WGS Rat Review

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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 40,000h21. 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 usedextensively in the industry for the manufacture of ammonia, methanol,hydrogen (for hydrotreating, hydrocracking of petroleum fractions and otherhydrogenations in the petroleum refining and petrochemical industry),hydrocarbons (by the Fischer-Tropsch process) and metals (by the reductionof 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=mol! ", !1"

CH4zH2O<COz3 H2 H2=CO~3; DH~206:3 kJ=mol! ", !2"

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, 2009Copyright # Taylor & Francis Group, LLCISSN: 0161-4940 print 1520-5703 onlineDOI: 10.1080/01614940903048661

325

Downloaded By: [Universidad de Sevilla] At: 09:41 9 August 2010

CH4z0:5 O2<COz2H2 H2=CO~2; DH~{35:6 kJ=mol! ", and !3"

CH4zCO2<2COz2H2 H2=CO~1; DH~247:4 kJ=mol! ": !4"

Reactions 1, 2, and 4 are endothermic while reaction 3 is exothermic. It may benoted that the molar ratio of H2 to CO varies depending on the source ofcarbon/oxygen.

Steam reforming [Eq. (2)] is the most popular mode of generating watergas, especially if the ultimate objective is the generation of pure hydrogensince it provides the highest molar ratio of H2/CO of 3. The exothermic partialoxidation [Eq. (3)] providing a H2/CO molar ratio of 2 is used in themanufacture 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 todifficulties in external heat supply, internal heat generation (autothermalreforming) 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 gaswith a very low H2/CO molar ratio of one. This process is used in themanufacture of water gas from natural gas for the reduction of iron orewherein 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 itcame into popular usage later, as a source of hydrogen for the Haber processfor the manufacture of ammonia:

CO g! "zH2O g! "<CO2 g! "zH2 g! " DH~{41:1 kJ=mol! ": !5"

In the initial stages of the Haber ammonia process, the hydrogen needed forthe process was obtained from the water gas generated by Eq. (1). In coal orcoke gasification [Eq. (1)], when steam is contacted with incandescent coke (atabout 1000uC), CO2 is an additional product [Eq. (6)] especially at lowertemperatures:

Cz2H2O<CO2z2H2 DH~90 kJ=mol! ": !6"

While CO2 could be easily removed from the products of the reaction (byabsorption in water), CO had to be removed by liquefaction or copper liquorscrubbing. A catalytic process to remove the CO from (CO + H2) mixtures wasneeded. In 1914, Bosch and Wild (2) discovered that the oxides of iron andchromium 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 theHaber process. Thenceforth, the water gas produced from the carbonaceoussource by steam reforming was passed over the iron-chromium catalyst to shift

326 Ratnasamy and Wagner


the CO to CO2 by the water gas shift reaction. Iron-based catalysts are stillused today industrially. There are four general types of water gas shiftcatalysts. One of them is the promoted iron oxide catalyst. Catalysts of thistype promote the shift reaction at moderately high temperatures (350–450uC)and are therefore called high temperature shift (HTS) catalysts. The secondtype is copper-zinc oxide catalyst and is called the low temperature shift (LTS)catalyst because it is used at relatively low temperatures (190–250uC). Thethird 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 wasinterest (in the past) in a fourth type of catalyst, medium temperature shift orMTS catalyst that operates at temperatures between the HTS and LTScatalysts. Normally, these are copper-zinc catalysts that are actually LTScatalysts modified (usually with iron oxide) to operate at slightly highertemperatures (275–350uC) than a standard LTS catalyst. In addition to theabove four, precious metal- based catalysts (mainly platinum and gold) havebeen under intensive investigation during the last decade for use in fuel cellapplications. Promoters, like Cu and Al2O3 are added to the conventional ironoxide - 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 (sectionII), however. Hence, to achieve higher conversions of CO at lowertemperatures (190–250uC), a second, more active catalyst, based on Cu-ZnO,was developed in the early 1960s and is used in the industry extensively. Itmay 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 adiabaticconverters where the effluent from the reformer system is converted in twoWGS reactors (HTS and LTS converters, respectively), with the second WGSreactor at a significantly lower temperature in order to shift the equilibriumtowards the favored hydrogen product (Fig. 1). The modern, two stage WGSconverter systems reduce the CO concentrations to about 0.3%(wt) from thehigh levels (10–50%) in the outlet from the reformers. During the last couple ofdecades, fuel cells, generating electricity from the reaction of hydrogen withoxygen, for stationary and mobile applications, have become popular. A crucialprerequisite for the techno-economic success of fuel cells, especially those thatoperate either at low temperatures (like the Polymer Electrolyte MembraneFuel Cells, PEMFC) or in mobile applications (as in automobiles), is thediscovery of improved reforming and WGS catalysts for the generation ofhydrogen which are much more active than those used in chemical plants. Therequirements of WGS catalysts for fuel cell applications are quite differentfrom those of the traditional Fe2O3-Cr2O3 or Cu-ZnO based catalysts. Thecatalyst bed must have a reduced volume and weight to be economical and

Water Gas Shift Catalysis 327


have sufficient durability to withstand rapid start-up and shut-downconditions. 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 twoHTS and LTS catalysts are extremely pyrophoric when activated (reduced)and, therefore, safety from runaway heat generation and fires cannot beensured upon air exposure. In response to these needs, noble metal- basedreforming and WGS catalysts are under intense development worldwide forfuel cell applications.

Newsome had provided an excellent review of the WGS literature up to1980 (4). Lloyd et al. have reviewed the industrial developments in this areaup to 1996 (5). A comprehensive review by Kochlofl in 1997 (6, 7) covers boththe fundamental and applied aspects of the field. More recently, in 2003,Ladebeck and Wagner have given a brief survey of the WGS catalystdevelopments, especially for fuel cell applications (8, 9).

This review, after highlighting the significant features of the conventionalHTS and LTS catalysts and processes based on the Fe2O3-Cr2O3 and Cu-ZnO-Al2O3 catalysts, respectively, critically examines the extensive results (both inthe journal and patent literature) from the study of noble metal based WGScatalysts 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).



structures, active sites, reaction intermediates and mechanisms. This subjecthas developed significantly, mainly, in the last decade due to its relevance tothe 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 theinlet to the WGS stage for some typical applications. The high N2 content inthe feed, for eventual application in ammonia synthesis, comes from theaddition of air in the secondary reformer to provide the N2 reactant for NH3.The water gas shift reaction is moderately exothermic (Eq. 5) and conversionsare 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=T! "{4:33# $, !7"

where T is in K. Accordingly, high conversions are favored at lowtemperatures and are not affected, significantly, by changes in total pressure.The reaction is reversible and the forward rate is strongly inhibited by thereaction 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.



inhibit, due to thermodynamic reasons, CO conversions. This limitation can,however, be mitigated by using two or more beds with heat removal betweenthem. CO levels at the exit of the HTS reactor are around 3–5wt% whilevalues around 0.3wt% can be achieved at an exit temperature of 200uC in theLTS reactor. The lower limit of the operating temperature in the LTS reactoris 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 originatingfrom condensation should be avoided, since leaching of the water-soluble Cr 6+

ion is equally unwelcome. Similarly, the lower limit on the pressure is theoperating pressure of the downstream units (10–60 bars). The water contenthas a strong influence on CO conversion. The water entering the WGS reactorcan be varied by controlling the amount added upstream at the reformingstage 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 HTSreactor 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).



determines the thermodynamic equilibrium conditions. The effect of waterconcentrations at various temperatures on the equilibrium CO concentrationis 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 andrepresents a syngas generated from autothermal reforming (ATR) andexcludes any residual hydrocarbons which may also be present. By increasingthe 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. Byoperating at 100uC higher temperature, a significant reduction of the reactorsize can be achieved by utilizing the more favorable kinetics at the highertemperature. The CO concentration at the inlet to the HTS reactor can varywidely in the range 12–40% (dry basis) depending on the raw material(natural gas or coal) and the reforming process (steam or autothermalreforming) utilized to generate the CO. The HTS exit (and LTS inlet)concentrations are in the range of 3–5% (dry basis) and depend on theoperating 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) whilevalues much higher than that stoichiometrically needed by Eq. (5) increase theenergy costs and adversely affect the process economy.

The method of producing the syngas will also affect the WGS equilibriumcompositions (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).



concentration (due to the dilution with nitrogen) compared to steamreforming. The lower H2 concentration increases the equilibrium COconversion whereas the high H2 concentrations expected with steamreforming lower the equilibrium CO conversions. Figures 5 and 6 (8, 9) showthe equilibrium CO composition as a function of H2 content at constant COand 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 COconcentration would be 1.66% for steam reforming at the same temperaturerequired to achieve 1% CO for a feedstock from autothermal reforming. Theeffect 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



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).



minimize the volume of the WGS reactor, especially in fuel processors for fuelcell 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 areavailable (4–11). The important structural and textural roles of Cr2O3 in thecatalyst formulation has also been investigated in detail (11). Two stage COconversion systems employing WGS using Fe2O3-Cr2O3 catalysts andmethanation using nickel-based catalysts for CO removal was the commonand economical design in ammonia synthesis up to the late 1950s. Most ofthose plants employed the Fe2O3-Cr2O3 HTS catalyst in the first, hightemperature reactor as well as in the second stage converter at temperaturesas low as 320uC. The conventional Fe2O3-Cr2O3 catalysts worked extremelywell for these high temperature applications but their relatively poorperformance in the lower temperature, second bed of these reactors motivatedfurther investigations. The early development of unsupported metallic coppercatalysts or copper supported on Al2O3, SiO2, MgO, pumice or Cr2O3 werecharacterized by relatively short life and low space velocity operations (400–1000h21). Important progress was made by the addition of ZnO or ZnO-Al2O3.These Cu-ZnO-Al2O3 catalysts exhibited not only a considerable increase inlifetime, but also an increase in the turnover number by an order ofmagnitude. Today, in industrial adiabatic converters, the syngas effluentfrom the reformer system is converted in two steps, with the second step at asignificantly lower temperature in order to shift the equilibrium towards thefavored hydrogen product. In the first step, catalysts based on the Fe2O3 –Cr2O3 oxides are applied at a reactor inlet temperature of 300–360uC and atotal pressure between 10 and 60 bars. Under normal operating conditions,the temperature rises, progressively, through the reactor bed and can increaseup to 500uC. At exit gas temperatures of 400 to 500uC, the CO content can bereduced in an industrial HTS converter to 5 vol% or lower. In this section, wereview the main features of the HTS catalyst/process and highlight thedevelopments 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 varybetween 30–100m2/g depending on the Cr2O3 and Al2O3 contents andcalcination temperatures. One of the major functions of Cr2O3 and Al2O3 isto prevent the sintering, and, consequent loss of surface area of the iron oxide



crystallites during the start – up and further operation (11). Pure Fe2O3, whenused as a HTS catalyst, deactivates fast due to sintering of the iron oxidecrystallites. In addition to being a textural promoter preventing the sinteringof iron oxide crystallites, Cr2O3 also functions as a structural promoter toenhance the intrinsic catalytic activity of Fe2O3. As supplied, the Fe2O3-Cr2O3

catalyst is a solid solution of a - Fe2O3 and Cr2O3, wherein the Cr3+ ionsubstitutes, isomorphously and partially, the Fe3+ ions in the a - Fe2O3 latticeframework. Even though most of the chromium ions in the fresh catalyst arepresent in the Cr3+ state, a small fraction, especially on the surface, is presentin 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=mol! ", and !8"

3Fe2O3zCO<2Fe3O4zCO2 DH~z24:8 kJ=mol! ": !9"

The reduction has to be done carefully and the reaction heat removed, to avoidfurther reduction of Fe3O4 (Eqs. 10–14).

Fe3O4zH2<3FeOzH2O DH~{63:8 kJ=mol! ", !10"

Fe3O4zCO<3FeOzCO2 DH~{22:6 kJ=mol! ", !11"

FeOzH2<FezH2O DH~{24:5 kJ=mol! ", !12"

FeOzCO<FezCO2 DH~{12:6 kJ=mol! ", and !13"

Fe3O4z4H2<3Fez4H2O DH~{149:4 kJ=mol! " !14"

Importantly, neither pure hydrogen nor H2-N2 mixtures should be used toreduce the HTS catalysts to avoid the occurrence of the strongly exothermicreduction to metallic Fe (Eq. 14). It is Fe3O4 that is the active phaseresponsible for the WGS reaction. The CrO3 phase present must also bereduced to Cr2O3 during start-up (Eqs. 15 and 16):

2 CrO3z3H2<Cr2O3z3 H2O DH~{684:7 kJ=mol! ", and !15"

2CrO3z3 CO<Cr2O3z3 CO2 DH~{808:2 kJ=mol! " !16"



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 operatingconditions 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 theseconditions. Formation of metallic Fe (due to low H2O/H2 ratios) can trigger thehighly 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=mol! ", and !17"

2CO<CzCO2 DH~{172:5 kJ=mol! " !18"

While maintaining sufficiently high H2O/H2 ratios is important, passingsteam, 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) andthereby lower catalytic activity:

2Fe3O4zH2O<3Fe2O3zH2 !19"

The Fe2O3-Cr2O3 catalysts are rugged and have a lifetime of 3–5 yearsdepending, mainly, on the temperature of operation. Unlike the Cu-ZnO (LTS)catalyst, the Fe2O3-Cr2O3 catalyst is not extremely sensitive to the presence ofsulfur and can tolerate the presence of substantial amounts of sulfur due tothe facile reversibility of the sulfidation reaction (Eq. 20):

Fe3O4z3H2SzH2<3FeSz4H2O DH~{75:0 kJ=mol! " !20"

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

4H2O) varies from

3 6 10210 to 45 6 10210 in the range of 300–450uC. The rate of the HTSreaction is limited by pore diffusion and linearly dependent on the steampartial pressures under industrial conditions (12). A power law – type rateEq.satisfactorily fits the experimental data (13). Apart from an increase in thepressure-drop across the catalyst bed during use due to inadequatemechanical crushing strength of the catalyst pellets, catalyst deactivation ismainly due to loss of iron oxide surface area by thermal sintering. In additionto the abovementioned risks of thermal sintering during start- up reduction ofthe catalyst, the exothermic nature of the WGS reaction (Eq. 5) also generatesa large amount of heat during the operational phase of the process. Forexample, it has been estimated (5) that the conversion of 1 vol% of CO resultsin a temperature rise of about 7–10uC in the catalyst bed. The syngasfeedstocks to the HTS reactor may contain anywhere from 8vol% (steamreformer) to 45% (partial oxidation/ autothermal reforming of methane or coal/coke) of CO. Hence, temperatures in the catalyst bed may rise by 500uC (to



about 800–850uC) if heat removal is inadequate. Distributing the catalyst intwo or three beds and providing inter-bed coolers can restrict the exittemperatures to about 450uC and outlet CO content to 3–5wt%. The Fe2O3-Cr2O3 catalysts can tolerate sulfur up to, about, even 1000 ppm. Their majordrawbacks are: (a) the toxicity of the water-soluble Cr6+ ions posing healthhazards during catalyst manufacture and handling, and (b) the low volumetriccatalytic activity (GHSV510,000 – 15,000h21), especially at low temperatures,when CO conversion is favored thermodynamically, necessitating the use oflarge catalyst bed volumes. The latter handicap is of crucial importance in fuelcell applications.

3.1.1. Influence of Catalyst Composition and Preparation Methods

The preparation method of the Fe2O3-Cr2O3 catalyst has a stronginfluence on their properties (14, 15). They are usually prepared bycoprecipitation of the hydroxides followed by drying and calcining them tothe corresponding oxides. The oxides are reduced in situ before use. Theprecipitation method involves the conventional coprecipitation of the mixediron and chromium nitrates with ammonium hydroxides. In an alternateimpregnation method, iron hydroxide gel is first prepared and thenimpregnated with chromium nitrate solution. Chromium retards the sinteringof the iron oxide crystallites both during activation and the WGS reaction (14,15). X-ray photoelectron spectroscopy revealed that there was surfaceenrichment of Cr ions in fresh samples prepared by both the coprecipitationand impregnation routes. The surface concentration of chromium was higherin the impregnated samples. However, after activation and running the WGSreaction, it was observed that the relative surface concentration of chromiumhad decreased significantly in both the samples, suggesting that, during theactivation and WGS reaction, some Cr ions had migrated from the surface intothe 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 latticeand occupies, preferentially, the octahedral sites because of its high crystalfield stabilization, in contrast to the Fe3+ ion (d5) which does not have anypreferred site. When the Cr ions occupy tetrahedral sites, they cause strain inthe magnetite lattice, thereby decreasing the average particle size of the ironoxide. The lower particle size, in turn, increases the surface area of the ironoxide and leads to enhanced catalytic activity. The presence of chromium alsocaused (16) an increase in the surface Fe2+/Fe3+ ratio (from XPS data) in thefresh samples and this effect was more pronounced for the impregnatedsample. After reaction, however, these ratios were smaller for both theprecipitated and impregnated samples than the fresh sample. The Mossbauerspectra of the pure magnetite sample indicated that the Fe3+ ions are in



tetrahedral sites (A sites) and Fe2+ ions are in octahedral sites (B sites) of Fe3O4.Chromium addition to magnetite (both by the precipitation and impregnationmethods) produced a decrease in the hyperfine magnetic fields in both sites dueto the partial replacement of iron ions by chromium. This decrease was morepronounced in B sites, indicating that Cr 3+ ions entered preferentially theoctahedral B sites. The sample prepared by the impregnation method was moreactive in the WGS reaction. The role of chromium is twofold: (a) as a texturalspacer/stabilizer for iron oxide crystallites (stabilization of their smallercrystallite size) and (b) as a structural promoter in increasing the intrinsiccatalytic activity of iron oxide crystallites due to (a) the increase in lattice straincaused by the substitution of Fe3+ by Cr3+ ions in themagnetite lattice and (b) anincrease in the surface area of the iron oxide crystallites as mentioned above.

During the last two decades, due to the rising cost of hydrocarbonfeedstocks, plants have been forced to keep operating costs low by being asenergy efficient as possible. One method used in improving energy efficiency isby reducing the overall steam to gas ratio of the plant (usually starting at theinlet to the reformer). Depending on the level, these lower steam-to-gas ratioscan cause over-reduction of the Fe2O3 in HTS catalyst and result in theformation of iron carbides. Iron carbides are very effective catalysts for theformation of hydrocarbons by the Fischer-Tropsch reactions. Products fromthe Fischer-Tropsch reactions would negatively impact both LTS catalystperformance and plant efficiency. The over-reduction can, also, result in avolume shrinkage within the catalyst pellet that weakens it and may lead toan increased rate of mechanical breakdown. To minimize the low steam to gasratio effects on HTS catalysts, the Sud-Chemie Inc. group developed andintroduced, in the late eighties, a copper promoted iron oxide- chromium oxideformulation that successfully suppressed the Fischer- Tropsch reactions incommercial 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.



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

oxide-chromium oxide catalyst and a copper - promoted, iron oxide- chromiumoxide catalyst. It may be seen that there is a significantly lower by-productformation in the latter. It is speculated that the presence of copper suppressesC-O cleavage (in CO), prevents the formation of iron carbides and therebyavoids the hydrogenation of the adsorbed carbon (to hydrocarbons) andfacilitates 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 influenceboth the equilibrium CO content and reaction kinetics. Other factors to beconsidered are pressure and catalyst activity.

