO2.

Thus, while there is sufficient experimental evidence to conclude that (a)

formate-like species are present, under WGS conditions, on the surface of Cu-

ZnO and precious metal-partially reducible, metal oxides; and (b) the

decomposition of these formate species under WGS conditions leads to the

products, CO2 and H2, it is not established that CO2 and H2 are derived only

from the surface formates and not, also, additionally, from other intermedi-

ates, like the carbonates/caboxylates or by a completely different mechanism,

like the redox mechanism, which does not involve any long-lived and

experimentally observable, surface intermediate. This viewpoint is further

supported by the investigations of Gokhale et al. (196–199) and Mhadeshwar

and Vlachos (209–210) (described later, Section 9.2.4).

Figure 41: Rate of CO2 production and rate of formate decomposition over the 0.6 AuCl atthree different temperatures under 2% 12CO + 7%H2O (60–62).


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9.2.2. The Redox Mechanism on Pt-based Catalysts

A redox mechanism is also being advocated (106, 148) for the LTS reaction

on Pt-ceria and other platinum group metals on ceria. According to this

picture, CO adsorbs on transition metal sites and reacts with oxygen from the

ceria, which, in turn, is reoxidised by H2O. In other words, it involves the

reaction of the reactants, CO and H2O, with the surface: CO with the oxide ion

of the ceria (to yield CO2) and H2O with the anion vacancies on ceria

(generating OH groups and, eventually, H). An important role of the metal is

to adsorb/activate CO and create of oxygen vacancies at the metal ceria

interface. In contrast to the formate theory, there is no postulate of a stable,

experimentally observable and kinetically relevant, surface intermediate.

Evidence for this mechanism came, initially, from kinetic studies (181). TPD

studies (182) have demonstrated that oxygen from ceria can react with CO

adsorbed on metals. It has also been established (183) that reduced ceria can

be oxidized by CO2. While the redox mechanism is well-established at high

temperatures in the case of the iron oxide-chromium oxide catalysts, its

applicability to LTS over Cu-ZnO and Pt-ceria catalysts is uncertain and

depends on confirmation of the ability of H2O to reoxidize the partially

reduced support oxide at temperatures below 250uC, especially in the presence

of significant amounts of hydrogen, as is the case for fuel cell applications (see

also sections 9.2 and 10). Such an unambiguous experimental confirmation is

desirable. Another feature of the ceria- based catalysts, namely, that high

temperature calcination lowers, not only the concentration of oxygen

vacancies and loosely – bound surface oxygen atoms, but also their WGS

activity lends additional support to the redox mechanism. Reaction orders on

Pd-ceria were, approximately, zeroth-order in CO, half-order in H2O, inverse-

half order in CO2 and inverse first order in H2 (106). The rate limiting step

was believed to be the dissociation of H2O on the ceria support. Diffuse

reflectance and FTIR spectroscopic measurements on Pd-ceria indicated that

the ceria existed (not surprisingly) in a reduced state under WGS conditions

and is covered by carbonate species that are removed only by reoxidation of

ceria (106). Such surface carbonates were also a cause of catalyst deactivation.

It may be noted, however, that, under their WGS reaction conditions, the Cu-

ZnO catalyst was much more active than all their metal-supported ceria

catalysts.

Azzam et al. (184–185) studied the WGS reaction on catalysts based on

ReO2-TiO2. Results pointed to contributions of an associative formate route

with redox regeneration and two classical redox routes involving TiO2 and

ReOx, respectively. Under their WGS reaction condition, rhenium was

present, at least partly, as ReOx providing an additional redox route for

WGS reaction in which ReOx is reduced by CO generating CO2 and re-oxidized

by H2O forming H2. The reaction between CO adsorbed on Pt and OH groups


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on titania was the rate-determining step. Gold nanoparticles supported on

reducible and non-reducible oxides with comparable gold particle size were

studied in the WGS reaction by Sandoval et al. (186). Not surprisingly, the

activity of Au on reducible oxides was much higher than the one observed on

non-reducible oxides. The optimum calcination temperature was 300uC. For

samples calcined at 300uC and reaction temperatures below 225uC the activity

varied (Fig. 42) as follows: TiO2 . CeO2 . Al2O3 . SiO2. A novel catalyst

consisting of platinum deposited over a cerium-modified titania substrate has

been, recently, reported by Gonzalez et al. (187). They showed better thermal

stability with respect to the bare TiO2 support and higher WGS activity than

those corresponding to individual titania or ceria supports. XPS and TPR

results revealed the intimate contact between Pt and cerium entities in the Pt/

CeO2–TiO2 catalyst that facilitates the reducibility of the support at lower

temperatures. The importance of CO adsorption on ceria in influencing the

rates of the WGS reaction was investigated by Li et al. (188). Au nanoparticles

on monoclinic ZrO2 showed much higher catalytic activity for the low-

temperature water gas shift reaction than those on tetragonal zirconia, mainly

due to the high CO adsorption capacity of monoclinic ZrO2. Formate species

formed by the reaction of adsorbed CO on gold nanoparticle with hydroxyl

groups on ZrO2 were postulated to be the reaction intermediates.

Figure 42: CO conversion over Au nanoparticles supported on TiO2, CeO2, Al2O3, and SiO2

(186).


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9.2.3. Mechanism over Cu- and Au- based Catalysts

One of the important issues in the mechanism of the LTS reaction is the

role of the copper and gold metal nanoparticles supported on ZnO or CeO2. Cu-

CeO2 and Au- CeO2 are two of the promising LTS catalysts. Cu and Au are

present, mostly as nanosized metallic particles, during the WGS reaction.

What is the intrinsic reactivity of these nanoparticles? Can nanoparticles of

Cu and Au catalyze the WGS reaction on their own without the aid of an oxide

support (such as ceria or ZnO)? Results from catalytic studies over bulk Cu

and Au may not be directly applicable to the nanoparticles. On the pure, bulk

metals, the WGS reaction, probably, proceeds by a redox mechanism (63, 173).

The mechanism may, however, be modified by the presence of the support

oxide (especially by a partially reducible one, like ceria) in intimate contact

with the metal nanoparticles and wherein metal-support interactions will be

more important. Rodriguez et al. (189–195) have addressed this issue. They

investigated the WGS reaction on Cu and Au nanoparticles supported on CeO2

(111) and ZnO (0001) surfaces (189–190). Pristine CeO2 (111) and ZnO (0001)

surfaces did not display any catalytic activity under their reaction conditions

(300–375uC, PCO 5 20 Torr, PH2O 5 10 Torr). Significant catalytic activity was

measured when Au or Cu particles (2–4 nm) were deposited (Fig. 43). The

deposition of Cu nanoparticles on ZnO (0001) produced a catalyst that was

clearly more active than the pure extended Cu surfaces. An even better

catalyst was obtained when the nanoparticles were supported on CeO2(111).

They found negligible WGS activity on Au (111) (Fig. 43) or polycrystalline Au.

Figure 43: Amounts of H2 produced during the WGS reaction on 0.5ML of gold or copperdeposited on CeO2 (111) and ZnO (0001). For comparison, the activities of Au (111) and Cu(100) are also included. The catalysts were exposed to a mixture of 20 Torr CO and 10 Torr H2Oat 625 K for 5 minutes in a batch reactor. A reaction time of 2–3 minutes was enough to reacha steady - state regime in the reactor (189–190).