Temperature: Since the reaction is exothermic, higher CO conversions canbe obtained by reducing the temperature at which the gas leaves the reactor.However, this principle applies only to a catalyst that is equilibrium- and notkinetically- limited. A reactor that is operating with the exit gas COconcentration above equilibrium (kinetically limited) may benefit from highergas/ bed temperatures. The exit gas temperature determines both the catalystreaction rate in the bottom of the bed and the CO equilibrium value of theoutlet gas. For a reactor loaded with a highly active catalyst, the exittemperature is determined primarily by the inlet temperature, CO concentra-tion, and steam/ gas ratio. Exit CO equilibrium is usually achieved in thesecases. The higher the inlet CO concentration and lower the inlet steam/gasratio, the larger will be the overall temperature rise through the bed. Thetemperature rise is also somewhat dependent on the other gas componentsand 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.



30–75uC are common across commercial reactors. For some of the recentcopper-promoted iron oxide-chromium oxide catalysts, the maximum operat-ing temperature is around 510uC.

Steam to gas ratio: Both laboratory and commercial data indicate thathigher steam/dry gas ratios in commercial ranges also increase the water gasshift reaction rate. As a result of the steam/ dry gas ratio effect on both thethermodynamic and kinetic properties of the process, higher values givehigher CO conversions and a lower exit CO content in the gas. In most plantconfigurations, the inlet steam/gas ratio cannot be independently controlled inthe HTS reactor. Other considerations, such as downstream gas purityrequirements and the overall site energy balance determine the inlet reformersteam/ gas ratio and, as a result, fix the value at the inlet of the HTSconverter. Commercial operating conditions are such that the equilibrium COconcentrations 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 exittemperature on the equilibrium CO concentration at the exit of the HTSreactor. They also illustrate the differences in achievable CO levels as afunction of the end-product (ammonia or hydrogen; Figs, 9 and 10,respectively) as well as the type of reformer feedstock (partial oxidation ofnatural gas or fuel oil; Figures 11 and 12, respectively) used to generate theHTS feed gas. In addition to CO conversions, the steam to gas ratio can alsoaffect the production of hydrocarbons (mainly methane) by the Fischer-Tropsch reaction. To minimize such undesirable reactions, a minimum steamto gas ratio of 0.4 and a maximum CO/CO2 ratio of 1.6 is ensured in the HTSreactor.

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



Pressure: The equilibrium CO concentration is virtually unaffected bysystem pressure. The pressure will, however, have an impact on the systemkinetics due to pore diffusion limitations and partial pressure effects of thereactants. 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.



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 afunction of time and operating temperature and is irreversible. Thermalsintering occurs more rapidly at higher temperatures. Hence, the maximumbed/exit gas temperatures are usually limited to less than 510uC. Somethermal sintering is unavoidable in the start-up and normal operation of thecatalyst. Commercial data suggest that there is approximately a 50% loss intotal surface area over the first few months of operation and, then, a further25% loss throughout the remaining life of the catalyst. As a result of thesechanges in surface area, the deactivation rate for the catalyst is faster duringthe first few months of operation and then stabilizes with very gradual agingafter the first year. Unlike LTS catalysts which can show distinct zones ofdeactivation (completely inactive, partially deactivated and essentially fresh,non-deactivated), the HTS catalyst undergoes a more gradual deactivationthat is spread throughout the bed. The more typical symptom of activity loss isa gradual spreading out of the reactor temperature profile and an increase inthe CO leakage. Figure 13 shows a typical change in the temperature profilewith time- on- stream in commercial reactors. For a relatively new catalyst,the temperature increases more sharply in the top 50–60% of the bed. Aftersome time-on-stream, the catalyst in the top is less active and the temperatureprofile changes. Throughout the catalyst aging period, however, therecontinues to be some catalyst activity through all of the beds. As the profilespreads more throughout the bed, it may become necessary to increase theinlet gas temperature in order to maintain acceptable exit CO levels.

Apart from thermal sintering, activity loss for HTS catalyst is mostcommonly due to the presence of poisons in the feedstock and the deposition ofsolids 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.



and/or silica) will coat the outer surface of the catalyst and block availablepores. In most modern plants, the HTS catalyst is downstream of equipmentoperating at higher temperatures such as reformers and heat exchangers. As aresult, there may be a gradual depositing of steam- volatile compounds intothe HTS bed. In time, these deposits can plug the pores in the catalyst and/ orthe void space between catalyst tablets. As a consequence, the activity declinesand the pressure drop may increase. Proper design of the pore size distributionand geometric shape of the catalyst pellets can minimize such effects. Thepresence of sulfur in the feed gas will affect the size of the converter, asallowances must be made for the adverse effect of sulfur on catalytic activity.The presence of oxygen (from the secondary reformer or the partial oxidationreactor) may also influence the design since the oxygen will be converted towater through an exothermic reaction. Thus, when the shift feed containsappreciable oxygen, an allowance may be necessary for the accompanyingtemperature rise due to this reaction. Any saturates, e.g. methane, ethane,propane or unsaturates (ethylene, propylene) in the process gas willessentially pass through the shift converter unchanged. There is no conclusiveevidence to indicate that the saturates will crack, or that the unsaturates willbe hydrogenated to any significant degree. Even if the unsaturates dohydrogenate, this side reaction apparently does not affect the catalyticactivity. Acetylene, on the other hand, can be troublesome, because it doeshydrogenate and impair the catalytic activity. When both diolefins and nitricoxide are present together in the feed stream, polymeric gums are usuallyformed 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 thecatalyst life. Although equilibrium CO leakages are often experienced at the

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



start of run and for some period of time, the reactor size and catalyst volume toachieve and maintain equilibrium throughout the total charge life would besubstantially higher and cost prohibitive. As a result of these cost considera-tions, the reactor is usually designed to be kinetically, rather thanequilibrium, limited. This means that any factor influencing the overallreaction kinetics will have a much more important impact on the requiredcatalyst volume for a given application. Typical design lives for a HTS catalystare 3–5 years before there is a need for catalyst replacement.

Since HTS catalysts/ reactors are usually designed to be kineticallylimited, the inlet gas temperature will have a significant impact on therequired catalyst volume. Lower inlet gas temperatures will requireincreased catalyst volumes to achieve similar levels of performance.Figure 14 shows the impact of temperature on reaction rates across a typicalFe2O3-Cr2O3 HTS catalyst. At the higher operating temperatures for HTSreactors, the WGS reaction is much more pore diffusion limited compared tothe LTS reaction. An increase in reaction rates can be achieved byincorporating a catalyst with high geometric surface area per unit loadedvolume of the reactor and/or increasing the size of the pores. Althoughpressure has no impact on the WGS equilibrium CO levels, there is asignificant influence on the reaction rate because of pore diffusionconsiderations. Figure 15 shows the relative influence of system pressureon reaction rates and the corresponding required catalyst volumes. The rateincreases with reactor pressure up to about 21 bar. Required catalystvolumes would correspondingly decrease with increasing pressure. Operatingpressure for a HTS plant is usually set more by an examination of overalleconomics (related to feedstock supply pressure and equipment costs) ratherthan catalyst reaction rate effects.

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



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 ofiron oxide (17–19):

COz O! "<CO2z!", and !21"

H2Oz!"<H2z O! ", !22"

where (O) is an oxygen atom on the oxide surface; and ( ) is a vacant site (anoxygen anion vacancy) on the surface caused by the removal of an oxygenatom. The surface is reduced by CO (Eq. 21), and subsequently, oxidized byH2O [Eq. (22)}. This mechanism is referred to, in subsequent literature, as the‘‘redox mechanism’’. The term redox mechanism denotes that the catalystitself undergoes changes in oxidation state during the course of themechanism. It does not refer to the oxidation state changes of the reactantsor products or associated intermediates. It should be pointed out, here, thatthe Eley-Rideal mechanisms refer to an adsorbed species reacting with a gas-phase species. The redox mechanism implied in Eqs. (21) and (22) isessentially the same as the Mars - Van Krevelen mechanism, with thedifference that the oxygen used to oxidize the catalyst [Eq. (22)] comes fromthe water rather than the gas-phase oxygen. Whether this oxygen atom(extracted from gas-phase H2O) is better referred to as an ‘‘adsorbed’’ speciesor, 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.



A multistep Langmuir- Hinshelwood type mechanism (Eqs. 23–27) wasproposed by Oki et al.(20–22) in 1973. From simultaneous exchange ratemeasurements, they concluded that while the evolution of gaseous H2 fromadsorbed H atoms [Eq. (27)] is the rate determining step at low COconversions, adsorption of CO [Eq. (23)} controls the overall reaction rate atsteady state, near-equilibrium, conditions prevalent in industrial reactors:

CO g! "za<CO a! ", !23"

H2O g! "z3a<2H a! "zO a! ", !24"

CO a! "zO a! "<CO2 a! "za, !25"

CO2 a! "<CO2 g! "za:, and !26"

2H a! "<H2 g! "z2a !27"

In the above Eqs., ‘‘a’ refers to an adsorption site. It can be located either onthe support or the metal oxide. The HTS reaction on Fe2O3-Cr2O3 catalystsprobably 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 amountsof 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 toincrease the reactivity of all the oxides with the promotional effect being mostpronounced with Cr2O3, U3O8 and CeO2-ZrO2 supports. Comparisons werealso 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 (BETarea 5 2.3m2/g) was more than 25 times that of Pt-Fe2O3-Cr2O3 (BET area 5

63m2/g) (Table 3). However, catalytic activity over this catalyst droppedquickly as temperature was reduced. Trimm (19, 21) has also compared thekinetics of the WGS reaction for the promoted and unpromoted Fe2O3-Cr2O3

catalysts. The general power rate law expression remained unchanged in theabsence and presence of the noble metal promoter indicating that it is thenumber of active sites that has increased by promotion (by noble metals).



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: 8wt% Cr2O3-Fe3O4, 8wt% CeO2 – Fe3O4,50wt% CeO2-ZrO2.

bNot measured. cNo measurement at 450uC; calculated from Arrheniusparameters.



Apparent activation energies were found to be similar (, 50kJ/mole) for Ptsupported on CeO2, Fe3O4-Cr2O3, CeO2-Fe3O4, Fe3O4, V2O5, MgO, MnO, andAl2O3, despite up to 15-fold differences in rates of reactions (Table 3;Figure 16). Since it is unlikely that surface diffusion of oxygen will havesimilar activation energies for such a variety of solids, the authors suggestedthat diffusion of oxygen on the surface or across the surface of the support toreact with CO adsorbed on the metal cannot be rate controlling in WGSreactions on Fe3O4-Cr2O3 promoted with noble metals. Rhodium was found tobe the most active promoter for Fe2O3-Cr2O3 oxides. More recently, this grouphad probed (24) the origin of rhodium promotion of Fe3O4-Cr2O3 catalysts forthe HTS reaction using various kinetic techniques and concluded that, of thetwo steps that may restrict the rate of the WGS reaction over iron- chromiumoxide catalysts (reduction by CO and H2 generation through reoxidation bywater), rhodium acts primarily by accelerating the latter.

Although the promoted catalysts are more efficient than the unpromotedFe2O3-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 toreduce CO concentrations still further.

During the last decade, attempts to develop improved HTS catalysts havebeen along two main lines:(A) replacing, at least partially, Fe by more activeelements (like noble metals), and (B) replacing Cr, partially or completely, bynon-toxic elements like Cu, Ca, Ce, Zr, La etc. (26–34). Promotion of theFe2O3-Cr2O3 catalysts by 2wt% Ag, Cu, Ba, Pb and Hg was explored byRhodes et al. (31). The catalysts were prepared by coprecipitation. Boron wasfound to poison the activity slightly whereas the others did increase theactivity between 350–440uC, with the relative order being Hg.Ag, Ba.Cu.Pb. unpromoted Fe2O3-Cr2O3 . B. From their results, the barium orsilver- promoted Fe2O3-Cr2O3 catalysts appear promising: a 10–15% increasein CO conversion was observed (when Ag or Ba was incorporated in theconventional Fe2O3-Cr2O3 catalysts) at their reaction conditions (400uC, 27bar, GHSV 5 1.2 6 10 6 h21; the volume of steam was 75 volume% of the drygas). The promoters decreased the activation energy of the reaction (Table 4).Based on the compensation effect (Fig. 17) seen when the activation energieswere plotted against the corresponding pre-exponential factors (in theArrhenius Eq.), the authors concluded that CO adsorption is an importantfactor controlling the relative catalytic activities of the various samples in theWGS reaction. Andreev et al.(29) studied the effect of the addition of CuO,CoO, and ZnO (5wt%) 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 (0.17–1.5wt%) in Fe2O3-Cr2O3



catalysts and found that Cu was in the metallic state under the WGS reactionconditions. 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 toguarantee 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 106h21. b ¡ 4kJ/mol.

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



catalyst, and health hazards cannot be fully ruled out despite considerableeffort. In addition, the spent catalyst ultimately poses a hazard to man and theenvironment and must be disposed of in accordance with the laws for thedisposal of toxic waste. HTS catalysts completely free from Cr and containingCa, 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 catalystsbased on alkali- promoted Co-, Cu-and Fe- manganese oxide systems werereported by Hutchings and coworkers (36, 37). The relative first order rateconstants in the WGS reaction for Fe-Cr, Fe-Mn, Cu-Mn and Co-Mn catalystswere found to be (36, 37) 1.0, 0.06, 0.75 and 1.75, respectively. Their Co-Mncatalyst, however exhibited significant methanation activity and the Cu-Mncatalysts were more sensitive to sulfur than the Fe-Cr formulations. Ladebeckand Kochloefl (38) had found that chromia-free, iron oxide catalysts containingabout 5wt% of Al2O3, 2wt% of Cu and 2.5wt% of CeO2 were very active for theHTS 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 (3wt%), prepared by thecoprecipitation (for Al and Fe) - impregnation (for Cu) method, in the HTSreaction. 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 ofchromium, in Fe- Cu – based catalysts for the HTS reaction. These Fe-Th-Cucatalysts were more active than the commercial Fe-Cr-Cu catalyst at H2O/CO5 0.6 and 370uC. Its high activity was attributed to an increase in surfacearea due to the presence of thorium. From a detailed study of chromium-free,iron-based HTS WGS catalysts, Natesakhawat et al. (26) concluded that acombination of copper and aluminum is a potential replacement for Cr in HTScatalysts. Further improvements in HTS activity of Fe-Al catalysts could beachieved by the addition of small amounts of copper or cobalt. The COconversions (at 400uC, CO/H2O/N2 5 1/1/18 (vol) and feed GHSV of 6000h21)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 activephase, magnetite (Fe3O4) by retarding its further reduction to FeO andmetallic Fe. The promotional effect of Cu was found to be strongly dependenton the preparation method. Fe-Cu-Al catalysts prepared by a one-step method(simultaneous coprecipitation of all the three component hydroxides) hadhigher CO conversions than those prepared by a two-step, coprecipitation –impregnation method (coprecipitation of the Fe and Al components followed byimpregnation of Cu on the precipitate obtained by coprecipitation) (Fig. 18).



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. Thebetter stabilization against sintering, at higher temperatures, of the coppercrystallites in the coprecipitated samples is probably the reason for itssuperior 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 fromthree different reduction sites. The reduction of hematite to magnetiteappears to have shifted from 300–290uC. The major peak from reduction ofCu species appears at 260uC with a very weak shoulder around 220uC. Thisshoulder is possibly due to reduction of the Cu species which are on theexternal surface of the catalyst and which can be reduced easily. The rest ofthe copper species appear to be more difficult to reduce as seen by a shift inreduction temperature from 220–260uC, possibly due to stronger interactionwith the hematite matrix. While the peak resulting from the reduction ofhematite to magnetite shifts to lower temperatures (from 300–290uC), thepeak from a further reduction of magnetite changes very little compared tothe non-promoted, Fe-Al sample. These results suggest that the preparationmethod makes a significant difference in the way Cu promoter isincorporated into the catalyst structure. During the last decade, moreactive and chromium-free, noble metal-based HTS catalysts are underdevelopment 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).



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 etal. (5). The development of highly efficient sulfur removal hydrodesulfurisa-tion technologies using Co(Ni)- MoO3- Al2O3 catalysts in the 1960s providedammonia manufacturers with syngas streams containing less than 1.0 ppmsulfur. This, in turn, enabled the use of the otherwise sulfur-sensitive Cu-ZnOcatalysts at sufficiently low temperatures (190–200uC) when the equilibriumCO concentrations, at the exit of the LTS converters, can be below 0.3%. Itmay be recalled that the Fe-Cr catalysts are not active enough below 350uC, atwhich temperature, the equilibrium CO concentrations are around 3–5%. As adirect consequence of having such low levels of CO (below 0.3%wt) from theCu-ZnO catalysts, it was economic to incorporate a methanation stage in theprocess 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).



was formerly used to remove residual CO, thereby enhancing the technoeco-nomic viability of large ammonia plants. The activity of metallic copper in theWGS reaction has, of course, been known for a long time (4–7). The problemwas the easy copper sintering and the subsequent loss of copper surface areaduring the activation (reduction of the precursor copper oxide) and use ofcopper catalysts. Various stabilizers, like SiO2, Cr2O3, Mn - Cr2O3 etc., wereevaluated for their ability to stabilize the copper surface area (39). Theintroduction 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 formulationsenabled the production of catalysts with high and stable copper surface areasand established the LTS process as a standard operation in any scheme ofhydrogen production from carbonaceous raw material. Currently, Cu-ZnO-Al2O3 based catalysts are used almost exclusively for industrial LTSoperations.

The feed gas to the LTS reactor is that exiting the HTS unit cooled, eitherby direct or indirect heat exchange, to approximately 200uC. LTS convertersare employed more frequently in hydrogen and ammonia - producing plantsthan in methanol or hydrocarbon (Fischer-Tropsch) plants. Hydrogen plantsnormally begin with primary steam/hydrocarbon reforming of natural gas tosyngas which is then water gas shifted over HTS and LTS catalysts tomaximize the hydrogen mole fraction of the effluent. Following CO2 scrubbingand methanation to remove unreacted CO, the product hydrogen is thenutilized for hydrocracking, hydrogenation or other service. In ammonia plants,there is a secondary reformer between the primary reformer and the HTSunits for the introduction of the requisite nitrogen. Process conditions in theWGS section are more severe than in hydrogen plants. This is because thedownstream ammonia process is considerably more sensitive to the purity ofthe hydrogen produced. Not only does a lower level of hydrogen reduceammonia production, but also the corresponding higher level of inerts (likeCH4 and CH3OH) increases the purge rate from the synthesis loop. Theprincipal deactivation mechanism for LTS catalysts is poisoning by sulfur andchlorides contained in the process gas. If only a single bed of LTS catalyst isemployed, this deactivation process begins as soon as the catalyst is placed onstream and, normally, within 6–12 months, a rise in CO leakage will bedetected. In an ammonia plant, a rise of 0.1% CO in the LTS converter effluentis roughly equivalent to a production loss of 30 T/day in a 3000 T/day ofammonia plant. To minimize this production loss and maintain a low COleakage for a long period of time, many plants have installed guard beds ofLTS catalysts immediately ahead of the main LTS unit. These beds areusually about 1/4 the size of the main LTS bed and serve to sacrificially screenpoisons from the main bed and to promote additional water gas shift. Ingeneral an ammonia process will tolerate up to 0.4% CO in the LTS effluent



before process economics dictate a catalyst change. Hydrogen plants maytolerate 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 natureof the precursor mixture and its evolution during the preparation steps seemto 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 theWGS reaction. Addition of zinc to copper increased the catalytic activity whichreached a maximum around a Cu/Zn ratio of 0.4. They observed that themethod of preparation of the Cu-ZnO catalyst is extremely important indetermining the catalytic activity. They also established, using x-raydiffraction, that the major constituents of a Cu-ZnO catalyst after use werecopper 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 catalystsare prepared by coprecipitation from the corresponding aqueous solutions ofmetal nitrates with sodium carbonate and having Cu/Zn atomic ratiosbetween 0.4 and 2.0 (42–45). To avoid extensive washing of the filter cakefor a reduction of its Na content, ammonium carbonate or hydroxide as aprecipitating agent was recommended by Sengupta et al. (46). The thermaldecomposition of the resultant, aqueous (Cu,Zn)(NH3)4(HCO3)2 complexes bysteam provides alkali-free Cu-Zn hydroxycarbonates. Petrini et al. (44, 45) hadalso noted that highly active catalysts can be prepared if Al(OH)3 is addedduring the Cu-Zn precipitation.

Gines et al. (47) reported a detailed study of the influence of preparationmethods on the activity and structure sensitivity of the Cu-ZnO-Al2O3 mixedoxide catalysts. Samples were prepared by coprecipitation from aqueoussolutions of the nitrates of Cu, Zn and Al with sodium carbonate at 60uC and aconstant pH around 7 in a stirred batch reactor. The precipitates were filtered,washed with distilled water at 60uC until no sodium ions were detectedand dried at 90–100uC overnight. Finally, the samples were decomposed inair for 8h at temperatures between 400–700uC. Depending on the ratio ofCu, Zn, and Al cations, different hydroxycarbonate phases were formed:malachite[Cu2(OH)2CO3], which is capable of substituting Cu by zinc andis 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 aurichalcitephase for Zn concentrations greater than 40mol%. No trace of copperhydroxynitrates like gherardite was observed. An important observationwas that hydrotalcite was selectively obtained as a single phase only inpreparations using a (Cu+Zn)/Al atomic ratio of 3, the stoichiometric, M2+/M 3+



metal cation ratio in hydrotalcite. The BET surface areas increased with Alcontent. On thermal decomposition, the mixed oxide, CuO-ZnO-Al2O3, wasobtained. X-ray diffraction data revealed the presence of crystalline CuO andZnO. 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 tothe hydrotalcite content in the hydroxycarbonate precursor: higher theamount of hydrotalcite in the precursor, the lower the CuO and ZnOcrystallite sizes in the resulting mixed, ternary oxide. The influence of thehydroxycarbonate precursor structure was preserved throughout the calcina-tion step and manifested itself in a different reducibility of the CuO/ZnOprecursors. After activating and reducing the samples, they were tested in theWGS reaction. Cu (I) oxide is a probable intermediate in the reduction of CuOto Cu metal. Completely reduced Cu clusters on ZnO constitute the active bulkphase 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 750mlSTP/ min (catalyst weight 5 0.5 g). The reaction was carried out at 230uC and1 bar. It should be pointed out that CO2, one of the products of the reversibleWGS reaction, was not included in the inlet gas mixture. A remarkable featureof their catalytic result is that the turnover frequency (number of CO2

molecules produced per surface copper atom per second) was essentiallyconstant, not only when the copper metal surface area was varied between 3–35m2/g Cu, but also when the CuO loading was varied between 30 and 50wt%,the Al/Zn atomic ratios between 0 and 2.5, the copper dispersion between 0.5and 5.0%, and the calcination temperature between 400–700uC, clearlysuggesting that the specific reaction rate is proportional to the copper metalsurface area. Based on these results, the authors concluded that (a) the WGSreaction is a structure insensitive reaction and linearly proportional to thesurface area of metallic copper; and that (b) both the metallic copperdispersion and catalytic activity were related to the amount of hydrotalcitecontained in the precursor precipitate; the higher the content of thehydrotalcite in the precursor, the higher the catalytic activity of the resultingcatalyst.

Contrary results, namely, that the turnover frequency does vary, by anorder of magnitude, when the copper metal surface area was changed from 10to 40m2 /g, for the Cu-ZnO-Al2O3 system had been reported earlier, byChinchen and Spencer (48). These authors had carried out the WGS reactionat 30 bar and their reaction mixture had included CO2 under conditions closerto 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 reactionsincreases with the surface area of metallic copper, there are no reports that,under industrial conditions, the rate of the reaction correlates linearly with



the metallic Cu area over the entire Cu-Zn composition range. While a high Cusurface area is a necessary prerequisite for catalytic activity, additionalfactors like the ‘‘microstrain’’ in the copper nanocrystallites due to thepresence of Zn ions probably affect catalytic activity. The hydroxycarbonateprecursors mentioned earlier probably influence the residual concentrations ofZn in the Cu metal crystallites in Cu-ZnO. Similarly, oxygen vacancies in ZnOformed, for example, during the reduction/activation of the catalyst or duringthe WGS reaction will also influence the catalytic activity indirectly byinfluencing the wetting behavior at the Cu/ZnO interface and, thereby, the‘‘microstrain’’ in the Cu crystallites. Hence, bulk structural changes in theZnO or Cu metal crystallites resulting from the preparation procedures cannotbe 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-Mnspinel oxides in the LTS reaction. They had originally found (50, 51) that Cu-Mn spinel catalysts which were prepared by coprecipitation with NH3, showeda WGS activity comparable to that of Cu-ZnO-Al2O3 catalysts in spite of theirlow surface area. Since Cu and Mn ions may not have coprecipitatedhomogeneously due to formation of the copper amine complex, [Cu(NH3)4]

2+

by the NH3 coprecipitation method, they later prepared (49) their catalysts bycitric acid complex, urea homogeneous coprecipitation or the Pechini method.The last method involves the polymerization accompanied with esterificationof ethylene glycol and citric acid during the precipitation of the hydroxides ofCu and Mn. Higher CO conversions were obtained for samples prepared by thecitric acid method. CO conversion was enhanced with a rise in the calcinationtemperature of the Cu-Mn spinel prepared by the citric acid method. Partialsubstitution of Fe or Al for Mn in the spinel lattice enhanced their COconversion 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 dataset can be obtained from an expression where all kinetic parameters are fitted,for example, to a power law. Such empirical kinetic expressions are essentialin reactor design calculations where it is necessary to have a very accuratedescription of the reaction rate. Different mechanisms, however, can lead to



the same overall kinetic expression. Hence, it is difficult to determine themechanism from an empirical kinetic expression alone. Microkinetic modelsare useful here as they are based on the knowledge about elementary stepsand their energetics. They enable us to estimate surface coverages, reactionorders, and activation enthalpy during reaction conditions. Ovesen et al. (58)had analysed the microkinetics of the WGS reaction under industrialconditions based on a model developed by them earlier (59). The reactionwas 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% Znand 5% Al. Ovesen et al.’s model (58, 59) is based on the surface redoxmechanism:

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, 6400h21 (49).



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 modelare shown in Table 5. The model describes the coverage of surface species inaddition to the overall rate. When this model was tested against measure-ments for an industrial (Cu-ZnO-Fe2O3) catalyst at 1 bar by Van Herwijnenand de Jong (52), a good agreement was found (58). From parallelphysicochemical measurements, it was deduced that the catalyst exposednanocrystallites of Cu (111) facets almost exclusively. The rate-determiningstep was dependent, critically, on the composition of the feed gas mixture. Itwas found that reaction step 2 above is rate limiting in a gas with a low ratio ofwater to carbon monoxide whereas reaction step 7 is rate limiting in a gas witha 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). Todescribe the kinetics of the water gas shift reaction at industrial conditions itwas necessary to include the synthesis and hydrogenation of formate (reactionsteps 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 (highpressure). 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{k2K2

HOHHH

K2H2OH~HH2OHO

r4~k4HOHHO{k4K4

HOHH

K5H2H~

PH2

P0H2

K6PCO

P0H~HCO

r7~k7HCOHO{k7K7

HCO2H

K8HCO2~PCO2

P0H

Note: ki is Forward RateConstant, Ki EquilibriumConstant, and hi SurfaceCoverage of Species i).



sequence for the industrial LTS reaction (58) consists of the steps 1 through11 with steps 2, 4, 7, and 10 as possible slow steps. The formate may bepresent on the surface but, it is not a species in the catalytic cycle for COconversion to CO2. These conclusions from kinetic studies are similar to thoseof a combined kinetic and DRIFTS study of Pt- and Au-based catalysts byMeunier et al. (60–62). It should be pointed out, here, that Ovesen et al.’s rateequations do not consider that co-adsorbed water molecules may influence therate of decomposition of the formate intermediate. It will be interesting toexplore the changes if this issue is taken into consideration (see sections 9 and10). The satisfactory agreement between the calculated exit mole fraction ofCO from the microkinetic model and the experimental exit mole fraction ofCO for Cu-ZnO-Al2O3 is shown in Fig. (21). This model was refined further in alater publication by Schumacher et al. (54). It was established that theadsorption energies for CO and oxygen (the latter arising from H2O) candescribe, to a large extent, changes in the remaining activation and adsorptionenergies through linear correlations. The model predicted well the order ofcatalytic activities for transition metals although it failed to describe theexperimental data quantitatively. The discrepancy was due to the neglect ofadsorbate-adsorbate interactions which play an important role at highcoverages. The model also predicted that the activity of copper can beimproved by increasing the strength with which CO and oxygen are bound tothe surface, thus suggesting possible directions for improving the LTScatalyst.

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



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 thecopper crystallites and poisoning by sulfur and chlorine compounds. Twiggand Spencer have reviewed the deactivation of copper-based catalysts inthe 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, catalyticactivity, when heated above 300uC. Indeed, one of the major roles of Al2O3

is to retard such growth of copper crystallites and function as a texturalpromoter. Details of the mechanism of the thermal sintering of Cucatalysts under hydrogen at elevated temperatures were studied by Tohjiet al. (65) using EXAFS techniques. As the temperature was increased inhydrogen, a quasi- two-dimensional layer of copper metal epitaxiallydeveloped over the ZnO support below 127uC. Between 127–230uC, smallcopper metal clusters dispersed over ZnO start to appear. Above 250–300uC, the small clusters fuse to give larger copper metal crystalsagglomerated on the support. Since these catalysts begin to lose coppersurface area and catalytic activity also above 250uC, it is reasonable toassume that the active sites for the WGS reaction are associated with thesmall copper clusters and their concentrations are diminished when thesmall crystallites grow into larger ones. Thermal sintering leads to theirgrowth and consequent catalytic deactivation (66).

4.4.2. Sulfur Poisoning

The second major cause of deactivation of these catalysts is poisoning bysulfur compounds present in the reaction gas stream. The LTS, Cu-ZnO-Al2O3

catalyst operates at 190–250uC, a temperature sufficiently low whereinthermodynamics favors strong adsorption of poisons. Sulfur is a powerfulpoison for Cu, as indicated by the change in enthalpy of the sulfidationreaction [Eq. (28)] (64):

2 CuzH2S[Cu2SzH2 DH~{59:4 kJ=mol !28"

Sulfiding of copper, hence, occurs. The corresponding equilibrium constant is(64) about 1 6 10+5. Sulfur, accumulating on the surface, blocks the pores andthe active sites leading to catalytic deactivation. To retain the long term



activity of copper catalysts, it is usual to maintain gas phase sulfurconcentrations below 0.1 ppm of S. In addition to keeping gas phaseconcentrations of sulfur low, the ZnO component of the catalysts is alsoengineered during catalyst manufacture to divert the sulfur away from thesmall Cu crystallites and absorb it in ZnO as ZnS. It is essential to keep thecrystallite size of ZnO as small as possible to accomplish this absorption. Thereaction of H2S with ZnO [Eq. (29)] is quite exothermic (66) and proceedsreadily:

ZnO s! "zH2S g! "[ZnS s! "zH2O g! " !29"

DHu 5 276.7 kJ /mol; DSu 5 23.0 J.mol21K21.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. Theequilibrium constant at 500K is 7.4 6 10 7. The reaction is strongly favoredthermodynamically.

4.4.3. Chloride Poisoning

Chlorine compounds, like HCl, form low-melting cuprous chloride (m.p. 5430uC) on reaction with copper in the Cu-ZnO catalyst. The ZnCl2 formed alsohas a low melting point (283uC). Their formation is favored thermodynami-cally (64) under the WGS reaction conditions [Eqs. (30, 31)]:

Cu s! "zHCl g! "[CuCl s! "z0:5 H2 g! " !30"

DHu 5 2 43.5 kJ /mol, and

ZnO s! "z2HCl g! "[ZnCl2 s! "zH2O g! " !31"

DHu 5 2 121.8 kJ /mol; DSu 5 117.2 J.mol21K21.These mobile chlorides facilitate the movement and sintering of copper as wellas the ZnO crystallites on the catalyst surface. The limits on HCl content toavoid catalyst poisoning are more severe than for H2S poisoning, on the orderof 1 ppb. Unlike the case of sulfur poisoning, the ZnO cannot offer anyprotection in the case of chloride poisoning.

In addition to the major poisons, sulfur and chloride, the Cu-ZnO catalystsare also deactivated by the presence of As, trivalent phosphorous, silica, andtransition metals like Fe, Co and Ni, in the feed stream. Due to their lowtemperature of operation, Cu-ZnO catalysts do not form significant amounts ofcoke when operated with purified feedstock (65).



5. SULFUR TOLERANT WGS CATALYSTS

The sulfur levels in natural gas or light petroleum naphtha are in the range of5–50 ppm and conventional hydrodesulfurisation of the feedstock with Co (Ni)-Mo- alumina catalysts is used before steam reforming them with nickel- basedcatalysts. The latter are deactivated in the presence of sulfur. H2S can beremoved from natural gas as well as hydrodesulfuriser effluents by reactionwith ZnO at 370uC. Other sulfur compounds can be removed from natural gasalso by absorption at ambient temperatures on activated charcoal (loaded withcopper) or molecular sieves. The efficiency of these absorption systemsdepends both on the type of sulfur compounds and on the amount of highmolecular weight hydrocarbons in the natural gas. Low boiling sulfurcompounds, like COS, are not strongly absorbed and condensable hydro-carbons can rapidly saturate the absorbent. Catalytic hydrodesulfurisationcan remove COS. The removal of H2S by absorption in a hot ZnO bed is usuallynot complete. Approximately 50 ppb H2S slips through and enters thereformer upstream of the WGS reactor. After the volume expansion due tothe reforming reaction, the resulting H2S concentration in the gas enteringthe WGS reactor is about 10 ppb. Due to the high temperatures in the reformerand the low capacity of modern Fe2O3-Cr2O3-Cu HTS catalysts for sulfurabsorption, nearly all this residual sulfur exits the HTS stage and is removedfrom the syngas by the Cu-ZnO-Al2O3 LTS catalyst located downstream.Unlike Fe2O3-Cr2O3 catalysts, the Cu-ZnO-Al2O3 catalysts are adverselyaffected by the presence of sulfur compounds in concentrations greater thanabout 0.1 ppm. The deactivation is irreversible even when the sulfur isremoved from the feed gas stream. Normally the Cu-ZnO-Al2O3, LTS catalystreactors are designed for a space velocity of 1000–2500 h21 to take into accountpoisoning by this sulfur. The actual catalyst volume in the reactor representsapproximately three times the volume needed by the kinetics. The Cu-ZnO-Al2O3 catalysts, thus, serve also as a total sulfur absorber protectingdownstream 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 thesyngas is generated by the gasification and partial oxidation of heavy fuel oil,tar sands, oil shale, coal, coke or biomass. Syngas from these raw materialscontain much larger concentrations of CO (up to 50%) (Table 1) and sulfur (upto 3% wt) (5). In such cases, the Fe2O3-Cr2O3 catalyst had to be used as theonly WGS catalyst; conventional, Cu-based LTS catalysts cannot be used atthe 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 itsactivity is much lower than in the oxide state. It is, hence, necessary to eitheroperate at higher H2O/CO ratios or remove the sulfur compounds from theprocess gas over sulfided Co-Mo-alumina catalysts before it enters the HTS



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 sulfurcompounds, these Co-Mo-alumina catalysts also serve to remove CO (by theWGS reaction) from the process gas and, thus, serve as sulfur-tolerant, sourgas shift catalysts. In fact, these catalysts are active mainly in the sulfidedform. When such sour gas catalysts based on Co-Mo sulfides are used, thepreferred minimum inlet sulfur in the feed for acceptable perfomance is about300 ppm. If these catalysts are adequately presulfided before use, then theycan operate satisfactorily even in feed streams that contain H2S at a level aslow as 35 ppm. Non-sulfided Co-Mo catalyst exhibits very little WGS activity.Commercial sour gas converters with Co-Mo catalysts operate in thetemperature range of 250–350uC and at pressures from atmospheric to 40bar. Typical process conditions in a Co-Mo- based sour gas shift catalyticreactor in a H2 plant using Texaco partial oxidation process to generate syngasfrom heavy oil are shown in Table 6. The syngas from the partial oxidationreactor contains 0.25% sulfur. The sulfided Co-Mo catalyst is deployed in 3beds. The CO content is reduced from 46% (vol) at the inlet to the first bed to1% at the exit of the third bed. It should be noted that all the cobalt moly-basedsour 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 atequilibrium. At high operating pressures and relatively high steam/dry gasratio, the resulting COS concentration is usually well below 0.1 ppmv.However, under certain circumstances, the COS concentration can be muchhigher and downstream COS hydrolysis has to be considered. One of theadvantages of the sour gas shift reaction using sulfided, cobalt molybdenumcatalysts is that they operate at much lower temperatures (250–350uC) thanconventional 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



Therefore, the water gas shift equilibrium is favored resulting in lower outletCO concentrations. The sour gas shift catalysts also need much less steam forthe same or even higher CO conversion since the possibility of metal formation(and accompanying methanation and Fischer-Tropsch reactions) from thesesulfided Co-Mo catalysts is remote. However, these catalysts operate at lowerspace velocities and, hence, need about 20% more catalyst than thecorresponding iron oxide- chromium oxide HTS catalysts. Additionally, theyalso 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 oilgasification.

Addition of alkali to these sulfided catalysts promotes their WGS activity(67, 68). It has also been reported (69, 70) that Co-Mo-based catalystspromoted by Ti improve the WGS activity of the former in the presence ofsulfur compounds. The sulfided Co-Mo catalysts are not affected by poisons,like NH3 or HCN when they are present in low concentrations (below about0.5%). Phenol is a catalyst poison but the rate of deactivation is relatively lowat low concentrations of phenol. Phenol poisoning is reversible and thecatalyst can be regenerated with steam - air regeneration. A high benzeneconcentration (above 10%) tends to decrease the catalyst’s activity. Chloride isa major poison for these catalysts. Even at a 1–2 ppb level, chlorides have anadverse effect on catalyst performance. The effect of chlorides is cumulativeand catalyst regeneration will not restore catalyst activity.