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The Au-ZnO(0001) system displayed catalytic activity that was worse than

that of Cu-ZnO(0001). In contrast, Au-CeO2 had an activity similar to that of

Cu- CeO2 (111). Their XPS data pointed to a lack of oxidation of the metals,

and, a reduction of the ceria support. They also evaluated the importance of

surface intermediates for the pure metal surface Cu (110), and the metal- ceria

catalyst. Importantly, in the case of pure, metallic Cu (110), analysis of the

surface after the WGS reaction showed it to be essentially free of formate and

carbonate species, suggesting that, on the pure metal surfaces, the WGS

reaction proceeds by the redox mechanism. In agreement with others, they

also identified adsorbed formate and carbonate-like species on the metal-ceria

surfaces after the WGS reaction. Using density functional calculations, they

had also investigated (193, 194) the WGS reaction on Cu29 and Au29 clusters

(representative of the metal nanoparticles formed on deposition on ceria or

ZnO supports) and on Cu (100) and Au (100) surfaces (representative of the

surface of the bulk metals). Figure 44 shows the calculated energy changes for

the WGS reaction on a Cu29 cluster. The reaction pathway with the minimum

energy barriers involves the following steps (Eqs. 36–41):

CO gð Þ<CO adsð Þ, ð36Þ

H2O gð Þ<H2O adsð Þ, ð37Þ

H2O adsð Þ?OH adsð ÞzH adsð Þ, ð38Þ

Figure 44: Reaction profile and structure for the WGS reaction on a Cu29 nanoparticle. Thezero energy is taken as the sum of the energies of the bare nanoparticles, gas-phase water,and carbon monoxide. The red bars represent the transition states, and the black barsrepresent reactants, intermediates, or products. Cluster side view: yellow - Cu, red -O, gray-C, white- H (189–190).


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CO adsð ÞzOH adsð Þ?OCOH adsð Þ, ð39Þ

OCOH adsð Þ?CO2 gð ÞzH adsð Þ, and ð40Þ

2H adsð Þ?H2 gð Þ: ð41Þ

The adsorption of CO or H2O on the Cu particles is exothermic. The first and

the most important energy barrier is for the dissociation of water into

adsorbed OH and H (reaction 38). Then, the reaction of OH and CO produces

an OCOH, carboxyl species. The final important energy barrier is for the

decomposition of this OCOH carboxyl intermediate into CO2 gas and adsorbed

H, which eventually yields the H2 gas. The DFT results indicated that a free,

metallic, nanoparticle of copper can catalyze the WGS reaction easily. A

comparison with the corresponding results on the Cu(100) surface of bulk

copper shows that the dissociation of H2O on the surface of bulk copper has a

larger activation energy barrier (1.13 ev vs. 0.94 ev on the nanoparticle) and

that no stable OCOH, carboxyl intermediate, is formed, as a redox mechanism

operates. The presence of corner or edge atoms in Cu29 favors the dissociation

of H2O. The Au29 nanoparticles and the bulk Au (110) surface could not

catalyze (169) the WGS action. Neither surface was able to adsorb and

dissociate water molecules. Figure 45 (189, 190) shows a correlation between

the calculated barrier (y axis) and the calculated energy (x axis) for water

dissociation on Au(100), Cu(100), as well as the ionic and neutral Au29 and

Cu29 particles. All the gold systems are characterized by a large activation

Figure 45: Correlation between the calculated barrier (DEa) and the calculated reactionenergy (DE) for water dissolution on several copper and gold systems (189–190).


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barrier and an endothermic DE value. There is significant improvement in

chemical reactivity in going from Au(100) to Au29, but not enough for

dissociation of the water molecule. These results, of course, cannot explain the

large catalytic activity of Au-ceria (Fig. 43) and highlight the important role of

ceria in the activation of the system. A perfect CeO2 (111) surface does not

dissociate water at low or, even, high temperatures. When O vacancies are

present, however, the H2O molecules dissociate on the partially reduced ceria

surface. Au and Cu particles facilitate the reduction of the ceria surface by the

CO/H2O mixture and, thereby, facilitate the most difficult step in the WGS

reaction, namely, the dissociation of H2O (189, 190).

9.2.4. The Carboxylate Mechanism

Surface carboxylic species had been observed earlier by spectroscopic

techniques on LTS catalysts (107, 108). Mhadeshwar and Vlachos (210) and

Gokhale et al. (196–199) proposed from theoretical calculations, that they are

important, reactive intermediates, which play a central role in the WGS

reaction. Gokhale et al. used periodic, self-consistent, density functional

theory (DFT-GGA) to investigate the WGS reaction mechanism on Cu(111),

the dominant facet of copper crystallites in the Cu-ZnO industrial WGS

catalysts. Their proposal for an alternate WGS reaction mechanism, involving

the oxidation of adsorbed CO by adsorbed OH, to form carboxyl (COOH)

species is compared with the conventional redox mechanism in Table 9. The

crucial difference is that, while in the conventional redox mechanism the

adsorbed CO is oxidized to CO2 by adsorbed O atoms, CO2 is formed by the

decomposition of an adsorbed carboxyl group (formed by the reaction of an

adsorbed CO with an adsorbed OH group) in the new proposal. CO2 may also

be generated by the reaction of the carboxyl with a second adsorbed OH group

(Table 9). They also suggested that although it is possible to form the carboxyl

group, COOH, in a single, elementary reaction step (reaction of CO with OH

Table 9: Redox and carboxyl mechanisms on Cu (111)a (175).

Redox mechanism Carboxyl mechanism

CO + * ) CO* CO + * ) CO*H2O+ * ) H2O* H2O+ * ) H2O*H2O + * ) H* + OH* H2O + * ) H* + OH*OH* + * ) O* + H* CO* + OH* ) COOH* + *OH* + OH* ) H2O* + O* COOH* + * ) CO2* + H*CO* + O* )CO2* + * COOH* + OH* ) CO2* + H2O*CO2* ) CO2 + * CO2* ) CO2 + *H* + H* ) H2 + 2* H* + H* ) H2 + 2*

aSteps in italics highlight differences between the two mechanisms.


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as postulated in the formate mechanism), that is less likely. That is because

OH binds to the surface through its O atom, and CO through its C atom,

whereas the formate binds through its two O atoms, not its C atom. Hence,

two O atoms of the formate will have to bind to the surface forming a bidentate

species, either sequentially (via unidentate formates) or, less likely,

simultaneously. Their calculations suggested that the easiest way to form

the formate, HCOO, is by reacting CO2 with atomic H. The 16 elementary

steps involved in their mechanistic model of the WGS reaction on Cu (111) are

shown in Table 10 and the corresponding reaction network in Figure 46. Using

the DFT-derived parameters as initial estimates for the microkinetic model

parameters, they fitted the 16- step model to the experimental WGS reaction

rate data published, earlier, by Koryabkin et al. (200). As may be seen from

Fig. 47, the agreement is satisfactory. Their model also tested well against the

kinetic data of Herwijnen and Jong on a Cu-ZnO-Al2O3 catalyst (52). Based on

the good ‘‘fit’’ between the calculated and observed data, they suggested that

Cu(111) may be a dominant active site for the WGS reaction on realistic

industrial catalysts. An alternate explanation may be that the WGS reaction

on these catalysts is not structure sensitive, and therefore, the reaction rate is

comparable on different Cu facets (196). From their model calculations they

also predicted that steps 5 (H2O* + * u H* + OH*) and 9 ( CO* + OH* u cis-

COOH * + *), in Table 10, are rate controlling under industrial conditions. In

the absence of CO2 and H2 co-feed, step 5 has a considerably stronger

influence on the overall reaction. To summarize their results on Cu(111): (a) H

abstraction from H2O appears to be the rate-controlling step for the entire

Table 10: Elementary Steps involved n the water gas shift reaction on Cu (111)(175).

Step No Elementary Step

1 CO + * R CO*2 H2 + 2* R 2H*3 H2O + * R H2O*4. CO2 + * R CO2*5 H2O* + * ROH* + H*6 OH* + * R O* + H*7 2OH* R O* + H2O*8 CO* + * R CO2* +*9 CO* + OH * R cis-COOH* + *10 cis-COOH* R COOH*11 COOH* +* R CO2* + H*12 COOH* + OH* R CO2* +H2O*13 CO2* + H* R HCOO* + *14 HCOO* + * R HCOO**15 CO2* + H2O*+ *R HCOO** + OH*16 CO2* + OH* + * R HCOO** + O*


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WGS reaction network, (b) carboxyl(COOH) is a very reactive, but short-lived ,

intermediate, and (c) formate (HCOO) formed, probably, from CO2 and H, is a

spectator species which tends to block active sites, and can reach substantial

surface coverages, particularly at high pressures. This site- blocking by

formate can, also, explain the observed negative WGS reaction order with

respect to CO2 (196–199).