Mellor et al. (72) reported novel Co-MnO and CoCr2O4 catalysts tolerant tosulfur up to levels of 220 ppm under WGS reaction conditions. However, incoal-derived process gases containing between 0.25 and 0.3mol% sulfur, andat a reaction temperature of 400uC, the Co-MnO catalyst deactivated rapidlyand irreversibly with formation of bulk Co9S8 and a surface manganese sulfidespecies. The CoCr2O4 catalyst deactivated only partially under similarconditions. Bulk sulfiding of the CoCr2O4 catalyst to CoCr2S4 occurred at550uC and this catalyst gave near equilibrium CO conversions in the WGSreaction. A pre-sulfided cobalt chromium catalyst demonstrated typical sulfurdependent mechanistic characteristics, with a maximum activity above 400uC(72). It may be noted that these sulfided Co-MnO and CoCr2O4 catalysts areactive in the WGS reaction only at temperatures considerably higher than thesulfided cobalt moly catalysts. They are, hence, under a thermodynamichandicap, 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 wasknown for many decades, their high price precluded their adoption in



commercial practice. The need for compact catalyst beds in automobileapplications of fuel cells provided an impetus for intense research in this fieldduring the last decade. Based on the experience from earlier studies ofautomotive exhaust catalysts, gold, platinum and other metal –on- partiallyreducible metal oxide supports have been the frontrunners in this area (73–115). It should, however, be noted that auto exhaust catalysis operates underan oxidative atmosphere above 400uC while the WGS fuel processingenvironment is a reducing atmosphere at temperatures between 180–450uCand wherein the partial pressures of H2 and CO2 are much greater. Inaddition, noble metals exhibit lower activity in WGS reactions below about250uC which limits the CO exit levels to about 0.5–1.0wt%. Of the manycatalysts that have been studied, precious metals (mainly Pt, Rh, Ru, Au, andPd) deposited on partially reducible oxides (ceria, zirconia, titania, iron oxides,and mixed oxides of ceria, like ceria- zirconia) have been the mostinvestigated. These catalysts are quite active in the 250–400uC range. Pre-reduction of these catalysts is not required and they can be safely exposed toair during cool down or start-up without significant loss of performance, acrucial requirement of fuel processor catalysts. The reaction rate for thesecatalysts is close to zero order for CO and, hence, advantageous in driving thereaction 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 differentformulations, combining precious metals with partially reducible oxides, havebeen 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 thatthe precious metal-based catalysts are bifunctional; both the metal andsupport have a significant influence on the overall performance. Ceria andceria- zirconia have been explored extensively as supports for LTS catalysts inthe last decade. The incorporation of Zr improves the thermal stability(against sintering), oxygen storage capacity and WGS activity of the ceriacrystallites. The bulk structure of zirconia-doped ceria is well known. Severaldifferent tetragonal phases of varying degrees of stability can be formeddepending upon the Zr doping level, preparation technique, crystallite size,shape and thermal history. The metastable ‘‘t’’ phase is commonly formedwhen Zr doping ranges from 15–28wt%. It has a tetragonal oxygen ionsublattice and a cubic Ce/Zr fluorite sublattice. Another metastable t’ phase,that can form between 28–63wt% Zr doping, is tetragonal (P42/nmc spacegroup) on both sublattices having a c/a lattice parameter ratio greater than



unity. Fluorite structures, such as CeO2, commonly occur as crystallites thatmaximize the most stable {111} surface face.

Apart from the noble metals and Au, other transition metals, such ascobalt and nickel have also been investigated as WGS catalysts. They,however, cause methanation of CO to CH4 under typical WGS reactionconditions, especially, below 350uC [Eq. (17)].

A meaningful comparison and rating of all the reported catalysts isdifficult since the various authors had prepared their catalysts by differentmethods, with different catalyst precursors, and had evaluated those usingdifferent compositions of the feedstocks and at different reaction conditions.

In an attempt to bring some order in the picture, Thinon et al. (93) haverecently screened about 20 metal-on-oxide catalysts for the WGS reactionunder identical conditions using a model reformate as the reaction mixture.They used commercial high-throughput equipment consisting of 16 parallelreactors set-up to compare the activity and selectivity of these bifunctionalWGS catalysts. The catalysts were prepared by impregnation of the supportswith a solution of the corresponding metal precursors. The supports used werecommercial metal oxide powders with surface areas between 30–80m2/gexcept Fe2O3 (7m2/g). The impregnated material were dried and calcined at400uC. 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 witha feed of 10% CO and 20% H2O diluted in Ar to investigate the forwardreaction 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 catalystswere found to promote methanation reactions. The salient features oftheir 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



defined as:

Activity mmol=kgcat:s! "~FCO| XCO=Wcat! ", !32"

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 thisreaction is the higher temperatures needed to partially reduce the former. Inan attempt to address this problem, Gonzalez et al. (94) were able to improvethe low temperature activity of Pt –TiO2 catalysts by incorporating ceria in thesupport. Pt supported on ceria - modified TiO2 catalyst showed better thermalstability and lower temperature reducibility compared to TiO2 and a higherWGS 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, GHSV5 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).



Pt-TiO2 (as well as those of Pd- and Ir-TiO2) was also improved by the additionof Re by Sato et al. (95) (Figure 23). Catalytic activities were evaluated bythem, in a closed gas circulation system, as an initial H2 formation rate in 10Torr 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.’sconclusions (81). Two important features were observed: (a) possible formationof bimetallic surface clusters with Re in the case of Pt and Pd; and (b)anchoring and ‘spacing’ of metal nanoparticles by highly dispersed Re overTiO2 in the case of Ir. Panagiotopoulou et al. (96) studied the influence of thesource of the TiO2 support in Pt-TiO2 catalysts and found that the WGSactivity depended strongly also on the phase composition and particle size ofthe TiO2 support; the activity increased with increasing reducibility of TiO2.Both TPR and Raman spectroscopy data indicated that the titania could bereduced by H2 or CO at temperatures as low as 150uC. Based on their results,the authors suggested that the titania surface undergoes successive reductionand oxidation by adsorbed CO and water, respectively, thereby cyclingbetween TiO22x and TiO2. Sato et al. (97), from CO adsorption, X-rayphotoelectron spectroscopy, in situ IR spectroscopy of adsorbed CO moleculesand catalytic studies of the Pt-Re-TiO2 system, observed that a bimetallic Pt-Re alloy is formed under reaction conditions and that an additional surfacecompound 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 2wt% Pt-Re/TiO2 (100uC), 1wt% Pd – Re/TiO2 (200uC) and 1wt% Ir-Re/TiO2 (100uC) (95).



mixture of Pt and ReOx. This surface compound, probably, accounts for thegreater activity of Pt-Re-TiO2. The binding energy of the Pt electrons is lowerin Pt-Re-TiO2 than in Pt-TiO2 suggesting that the Pt crystallites are slightlynegatively charged. CO is adsorbed more strongly on Pt in Pt-Re-TiO2 than onPt-TiO2. This is understandable since a more negative Pt will transferelectrons more easily to the antibonding orbitals of CO, thereby stabilizing COin the adsorbed state. IR spectra of CO on Pt-TiO2 reveals only linearlyadsorbed CO. On Pt-Re-TiO2, bridged CO as well as formate ions are seenadditionally 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 toCO2. 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 avery active and selective catalyst for the WGS reaction, they deactivated withtime on stream (Fig. 24). Catalyst deactivation during the WGS reaction wasalso a problem with Pt- and Pd- ceria catalysts (86–87, 101). Wang et al. (86)investigated the mechanism responsible for the irreversible deactivation ofceria- supported precious metals for the WGS reaction through acceleratedaging tests. They showed that deactivation of Pd- ceria occurs more rapidly at400uC than 250uC when operating with an integral reactor in 25 Torr each ofCO and H2O. By heating a fresh catalyst in H2, H2O, CO or CO2, it wasdiscovered that deactivation occurs due to the presence of CO. Similarconclusions 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 22h 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 GHSV5410,000h21, mcat 551mg (98).



metal dispersion by CO adsorption and by X-ray diffraction show (86) thatdeactivation on Pd- ceria and Pt - ceria catalysts was due to loss of noble metalsurface area. Pd dispersion values, for example, decreased from 23% (freshcatalyst) to 3% after a 10hr treatment in CO at 400uC causing the COconversion to decrease from 25% (fresh catalyst) to 3% after treatment for only2hr in CO at 400uC. The corresponding dispersion values after similartreatments in H2O and CO2 were 25 and 16, respectively. The CO conversionswere 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 withceria crystallite sizes ranging from 7.2 to 40nm and Pd loadings of either 1 or6wt% demonstrated that the rates were strictly proportional to the COadsorption capacity and, hence, Pd surface area (Fig. 25). Later, using in-situFTIR spectroscopy in the OCO stretching region, Gorte et al. (106) observedstrongly – held, carbonate-like species on the surface, formed from CO. Thesewere 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 notreduce catalytic activity, and (b) high temperature oxidation also did notrestore the activity of deactivated catalysts, they ruled out the over-reductionof ceria as a contributory factor to the deactivation of the catalyst. Justbecause 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, 1wt%Pd/ceria, withceria calcined at 600uC, &, 6wt%Pd/ceria, with ceria calcined at 600uC, D, 1wt% Pd/ceria,with ceria precipitated and calcined at 350uC; $, 1wt% Pd/ceria, with ceria calcined at800uC; m, 1wt% Pd/ceria, with ceria calcined at 950uC (86–87).



species in question is causing deactivation. First, one has to see the increase ofthat 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 theintermediate (e.g., formate, carbonate) to increase because the metal mayassist in decomposing that intermediate and any loss in the metal’s interactionwith the support (e.g., ceria) would cause the inventory of the intermediate tobuild up (since it would be reacting more slowly with loss of the metal’sinteraction). So, while the first point is necessary, the second point explainswhy the first point is not sufficient. Zalc et al. (109) observed a strongdependence of the deactivation rate on the presence of hydrogen in the feedand 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 potentialcause of deactivation was the growth of ceria crystallites and occlusion of Ptcrystallites, and the consequent decrease of the BET surface area during thereaction. While all the above-mentioned factors may potentially lead tocatalyst deactivation, it is difficult to extrapolate the validity and relevance ofthese conclusions to the WGS reaction in industrial reactors in view of thewidely different methods of catalyst preparation, activation and reactionconditions used by the various authors. Noble metal- based catalystscontaining a combination of Pt, CeO2 and TiO2 have, recently, been claimedto be superior WGS catalysts (111). Baidya et al. (112) have compared thestructure, reducibility and catalytic activity of various solid solution oxidescontaining 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 stepsolution combustion method. Their fluorite structure and solid solutionformation were confirmed by XRD Rietveld calculations. Temperatureprogrammed reduction and XPS study of Ce12x TixO2 (x 5 0.00–0.04) showedcomplete 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 thereducibility of CeO2. Ce0.84Ti0.15 Pt0.01O22d crystallized with a fluoritestructure and Pt was ionically substituted with 2+ and 4+ oxidation states.The amount of hydrogen adsorbed at 30uC over Ce0.84Ti0.15Pt0.01O22d was twoorders of magnitude larger than that over pure 8nm Pt metal crystallites. COand 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 thePt2+/Pt0 and Ti4+/Ti3+ redox couples in addition to Ce4+/Ce3+ were heldresponsible for the higher reducibility and catalytic activity in the oxidation ofCO (Figure 26). As may be seen from the figure, the Ce0.84Ti0.15Pt0.01O2-d has a



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 theoxygen ion vacancy concentration and the consequent increase in the oxygenstorage capacity of the material. Parallel TPR and XPS measurements alsoconfirmed the greater reducibility of the Pt-titania-ceria samples. XPS dataalso indicated the presence of Pt in the ionic (Pt2+) state. In view of the knowndeactivation of both the Pt-TiO2 and Pt-CeO2 samples during prolonged WGSreactions, 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 fuelcell applications was investigated by Zalc et al. (109) who prepared a variety ofPt-ceria WGS catalysts and tested them in the range 250–450uC under feedand reaction conditions typical of a reformer outlet. They observed first orderdeactivation. Virtually identical deactivation rates were found for all the Pt-ceria catalysts tested. Significantly lower deactivation rates were observedwhen hydrogen was not present in the feed. Attempts to rejuvenate thecatalyst by heating under steam and under air were unsuccessful (109).Catalyst deactivation is still a major obstacle in the commercialization of WGScatalysts for fuel processing. The goal for most programs is 40,000 hours ofcatalyst life. This is an ambitious goal of about 4.5 years of continuousoperation. There are indications, from studies in the industry that addition ofother rare earth elements like lanthanum, praseodymium, etc., to the ceria-zirconia support can reduce, to some extent, the agglomeration of ceriacrystallites 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).



catalysts is the relatively larger formation of hydrocarbons, includingmethane, 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 lowertemperatures? The known higher Fischer-Tropsch activity of the Pt groupmetals (compared to copper) for the synthesis of hydrocarbons from CO andH2, at 200–350uC, is a handicap in the WGS reaction. We may, perhaps, haveto reduce the Fischer-Tropsch activity of the noble metal component to thelevel of copper without, however, sacrificing its greater catalytic activity inWGS reactions. Incorporation of the Pt in the lattice sites of the partiallyreducible cerium oxide (as in the work of Baidya et al. (99)) and preserving thePt in the ionic state under the WGS reaction conditions, may be one potentialsolution since ionic Pt, while active in redox reactions (109), is not known topossess Fischer –Tropsch activity.

7. Au-BASED WGS CATALYSTS

The low WGS activity of Pt-, Rh- and Pd-based catalysts below 250uC had ledto increased interest in more active catalysts to take advantage of thefavorable thermodynamics at these temperatures. In the past 10 years,supported gold catalysts with remarkably high activity for the WGS reactionhave been discovered (116, 117). Gold catalysts can offer some advantages inthe 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 noexceptional pre-treatment before use. Figure 27 illustrates the high activity ofgold when compared to the Pt and the Cu-Zn-based catalysts. First developedas a low temperature catalyst for the preferential oxidation of carbonmonoxide (in a mixture of CO and H2) by Haruta et al. (118), it was soonrecognized that the catalytic activity was high only when the particle size ofgold was very small, of the order of 1–5 nanometers (119). Extension of thestudies 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 2nmparticles) (120). The dissociative adsorption of water on the nano gold particlesfollowed 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 particlessupported on other supports including TiO2 (123–125), and ceria (126–128)were also found to lead to catalysts active at temperatures below 200uC. Basedon their studies of Au-TiO2, Andreeva et al.(120) postulated the existence of



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 orderto estimate the relative contributions of the metallic and ‘‘ionic’’ gold tocatalytic activity, Fu et al. (137) measured the rates of the WGS reaction afterleaching out the metallic Au from Au-ceria with NaCN. The rates andapparent activation energies were the same, before and after leaching withNaCN, highlighting the importance of the fraction of the gold (presumablyionic) that was not leached out by the NaCN treatment. The amount of such‘‘ionic’’ gold inserted in ceria was found to increase with decreasing crystallitesize of ceria. Large crystallites of ceria did not retain any gold. Incorporation ofgold also increased the stability of the ceria microcrystallites. Catalyticactivity in the WGS reaction was also reasonably stable. When this idea wasextended and gold ions were stabilized in the framework of an ionic lattice, asin Au2Sr5O8 or La2Au0.5O4 (128), not only was the sintering reduced and thethermal stability of the catalyst increased but the catalytic activity was alsoenhanced. It should be mentioned here that recent studies indicate that ionicgold is unlikely to be present, in the steady state reducing conditions duringWGS reaction, especially, at higher pressures (129). In-situ time-resolved X-ray diffraction and X-ray absorption spectroscopy were used by Rodriguez etal. (129) to monitor the behavior of nanostructured Au-CeO2 catalysts underthe WGS reaction. Above 250uC, a complete AuOx ) Au transformation wasobserved with high catalytic activity. Photoemission results for the oxidationand 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).



sites responsible for the WGS activity at high temperatures. They suggestedthat the active sites in Au - ceria catalysts involve pure gold nanoparticles incontact with O vacancies on the ceria. The role of cationic Au3+ and nonionicAu0 species in the LTS reaction over Au-ceria catalysts was also studied byKarpenko et al. (130–131) by comparing the reaction behavior of a cyanide-leached catalyst with that of non-leached catalysts. Using rate measurementsas well as in situ spectroscopic and structure-sensitive techniques, they foundthat, based on the Au mass balance, cyanide leaching removed all the Auexcept for ionic Au3+ species, and that leaching resulted in pronounced decayof the catalyst mass- normalized activity to 1–25% of that of a non-leachedcatalyst. The extent of the activity loss strongly depended on the post- leachingtreatment of the leached catalyst. Both the catalyst pretreatment after theleaching and, in particular, the WGS reaction resulted in considerablereformation of Au0 aggregates and metallic Au0 nanoparticles as indicatedby 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 nanoparticlescontribute predominantly to the WGS activity, and (b) cationic gold has anegligible contribution to the WGS activity in the steady state. Au ions are,expectedly reduced to Au0 atoms in the reducing atmosphere during the WGSreaction. Combining TEM, XRD, XPS, DRIFTS and activity studies, theyconcluded (131), further, that for reaction up to 200uC, catalyst deactivationwas dominated by the formation of stable adsorbed monodentate carbonatespecies. The influence of other effects, such as catalyst reduction/ oxidationwere less significant.

Au-CeO2 and Au-CeO2-Al2O3 catalysts were also investigated by Andreevaet al. (132) who compared samples, prepared by a mechanochemicalactivation, with those prepared by a conventional coprecipitation; the formerwere more active. This was attributed to the smaller size of the Au and ceriacrystallites in the former. The main role of alumina was that of a texturalpromoter in stabilizing the Au and ceria crystallites against agglomerationduring the WGS reaction and, thereby, maintaining a high catalytic activity inthe steady state. The addition of alumina to ceria results in smaller ceriacrystallites and, consequently, an increase in the number of oxygen vacanciesand oxygen storage capacity of ceria, as estimated from temperature-programmed reduction experiments. A correlation was found between WGSactivity and the oxygen storage capacity of the samples (132).

In an attempt to gain insights into the reactivity of supported Aunanoparticles, Janssens et al. (133) applied density – functional calculations,adsorption studies of CO and oxygen on single crystal surfaces and WGSactivity measurements on well-characterised, supported gold particles. Theyattributed the increasing activity of supported Au catalysts, with decreasing



Au particle size, to the increasing number of low - coordinated Au atomspresent in such small particles. Their DFT calculations indicate thatadsorption of CO and oxygen on the densely–packed surfaces (which exposeAu atoms with high coordination numbers, like 8) is generally difficult orthermodynamically not possible. On the other hand, adsorption was favoredon Au atoms with a lower coordination number. The effect of the Aucoordination number on the adsorption strength of CO and oxygen was foundto be larger than other electronic effects or strain and was, therefore, a crucialparameter for the catalytic activity. The smaller particle size and supporteffects influence the catalytic activity only indirectly through their influenceon 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 thelowest coordination numbers) were the most reactive. During the synthesis ofthe various supported Au catalysts, the properties of the support surface (i.e.,quality and number of nucleation sites) influence the size, dispersion andmorphology of the Au nanoparticles, and, thereby, the concentration of active,low coordinated sites. Moreover, during catalytic operation, the metal-supportinterface energy, which is influenced largely by the support, has a significantinfluence 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 thesemolecules 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).



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 electronicstructure. The low- coordinated Au atoms have high-lying metal d states,which are in a better position to interact with the adsorbate valence state thanthe 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 incatalytic reactions. The trend that the CO adsorption strength on Au increaseswith 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 desorptionaround2103 to283uC (in Fig. 29) to CO adsorbed on defect or corner sites, thedesorption around 2 123uC to CO adsorption on the (110)-(162) surface, and

Figure 29: Temperature programmed desorption of CO on various Au samples (133).



the desorption around 2163 to 2143uC to CO adsorbed on the edge of thenanoparticles, or step sites on the single crystal surfaces. For particles with agiven shape, the number of corner atoms per gold particle is independent ofthe particle diameter. Hence, for a given amount of gold in the catalyst, thelarger the number of metal particles (i.e., larger the dispersion), the larger willbe the number of corner atoms and, hence, the catalytic activity. Thisexplanation (133) is fundamentally different from the quantum size effectsconventionally invoked and which ascribes the higher catalytic activity ofsmall Au particles to changes in electronic structure as the particle sizedecreases. Though quantum size effects are important for very small particlescontaining only a few atoms (134), they seem not to be necessary to explain thecatalytic effect for supported Au particles larger than about 1.5 nm. It may benoted that Bond and Thompson had also pointed out earlier that significantchemisorption of molecules like CO occurs only when an adequate number oflow-coordination surface Au atoms are present (135).

One of the major drawbacks of the gold catalyst is catalytic deactivationduring 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 WGSreaction and physical blocking of those sites at which participation of thesupport is essential for high activity. Such species include carbonates,bicarbonates, formates etc. In attempts to prolong the catalytic activity ofsupported gold catalysts (for CO oxidation), Moreau and Bond (136) haverecently found that inclusion of Fe(OH)3 or lanthanum oxide during thepreparation of Au catalysts supported on ceria and zirconia, gave betteractivity and much improved stability with time-on-stream. This effect waslinked to the ability of the FeOx phase to provide hydroxyl groups, stable atthe reaction temperatures, that are needed for the catalytic action and to formanion vacancies (by replacement of a tetra- by a trivalent metal cation) atwhich O2 or H2O molecules can chemisorb. The effect is similar to those seenwhen La3+ or, Fe3+ are dispersed in the ceria lattice (136).

The relatively high activity of gold catalysts has been challenged. Jacobset al. (138) reported that a 5%Pt- ceria catalyst was much more active than a5% Au-ceria catalyst. They attributed their distinctive results, essentially, to(a) the higher content of Pt in their catalysts, (b) the complete reduction ofplatinum oxide, and (c) the ‘‘careful activation’’ of their samples. Differences inthe details of catalyst preparation and determination of the dispersion of noblemetals on partially reducible oxides (see Section 10), adopted by the variousresearchers, perhaps, explain such differences. In addition, differences in thecomposition of the feedstocks used to evaluate the catalytic activities of theAu- and Pt-based catalysts and the reaction conditions will also influence theconclusions. The two metals, Pt and Au, respond differently (116) to changes



in the concentrations of the reactants. The power law dependencies, on theconcentrations of the reactants and products, are different for Pt and Au (116).Thus, under certain conditions, the order with respect to carbon monoxide isnegative for Pt but positive for Au. The rate for Au, for example, was given by:

R~k CO# $0:7 H2O# $0:6 CO2# ${0:3 H2# ${0:9: !33"

The positive order with respect to CO reflects the weak adsorption of CO on Auin contrast to Pt where CO is much more strongly adsorbed, at least, up to200–250uC, accounting for the negative order with respect to CO on Pt. Inaddition to catalytic activity, the stability of the catalyst during prolonged useunder various process conditions is also of major importance. Here, Pt-basedcatalysts are quite rugged and have a distinct advantage over Au-basedcatalysts. The performance of the latter is more sensitive to conditions ofstorage and operation. In view of the importance of resolving this issue (therelative superiority of Pt- and Au- based catalysts) in the design of fuelprocessors for fuel cells, a quantitative comparison of their kinetic behaviorusing industrial feedstock and under identical, but, realistic conditions, isdesirable.

Can the performance of gold-based catalysts be improved? Two approacheshave been taken in the last few years: (a) Improving the metal function bycombining Au with another metal (like Pt) to form bimetallic catalysts, and (b)incorporating promoters in the ceria support. Juan et al. (139) have reportedthat when Pt, Pd, W, or Ni is added to Au – ceria, there is a synergistic effectand the resultant bimetallic catalysts are more active than Au-Ceria or Pt-Ceria. The catalysts were tested in the temperature range 150–500uC, with aH2O/ CO ratio of 13.5 and GHSV of 52,000h21 (Fig. 30). The Au-Pt -Ceriaclearly displayed a much higher activity compared to Au-Ceria at the same

Figure 30: CO conversion over Au-M bimetallic promoters on CeO2 (139).



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 thesynergy between the two metallic functions (Au and Pt), the atomic loadingsas 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 bimetallicgold and other metal- supported catalysts are expressed in terms of COfractional conversion without using their normalized specific activities(activity per metal site) since the interactions between different metals overthese bimetallic gold catalysts is not, sufficiently, clear to differentiatebetween CO adsorption on Au or Pt or both. The total number of metal sitesper gram of the catalyst (in mmoles per g), evaluated via CO chemisorptionwith 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, a3wt%(Au-Pt)- CeO2 displayed the best catalytic activity in the WGS reaction.Very interestingly, the WGS activity was strongly correlated with the surfacereducibility data from temperature-programmed reduction experiments(Figure 31). The reducibility of the catalysts, in turn, depended on themodified local electronic band structure of the promoted ceria. CeO2 shows twodistinct reduction peaks, one at 440uC (assigned to reduction of surfaceoxygen) and another at 800uC (reduction of bulk ceria oxygen).The incorpora-tion of gold and the metallic promoters in the ceria catalyst facilitated thereduction of surface oxygen at lower temperatures while the reduction of bulkoxygen remained unchanged. Among the different promoters studied by them,the Au-Pt-Ceria combination was more effective than Pt, Pd, or Au alone onceria in giving higher WGS activity at lower temperatures. The temperature ofthe surface oxygen reduction peaks, in Au-Pt, was 120uC. The correspondingtemperatures in the case of Pt and Pd supported on ceria were 130uC and135uC, respectively. It may be noted that Fu et al. (140) and Andreeva et al.(141) had also reported similar low temperatures (around 150uC) in theTPRprofile 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 temperatureslower than even noble metals. The ranked order (139) of the lowering intemperature (Fig. 31) of the first reduction peak (from surface oxygen lossfrom 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 ofthese samples.

Ceria is a n-type semiconductor whose electronic band structure can bemodified by promoters. From the UV diffuse reflectance spectra of thesamples, a clear, blue shift of the absorption edge (O2p R Ce 4f) of the ceriaupon doping with Au or Au-Pd was observed. The degree of band gap wideningwas found to relate to different bimetallic promoters (139). The order of



bandgap widening was remarkably similar to both the orders of surfaceoxygen reducibility and WGS activity suggesting an electron transfermechanism at the interface between ceria and the metallic componentsfacilitating the redox transformations occurring in ceria (139):

2Mz2 O# $u2 Mzz2# $zO2z2e, and !34"

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).



2ez2Ce4zu2 Ce3z, !35"

where M represents a metal, like Pt, Pd, Au etc., [ ] an oxygen anion vacancyand [O] denotes an oxygen anion on the surface.

One of the problems in comparing, quantitatively, the specific WGSreaction rate data over precious metal- reducible oxide catalysts from differentlaboratories is the absence of a reference method for determining thedispersion of the metal on the support. In the case of a metal dispersed overa non-reducible oxide, like Pt-silica or Pt-alumina, the experimentalprocedures (temperature of reduction, temperature of H2 or CO adsorption,etc.) and the stoichiometry of the chemisorption for determination ofreproducible and accurate metal dispersion values are accepted and havebeen standardized. An important difficulty originates when the support canalso adsorb (and, even react with) the probe molecules, H2 or CO in quantitiescomparable or even more than the metal itself. Such is the case for redoxsupports such as ceria, titania or ceria-zirconia on which hydrogen spilloverprocesses (from the metal to the oxide) occur easily in the presence of ametallic phase, like Pt or Au. Perrichon et al. (144) determined the Ptdispersion by chemisorption of H2 and CO in a series of Pt-ceria-zirconiacatalysts covering the full range of composition between ceria and zirconiausing volumetric techniques and FTIR spectroscopy. Using IR spectroscopy ofadsorbed CO to distinguish the CO adsorption on the Pt surface from that onthe ceria-zirconia support allowed them to validate a protocol of hydrogenchemisorption for measuring the metal dispersion. A first method is based onthe CO adsorption isotherm analysis, using IR spectroscopy as the detectiontool. Apart from the quantitative analysis of the adsorbed/desorbed gas phase,this method also gives information about the coordination mode of the COmolecule on the metal particle, linear or bridged. The second hydrogenchemisorption method is based on the use of a double isotherm of hydrogenadsorption at 278uC , this low temperature being required to suppress thehydrogen spillover from the metal to the ceria- zirconia support. Theirreversible adsorption of H2, measured either at saturation or by extrapola-tion to zero pressure, leads to the most reliable metal dispersion values whichcan be independently confirmed by FTIR spectroscopy of adsorbed CO. Ptdispersion, 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



electricity. The WGS reaction is a critical step in reducing the COconcentration in such H2 feed streams, especially in low temperature fuelcells 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 lessconstrained by the need for compaction and fast response as it is forautomotive applications. Compared to conventional industrial WGS plants forthe generation of hydrogen, however, a reduction in reformer and water gascatalytic reactor sizes by over two orders of magnitude is necessary before fuelcells can compete techno-economically with other modes of electricitygeneration in automobile applications.

Among the various fuel cells, the Proton Exchange Membrane Fuel Cells(PEMFC) offer great promise as an alternative to traditional fuel combustionfor generation of electricity for mobile and stationary applications. This H2

must have a CO concentration lower than 50ppm. In a typical fuel processorfor a PEMFC, the hydrocarbon undergoes reforming by the steam reforming(SR), autothermal reforming (ATR) or catalytic partial oxidation (CPO). Thereformate then undergoes a series of reactions with the goal of reducing theconcentration of CO and increasing the concentration of H2. The first is thewater gas shift reaction which reduces the concentration of CO in thereformate from about 10% to less than 1% while increasing the hydrogenconcentration. Further CO clean-up methods, such as preferential CO

Figure 32: Water gas shift in a Fuel Processor for fuel cells (8–9).



oxidation or selective CO methanation are needed to reduce CO levels to below50ppm.

Most of the papers on precious Pt- group or gold catalysts, described in theearlier sections, report results with powder catalysts or those in the form ofextrudates or tablets. However, in many fuel cell applications, due torequirements of very high space velocities (to reduce the volume of thecatalyst bed), low pressure drops and mechanical strength, the use of monolithcatalysts is almost mandatory. In response to this requirement, during thelast decade, many publications and patents (especially from industrialresearch 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 arequite different from those needed in NH3 or H2 plants (Table 8). Thedevelopment of robust WGS catalysts that can operate in such demandingconditions is critical in the development of hydrogen generators for fuel cells.Furthermore, much more active catalysts are sought to make the fuelprocessor as compact as possible. The challenge is formidable to achieve suchhigh catalytic activity at low temperatures (111, 145–148). It is also desirableto replace the two, HTS and LTS, reactors operating in the 350–450uC and200–300uC , respectively, by a single, medium temperature, shift reactor inthe 200–350uC range. The activity of the Pt-group catalysts is inadequatebelow 250uC. Since CO concentrations as high as1% are tolerated by recentpreferential oxidation (PROX) catalysts used in mobile fuel processors withoutsacrificing too much efficiency, this enables one to run monolithic WGScatalysts at temperatures as high as 300–350uC to reduce the residual COcontent to about 1% in the reformate gas. Even for stationary applications,this concept of replacing the two HTS and LTS reactors by a single shiftreactor 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



ruggedness of the monolith catalysts. While the efficiency of the monolithcatalysts is slightly lower (about 2–5%), the catalyst bed volume is about 90%smaller than a comparable system with a particulate catalyst. At the presentstage of development of the Pt-group metal-based, single stage shift reactor, atwo-stage PROX or selective methanation catalyst may be necessary to reducethe CO content of the reformate to , 10 ppm. The discovery of PEMFC anodesthat can tolerate higher amounts of CO in the H2 stream (Pt-Ru instead of Pt)may improve the situation further. Additionally, replacement of the lowtemperature (100uC) polyvinyl styrene – based electrolyte membrane by thehigh temperature (200uC) polybenzimidazole - based membranes in thePEMFC (enabling the fuel cell operation at 200uC) can also lead to a greatertolerance of CO by the fuel cell Pt anodes since the poisoning by strongly-heldCO 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) theactive metal (the metals Pt, Pd, Rh, Au or their mixtures), and (c) the supportmetal oxide (ceria, zirconia, lanthana, titania, alumina or their mixtures)powder (the ‘washcoat’). The high geometric surface area of honeycombmonoliths combined with their good mechanical strength and low pressuredrop make them particularly attractive for vehicular applications. Theperformance and durability of the finished catalyst depends significantly onthe quality of the washcoat. It is extremely important that the washcoatingprocess produces a reproducible, uniform, layer of the washcoat. Apart fromthe active metals and the support metal oxide, there are several othermaterials which can act as additive, binder or adhesive to the washcoat slurrywhich is deposited on the monolith prior to impregnation of the noble metal.Acetic acid is added to most washcoating slurries as a peptizing and dispersingagent to maintain an adequately low viscosity of the washcoating solution.Major steps in slurry preparation are particle size reduction of the supportmetal oxide powder and addition of appropriate acid/sol to adjust the pH,viscosity and homogeneity of the support metal oxide slurry in water. Sizereduction of powder particles in water down to a few microns to achieve welldispersed, homogeneous aqueous slurry can be accomplished by ball milling.The particle size distribution of the washcoat affects the mechanical strengthof the finished washcoat and its adhesion to the monolith, as well as therheological properties of the slurry during the washcoating process. In thenext step, the materials are dispersed in an acidic medium in a tank with ahigh-shear mixer. The solids content in the slurry is typically 30–50%wt. Afterprolonged mixing, the slurry suspension becomes a stable colloidal system.



The amount of washcoat that can be deposited on the monolith depends on theproperties of the monolith and the slurry. The catalyst samples are thenprepared by dipping the cordierite monolith in this slurry until the desiredloading is reached. The monolith is then air-knifed to remove excess slurryfollowed by drying at 110uC and calcining at 550uC. The monolith is thenimpregnated, with an aqueous solution containing the precious metal complex,dried and calcined. The physicochemical properties of the washcoated monolithcatalyst, like chemical composition, BET surface area, pore volume, metaldispersion and washcoat adhesion are then measured. When ceria-zirconia isused as the support oxide for the platinum groupmetals, their content, in g/litervolume of the final monolith catalyst, is between 200–500. The noble metalcontent is between 2–10 g/liter of the monolith. The costs of the noble metal andthe rare earth oxides constitute a significant part of the cost of the fuel processorin a fuel cell and efforts are in progress to reduce them further.

The ceramic honeycomb monolith, which is generally used, has somedisadvantages. It has a non-uniform flow distribution due to unidirectionalchannels and a closed structure between channels and slow diffusion rate ofreactants to the catalyst surface due to low turbulence in the channels.Further, when the catalyst particles are washcoated into the channels, thecatalyst particles are not uniformly deposited and are mainly deposited atcorners of square-shaped channels in the monolith, thus leading to lowercatalytic activity. In addition, the low thermal conductivity of the ceramicmaterial is also a disadvantage in dissipating away the heat generated inexothermic reactions like the WGS reaction. Metallic monoliths have highthermal conductivity, larger geometric surface area per unit volume, easyfabrication and have uniformly deposited catalyst particles. The gas flux flowsin the channel direction as well as in the direction perpendicular to thechannels. Thus, a turbulent flow of the reactant gases results, leading to highmass transfer rates. Consequently, the required reactor volumes aredecreased. The metallic monolith can be made of a refractory metal likestainless steel or other iron-based corrosion resistant alloys (e.g., iron-chromium alloy). They are typically fabricated from such materials by placinga flat and corrugated metal sheet, one over the other, and rolling the stackedsheets into a tubular configuration about an axis parallel to the configuration,to provide a cylindrical –shaped body having a plurality of fine, parallel gasflow passages, which can range from 200 to 1200 per square inch of face areacompared 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



a few examples will be given here to illustrate the advantages of the monolithcatalyst. 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 catalystreaches equilibrium conversion levels around 275uC. The high space velocityof operation (GHSV 5 20,000 h21) should especially be noted. The advantagesof the noble metal-washcoated, monolith catalysts are apparent from theseresults. A conventional Cu-ZnO-Al2O3 LTS catalyst does not give equilibriumconversions 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 compositionand ceria-zirconia is a superior support to ceria or zirconia alone. Theinfluence of monolith geometry and external geometric surface area on COconversion 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 alower activity suggesting that mass transfer from the bulk fluid to the catalystsurface, and, not the surface reaction, is controlling the rate of the reaction inmonoliths with lower geometric surface areas (like the 225 cpsi monolith). Asexpected, 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 washcoatedper unit volume of the monolith (keeping the Pt content constant) varies,usually, between 200–500 g/L of monolith. Ceria-zirconia amounts beyond500 g/liter of the monolith do not increase the catalytic activity. The superior

Figure 33: CO conversion over supported Pt, Re and PtRe catalysts.



Figure 35: Influence of gas velocities on CO conversion over monolith catalysts.

Figure 34: Influence of monolith geometry (CPSI) on CO conversion.



activity of TiO2- based supports, illustrated earlier (Section 6), is alsoreproduced 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 bementioned that the catalytic activities of the monoliths shown in Figs. 33–36above 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).



the monolith catalysts also undergo deactivation on prolonged use.Deactivation of Pt group metal–ceria catalysts with time-on-stream in realisticreformate gas streams have been reported (149, 150) and attributed to variousreasons, such as metal-induced over- reduction of ceria, precious metalsintering after high temperature reaction aging of Pd and Pt- CeO2 catalysts(86, 87) and carbonate formation (106). Deactivation due to carbonateformation has also been reported for Pt-CeO2 (151) and Au-CeO2 (137).Apart from their high cost, some of the technical drawbacks of noble metalcatalysts in WGS applications at low temperatures include (a) lower catalyticactivity (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 carbonatespecies on the surface, which, eventually cause catalyst deactivation. Theknown Fischer-Tropsch activity of the noble metals in the 200 – 300uC range(153) is, perhaps, relevant in accounting for their methanation activity underWGS conditions. In 2007, researchers from the Honda Motor Companyreported (152) results of combinatorial catalysis for over 250,000 materialsand claimed that catalysts containing a combination of (a) one noble metal likePt or Rh, (b) one group 11 metal like Cu, Ag, Au, and (c) one partially reducibleoxide like ceria, zirconia, titania, lanthana,vanadia or their mixtures, formsuperior WGS catalysts active and stable at low temperatures. While most ofthe other elements of the above combination were well-known for theirimportance in WGS catalysis, the inclusion of vanadia as a promoting supportfor the WGS reaction is interesting and may open future possibilities.Suppression of methanation activity of monolith catalysts during WGSreactions at low temperatures by inclusion of basic oxides, like ZnO, MgO,CaO, SrO, and BaO, in the catalyst support, has been claimed by Farrauto’sgroup (154). Inclusion of any of these basic oxides in the catalyst formulationhas been claimed to reduce the methane content in the outlet of the WGSreactor to less than 5ppm.

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 formulationhas been published (155). The authors prepared their Pt-CeO2-Al2O3 catalystby a sol-gel method, washcoated it in the micro channels of stainless steelplatelets and evaluated them for catalytic activity in the WGS reaction.Microstructured metal platelets offer excellent temperature control due totheir good thermal conductivity and small dimensions. Moreover, the use ofthin washcoat layers of catalysts eliminates the intraparticle diffusionlimitations that occur for fast reactions. The superior catalytic activity ofthe platelets compared to the pellet samples (Fig. 38) was attributed to thediffusion limitations inside the pellet samples. The conversion over the pelletsamples above 290uC is lower than that of the platelet samples indicating thatthere might be diffusion limitations inside the pellets, as the pellet size



(250 mm) is substantially larger than the equivalent particle diameter of thecoated, catalyst layer (37mm) inside the micro channel. To verify thishypothesis, a simulation was performed that took into account explicitly thediffusion of matter inside the pellets. The simulation was based on a value ofthe average particle diameter of 250mm, a tortuosity factor of 5 and the meanpore diameter and pellet porosity as calculated from the N2 adsorption data.Figure 38 compares the experimental data and the simulation. It can beobserved that at 260uC, the powder and the platelet indeed give similar initialrates. Above this temperature, the rates (for the pellet) are lower due todiffusion limitations inside pores, until the thermodynamic equilibrium isreached, at which point, the rate for the reverse water gas shift is close to theforward WGS rate and, hence, the overall rate is lower. They concluded that,catalysts deposited on micro structured platelets lead to a better utilization ofthe Pt metal.

9. SURFACE STRUCTURES, ACTIVE SITES AND REACTIONMECHANISMS

The mechanism of the WGS reaction has been studied thoroughly for manydecades. Even though there is some consensus on the redox mechanismprevailing over the iron-chromia catalysts at high temperatures, there isconsiderable 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).