This approach has, recently, been extended by the same group (197) to the

water gas reaction on Pt(111) surface of bulk Pt metal. It should be borne in

Figure 47: Experimental WGSR rates versus rates predicted by the microkinetic model(196–199).

Figure 46: Reaction rate for the water gas reaction. A reaction scheme including both thesurface redox mechanism and the carboxyl mechanism is outlined. The thermochemistry andthe kinetic barriers for all the elementary steps are given in electron volts. For reactionsinvolving bond making, the activation barriers are reported with respect to the adsorbedreactants at infinite separation from each other. The minimum energy pathway for the WGSRis highlighted with green (196–199).


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mind that the influence of the support on the electronic and textural

characteristics of the metal surface OH groups or reaction intermediates

was not explicitly taken into account in these calculations. The results are,

broadly, similar to those described above for copper (196). The contribution of

the surface redox mechanism to the water gas reaction on Pt, involving CO

oxidation by atomic O, is negligible under the range of conditions studied

(250–300uC, 1 bar, feed 5 (CO+H2O+CO2+ H2)). The lowest energy path

involves the formation of the carboxyl (COOH) intermediate, which is

subsequently, decomposed by reaction with OH (COOH + OH R CO2 +H2O). The OH species is, then, regenerated by dissociation of the formed H2O.

When the concentration of the OH groups is limited, the direct decomposition

route (COOH R CO2 + H) dominates. Additional H2O in the feed increases the

OH coverage and makes the OH 2 mediated COOH + OH, low energy

decomposition path more kinetically accessible, thus, enhancing the water gas

reaction rates.

9.2.5. Alkali-doped, Pt-based, LTS Catalysts

One of the key, rate- determining, steps in the formate mechanism is the

C-H bond scission in the surface formate intermediate. Evin et al., have

recently found in accord with the results of Pigos et al. (178), that alkali

doping weakens the C-H bond, as demonstrated by a shift to lower

wavenumbers of the n(CH) vibrational mode, and enhances the LTS reaction

over Pt-Ceria catalysts significantly (201). However, with high alkalinity (,2.5% Na or equimolar amounts of K, Rb, or Cs), a trade- off was observed such

that while the formate became more reactive, the stability of the adsorbed

carbonate species, which arises from the decomposition of the initially- formed

formate intermediate, was found to increase. This was observed by TPD-MS

measurements of the adsorbed CO2 probe molecule. Increasing the amount of

alkali to levels that were too high also led to (a) lower catalyst BET surface

area, (b) the blocking of the Pt surface sites as observed in infrared

measurements, and (c) a shift to higher temperature of the surface shell

reduction step of ceria during TPR. When the alkalinity was too high, the CO

conversion rate during the water-gas shift reaction also decreased in

comparison with the undoped Pt-ceria catalyst. However, at lower levels of

the alkali, the above-mentioned inhibiting factors on the water-gas shift rate

were suppressed such that the weakening of the formate C-H bond could be

utilized to improve the overall turnover efficiency during the water-gas shift

cycle. This was demonstrated at 0.5% Na (or equimolar levels of K) doping

levels. Not only was the formate turnover rate found to increase significantly

during both transient and steady state DRIFT tests, but this effect was

accompanied by a notable increase in the CO conversion rate during low


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temperature water-gas shift. Evin et al., (202) had also observed, using in-situ

DRIFTS spectroscopy, that, on steaming pre-adsorbed formate species, these

species were more reactive in the decomposition step in the catalytic cycle (i.e.,

decomposition of the formate to H2 + CO2) for the Li- and Na- doped catalysts

relative to undoped Pt- ceria (202). For example, the relative formate

decomposition rates were 1.0, 1.2, and 1.41 for the undoped, 0.15% Li- doped

and 0.5% Na- doped Pt-Ceria, catalysts respectively. The CO conversion at

225uC, correspondingly, increased from 12% for the undoped 2% Pt- ceria to

14% for 2%Pt- 0.15%Li-ceria and 24.3% for the 2% Pt-ceria sample doped with

0.5% Na. However, with increasing atomic number over the series of alkali –

doped catalysts, the stability of the carbonate species (another surface species

formed during the WGS reaction) was also found to increase. This was

observed during TPD-MS measurements of the adsorbed CO2 probe molecule

by a systematic increase of a high temperature peak for a fraction of the CO2

desorbed. This result indicates that alkali-doping is an optimization problem-

that is, while improving the decomposition rates of formate species, the

carbonate intermediate stability also increases, making it difficult to liberate

the CO2. An optimal amount of basicity, sufficient to decompose the formate,

but, not enough to stabilize too much the carbonate, is needed. Infrared

spectroscopy results of CO adsorbed on Pt and ceria suggested that the alkali

dopant is located on, and electronically modifies, both the Pt and ceria

components. Alkali doping may, thus, provide a path forward for improving

the WGS rate by means other than resorting to higher noble metal loadings.

It has, of course, been known for a long time that alkali metals promote

the WGS reaction rate. In 1981, Sato and White (203) doped Pt- TiO2 with

NaOH and found an improvement in the photocatalyzed water gas shift rate.

Klier (204) also highlighted the promoting influence of alkali dopants, their

relative efficiency being, Cs. Rb. K.Na, Li. Klier also suggested that the

alkali should be present at concentrations less than a monolayer. Campbell et

al., (205), observed a promotion of the WGS activity of Cu (110) by Cs ions. Cs

1.5–2.0 CO3 was found after the reaction by surface analysis (TDS, XPS, AES) of

the catalyst. In kinetic studies, using a low H2O/CO ratio, they found that on

the optimally Cs- promoted surface, the reactant orders were zero order for

H2O and 0.5 order for CO, suggesting that H2O dissociation was not rate

controlling on the alkali-promoted catalyst. They proposed a redox mechanism

to describe the catalysis of both the clean and Cs- doped surfaces with Cs

playing the role of O mediator among CO2, H2O, and CO, where Cs is,

primarily, in the form of a carbonate. Honda Research Inc., has also claimed

(152) a remarkable improvement in the WGS activity of Pt- ZrO2 catalysts for

fuel processors for use in fuel cell applications by doping the catalysts with

alkali. Among the promising compositions discovered was an important

improvement when Pt- ZrO2 was doped with Na alone or in combination with


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vanadium. In their DRIFTS spectroscopic study of metal- ceria catalysts,

Pigos et al. (178) found that formate C-H stretching bands were strongly

shifted to lower wavenumbers upon CO adsorption indicative of a weakening

of the C-H bond. Further, in transient formate decomposition experiments,

both in the presence and absence of steam, they reported that formates, over

Pt-Na- ZrO2, decomposed at twice the rate of those observed on Pt-ZrO2

without Na. Among the alkali dopants, Na was found to provide the most

benefit.

9.2.6. LTS over Pt Supported on Non-reducible Oxides

What is the mechanism prevailing over catalysts comprising of noble

metals supported on non-reducible oxide supports, like alumina or silica? The

fact that gamma alumina is an ‘‘irreducible’’ oxide at the WGS reaction

conditions will seem to exclude the redox mechanism involving oxygen ion

vacancies on the support, extensively discussed in the literature for partially

reducible metal oxides and supported metal catalysts on such carriers.