Newsome (4) and Kochloffl (6, 7) have reviewed the literature in this field upto1975 and 1996, respectively. The mechanism over the Au based catalysts hasbeen reviewed by Burch (116). The WGS reaction involves the removal of anoxygen atom from the H2O molecule to liberate H2 and addition of the oxygento CO to form CO2. The dissociation of H2O can occur on the metal, the supportor 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) asexemplified in Eqs. 21 and 22 or (b) those that involve a rate – determiningstep in which the reaction takes place between two adsorbed species (theLangmuir-Hinshelwood mechanism), illustrated, for example, in Eqs. 23–27.The first is exemplified by the redox mechanism proposed long ago, in 1949, byTemkin (17–19) and developed further during the past decades. The multistepmechanism (Eqs. 23–27) proposed by Oki et al., in 1973 (20–22), as well as theformate mechanism fall in the second category (the L-H mechanism). Theredox mechanism for the HTS reaction was supported by the results ofBoreskov et al., (156–157) who established that the rates of reduction andoxidation of an iron oxide-based WGS catalyst were in good agreement withreaction rates calculated from Eqs. 21 (surface oxidation of the catalyst) and22 (surface reduction). Additional support for this mechanism was derivedfrom the results of Tinkle and Dumesic (158) who, from rates of adsorption/desorption and interconversion of CO and CO2 (using isotope exchangetechniques) over iron oxide – chromium oxide catalysts, concluded that CO/CO2 interconversion is fast compared with adsorption/desorption of CO andCO2. 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 themechanism 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- Feand Fe3O4. Below 570uC, the reduction of Fe2O3 to Fe metal proceeds in twosteps via an Fe3O4 intermediate. The reduction of Fe2O3 to Fe3O4 isexothermic, whereas further reduction to the metal is endothermic.Hematite crystallizes in the Al2O3 (corundum) structure with a closely packedoxygen lattice, with Fe3+ cations occupying octahedral sites. Its structure canbe visualized (159) as being composed of Fe-O3-Fe units (triplets) of closelypacked oxygen atoms with Fe(III) on either side. The Fe(III) atoms in each of



these Fe-O3-Fe units have opposite spins, being antiferromagnetically coupledas a result of superexchange interaction through the triad of oxygen atoms.Substitution of Fe cation sites by other metal (M) ion substituents in thestructure 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- typestructures, octahedral sites are occupied by 2+ and 3+ ions, whereastetrahedral sites are occupied only by 3+ ions. Boreskov et al. (156–157)had, earlier, demonstrated that the octahedral Fe2+ and Fe3+ ions located inthe magnetite-based structure function as a redox couple, and that magnetite-based catalysts can be highly effective for the complete dissociation of H2Ointo H2 and adsorbed oxygen under HTS reaction conditions. Waterdissociation causes the oxidation of Fe2+ to Fe3+ and liberates H2. The Fe3+

centers may, subsequently, be reduced to Fe2+ by CO or H2, thereby producingCO2 (or H2O) to complete the reaction loop. Many of these substitutionsimprove the thermal and textural stability of the structure while promotingthe reducibility of Fe2O3 to Fe3O4. Inverse and mixed spinel structures readilyundergo rapid electron exchange between the 2+ and 3+ states, therebycatalyzing the WGS reaction. A detailed investigation of the structuralproperties of the magnetite (Fe3O4) - based HTS catalyst system has beenpublished, recently, by Khan et al. (159). These authors prepared metal-doped,iron oxide-based catalysts with nominal composition of Fe1.82 M0.18O3, (whereM5 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. Theresulting cake was dried at 80uC and further calcined at 500uC in an inertenvironment. The structure of the materials was studied by various structuraltechniques and evaluated for their catalytic activity in the WGS reaction at350 – 550uC, different steam to CO ratios (1, 3.5, and 7) and a GHSV of60,000 h 21. On activation, the hematite- like Fe1.82M0.18O3 phase transformedinto either an inverse or mixed Fe2.73M0.27O4 magnetite - like spinel phase.The activity of Cr- and Ce- substituted Fe3O4 materials approachedequilibrium levels at high temperatures. At lower temperatures, the activityof these magnetite- based catalysts was limited by the dissociation of steam.Interestingly, they discovered that Ce- substituted Fe3O4 spinels are quitepromising HTS catalysts under steam-rich and high temperature conditions.Khan et al. also carried out the temperature- programmed reduction inhydrogen of their various promoted iron catalysts. Each promoter ioninfluenced the reduction profile of iron oxide in a unique manner. In the caseof the Fe2O3-Cr2O3 catalyst, the first reduction peak at 225uC corresponded tothe reduction of Cr6+ to Cr3+. Further partial reduction of Cr3+ to Cr2+, whichwould be expected at 490uC, was not observed. Reduction of Fe2O3 to Fe3O4

was observed at Tmax 350uC, whereas further reduction to FeO occurred at



higher temperatures. Adding chromium to Fe2O3 did not improve thereducibility of hematite to magnetite. Based on XRD, TPR and Mossbauerstudies 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 thedisplaced Fe2+, consequently located in tetrahedral sites. In the Fe2O3-CuOcatalyst, the reduction of Cu2+ to Cu+ occurred at 143uC. An interestingobservation was that the addition of Cu to Fe2O3 decreased the reductiontemperature of hematite to magnetite considerably to 190uC, compared to348uC for the pristine hematite sample. The addition of Cu to iron oxide thusimproved reducibility of Fe3+ to Fe2+ species. On doping with Cu, the mobilityof lattice oxygen and hydroxyl groups increased, due to the greaterelectronegativity of Cu (1.9) compared to Fe (1.8), thereby perhaps improvingcatalytic activity. Promoting iron oxide with cerium causes a shift in thereduction temperature peak maxima of both hematite - to - magnetite andmagnetite – to - wustite to lower temperatures. In the Fe2O3-Cr2O3, the ceriasurface shell reduction occurs at 380uC, instead of 485uC as in pure ceria. Thetemperature of bulk reduction of ceria, however, was not affected by thepresence of iron. These results are significant in the development of Cr-freeiron oxide-based, HTS catalysts.

The propensity of metallic iron and the lower oxides of iron to be oxidizedby steam and evolve hydrogen at high temperatures are well known.Additionally, the adsorption and surface concentrations of CO at hightemperatures will also be low. The redox mechanism will thus be favored athigh temperatures and over those catalysts which can dissociate H2O into H2

and O2, a crucial requirement of the redox mechanism. Reviewing this area in1996, Kochloffl (6, 7) concluded that the WGS reaction at high temperaturesover 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 iscurrently used as the standard, industrial catalyst, (b) the Pt group metalssupported on partially reducible oxides, like ceria, titania, zirconia or theirmixtures, and (c) Au supported on the same, above-mentioned oxides. Thereaction mechanism on oxide-supported Au catalysts may be different fromthat on supported platinum metal catalysts, because of (a) the loweradsorption energy of CO on the Au nanoparticles; and (b) the inactivity ofAu (unlike the Pt- group of metals) for H2O dissociation. Some of the featuresof the LTS reaction over Cu-ZnO catalysts that distinguish them from the HTSreaction over iron oxide – chromia are: (1) the dissociation of H2O to H2 and O



over copper metal or ZnO, at these low temperatures, is less documentedcompared to that on the iron oxide, Fe3O4, at the higher temperatures (120);(2) the amount of CO adsorbed on metallic copper, even at the lowtemperatures, is less than that on the platinum group of metals at similartemperatures and (3) the WGS rate is proportional to the CO partial pressureto the first order over the Cu-ZnO compared to the zero order observed overthe Pt-based catalysts. The last feature implies that, reducing the COconcentration from 1.0% (a typical value in the LTS reactor) to 0.5% requirestwice as much catalyst as reducing it from 2% to 1% thus leading to largesecond stage LTS reactors. Hence, major efforts have been made in the lasttwo decades, in the field of fuel processors, to develop cost-effective catalyststhat are tolerant to oxygen exposure, have robust, high volumetric activities atlow temperatures and whose CO conversion rate is independent of the COconcentration (zero order) in the range 3–0.3%. As mentioned earlier (Section6), noble metal- based catalysts, like Pt- ceria, Pt-ceria-zirconia, Pt- titaniaand 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 thesecatalysts so that they can be used successfully in fuel cells. It is in this contextthat a better understanding of the basic mechanism of the LTS reaction overthese noble metal- based catalysts is of importance. A redox mechanism,involving the reduction of the catalyst by CO and reoxidation (and, thereby theregeneration of the catalyst by H2O), similar to the one described above for theHTS reaction over the iron oxide- chromium oxide catalyst, has also beenproposed 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-ceriainterface. Two Ce4+ ions are reduced to the Ce3+ state in this process. Theresulting reduced ceria lattice is then reoxidised through the dissociation ofthe incoming H2O. The O vacancy is, thereby, refilled, formally oxidizing twoCe3+ to Ce4+ and releasing molecular H2 in the process. If the sites for theadsorption of CO and H2O are different, there will be no competition betweenCO and H2O for adsorption and the zero order rate dependence for CO may beobserved. If the sites for the adsorption of CO or H2O is occupied by the otherreactant (H2O and CO, respectively), or, by strongly adsorbed species (like theformates, carbonates or carboxylates), the reaction order, may be different.However, it should be noted that at the lower temperatures characteristic ofthe LTS reactions over catalysts like Cu-ZnO, and Pt-ceria (190–250uC), theability of steam to reoxidise the partially reducible oxide supports (with orwithout the presence of the noble metals) has not been demonstrated so far; Inaddition, in the case of fuel cell conditions (for Pt-Ceria), this reoxidation ofCe3+ to Ce4+ has to occur in the presence of a considerable amount of hydrogen.



Techniques, like XANES, can, perhaps, be used to resolve this issue. Inaddition to the redox mechanism, an associative mechanism involving asurface 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 thereaction involved the adsorption of CO and H2O on the catalyst surface to forma surface intermediate that subsequently decomposed to CO2 and H2. Toisolate and identify the ‘‘surface intermediate’’, they reacted it with NH3 andobtained ammonium formate. The formation and decomposition of a formate-or formic acid – type surface species was speculated to lead to the formation ofthe products, H2 and CO2, over copper – based catalysts, like Cu-Cr2O3.Similarly, using dimethyl sulphate as a methylating agent to trap the surfaceformate, Deluzarche et al. (161) obtained dimethyl formate further supportingthe presence of formates on the surface during the WGS reaction. During thelast few decades, extensive research using a variety of techniques has,conclusively, proven the existence of a formate species at low temperaturesover Cu-ZnO as well as precious metal- reducible metal oxide catalysts (162).Boreskov and Davydov (163–164) had earlier carried out pioneering IRspectroscopic studies supporting the associative mechanism over a widenumber of copper- based catalysts including the industrially important Cu-ZnO. Additional early work supporting the formate associative mechanism forCu-ZnO include those of Herwijnen et al.(165–166) who observed the nearlyidentical conversion rates for the water gas shift reaction and formatedecomposition. Rhodes et al.(167) have, however, raised some doubts as towhether these surface formates constitute the only, or, even the mainintermediates in the reaction path of CO and H2O to H2 and CO2 or are merelyspectators 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. Theirresults indicated that surface OH groups, formed by reaction of H2O withoxygen vacancies on partially reduced CeO2, reacted with CO to form bridgedformates. The bridged formates were converted to bidentate formates above170uC. This transformation occurred at room temperatures in the presence ofwater. About 30% of these adsorbed bidentate formates were, in turn,decomposed to the final products (H2 and CO2) and adsorbed, unidentatecarbonates. The rest decomposed back to the reactants, CO and H2O. Thesetransformations of the bidentate formates were also influenced by thepresence of H2O. Coadsorbed H2O also promoted the decomposition ofthe unidentate carbonates to CO2. In addition to the unidentate carbonates,



the presence of surface carboxylates and bidentate carbonates were alsoidentified 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 steamin about 6min at 300uC , while in the presence of steam, they decomposedcompletely and much more rapidly in 8min even at 140uC to produce H2 andunidentate carbonates. Based on mechanistic studies, including kineticisotope effect (170) and isotope tracer studies (172), they, further, suggestedthat the rate-limiting step in the LTS reaction is the cleavage of the C-H bondof the surface formate. Bridging OH groups and surface formates have alsobeen identified, earlier, by IR spectroscopy over thoria and zirconia-basedcatalysts (173–175). Jacobs et al. have also observed that Pt-thoria (176) andPt-zirconia (177) possess much higher WGS activity than the correspondingoxides without the precious metals and attributed it to the more facileformation of the Type II bridged OH groups, and the surface formates derivedfrom them, over the Pt-loaded catalysts. In addition, kinetic isotope effectssimilar to those observed in the case of Pt-ceria were also observed for Ptsupported on thoria and zirconia suggesting that the rate-limiting step waslikely to be the C-H bond scission of the formate intermediate in these casesalso. To summarize, in the formate mechanism, a bidentate formate producedfrom CO and surface OH groups acts as an intermediate. The bidentateformate, then, decomposes to gaseous hydrogen and a surface unidentatecarbonate, which further decomposes to gaseous CO2. One of the majorcontributions of Jacob et.al., is the discovery that co-adsorbed water plays acrucial 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, leadingto the formation of surface, terminal OH groups on ceria and (2) to catalyse thedecomposition of the formate to H2 and CO2. The rate determining step islikely 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-limitingstep, then, factors that facilitate the cleavage of the formate C-H bond shouldenhance the reaction rate. More specifically, addition of bases, like the alkaliions, which are known to accelerate the decomposition of formates, shouldenhance the WGS rates. Pigos et al. (178) have recently observed that theincorporation of Na in Pt- zirconia catalysts does, indeed, enhance the WGSreaction rates. Interestingly, their DRIFTS investigations suggest thatincorporation of Na in Pt- zirconia modifies the electronic properties of thesurface formate and weakens its C-H bond. Two significant features of theirDRIFTS results (178) are (a) The C-H band of the formate species wasshifted to lower wavenumbers from 2880 cm21 (Pt-ZrO2) to 2842 and



2804 cm2 1 (Pt-Na-ZrO2), respectively, indicating a weakening of the C-H bondon Na incorporation; and (b) The ratio of the intensity of the bridged to linearPt carbonyls increased from 1:5 to 4:5, thus, favoring bridge-bonded CO in thePt-Na-ZrO2 catalyst. The influence of catalyst basicity in increasing theconcentration of the bridged carbonyls had, already, been reported by Mojet etal., who found (179) that, in Pt-SiO2 and Pt-K-L (LTL) zeolites, increasing theK+ ion content also increased the concentration of the bridged carbonyls. Thefaster decomposition of the surface formate over Pt-Na-ZrO2 is illustrated inFig. 39. 20% of the initial intensity of the IR band of formate is decreased in5.2min for the Pt-ZrO2 and 2.4min over Pt-Na-ZrO2. CO2 was also seen in thepresence of steam. In addition to the formate, carbonates and carboxylateswere also seen after steaming. Pigos et al. (178) also compared the stability ofthe formate species, on Pt-ZrO2 and PtNa-ZrO2, in the absence of steam, bymonitoring their thermal decomposition. It may be recalled that Shido andIwasawa (107, 108) had, earlier, observed that in the absence of steam,formates on metal-ceria catalysts decompose, primarily, in the reversedirection, back to CO and OH. To follow the thermal decomposition of theformate 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 formatebands. The areas of the formate C-H bands (at 2880 and 2966 cm21) and theircorresponding 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 overallformate 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 bondbreaking in the formate is more facile for the Na-promoted catalyst. It may beobserved 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).



decomposition rate constant and surface coverage determined simultaneouslyat the steady state under WGS reaction conditions. Based on DRIFTS analysescombined with the utilization of isotopic tracers, it had been shown (180) thatformates were less reactive than carbonyl and carbonate species under steadystate 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 byMeunier et al.(60–62) using simultaneous DRIFTS and kinetic analysis. Theexchange of the product CO2 and formate and carbonate surface species werefollowed during an isotope exchange of the reactant, CO, using a DRIFTS cellas the reactor. In independent experiments, the DRIFTS cell yielded identicalreaction rates to that measured in a quartz, plug-flow reactor. The DRIFTSsignal was used to quantify the relative concentrations of the surface speciesas 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).



155–220uC suggested the presence of two types of surface formates: ‘‘slowformates’’ that display an exchange rate constant 10- to 20 fold slower thanthat of CO2 and ‘‘fast formates’’ that exchanged on a time scale similar to thatof CO2. Figure 41 compares the molar rate of formate decomposition to that ofCO2 formation per unit mass of catalyst. The specific rate of CO2 formationwas determined from the CO2 concentration in the DRIFTS cell exhaust gas(by gas chromatography), the sample weight in the DRIFTS crucible and theflow rate of the reactants. The rate of formate decomposition was calculatedfrom 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 wasmore than an order of magnitude (about 60-fold) higher than the rate ofdecomposition of the (slow + fast) formates, indicating that the formates,detected by DRIFTS, cannot be the only reaction intermediates in theproduction 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) thedecomposition of these formate species under WGS conditions leads to theproducts, CO2 and H2, it is not established that CO2 and H2 are derived onlyfrom 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 andexperimentally observable, surface intermediate. This viewpoint is furthersupported by the investigations of Gokhale et al. (196–199) and Mhadeshwarand 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).



9.2.2. The Redox Mechanism on Pt-based Catalysts

A redox mechanism is also being advocated (106, 148) for the LTS reactionon Pt-ceria and other platinum group metals on ceria. According to thispicture, CO adsorbs on transition metal sites and reacts with oxygen from theceria, which, in turn, is reoxidised by H2O. In other words, it involves thereaction of the reactants, CO and H2O, with the surface: CO with the oxide ionof 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 isto adsorb/activate CO and create of oxygen vacancies at the metal ceriainterface. 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). TPDstudies (182) have demonstrated that oxygen from ceria can react with COadsorbed on metals. It has also been established (183) that reduced ceria canbe oxidized by CO2. While the redox mechanism is well-established at hightemperatures in the case of the iron oxide-chromium oxide catalysts, itsapplicability to LTS over Cu-ZnO and Pt-ceria catalysts is uncertain anddepends on confirmation of the ability of H2O to reoxidize the partiallyreduced support oxide at temperatures below 250uC, especially in the presenceof significant amounts of hydrogen, as is the case for fuel cell applications (seealso sections 9.2 and 10). Such an unambiguous experimental confirmation isdesirable. Another feature of the ceria- based catalysts, namely, that hightemperature calcination lowers, not only the concentration of oxygenvacancies and loosely – bound surface oxygen atoms, but also their WGSactivity lends additional support to the redox mechanism. Reaction orders onPd-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 stepwas believed to be the dissociation of H2O on the ceria support. Diffusereflectance and FTIR spectroscopic measurements on Pd-ceria indicated thatthe ceria existed (not surprisingly) in a reduced state under WGS conditionsand is covered by carbonate species that are removed only by reoxidation ofceria (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 ceriacatalysts.

Azzam et al. (184–185) studied the WGS reaction on catalysts based onReO2-TiO2. Results pointed to contributions of an associative formate routewith redox regeneration and two classical redox routes involving TiO2 andReOx, respectively. Under their WGS reaction condition, rhenium waspresent, at least partly, as ReOx providing an additional redox route forWGS reaction in which ReOx is reduced by CO generating CO2 and re-oxidizedby H2O forming H2. The reaction between CO adsorbed on Pt and OH groups



on titania was the rate-determining step. Gold nanoparticles supported onreducible and non-reducible oxides with comparable gold particle size werestudied in the WGS reaction by Sandoval et al. (186). Not surprisingly, theactivity of Au on reducible oxides was much higher than the one observed onnon-reducible oxides. The optimum calcination temperature was 300uC. Forsamples calcined at 300uC and reaction temperatures below 225uC the activityvaried (Fig. 42) as follows: TiO2 . CeO2 . Al2O3 . SiO2. A novel catalystconsisting of platinum deposited over a cerium-modified titania substrate hasbeen, recently, reported by Gonzalez et al. (187). They showed better thermalstability with respect to the bare TiO2 support and higher WGS activity thanthose corresponding to individual titania or ceria supports. XPS and TPRresults revealed the intimate contact between Pt and cerium entities in the Pt/CeO2–TiO2 catalyst that facilitates the reducibility of the support at lowertemperatures. The importance of CO adsorption on ceria in influencing therates of the WGS reaction was investigated by Li et al. (188). Au nanoparticleson monoclinic ZrO2 showed much higher catalytic activity for the low-temperature water gas shift reaction than those on tetragonal zirconia, mainlydue to the high CO adsorption capacity of monoclinic ZrO2. Formate speciesformed by the reaction of adsorbed CO on gold nanoparticle with hydroxylgroups on ZrO2 were postulated to be the reaction intermediates.