Olympiou et al. (206) studied the mechanism of the WGS reaction over

alumina-supported Pt, Pd, and Rh catalysts, using steady state isotopic

transient kinetic analysis (SSITKA) techniques coupled with mass spectro-

metry. In particular, the concentrations (mmol g21) of active intermediate

species found in the carbon-path from CO to the CO2 product (using 13CO),

and in the hydrogen- path from H2O to H2 (using D2O) were determined

(Table 11). It was found that by increasing the reaction temperature from 350

to 500uC, the concentrations of the active species in both the carbon and

hydrogen paths increased significantly. Based on (a) the large concentration of

the active species present in the hydrogen- path (OH/H located on the alumina

support), which was larger than six equivalent monolayers (based on the

exposed platinum metal surface area), (b) the small concentration of OH

groups along the periphery of the metal-support interface, and (c) the

Table 11: Concentration of active ‘‘H-containing’’ (H-pool) and ‘‘C-containing’’(C-pool) surface species at water gas shift conditions (181).

Catalyst T (uC)H-pool

(mmol gcat21) or (h)a

C-pool(mmol gcat

21) or (h)a

0.5 wt% Pt/c-Al2O3 350 350 (28.5) 1.3 (0.1)500 1664 (135.6) 31.7(2.6)

0.5 wt% Pd/c-Al2O3 350 235 (9.8) 0.5 (0.02)500 3194 (132.7) 28.6(1.2)

0.5 wt%Rh/c-Al2O3 350 138 (6.2) 2.4(0.1)500 1093 (49.0) 27.3(1.2)

aCoverage in monolayers of exposed surface metal area.


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significantly smaller concentration (mmol g21) of active species present in the

carbon-path (adsorbed CO on the noble metal and formate species on the

alumina support and/ or at the metal- support interface), the authors

suggested that the diffusion of OH/H species on the alumina support towards

catalytic sites present in the hydrogen pathway may be the slow step in the

reaction mechanism. The OH/H species were considered to be formed by the

dissociation of H2O on the alumina support. The role of the noble metal was (a)

to activate the CO molecule by chemisorption, and (b) to promote formate

decomposition into CO2 and H2 products. There was also a correlation between

catalytic activity and the surface concentration and binding energy of CO on

the noble metals. Among the alumina-supported noble metals, the order of

activity was found to be Pt . Rh . Pd. It may be remarked that these results

lend strong support to the WGS mechanism proposed by Grenoble et al. (53)

for Pt-alumina in 1981. For the formation of the formate entity, the CO

adsorbed on the Pt metal must react with the OH group adsorbed,

presumably, on the alumina. Whether the formation of the formate is the

result of diffusion of CO from the Pt surface to the Al-OH sites or the diffusion

of the – OH groups from alumina towards CO adsorbed along the

circumference of the metal-support interface is not clear. Duprez (207) has

discussed the mechanism of migration of the OH/H species on metal oxide

surfaces with basic and weak Bronsted acidic character (like gamma alumina).

The problem is still unresolved.

A mechanism based on the interaction of CO with Pt and H2O with ceria

and derived from a kinetic study using a microstructured reactor has been,

recently, proposed for the WGS reaction over a Pt-CeO2-Al2O3 catalyst by

Germani and Schuurman (208). The use of a microstructured reactor, rather

than a packed bed reactor, enabled the measurement of the intrinsic kinetics

of the reaction. The reaction rate was almost zero order in CO and was

strongly inhibited by the partial pressure of hydrogen, and, to a lesser extent

by that of CO2. The rate equation that fitted the data best was based on a dual-

site mechanism with a rate-determining step that involved a species adsorbed

on Pt, a species adsorbed on ceria and a free Pt site. Based on this observation,

a reaction mechanism was proposed where CO, adsorbed on Pt, reacts with

water, dissociatively chemisorbed on ceria, to yield a carboxyl species as an

intermediate. This carboxyl species reacts with a second hydroxyl group and

decomposes over a free Pt site into carbon dioxide and hydrogen as shown

below:

COz�uCO�, ð42Þ

H2OzCe-OuHO-Ce-OH, ð43Þ


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CO�zHO-Ce-OHuCOOH�zCe-OH, ð44Þ

COOH�zCe-OHz�u2H�zCe-O-CO2, ð45Þ

2H�uH2z2�, and ð46Þ

Ce-O-CO2uCO2zCe-O: ð47Þ

In the above equations * represents a Pt site. The rate determining step is the

reaction between the carboxyl species, and the second hydroxyl group on the

ceria (Eq. 45). Once an adjacent Pt site becomes free, this carboxyl complex

decomposes into the reaction products. Hydrogen competes with CO for Pt

adsorption sites and, therefore, retards the rate. Similarly, CO adsorbs

strongly on ceria and has a negative influence on the rate. This mechanism

differs from the mechanism proposed by Shido and Iwasawa (107–108) in that

the surface intermediate is postulated to be a carboxylate and not a formate. It

may be mentioned here that the reaction of CO with Type II bridging OH

groups on ceria to form carboxy species has not been confirmed unambigu-

ously by experiment. We may recall that a carboxyl surface intermediate has

also been postulated, more recently, by Gokhale et al. (196–199). A key feature

of the carboxyl mechanism is the conversion of adsorbed CO to adsorbed

COOH. From microkinetic studies of the WGS reaction system, Mhadeshwar

and Vlachos (210) had, earlier, made an important observation that while the

formation of the carboxyl intermediate from H2O, namely, CO* + H2O* )COOH* + H*, can be the rate-determining step (as per their model), competing

paths for CO* oxidation on Pt by OH* (i.e., CO* + HO*) COOH* + *) instead

of H2O*, cannot completely be ruled out as being important (with the H2O*

decomposition being the rate determining step), due to the relatively small

differences in activation energies of these parallel oxidation paths. They have

also provided a very useful and comprehensive compilation of all the

significant rate expressions and reaction orders for the WGS reaction on

different catalysts postulated in the literature up to 2005 [Table 1 of ref. 185].

One of the drawbacks of the Pt group of metals is that they are less active

in the WGS reaction below about 250uC. On the other hand, Cu-ZnO is an

outstanding WGS catalyst in the range of 200–250uC. Some of the latter’s

drawbacks are (a)the necessity to operate at low GHSV values, (b) its complex

and time-consuming activation protocol before use, and (c) its instability on

contact with air. The Pt group metals do not have these disadvantages. In an

attempt to combine the advantages of both the copper and Pt- group metals in


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a single catalyst formulation, Fox et al. (211) have investigated the catalytic

properties of Pd- promoted Cu-Ceria catalysts for the oxygen-assisted, WGS

reaction. It may be noted that the conventional ZnO support has been replaced

by ceria in this formulation. Cu-CeO2, Pd-CeO2, and Cu-Pd-CeO2 catalysts

were prepared and their reduction followed by in-situ XPS to explore the metal

– metal and metal - support interactions in the bimetallic Cu-Pd-CeO2.

Addition of only 1 wt% Pd to 30 wt% Cu-CeO2 greatly enhances the

reducibility of both dispersed CuO as well as the ceria support, presumably

by hydrogen spillover from Pd. In-vacuo reduction (inside the XPS chamber)

up to 400uC results in a continuous growth of metallic copper and Ce3+ surface

species. Support copper, in turn, destabilizes palladium metal (Pdu) with

respect to PdO, this mutual perturbation indicating a strong, intimate

interaction between the Cu-Pd components. The presence of Pd, apparently,

increased the fraction of copper that remains in the metallic state thereby

enhancing its catalytic activity. It may be recalled that the increase of

catalytic activity with metallic copper surface area is well known. Palladium

addition at only 1 wt% significantly improved CO conversion at 180uC,

compared to a monometallic 30 wt% Cu- CeO2 catalyst (Fig. 48). As

anticipated, the Pd-Ceria was the least active compared to Cu-CeO2 and Cu-

Pd-CeO2 at low temperatures. It should be noted that the feedstock used by

Fox et al. (211) (Fig. 48) contained oxygen (2% air) and is not representative of

the conventional feedstock from a steam reformer to the WGS reactor. Reactor

design considerations will be crucial in such a situation to avoid the oxidation

of the hydrogen or CO over the precious metal and the consequent exothermic

temperature rise. In a similar vein, but combining the relative advantages of

two support components, ceria and titania, Gonzalez et al. (212) have recently

Figure 48: CO conversion over 1 wt% Pd, and 30 wt% Cu catalysts and 1 wt% Pd – 30 wt%Cubimetallic catalyst at 180uC. Feed gas: 4% CO, 10% CO2, 26% Ar, 2% air, balance H2. H2O/CO510 (211).