Figure 42: CO conversion over Au nanoparticles supported on TiO2, CeO2, Al2O3, and SiO2

(186).



9.2.3. Mechanism over Cu- and Au- based Catalysts

One of the important issues in the mechanism of the LTS reaction is therole 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 arepresent, mostly as nanosized metallic particles, during the WGS reaction.What is the intrinsic reactivity of these nanoparticles? Can nanoparticles ofCu and Au catalyze the WGS reaction on their own without the aid of an oxidesupport (such as ceria or ZnO)? Results from catalytic studies over bulk Cuand Au may not be directly applicable to the nanoparticles. On the pure, bulkmetals, the WGS reaction, probably, proceeds by a redox mechanism (63, 173).The mechanism may, however, be modified by the presence of the supportoxide (especially by a partially reducible one, like ceria) in intimate contactwith the metal nanoparticles and wherein metal-support interactions will bemore important. Rodriguez et al. (189–195) have addressed this issue. Theyinvestigated 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 wasmeasured when Au or Cu particles (2–4 nm) were deposited (Fig. 43). Thedeposition of Cu nanoparticles on ZnO (0001) produced a catalyst that wasclearly more active than the pure extended Cu surfaces. An even bettercatalyst 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 amixture of 20 Torr CO and 10 Torr H2Oat 625K 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).



The Au-ZnO(0001) system displayed catalytic activity that was worse thanthat of Cu-ZnO(0001). In contrast, Au-CeO2 had an activity similar to that ofCu- 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 ofsurface intermediates for the pure metal surface Cu (110), and the metal- ceriacatalyst. Importantly, in the case of pure, metallic Cu (110), analysis of thesurface after the WGS reaction showed it to be essentially free of formate andcarbonate species, suggesting that, on the pure metal surfaces, the WGSreaction proceeds by the redox mechanism. In agreement with others, theyalso identified adsorbed formate and carbonate-like species on the metal-ceriasurfaces after the WGS reaction. Using density functional calculations, theyhad also investigated (193, 194) the WGS reaction on Cu29 and Au29 clusters(representative of the metal nanoparticles formed on deposition on ceria orZnO supports) and on Cu (100) and Au (100) surfaces (representative of thesurface of the bulk metals). Figure 44 shows the calculated energy changes forthe WGS reaction on a Cu29 cluster. The reaction pathway with the minimumenergy barriers involves the following steps (Eqs. 36–41):

CO g! "<CO ads! ", !36"

H2O g! "<H2O ads! ", !37"

H2O ads! "?OH ads! "zH ads! ", !38"

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).



CO ads! "zOH ads! "?OCOH ads! ", !39"

OCOH ads! "?CO2 g! "zH ads! ", and !40"

2H ads! "?H2 g! ": !41"

The adsorption of CO or H2O on the Cu particles is exothermic. The first andthe most important energy barrier is for the dissociation of water intoadsorbed OH and H (reaction 38). Then, the reaction of OH and CO producesan OCOH, carboxyl species. The final important energy barrier is for thedecomposition of this OCOH carboxyl intermediate into CO2 gas and adsorbedH, which eventually yields the H2 gas. The DFT results indicated that a free,metallic, nanoparticle of copper can catalyze the WGS reaction easily. Acomparison with the corresponding results on the Cu(100) surface of bulkcopper shows that the dissociation of H2O on the surface of bulk copper has alarger activation energy barrier (1.13 ev vs. 0.94 ev on the nanoparticle) andthat no stable OCOH, carboxyl intermediate, is formed, as a redox mechanismoperates. The presence of corner or edge atoms in Cu29 favors the dissociationof H2O. The Au29 nanoparticles and the bulk Au (110) surface could notcatalyze (169) the WGS action. Neither surface was able to adsorb anddissociate water molecules. Figure 45 (189, 190) shows a correlation betweenthe calculated barrier (y axis) and the calculated energy (x axis) for waterdissociation on Au(100), Cu(100), as well as the ionic and neutral Au29 andCu29 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).



barrier and an endothermic DE value. There is significant improvement inchemical reactivity in going from Au(100) to Au29, but not enough fordissociation of the water molecule. These results, of course, cannot explain thelarge catalytic activity of Au-ceria (Fig. 43) and highlight the important role ofceria in the activation of the system. A perfect CeO2 (111) surface does notdissociate water at low or, even, high temperatures. When O vacancies arepresent, however, the H2O molecules dissociate on the partially reduced ceriasurface. Au and Cu particles facilitate the reduction of the ceria surface by theCO/H2O mixture and, thereby, facilitate the most difficult step in the WGSreaction, namely, the dissociation of H2O (189, 190).

9.2.4. The Carboxylate Mechanism

Surface carboxylic species had been observed earlier by spectroscopictechniques on LTS catalysts (107, 108). Mhadeshwar and Vlachos (210) andGokhale et al. (196–199) proposed from theoretical calculations, that they areimportant, reactive intermediates, which play a central role in the WGSreaction. Gokhale et al. used periodic, self-consistent, density functionaltheory (DFT-GGA) to investigate the WGS reaction mechanism on Cu(111),the dominant facet of copper crystallites in the Cu-ZnO industrial WGScatalysts. Their proposal for an alternate WGS reaction mechanism, involvingthe oxidation of adsorbed CO by adsorbed OH, to form carboxyl (COOH)species is compared with the conventional redox mechanism in Table 9. Thecrucial difference is that, while in the conventional redox mechanism theadsorbed CO is oxidized to CO2 by adsorbed O atoms, CO2 is formed by thedecomposition of an adsorbed carboxyl group (formed by the reaction of anadsorbed CO with an adsorbed OH group) in the new proposal. CO2 may alsobe 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 carboxylgroup, 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.



as postulated in the formate mechanism), that is less likely. That is becauseOH 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 bidentatespecies, either sequentially (via unidentate formates) or, less likely,simultaneously. Their calculations suggested that the easiest way to formthe formate, HCOO, is by reacting CO2 with atomic H. The 16 elementarysteps involved in their mechanistic model of the WGS reaction on Cu (111) areshown in Table 10 and the corresponding reaction network in Figure 46. Usingthe DFT-derived parameters as initial estimates for the microkinetic modelparameters, they fitted the 16- step model to the experimental WGS reactionrate data published, earlier, by Koryabkin et al. (200). As may be seen fromFig. 47, the agreement is satisfactory. Their model also tested well against thekinetic data of Herwijnen and Jong on a Cu-ZnO-Al2O3 catalyst (52). Based onthe good ‘‘fit’’ between the calculated and observed data, they suggested thatCu(111) may be a dominant active site for the WGS reaction on realisticindustrial catalysts. An alternate explanation may be that the WGS reactionon these catalysts is not structure sensitive, and therefore, the reaction rate iscomparable on different Cu facets (196). From their model calculations theyalso predicted that steps 5 (H2O* + * u H* + OH*) and 9 ( CO* + OH* u cis-COOH * + *), in Table 10, are rate controlling under industrial conditions. Inthe absence of CO2 and H2 co-feed, step 5 has a considerably strongerinfluence on the overall reaction. To summarize their results on Cu(111): (a) Habstraction 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*



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 aspectator species which tends to block active sites, and can reach substantialsurface coverages, particularly at high pressures. This site- blocking byformate can, also, explain the observed negative WGS reaction order withrespect to CO2 (196–199).

This approach has, recently, been extended by the same group (197) to thewater 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).



mind that the influence of the support on the electronic and texturalcharacteristics of the metal surface OH groups or reaction intermediateswas not explicitly taken into account in these calculations. The results are,broadly, similar to those described above for copper (196). The contribution ofthe surface redox mechanism to the water gas reaction on Pt, involving COoxidation by atomic O, is negligible under the range of conditions studied(250–300uC, 1 bar, feed 5 (CO+H2O+CO2+ H2)). The lowest energy pathinvolves the formation of the carboxyl (COOH) intermediate, which issubsequently, 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 decompositionroute (COOH R CO2 + H) dominates. Additional H2O in the feed increases theOH coverage and makes the OH 2 mediated COOH + OH, low energydecomposition path more kinetically accessible, thus, enhancing the water gasreaction rates.

9.2.5. Alkali-doped, Pt-based, LTS Catalysts

One of the key, rate- determining, steps in the formate mechanism is theC-H bond scission in the surface formate intermediate. Evin et al., haverecently found in accord with the results of Pigos et al. (178), that alkalidoping weakens the C-H bond, as demonstrated by a shift to lowerwavenumbers of the n(CH) vibrational mode, and enhances the LTS reactionover 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 suchthat while the formate became more reactive, the stability of the adsorbedcarbonate species, which arises from the decomposition of the initially- formedformate intermediate, was found to increase. This was observed by TPD-MSmeasurements of the adsorbed CO2 probe molecule. Increasing the amount ofalkali to levels that were too high also led to (a) lower catalyst BET surfacearea, (b) the blocking of the Pt surface sites as observed in infraredmeasurements, and (c) a shift to higher temperature of the surface shellreduction step of ceria during TPR. When the alkalinity was too high, the COconversion rate during the water-gas shift reaction also decreased incomparison with the undoped Pt-ceria catalyst. However, at lower levels ofthe alkali, the above-mentioned inhibiting factors on the water-gas shift ratewere suppressed such that the weakening of the formate C-H bond could beutilized to improve the overall turnover efficiency during the water-gas shiftcycle. This was demonstrated at 0.5% Na (or equimolar levels of K) dopinglevels. Not only was the formate turnover rate found to increase significantlyduring both transient and steady state DRIFT tests, but this effect wasaccompanied by a notable increase in the CO conversion rate during low



temperature water-gas shift. Evin et al., (202) had also observed, using in-situDRIFTS spectroscopy, that, on steaming pre-adsorbed formate species, thesespecies 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 catalystsrelative to undoped Pt- ceria (202). For example, the relative formatedecomposition rates were 1.0, 1.2, and 1.41 for the undoped, 0.15% Li- dopedand 0.5% Na- doped Pt-Ceria, catalysts respectively. The CO conversion at225uC, correspondingly, increased from 12% for the undoped 2% Pt- ceria to14% for 2%Pt- 0.15%Li-ceria and 24.3% for the 2% Pt-ceria sample doped with0.5% Na. However, with increasing atomic number over the series of alkali –doped catalysts, the stability of the carbonate species (another surface speciesformed during the WGS reaction) was also found to increase. This wasobserved during TPD-MS measurements of the adsorbed CO2 probe moleculeby 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, thecarbonate intermediate stability also increases, making it difficult to liberatethe CO2. An optimal amount of basicity, sufficient to decompose the formate,but, not enough to stabilize too much the carbonate, is needed. Infraredspectroscopy results of CO adsorbed on Pt and ceria suggested that the alkalidopant is located on, and electronically modifies, both the Pt and ceriacomponents. Alkali doping may, thus, provide a path forward for improvingthe 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 promotethe WGS reaction rate. In 1981, Sato and White (203) doped Pt- TiO2 withNaOH and found an improvement in the photocatalyzed water gas shift rate.Klier (204) also highlighted the promoting influence of alkali dopants, theirrelative efficiency being, Cs. Rb. K.Na, Li. Klier also suggested that thealkali should be present at concentrations less than a monolayer. Campbell etal., (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) ofthe catalyst. In kinetic studies, using a low H2O/CO ratio, they found that onthe optimally Cs- promoted surface, the reactant orders were zero order forH2O and 0.5 order for CO, suggesting that H2O dissociation was not ratecontrolling on the alkali-promoted catalyst. They proposed a redox mechanismto describe the catalysis of both the clean and Cs- doped surfaces with Csplaying 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 forfuel processors for use in fuel cell applications by doping the catalysts withalkali. Among the promising compositions discovered was an importantimprovement when Pt- ZrO2 was doped with Na alone or in combination with



vanadium. In their DRIFTS spectroscopic study of metal- ceria catalysts,Pigos et al. (178) found that formate C-H stretching bands were stronglyshifted to lower wavenumbers upon CO adsorption indicative of a weakeningof the C-H bond. Further, in transient formate decomposition experiments,both in the presence and absence of steam, they reported that formates, overPt-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 mostbenefit.

9.2.6. LTS over Pt Supported on Non-reducible Oxides

What is the mechanism prevailing over catalysts comprising of noblemetals supported on non-reducible oxide supports, like alumina or silica? Thefact that gamma alumina is an ‘‘irreducible’’ oxide at the WGS reactionconditions will seem to exclude the redox mechanism involving oxygen ionvacancies on the support, extensively discussed in the literature for partiallyreducible metal oxides and supported metal catalysts on such carriers.Olympiou et al. (206) studied the mechanism of the WGS reaction overalumina-supported Pt, Pd, and Rh catalysts, using steady state isotopictransient kinetic analysis (SSITKA) techniques coupled with mass spectro-metry. In particular, the concentrations (mmol g21) of active intermediatespecies 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 350to 500uC, the concentrations of the active species in both the carbon andhydrogen paths increased significantly. Based on (a) the large concentration ofthe active species present in the hydrogen- path (OH/H located on the aluminasupport), which was larger than six equivalent monolayers (based on theexposed platinum metal surface area), (b) the small concentration of OHgroups 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

0.5wt% Pt/c-Al2O3 350 350 (28.5) 1.3 (0.1)500 1664 (135.6) 31.7(2.6)

0.5wt% Pd/c-Al2O3 350 235 (9.8) 0.5 (0.02)500 3194 (132.7) 28.6(1.2)

0.5wt%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.



significantly smaller concentration (mmol g21) of active species present in thecarbon-path (adsorbed CO on the noble metal and formate species on thealumina support and/ or at the metal- support interface), the authorssuggested that the diffusion of OH/H species on the alumina support towardscatalytic sites present in the hydrogen pathway may be the slow step in thereaction mechanism. The OH/H species were considered to be formed by thedissociation 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 formatedecomposition into CO2 and H2 products. There was also a correlation betweencatalytic activity and the surface concentration and binding energy of CO onthe noble metals. Among the alumina-supported noble metals, the order ofactivity was found to be Pt . Rh . Pd. It may be remarked that these resultslend 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 COadsorbed on the Pt metal must react with the OH group adsorbed,presumably, on the alumina. Whether the formation of the formate is theresult of diffusion of CO from the Pt surface to the Al-OH sites or the diffusionof the – OH groups from alumina towards CO adsorbed along thecircumference of the metal-support interface is not clear. Duprez (207) hasdiscussed the mechanism of migration of the OH/H species on metal oxidesurfaces 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 ceriaand derived from a kinetic study using a microstructured reactor has been,recently, proposed for the WGS reaction over a Pt-CeO2-Al2O3 catalyst byGermani and Schuurman (208). The use of a microstructured reactor, ratherthan a packed bed reactor, enabled the measurement of the intrinsic kineticsof the reaction. The reaction rate was almost zero order in CO and wasstrongly inhibited by the partial pressure of hydrogen, and, to a lesser extentby 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 adsorbedon 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 withwater, dissociatively chemisorbed on ceria, to yield a carboxyl species as anintermediate. This carboxyl species reacts with a second hydroxyl group anddecomposes over a free Pt site into carbon dioxide and hydrogen as shownbelow:

COz%uCO%, !42"

H2OzCe-OuHO-Ce-OH, !43"



CO%zHO-Ce-OHuCOOH%zCe-OH, !44"

COOH%zCe-OHz%u2H%zCe-O-CO2, !45"

2H%uH2z2%, and !46"

Ce-O-CO2uCO2zCe-O: !47"

In the above equations * represents a Pt site. The rate determining step is thereaction between the carboxyl species, and the second hydroxyl group on theceria (Eq. 45). Once an adjacent Pt site becomes free, this carboxyl complexdecomposes into the reaction products. Hydrogen competes with CO for Ptadsorption sites and, therefore, retards the rate. Similarly, CO adsorbsstrongly on ceria and has a negative influence on the rate. This mechanismdiffers from the mechanism proposed by Shido and Iwasawa (107–108) in thatthe surface intermediate is postulated to be a carboxylate and not a formate. Itmay be mentioned here that the reaction of CO with Type II bridging OHgroups on ceria to form carboxy species has not been confirmed unambigu-ously by experiment. We may recall that a carboxyl surface intermediate hasalso been postulated, more recently, by Gokhale et al. (196–199). A key featureof the carboxyl mechanism is the conversion of adsorbed CO to adsorbedCOOH. From microkinetic studies of the WGS reaction system, Mhadeshwarand Vlachos (210) had, earlier, made an important observation that while theformation of the carboxyl intermediate from H2O, namely, CO* + H2O* )COOH* +H*, can be the rate-determining step (as per their model), competingpaths for CO* oxidation on Pt by OH* (i.e., CO* + HO* ) COOH* + *) insteadof H2O*, cannot completely be ruled out as being important (with the H2O*decomposition being the rate determining step), due to the relatively smalldifferences in activation energies of these parallel oxidation paths. They havealso provided a very useful and comprehensive compilation of all thesignificant rate expressions and reaction orders for the WGS reaction ondifferent 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 activein the WGS reaction below about 250uC. On the other hand, Cu-ZnO is anoutstanding WGS catalyst in the range of 200–250uC. Some of the latter’sdrawbacks are (a)the necessity to operate at low GHSV values, (b) its complexand time-consuming activation protocol before use, and (c) its instability oncontact with air. The Pt group metals do not have these disadvantages. In anattempt to combine the advantages of both the copper and Pt- group metals in



a single catalyst formulation, Fox et al. (211) have investigated the catalyticproperties of Pd- promoted Cu-Ceria catalysts for the oxygen-assisted, WGSreaction. It may be noted that the conventional ZnO support has been replacedby ceria in this formulation. Cu-CeO2, Pd-CeO2, and Cu-Pd-CeO2 catalystswere 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 1wt% Pd to 30wt% Cu-CeO2 greatly enhances thereducibility of both dispersed CuO as well as the ceria support, presumablyby hydrogen spillover from Pd. In-vacuo reduction (inside the XPS chamber)up to 400uC results in a continuous growth of metallic copper and Ce3+ surfacespecies. Support copper, in turn, destabilizes palladium metal (Pdu) withrespect to PdO, this mutual perturbation indicating a strong, intimateinteraction between the Cu-Pd components. The presence of Pd, apparently,increased the fraction of copper that remains in the metallic state therebyenhancing its catalytic activity. It may be recalled that the increase ofcatalytic activity with metallic copper surface area is well known. Palladiumaddition at only 1wt% significantly improved CO conversion at 180uC,compared to a monometallic 30wt% Cu- CeO2 catalyst (Fig. 48). Asanticipated, 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 byFox et al. (211) (Fig. 48) contained oxygen (2% air) and is not representative ofthe conventional feedstock from a steam reformer to the WGS reactor. Reactordesign considerations will be crucial in such a situation to avoid the oxidationof the hydrogen or CO over the precious metal and the consequent exothermictemperature rise. In a similar vein, but combining the relative advantages oftwo support components, ceria and titania, Gonzalez et al. (212) have recently

Figure 48: CO conversion over 1wt% Pd, and 30wt% Cu catalysts and 1wt% Pd – 30wt%Cubimetallic catalyst at 180uC. Feed gas: 4% CO, 10% CO2, 26% Ar, 2% air, balance H2. H2O/CO510 (211).



demonstrated the performance enhancement in the WGS reaction when Pt isdispersed over a mixed oxide support material containing both ceria andtitania. TiO2 and CeO2 have complementary physical properties which can besynergized to improve the performance of a catalyst support in the WGSreaction. For example, the redox properties and thermal stability of titaniacan be improved by replacing, partially, Ti4+ ions by the Ce4+ ions in thetitania lattice (213–214). Gonzalez et al.(212) have found that Pt supported onCe modified TiO2 catalyst shows better thermal stability (with respect to bareTiO2 support) and higher WGS activity than those corresponding to individualtitania and ceria supports, indicating a synergistic effect between Pt and theCe- modified TiO2 support (Fig. 49). XPS and TPR results revealed theintimate contact between Pt and cerium entities in the Pt-CeO2-TiO2 catalystthat facilitates the reducibility of the support at lower temperatures (Fig. 50)while the Ce-O-Ti interactions decrease the overreduction of TiO2 at hightemperatures. It may be noted that the addition of cerium to TiO2 had alsoincreased the hydroxyl concentration on the support. This is probably one ofthe 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 OHgroups in the mechanism of the LTS reaction.