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demonstrated the performance enhancement in the WGS reaction when Pt is

dispersed over a mixed oxide support material containing both ceria and

titania. TiO2 and CeO2 have complementary physical properties which can be

synergized to improve the performance of a catalyst support in the WGS

reaction. For example, the redox properties and thermal stability of titania

can be improved by replacing, partially, Ti4+ ions by the Ce4+ ions in the

titania lattice (213–214). Gonzalez et al.(212) have found that Pt supported on

Ce modified TiO2 catalyst shows better thermal stability (with respect to bare

TiO2 support) and higher WGS activity than those corresponding to individual

titania and ceria supports, indicating a synergistic effect between Pt and the

Ce- modified TiO2 support (Fig. 49). XPS and TPR results revealed the

intimate contact between Pt and cerium entities in the Pt-CeO2-TiO2 catalyst

that facilitates the reducibility of the support at lower temperatures (Fig. 50)

while the Ce-O-Ti interactions decrease the overreduction of TiO2 at high

temperatures. It may be noted that the addition of cerium to TiO2 had also

increased the hydroxyl concentration on the support. This is probably one of

the contributing factors to the greater catalytic activity of the Pt-CeO2-TiO2

compared to Pt-TiO2. This data also underlines the important role of the OH

groups in the mechanism of the LTS reaction.

While discussing the active sites and mechanism of the WGS reaction, it is

instructive to recall two important features of the landmark postulate of Hugh

Taylor in 1926 (215) on active sites over solid catalysts: (a) particular atoms or

groups of atoms on the surface of solids are the active sites responsible for the

Figure 49: CO conversion for the WGS reaction on supported Pt catalysts: (m) Pt/TiO2, (&)Pt/Ce-TiO2, ($) Pt/CeO2 (reference). Reaction conditions: total pressure 1 atm, GHSV521200liter.h21kgcat

21, feed gas composition (mol% ): H2 28%, CH4 0.1%, CO 4.4%, CO2 8.7%, N2

29.2%, H2O 29.6%. Dotted line shows thermodynamic equilibrium limit (212).


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catalytic activity and selectivity and, importantly in the present context, (b)

the identity and concentration of the active sites on a catalyst are dependent,

not only on the procedures adopted during its preparation, but also on the

particular reaction conditions, i.e., the relative concentration of the reactants,

temperature, pressure etc. If these conditions are changed, then, the identity

and concentration of the active sites are also likely to change and,

consequently, the dominant reaction mechanistic path from the reactants

(CO and H2O) to the products (CO2 and H2) will be different and depend on

particular reaction conditions. The CO concentration at the inlet to the WGS

section can vary widely in the range 10–40% (dry basis) depending on the raw

material (natural gas or coal) and the reforming process utilized to generate

the syngas. The H2O concentrations will also vary depending on the type of

reformer (steam, partial oxidation or autothermal reformer) that is utilized

upstream of the WGS reactor. Steam reformers utilize higher H2O/ carbon

molar ratios (3–5) than partial oxidation or autothermal reformers (0.5–2.0).

Consequently, the concentration of H2O at the WGS reactor inlet will be

higher when steam reformers are used. Similarly, the concentration of CO2

will be higher when the syngas is generated in a partial oxidation or

Figure 50: Temperature-programmed reduction-MS profiles of (a) Pt/TiO2 and (b) Pt/Ce-TiO2

calcined catalysts (212).


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autothermal reformer than in a steam reformer. It is normal to expect that the

different concentrations of CO, CO2 and H2O in the feedstocks, from these

various H2 generation system configurations, will influence, differently, the

nature and, especially, the concentration of chemical species present on the

catalyst surface (OH groups, H atoms, anion vacancies etc.). Hence, it should

not be surprising that different WGS mechanisms can prevail on the same

catalyst under different reactants concentration and temperature/pressure

conditions, especially in the WGS reaction that is equilibrium-limited at high

temperatures and kinetically limited at low temperatures.

Burch (116) has published a critical discussion of the relative merits of the

various mechanisms that have been proposed for the LTS reaction over metal

– partially reducible oxide supports. He has also presented a ‘‘Universal

mechanism’’ for the WGS reaction that seeks to integrate the salient features

of the formate and redox mechanisms into a single model that is consistent

with all the experimental observations (Fig. 51). Figure 51a shows the

Figure 51: (a) ‘‘carbonate/carboxylate’’mechanism for the reverse WGS reaction (b)‘‘carbonate/carboxylate’’ mechanism for the WGS reaction. (c) ‘‘universal’’ mechanism forthe WGS reaction (116).


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mechanism for the reverse water gas reaction, CO2 + H2) CO + H2O (RWGS);

Fig. 51b shows the corresponding mechanism for the forward WGS reaction,

CO + H2O) CO2 + H2. In both cases, the importance of the carbonates and/or

carboxylates is emphasized. Figure 51c shows Burch’s ‘‘universal’’ mechanism.

One crucial feature of Burch’s universal mechanism is that, while in the

formate mechanism, postulated originally by Shido and Iwasawa (107–108)

and elaborated upon by Jacobs et al. (202), the formate intermediate is formed

from insertion of CO into an OH bond, both being adsorbed on the support, it is

formed (in Burch’s postulated mechanism) from the addition of an H to a CO,

both being adsorbed on the metal particle (116). However, there is no

unambiguous experimental evidence for this assumption. In fact, in studies on

transient formate decomposition either with the unpromoted catalyst or the

catalyst promoted with different metals (e.g., Pt, Au) and loadings of metal,

once the surface shell of ceria is reduced, the formate concentration from

reaction of CO with the Type II bridging OH groups is high and, for the most

part, have the same intensity over all the catalysts before the H2O is added to

decompose them. If the formate had been anchored on the metal, a variation of

the intensity with the type and concentration of the metal would have been

observed (223). Both modes of formation of the surface formate intermediate

are shown in Figure 51c. It is important to note that a carbonate–like species

is also involved in the reaction path in all the three postulated mechanisms,

including the formate mechanism, wherein the formate formed initially reacts

with H2O to form a carbonate which finally decomposes to yield CO2. The

dominant mechanism will depend (116) on the reaction conditions, specifically

the temperature and the H2O/CO2 ratio. It can change from a redox-type

process to one dominated by surface intermediate species, including formates,

carbonates and carboxylates. We envision three situations (116):

N At high temperatures, where desorption and/ or decomposition ofintermediates, like formate and carbonate species will be very fast, theredox processes would be expected to be important in determining the rateof the reaction. This is particularly valid in the presence of a highconcentration of H2O when the surface is covered, to a significant extent,by OH groups;

N At low temperatures, and, especially in the presence of a substantialamount of CO2, the final carbonate decomposition step in the mechanismwill be the slow, rate-determining step; and,

N At intermediate temperatures, especially in the presence of a largeconcentration of water, and a low concentration of CO2, the formatedecomposition step in the mechanism would be slow and rate- determining.

As depicted in Fig. 51b, the active sites in the WGS reaction are the oxygen

vacancies which dissociate water into OH adsorbed on the support. Metal sites


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also adsorb and activate CO. In the next step, the OH groups, located on the

support oxide, react with the CO adsorbed on the metal particles to form the

surface intermediates, the formates and carbonates. The latter decompose to

CO2. It must, however, be emphasized that the above mechanism involving

long-lived surface intermediates, like formates or carbonates/ carboxylates, is

valid only below, say, 350uC. At higher temperatures, over HTS catalysts, like

Fe2O3- Cr2O3, the direct oxidation and reduction of the catalyst by H2O and

CO, respectively, by Eley-Rideal-type of reactions are well known, and, hence,

the redox mechanism will, probably be the dominant mechanism. These Eley-

Rideal processes are less favored at low temperatures and, hence, the rates of

formation or decomposition of surface intermediates, like formates and

carbonates/ carboxylates, assume critical importance. One further point in

relation to this mechanism is its relative importance in the case of catalysts,

like Pt-Al2O3, wherein a significant loss of the support hydroxyl groups occurs

only above 400uC (216). Surface oxygen vacancies, that play such an important

role in the above-mentioned mechanisms that invoke the formation of stable

surface intermediates located at these oxygen vacancies, are unlikely to be

present in significant concentrations on the alumina surface in the 190–300uCrange, typical of the LTS reaction, especially in the presence of significant

amounts of H2O. It may be noted that an oxygen anion vacancy is the starting

point in all the mechanisms depicted in Figures 51. In such cases, both the

activation of CO and the dissociation of H2O occur, perhaps, on the Pt metal. It

may, however, be relevant to mention here, that, Chenu et al. (177) had

reported the observation of surface defects, oxygen vacancies and type II

bridging OH groups on non-easily-reducible oxides, like MgO and ZrO2 under

reducing conditions similar to those that prevail during the WGS reaction,

when they were promoted with Pt. After activating the catalyst, DRIFTS of

CO adsorption was used to probe the active OH groups via the generation of

formate species. It is important to note that in the absence of H2O, formates

are quite stable, such that their intensity upon CO adsorption gives a good

qualitative indication of the number of active OH group defect sites. While

formates were observed over both the Pt-ZrO2 and Pt-MgO, the band

intensities were lower as compared with Pt- ceria, suggesting a lower

concentration of defect-associated active OH groups on ZrO2 and MgO. The

WGS rates and formate band intensities from CO adsorption (used to probe

the active OH groups) followed the same trend: Pt-Ceria . Pt-monoclinic ZrO2

. Pt-tetragonal ZrO2 . Pt-magnesia. The Pt content was 1%(wt) in all the

catalysts. On lowering the temperatures, Pt-magnesia was inactive below

400uC, while Pt-Ceria was active up to 250uC, under their reaction conditions

(Feed: 3.75 ml/min CO, 125 ml/min H2O, 100 ml/min H2, 10 ml/min N2; 33 ml of

catalyst). Similar results of generating active OH groups on non-easily

reducible oxides, like ZrO2, on promotion with Pt was also reported by Pigos et


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al. (178). They found that the catalytic activity of Pt-ZrO2 was improved,

significantly, when Na was added, either alone or in combination with

vanadium oxide. The results of parallel, DRIFTS spectroscopic experiments

indicated that formate species were more reactive on the Na - promoted

catalysts; The formate C - H stretching bands were shifted to lower

wavenumbers upon CO adsorption. In transient formate decomposition

experiments, both in the presence and absence of steam, the formates over

Pt-Na-ZrO2 decomposed at twice the rate of those observed on Pt-ZrO2 without

Na. This was further supported by steady state WGS rates which confirmed

that the formate species were more reaction rate- limited in DRIFTS for the

Na - promoted catalysts relative to those without Na. These results of Chenu

et al. (177) and Pigos et al. (178) suggest that the active sites for the WGS

reaction, namely, the oxygen anion vacancies and associated OH groups, can

be generated also on otherwise non-reducible oxide supports under reaction

conditions when the catalyst formulation contains elements like the noble

metals.

9.2.7. Catalyst Deactivation

Deactivation during long-term tests has been one of the major drawbacks

of the noble metal- based catalysts with ceria or titania supports. The

facilitating role of bases, like the alkali ions, in the decomposition of formates

had been noted earlier (178). The deactivation of WGS activity along with

strongly held, carbonate surface intermediates had been observed by Gorte et

al. (86–87). At high temperatures (above 350uC) and over the iron oxide-

chromium oxide catalysts, these carbonates decompose more easily and no

deactivation is observed. Do the strongly-held carbonates impede the

reoxidation of the oxide surface or the release of the H2 molecule? The use

of acidic oxides, like those of Nb, Mo, Ta, and W, to enhance the WGS activity

of Pt-Ceria-Zirconia has been reported, recently, by Opalka et al. (217). Aided

by density functional simulations, these authors observed that doping Pt-

ceria-zirconia with acidic, transition metal dopants such as Nb, Mo, Ta and W

oxides increased the oxide surface affinity for water and the turnover rate of

the WGS reaction. The Pt/ Mo-doped –ceria-zirconia, Pt/ Mo0.1Ce 0.7Zr 0.2O2,

for example, was significantly more active (by 15–20% in CO conversion) than

the undoped sample in the temperature range, 200–300uC. The composition of

their feedstock was: 4.9% CO, 10.5% CO2, 33% H2O and 30.3% H2, and the

GHSV was 300,000 h21, simulating the environment in a fuel processor of a

fuel cell. Only initial catalytic activities were reported. Kinetic rate analysis

for the CO conversion yielded reaction orders approaching 0 for CO and 1 for

H2O. They characterised the nature of adsorbed CO and the formate and

carbonate intermediates, formed during the WGS reaction at 200uC over


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catalysts without and with Pt loading, by in-situ infrared spectroscopy. After

oxidizing the nanocrystallites in dry air for 2 h and, then, exposing them to low

CO pressures, only weakly adsorbed CO was detected. When the CO pressures

were increased to 4 bar or higher, small amounts of adsorbed CO2, formates

and carbonates were also observed, indicating that the catalyst was partially

reduced by CO. In the presence of H2O and, under WGS conditions, linearly

adsorbed CO and significant amounts of formates and carbonates were also

observed. The larger, more basic Ce3+ ions formed under reducing conditions

enhance the further reactions of adsorbed CO to form formates and

carbonates. The formates were weakly bonded and could be removed by

outgassing the catalysts in dry nitrogen. The carbonates, on the other hand,

were removed only on oxidation of the catalyst above 270uC. Under the wet,

reducing WGS conditions, on the same oxides with Pt loading, CO linear

adsorption was observed only on the Pt metal (not ceria). Formate and

carbonate formation was observed on the ceria- zirconia oxides. If ligand (CO)

complexation of the oxide surface leading to strongly-held formates and

carbonates is locally specific to the reduced sites, then, CO associative

reactions will compete with or impede the reduction (by CO) or oxidation (by

H2O) of the catalyst and, hence, influence the redox mechanism. Based on

their density functional simulations and IR spectroscopic/kinetic experimental

results, the authors suggested (217) that the associative formation of formates

and carbonates was, indeed, coupled with the bifunctional redox mechanisms

that lead to the reduction of the oxide surface. They observed, further, that the

rate at which O could be removed by CO from the lattice, outstripped the rate

at which H2O could chemisorb and react to replenish the lost O. Hence, the

rate- limiting step, over Pt-ceria and Pt-ceria-zirconia, was not lattice

reduction but rather reoxidation. While the reduction of ceria by CO/H2 is

not in question, its reoxidation by H2O at low temperatures needs to be

verified experimentally by a direct method. To shift the WGS oxidation –

reduction balance towards reoxidation, they added more acidic, less reducible

dopants like Nb, Mo, Ta and W oxides, to make the reoxidation more

favorable. These are strong electron acceptors and are fully oxidized in their

Lewis acid- like oxide phases with generally empty d orbitals (d0 oxides). On

ceria and ceria-zirconia, formate formation was very close in energy to

reoxidation of the reduced surface by H2O. In the presence of acidic transition

metal dopants, however, surface reoxidation was significantly more favorable

than the reaction of adsorbed H2O with CO to form stable formate and

carbonate complexes. For example, while the enthalpy for reoxidation (by

H2O) of the oxide surface of ceria- zirconia was 2134 (eV), it changed to 2340

(eV) on doping with Mo indicating that Mo facilitates the reoxidation of the

surface (217). The transition metal oxide dopants, apparently, shifted the

relative balance of the reaction steps, enhancing the refilling of oxide


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vacancies and H2 generation, thereby minimizing the blocking of active sites

by formate and carbonate formation. As a consequence, both catalytic activity

and life were improved (217).

How do dopant ions, like Mo, facilitate surface reoxidation vis-a-vis

formate/carbonate formation? An understanding of the interaction of H2O

with the catalyst surface holds the key to the answer. Atomic simulation

calculations indicated (217) that the H2O molecule dissociated, preferentially,

over the dopant ion to hydroxylate the dopant and to protonate the surface

oxygen ions in the adjacent fluorite lattice. The average H2O adsorption

enthalpy of 257.7 for ceria- zirconia increased to 2 82.3 KJ/mole on doping

with Mo and to 2110.4 KJ/mole on introduction of W in the fluorite lattice

(217).

9.2.8. The LTS Reaction over Non-oxide Supports

While the role of the oxygen anion vacancies and surface OH groups on

metal- metal oxide-based catalysts in the WGS mechanism has been studied

extensively, there is another group of WGS catalysts based on molybdenum

carbide wherein the mechanistic picture is less clear (218–221). Patt et al.

(218) reported high activity for LTS over Mo2C catalysts and obtained higher

activity than over a commercial Cu-ZnO-Al2O3 catalyst. The precursor for

molybdenum was ammonium paramolybdate. The salt was dissolved in warm

water. Then, the liquid was, slowly, evaporated and the solid was calcined in

dry air at 500uC. The oxide was carburized using a temperature-programmed

treatment with CH4 and H2. The LTS activities of the resulting solid (BET

surface area 5 61 m2/g) as well as that of a commercial Cu-ZnO-Al2O3 catalyst

of similar surface area were compared at various temperatures in the range,

220–295uC, using a feed containing 62.5% H2, 31.8% H2O, and 5.7% CO. The

divergence of this feedstock composition from those in commercial practice,

especially the absence of CO2, may be noted. Under these conditions, the CO

conversion over the Mo2C catalyst was at least 50% higher than that over the

Cu-ZnO-Al2O3 sample. Moon and Ryu (219) found that the optimum

carburization temperature was 640–650uC. After repeated thermal cycling

in reductive and oxidative atmospheres, the authors found that even though

there was a decay in catalytic activity of both the Mo2C and a Cu-ZnO-Al2O3

catalysts, the Mo2C was, relatively, more stable. XPS results indicated that

the deactivation of the Mo2C catalyst was linked to the formation of the MoO3

oxide on reaction with H2O. Based on a Density Functional Theory study of

the WGS reaction over Mo2C, combined with infrared spectroscopy experi-

mental results, a redox mechanism was proposed by Tominaga and Nagai

(220). Upon CO adsorption, the authors observed two bands at 1626 cm21 and

1450 cm21. Instead of assigning these to formate species, the authors


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concluded that, since no accompanying symmetric band was observed in the

1350–1365 cm21 range, the 1626 cm21 band was not due to formate but

another vibration. The 1450 cm21 band was assigned to unidentate carbonate.

Based on DFT calculations, the authors assigned the 1626 cm21 band to CO

adsorption on a 3-fold Mo site. A summary of their proposed redox mechanism

is given below where * denotes a free adsorption site:

COz�[CO�, ð48Þ

H2Oz�[H2O�, ð49Þ

H2O�z�[HO�zH�, ð50Þ

HO�z�[O�zH�, ð51Þ

CO�zO�[CO2z2�, and ð52Þ

2H�[H2z2�: ð53Þ

The rate-limiting barrier was found to be the reaction of O with CO to form

CO2 (Eq. 52). This group has, also, extended their study of molybdenum

carbides to include cobalt-containing samples (221). Catalysts were prepared

by combining aqueous solutions of Co(NO3)2 and ammonium heptamolybdate

(NH4)6Mo7O24 and stirring at 80uC to produce a viscous mixture. Solids were

dried in an oven and calcined at 500uC. Carburization was carried out using

20% CH4/H2 mixtures and a temperature-programmed procedure. The

optimum Co content (for catalytic activity using a feed containing 10.5%

CO, 21% H2O the balance being He) was 50% (i.e., Co0.5Mo0.5C). Both the

initial activity and long-term stability of the Co0.5Mo0.5C catalyst was superior

to that of Mo2C. Even though its initial activity was superior to that of a Cu-

ZnO-Al2O3 catalyst, the latter’s long-term stability was better. In view of its

potential as a sulfur-tolerant LTS catalyst, similar to the Co-Mo sour gas shift

catalysts (Section 5) that operate at higher temperatures, further investiga-

tions on this system seem warranted.

10. CONCLUSIONS AND CHALLENGES

The WGS reaction is one of the primary industrial reactions that produce

hydrogen. The utilization of coal and biomass for the production of electrical


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power, chemicals, petrochemicals, and hydrogen and transportation fuels is

gaining importance. Declining resources and increasing prices of crude oil are

some of the major driving forces. In future, due to considerations of global

warming, coal may have to be used only in CO2- free power plants, which can

only work in combination with CO2 sequestration. This is only possible when

all carbon compounds in the feedstock are converted to CO2. CO conversion

processes (like the water gas shift reaction) play a key role here. Some sources

predict that by the year 2030, 10% of the yearly consumption of energy will

originate from the WGS reaction (222). The WGS reaction is a well-established

process in conventional chemical plants for the manufacture of ammonia,

methanol, refinery hydrogen, hydrocarbons (by the Fischer - Tropsch process),

etc. In current commercial practice, the WGS conversions are kinetically

limited at low temperatures and thermodynamically limited at high

temperatures. Due to intensive efforts during the last two decades, significant

progress has been made in the study of the mechanism of the WGS reaction

(223). At high temperatures (above 350uC) and over the iron oxide-based

catalysts, the redox mechanism, involving the reduction of the catalyst by CO

and H2 and its reoxidation by H2O probably prevails. At lower temperatures,

even though the detailed mechanism is not established definitively, the

following picture is beginning to emerge: The mechanism and the rate

determining step depends on the nature of the catalyst and process conditions.

On precious metal (Pt, Au) – partially reducible metal oxides (ceria, ceria-

zirconia, titanium oxide), the CO adsorbed, mostly on the metal, reacts with

the surface OH groups on the support to form surface species, like the

formates, carbonates and carboxylates. The concentration and stability of

these species depend on the support oxide, temperature and the partial

pressures of the reactants, especially H2O. Some of these surface species (like

the formates and carbonates) are also intermediates in the reaction path and

decompose to CO2 at higher temperatures and/or partial pressures of H2O.

The decomposition of these various surface species to CO2 is a crucial and slow

step in the reaction path. It is faster at higher temperatures and partial

pressures of H2O. The nature of the catalyst (metal type and loading) and

process conditions have a profound influence on the decomposition of these

intermediates. The accumulation of these species on the surface and, the

consequent, blocking of the catalytically active sites by them lead to loss of

catalytic activity. Factors that hasten their removal, by conversion to CO2, will

improve the catalytic performance of these catalysts. The primary reason for

the accumulation of these intermediates on the surface (and indirectly to loss

of catalytic activity) must be sought in the physicochemical changes under-

gone by the catalyst (e.g., loss of metal-support interface area). Basic additives

like the alkali ions facilitate the decomposition of formates and improve the

low temperature performance of catalysts, like Pt-Ceria, Pt-Ceria-Zirconia


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and Pt-Titania. Similarly, acidic additives, like the oxides of Nb, Mo, Ta and W

facilitate the decomposition of the carbonates to CO2 and accelerate the rate.

The strong adsorption and binding of CO on metals like Pt, at low

temperatures below 250uC, is, also, a rate – inhibiting factor and contribute

to the low activity and slow deactivation of these catalysts at low

temperatures. Development of WGS catalysts kinetically more active in the

190–250uC range is a general challenge. While the Cu-ZnO catalysts are active

in this range, their activity is low necessitating the use of low GHSVs (3000–

5000 h21). A more precise understanding of the mechanism will lead to the

development of better WGS catalysts for hydrogen generation in fuel cells.

There are other challenges to develop improved WGS catalyst and process

versions. Some of them are:

Challenges in High Temperature Shift: (1) Replacing chromium in the

iron oxide- based catalyst by a non-toxic promoter; even though many

chromium-free formulations(containing copper, for example) are in the

pipeline, they are not, yet, proven in commercial usage. (2) Discovery of a

novel, non- noble metal catalyst with higher catalytic activity that will enable

operation at GHSV 5 40,000 h21 and above; the iron oxide- based catalysts

operate at , 15000 h21. This requirement is especially relevant to fuel cell

applications; (3) Discovery of catalysts that can function successfully at low

steam to gas ratios will reduce, significantly, the energy costs associated with

hydrogen generation.

Challenges in Low Temperature Shift: There are, at least, three major

drawbacks in the use of Cu - ZnO – Al2O3 catalysts even in conventional,

stationary applications: (1) their relatively, low catalytic activity (GHSVs of

around only 3000 – 5000 h 21 leading to large – volume catalyst beds), (2) their

elaborate start-up, activation procedures and, (3) their high sensitivity to

sulfur and chlorine compounds as well as to steam below the dew point of H2O

at the operating temperature and pressure. Current pressures for many

downstream applications, for example, restrict the minimum LTS tempera-

tures to 190 – 200uC. Operation at lower pressures in fuel cells, for example,

can benefit from favorable thermodynamics at lower temperatures if a

suitable catalyst can be discovered. Even though well- formulated, modern,

Cu-ZnO-Al2O3 catalysts produce only minor amounts of methanol, it is

desirable to reduce this quantity still further or, even better, eliminate

methanol formation altogether. The above constraints are much more

important for LTS catalysts in mobile fuel cell applications. It is mainly

because of these and other reasons (like the pyrophoricity of the copper –

based catalysts) that the noble metal – reducible oxide catalysts are being

investigated as potential alternatives.

Challenges in Fuel Cell Applications: Even though noble metal –

based catalysts for fuel processors are already in the market, a completely


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satisfactory catalyst for WGS applications in fuel cells is not yet in commercial

operation. Ceria and titania- based platinum catalysts are the front runners as

potential water gas catalysts in fuel cell applications. Apart from their high

costs, some of their major drawbacks include their low activity below 250uCand deactivation in long – term operations, especially at lower temperatures

and high pressures. Formation of hydrocarbons at low temperatures and high

pressures (Fischer – Tropsch activity) is yet another drawback of these

catalysts. The present copper – based catalysts do not form significant

amounts of hydrocarbons (like methane) under LTS conditions. Attempts to

improve the long-term life of noble metal catalysts by incorporating acid or

basic additives in the oxide support have been described above (see Section 9).

Modifying the electronic properties of the noble metal, by alloying surface Pt

atoms with those, like Re, Au, Ag and Cu, which do not adsorb CO so strongly,

may, perhaps, be necessary to prevent poisoning by strongly held CO and,

thereby, increase their catalytic activity at low temperatures. Recent

theoretical calculations, by Knudsen et al. (224) indicate that this may indeed

be a promising approach. These authors find that a Cu-Pt surface-alloy binds

CO more weakly than pure Pt (Figure 52) and, hence, CO poisoning at low

temperatures is less likely with the alloy than with the pure metal; in a

temperature programmed desorption experiment, adsorbed CO desorbs at

lower temperatures from Cu-Pt surface alloys than from a pure Pt surface

(Figure 52). Interestingly, the Cu/Pt is also able to activate and dissociate H2O

more easily, the latter being the usual rate-determining step for the WGS on

Figure 52: CO TPD spectra for CuPt (111) surface alloys with varying amounts of copper(ML5 monolayer of Cu) after exposure of 10 Langmuir of CO at 2107uC (224).


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several metal surfaces. The higher WGS activity of Pt-Cu compared to Pt-

alone catalysts had already been claimed by researchers from the Honda

Motor Company (152). Similarly, the higher activity of Pt-Re and Pt-Ag

compared to Pt catalysts has also been claimed by Sud-Chemie Inc. (225).In

addition to the above challenges, development of a single catalyst formulation

for both the hydrocarbon reforming and water gas shift reactions will go a long

way in simplifying the inventory and design of fuel processors. Noble metals,

like Pt, Rh, and Re are the prime candidates for the composite catalyst.

However, materials, that can function as efficient catalyst supports for the

high temperature reforming/ partial oxidation as well as the low temperature

water gas shift reaction, will have to be discovered and developed to meet this

challenge. Major efforts in this direction are in progress worldwide.

Challenges in Fundamental studies: There are also more basic

fundamental problems and key issues that remain to be addressed and

clarified on the mechanistic aspects of the water gas shift, especially the LTS

reaction; (a) The role of oxygen mobility in the oxide component of the noble

metal- partially reducible oxide catalyst needs to be investigated further; in

the redox mechanism, the role of oxygen mobility is very clear and obvious.

However, the role of surface oxygen mobility is also crucial in the associative

mechanism because intermediates (e.g., formates and carbonates) are bound

to the oxide by their surface oxygen atoms and presumably move across the

oxide surface to the metal; information about the nature of the mobile species

(O22/OH 2) and the kinetics of their mobility will benefit both the redox and

associative mechanisms; (b) The need to confirm, unambiguously, that H2O

can indeed reoxidize partially reduced cerium oxide at low temperatures (150–

250uC). Such a reoxidation is a fundamental assumption in the redox

mechanism and its occurrence at high temperatures is well established in

the case of the Fe2O3-Cr2O3 catalysts. There has been no confirmatory

evidence of such a reoxidation process at low temperatures in the case of

partially reduced cerium oxide. Does the presence of a metal (like platinum)

facilitate the reoxidation of partially reduced ceria by water molecules at low

temperatures? While it is known that Pt enhances the reduction of ceria by

hydrogen and creation of surface oxygen vacancies and OH groups, it is not

confirmed that Pt also facilitates the reoxidation of the reduced ceria by H2O

at low temperatures. These experiments are crucial to confirm the redox

mechanism at low temperatures; (c) The need to establish a standard protocol

to estimate the noble metal dispersion when they are supported on partially

reducible oxides like ceria, titania, ceria-zirconia, chromia etc.; while many

labs(especially in the industry) have evolved in-house empirical methods for

catalyst screening purposes, a more scientific foundation is desirable;(d) The

need to model, more accurately, the transition state of formate decomposition

in the presence of co-adsorbed water. In the absence of water, formate is quite


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stable and decomposes (reverse thermal decomposition) to CO and OH only at

. 300uC. But, and this is important to note, in the presence of steam (as

during the WGS reaction), the formate decomposes very rapidly even below

150uC in the forward direction to a carbonate, the precursor of CO2 and H2,

picking up the second H, probably from a bridging OH group; most of the

theoretical models developed so far have not explicitly considered this major

influence of co-adsorbed water molecules in enhancing the decomposition of

formate ions to CO2 and H2 at temperatures typical of the WGS reaction and

have treated the decomposition only as a thermal decomposition; hence,

theoretical calculations taking into account the original transition state

picture of Shido and Iwasawa (108) and further elaborated by Jacobs et.al.,

which involves a ‘‘reactant-promoted decomposition of the formate’’ by co-

adsorbed water molecules, are desirable.

ACKNOWLEDGMENTS

We thank the reviewers for useful comments.

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Water Gas Shift Catalysis

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