While discussing the active sites and mechanism of the WGS reaction, it isinstructive to recall two important features of the landmark postulate of HughTaylor in 1926 (215) on active sites over solid catalysts: (a) particular atoms orgroups 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).



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 theparticular reaction conditions, i.e., the relative concentration of the reactants,temperature, pressure etc. If these conditions are changed, then, the identityand 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 onparticular reaction conditions. The CO concentration at the inlet to the WGSsection can vary widely in the range 10–40% (dry basis) depending on the rawmaterial (natural gas or coal) and the reforming process utilized to generatethe syngas. The H2O concentrations will also vary depending on the type ofreformer (steam, partial oxidation or autothermal reformer) that is utilizedupstream of the WGS reactor. Steam reformers utilize higher H2O/ carbonmolar ratios (3–5) than partial oxidation or autothermal reformers (0.5–2.0).Consequently, the concentration of H2O at the WGS reactor inlet will behigher 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).



autothermal reformer than in a steam reformer. It is normal to expect that thedifferent concentrations of CO, CO2 and H2O in the feedstocks, from thesevarious H2 generation system configurations, will influence, differently, thenature and, especially, the concentration of chemical species present on thecatalyst surface (OH groups, H atoms, anion vacancies etc.). Hence, it shouldnot be surprising that different WGS mechanisms can prevail on the samecatalyst under different reactants concentration and temperature/pressureconditions, especially in the WGS reaction that is equilibrium-limited at hightemperatures and kinetically limited at low temperatures.

Burch (116) has published a critical discussion of the relative merits of thevarious mechanisms that have been proposed for the LTS reaction over metal– partially reducible oxide supports. He has also presented a ‘‘Universalmechanism’’ for the WGS reaction that seeks to integrate the salient featuresof the formate and redox mechanisms into a single model that is consistentwith 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).



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/orcarboxylates is emphasized. Figure 51c shows Burch’s ‘‘universal’’ mechanism.One crucial feature of Burch’s universal mechanism is that, while in theformate mechanism, postulated originally by Shido and Iwasawa (107–108)and elaborated upon by Jacobs et al. (202), the formate intermediate is formedfrom insertion of CO into an OH bond, both being adsorbed on the support, it isformed (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 nounambiguous experimental evidence for this assumption. In fact, in studies ontransient formate decomposition either with the unpromoted catalyst or thecatalyst promoted with different metals (e.g., Pt, Au) and loadings of metal,once the surface shell of ceria is reduced, the formate concentration fromreaction of CO with the Type II bridging OH groups is high and, for the mostpart, have the same intensity over all the catalysts before the H2O is added todecompose them. If the formate had been anchored on the metal, a variation ofthe intensity with the type and concentration of the metal would have beenobserved (223). Both modes of formation of the surface formate intermediateare shown in Figure 51c. It is important to note that a carbonate–like speciesis also involved in the reaction path in all the three postulated mechanisms,including the formate mechanism, wherein the formate formed initially reactswith H2O to form a carbonate which finally decomposes to yield CO2. Thedominant mechanism will depend (116) on the reaction conditions, specificallythe temperature and the H2O/CO2 ratio. It can change from a redox-typeprocess 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 oxygenvacancies which dissociate water into OH adsorbed on the support. Metal sites



also adsorb and activate CO. In the next step, the OH groups, located on thesupport oxide, react with the CO adsorbed on the metal particles to form thesurface intermediates, the formates and carbonates. The latter decompose toCO2. It must, however, be emphasized that the above mechanism involvinglong-lived surface intermediates, like formates or carbonates/ carboxylates, isvalid only below, say, 350uC. At higher temperatures, over HTS catalysts, likeFe2O3- Cr2O3, the direct oxidation and reduction of the catalyst by H2O andCO, 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 offormation or decomposition of surface intermediates, like formates andcarbonates/ carboxylates, assume critical importance. One further point inrelation to this mechanism is its relative importance in the case of catalysts,like Pt-Al2O3, wherein a significant loss of the support hydroxyl groups occursonly above 400uC (216). Surface oxygen vacancies, that play such an importantrole in the above-mentioned mechanisms that invoke the formation of stablesurface intermediates located at these oxygen vacancies, are unlikely to bepresent in significant concentrations on the alumina surface in the 190–300uCrange, typical of the LTS reaction, especially in the presence of significantamounts of H2O. It may be noted that an oxygen anion vacancy is the startingpoint in all the mechanisms depicted in Figures 51. In such cases, both theactivation of CO and the dissociation of H2O occur, perhaps, on the Pt metal. Itmay, however, be relevant to mention here, that, Chenu et al. (177) hadreported the observation of surface defects, oxygen vacancies and type IIbridging OH groups on non-easily-reducible oxides, like MgO and ZrO2 underreducing conditions similar to those that prevail during the WGS reaction,when they were promoted with Pt. After activating the catalyst, DRIFTS ofCO adsorption was used to probe the active OH groups via the generation offormate species. It is important to note that in the absence of H2O, formatesare quite stable, such that their intensity upon CO adsorption gives a goodqualitative indication of the number of active OH group defect sites. Whileformates were observed over both the Pt-ZrO2 and Pt-MgO, the bandintensities were lower as compared with Pt- ceria, suggesting a lowerconcentration of defect-associated active OH groups on ZrO2 and MgO. TheWGS rates and formate band intensities from CO adsorption (used to probethe 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 thecatalysts. On lowering the temperatures, Pt-magnesia was inactive below400uC, while Pt-Ceria was active up to 250uC, under their reaction conditions(Feed: 3.75ml/min CO, 125ml/min H2O, 100ml/min H2, 10ml/min N2; 33ml ofcatalyst). Similar results of generating active OH groups on non-easilyreducible oxides, like ZrO2, on promotion with Pt was also reported by Pigos et



al. (178). They found that the catalytic activity of Pt-ZrO2 was improved,significantly, when Na was added, either alone or in combination withvanadium oxide. The results of parallel, DRIFTS spectroscopic experimentsindicated that formate species were more reactive on the Na - promotedcatalysts; The formate C - H stretching bands were shifted to lowerwavenumbers upon CO adsorption. In transient formate decompositionexperiments, both in the presence and absence of steam, the formates overPt-Na-ZrO2 decomposed at twice the rate of those observed on Pt-ZrO2 withoutNa. This was further supported by steady state WGS rates which confirmedthat the formate species were more reaction rate- limited in DRIFTS for theNa - promoted catalysts relative to those without Na. These results of Chenuet al. (177) and Pigos et al. (178) suggest that the active sites for the WGSreaction, namely, the oxygen anion vacancies and associated OH groups, canbe generated also on otherwise non-reducible oxide supports under reactionconditions when the catalyst formulation contains elements like the noblemetals.

9.2.7. Catalyst Deactivation

Deactivation during long-term tests has been one of the major drawbacksof the noble metal- based catalysts with ceria or titania supports. Thefacilitating role of bases, like the alkali ions, in the decomposition of formateshad been noted earlier (178). The deactivation of WGS activity along withstrongly held, carbonate surface intermediates had been observed by Gorte etal. (86–87). At high temperatures (above 350uC) and over the iron oxide-chromium oxide catalysts, these carbonates decompose more easily and nodeactivation is observed. Do the strongly-held carbonates impede thereoxidation of the oxide surface or the release of the H2 molecule? The useof acidic oxides, like those of Nb, Mo, Ta, and W, to enhance the WGS activityof Pt-Ceria-Zirconia has been reported, recently, by Opalka et al. (217). Aidedby density functional simulations, these authors observed that doping Pt-ceria-zirconia with acidic, transition metal dopants such as Nb, Mo, Ta and Woxides increased the oxide surface affinity for water and the turnover rate ofthe 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) thanthe undoped sample in the temperature range, 200–300uC. The composition oftheir feedstock was: 4.9% CO, 10.5% CO2, 33% H2O and 30.3% H2, and theGHSV was 300,000h21, simulating the environment in a fuel processor of afuel cell. Only initial catalytic activities were reported. Kinetic rate analysisfor the CO conversion yielded reaction orders approaching 0 for CO and 1 forH2O. They characterised the nature of adsorbed CO and the formate andcarbonate intermediates, formed during the WGS reaction at 200uC over



catalysts without and with Pt loading, by in-situ infrared spectroscopy. Afteroxidizing the nanocrystallites in dry air for 2 h and, then, exposing them to lowCO pressures, only weakly adsorbed CO was detected. When the CO pressureswere increased to 4 bar or higher, small amounts of adsorbed CO2, formatesand carbonates were also observed, indicating that the catalyst was partiallyreduced by CO. In the presence of H2O and, under WGS conditions, linearlyadsorbed CO and significant amounts of formates and carbonates were alsoobserved. The larger, more basic Ce3+ ions formed under reducing conditionsenhance the further reactions of adsorbed CO to form formates andcarbonates. The formates were weakly bonded and could be removed byoutgassing 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 linearadsorption was observed only on the Pt metal (not ceria). Formate andcarbonate formation was observed on the ceria- zirconia oxides. If ligand (CO)complexation of the oxide surface leading to strongly-held formates andcarbonates is locally specific to the reduced sites, then, CO associativereactions will compete with or impede the reduction (by CO) or oxidation (byH2O) of the catalyst and, hence, influence the redox mechanism. Based ontheir density functional simulations and IR spectroscopic/kinetic experimentalresults, the authors suggested (217) that the associative formation of formatesand carbonates was, indeed, coupled with the bifunctional redox mechanismsthat lead to the reduction of the oxide surface. They observed, further, that therate at which O could be removed by CO from the lattice, outstripped the rateat which H2O could chemisorb and react to replenish the lost O. Hence, therate- limiting step, over Pt-ceria and Pt-ceria-zirconia, was not latticereduction but rather reoxidation. While the reduction of ceria by CO/H2 isnot in question, its reoxidation by H2O at low temperatures needs to beverified experimentally by a direct method. To shift the WGS oxidation –reduction balance towards reoxidation, they added more acidic, less reducibledopants like Nb, Mo, Ta and W oxides, to make the reoxidation morefavorable. These are strong electron acceptors and are fully oxidized in theirLewis acid- like oxide phases with generally empty d orbitals (d0 oxides). Onceria and ceria-zirconia, formate formation was very close in energy toreoxidation of the reduced surface by H2O. In the presence of acidic transitionmetal dopants, however, surface reoxidation was significantly more favorablethan the reaction of adsorbed H2O with CO to form stable formate andcarbonate complexes. For example, while the enthalpy for reoxidation (byH2O) 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 thesurface (217). The transition metal oxide dopants, apparently, shifted therelative balance of the reaction steps, enhancing the refilling of oxide



vacancies and H2 generation, thereby minimizing the blocking of active sitesby formate and carbonate formation. As a consequence, both catalytic activityand life were improved (217).

How do dopant ions, like Mo, facilitate surface reoxidation vis-a-visformate/carbonate formation? An understanding of the interaction of H2Owith the catalyst surface holds the key to the answer. Atomic simulationcalculations indicated (217) that the H2O molecule dissociated, preferentially,over the dopant ion to hydroxylate the dopant and to protonate the surfaceoxygen ions in the adjacent fluorite lattice. The average H2O adsorptionenthalpy of 257.7 for ceria- zirconia increased to 2 82.3KJ/mole on dopingwith Mo and to 2110.4KJ/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 onmetal- metal oxide-based catalysts in the WGS mechanism has been studiedextensively, there is another group of WGS catalysts based on molybdenumcarbide wherein the mechanistic picture is less clear (218–221). Patt et al.(218) reported high activity for LTS over Mo2C catalysts and obtained higheractivity than over a commercial Cu-ZnO-Al2O3 catalyst. The precursor formolybdenum was ammonium paramolybdate. The salt was dissolved in warmwater. Then, the liquid was, slowly, evaporated and the solid was calcined indry air at 500uC. The oxide was carburized using a temperature-programmedtreatment with CH4 and H2. The LTS activities of the resulting solid (BETsurface area 5 61m2/g) as well as that of a commercial Cu-ZnO-Al2O3 catalystof 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. Thedivergence of this feedstock composition from those in commercial practice,especially the absence of CO2, may be noted. Under these conditions, the COconversion over the Mo2C catalyst was at least 50% higher than that over theCu-ZnO-Al2O3 sample. Moon and Ryu (219) found that the optimumcarburization temperature was 640–650uC. After repeated thermal cyclingin reductive and oxidative atmospheres, the authors found that even thoughthere 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 thatthe 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 ofthe 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 and1450 cm21. Instead of assigning these to formate species, the authors



concluded that, since no accompanying symmetric band was observed in the1350–1365 cm21 range, the 1626 cm21 band was not due to formate butanother vibration. The 1450 cm21 band was assigned to unidentate carbonate.Based on DFT calculations, the authors assigned the 1626 cm21 band to COadsorption on a 3-fold Mo site. A summary of their proposed redox mechanismis given below where * denotes a free adsorption site:

COz%[CO%, !48"

H2Oz%[H2O%, !49"

H2O%z%[HO%zH%, !50"

HO%z%[O%zH%, !51"

CO%zO%[CO2z2%, and !52"

2H%[H2z2%: !53"

The rate-limiting barrier was found to be the reaction of O with CO to formCO2 (Eq. 52). This group has, also, extended their study of molybdenumcarbides to include cobalt-containing samples (221). Catalysts were preparedby combining aqueous solutions of Co(NO3)2 and ammonium heptamolybdate(NH4)6Mo7O24 and stirring at 80uC to produce a viscous mixture. Solids weredried in an oven and calcined at 500uC. Carburization was carried out using20% CH4/H2 mixtures and a temperature-programmed procedure. Theoptimum 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 theinitial activity and long-term stability of the Co0.5Mo0.5C catalyst was superiorto 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 itspotential as a sulfur-tolerant LTS catalyst, similar to the Co-Mo sour gas shiftcatalysts (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 producehydrogen. The utilization of coal and biomass for the production of electrical



power, chemicals, petrochemicals, and hydrogen and transportation fuels isgaining importance. Declining resources and increasing prices of crude oil aresome of the major driving forces. In future, due to considerations of globalwarming, coal may have to be used only in CO2- free power plants, which canonly work in combination with CO2 sequestration. This is only possible whenall carbon compounds in the feedstock are converted to CO2. CO conversionprocesses (like the water gas shift reaction) play a key role here. Some sourcespredict that by the year 2030, 10% of the yearly consumption of energy willoriginate from the WGS reaction (222). The WGS reaction is a well-establishedprocess 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 kineticallylimited at low temperatures and thermodynamically limited at hightemperatures. Due to intensive efforts during the last two decades, significantprogress has been made in the study of the mechanism of the WGS reaction(223). At high temperatures (above 350uC) and over the iron oxide-basedcatalysts, the redox mechanism, involving the reduction of the catalyst by COand H2 and its reoxidation by H2O probably prevails. At lower temperatures,even though the detailed mechanism is not established definitively, thefollowing picture is beginning to emerge: The mechanism and the ratedetermining 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 withthe surface OH groups on the support to form surface species, like theformates, carbonates and carboxylates. The concentration and stability ofthese species depend on the support oxide, temperature and the partialpressures of the reactants, especially H2O. Some of these surface species (likethe formates and carbonates) are also intermediates in the reaction path anddecompose to CO2 at higher temperatures and/or partial pressures of H2O.The decomposition of these various surface species to CO2 is a crucial and slowstep in the reaction path. It is faster at higher temperatures and partialpressures of H2O. The nature of the catalyst (metal type and loading) andprocess conditions have a profound influence on the decomposition of theseintermediates. The accumulation of these species on the surface and, theconsequent, blocking of the catalytically active sites by them lead to loss ofcatalytic activity. Factors that hasten their removal, by conversion to CO2, willimprove the catalytic performance of these catalysts. The primary reason forthe accumulation of these intermediates on the surface (and indirectly to lossof catalytic activity) must be sought in the physicochemical changes under-gone by the catalyst (e.g., loss of metal-support interface area). Basic additiveslike the alkali ions facilitate the decomposition of formates and improve thelow temperature performance of catalysts, like Pt-Ceria, Pt-Ceria-Zirconia



and Pt-Titania. Similarly, acidic additives, like the oxides of Nb, Mo, Ta and Wfacilitate the decomposition of the carbonates to CO2 and accelerate the rate.The strong adsorption and binding of CO on metals like Pt, at lowtemperatures below 250uC, is, also, a rate – inhibiting factor and contributeto the low activity and slow deactivation of these catalysts at lowtemperatures. Development of WGS catalysts kinetically more active in the190–250uC range is a general challenge. While the Cu-ZnO catalysts are activein this range, their activity is low necessitating the use of low GHSVs (3000–5000h21). A more precise understanding of the mechanism will lead to thedevelopment of better WGS catalysts for hydrogen generation in fuel cells.There are other challenges to develop improved WGS catalyst and processversions. Some of them are:

Challenges in High Temperature Shift: (1) Replacing chromium in theiron oxide- based catalyst by a non-toxic promoter; even though manychromium-free formulations(containing copper, for example) are in thepipeline, they are not, yet, proven in commercial usage. (2) Discovery of anovel, non- noble metal catalyst with higher catalytic activity that will enableoperation at GHSV 5 40,000 h21 and above; the iron oxide- based catalystsoperate at , 15000h21. This requirement is especially relevant to fuel cellapplications; (3) Discovery of catalysts that can function successfully at lowsteam to gas ratios will reduce, significantly, the energy costs associated withhydrogen generation.

Challenges in Low Temperature Shift: There are, at least, three majordrawbacks in the use of Cu - ZnO – Al2O3 catalysts even in conventional,stationary applications: (1) their relatively, low catalytic activity (GHSVs ofaround only 3000 – 5000 h 21 leading to large – volume catalyst beds), (2) theirelaborate start-up, activation procedures and, (3) their high sensitivity tosulfur and chlorine compounds as well as to steam below the dew point of H2Oat the operating temperature and pressure. Current pressures for manydownstream 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 asuitable catalyst can be discovered. Even though well- formulated, modern,Cu-ZnO-Al2O3 catalysts produce only minor amounts of methanol, it isdesirable to reduce this quantity still further or, even better, eliminatemethanol formation altogether. The above constraints are much moreimportant for LTS catalysts in mobile fuel cell applications. It is mainlybecause of these and other reasons (like the pyrophoricity of the copper –based catalysts) that the noble metal – reducible oxide catalysts are beinginvestigated as potential alternatives.

Challenges in Fuel Cell Applications: Even though noble metal –based catalysts for fuel processors are already in the market, a completely



satisfactory catalyst for WGS applications in fuel cells is not yet in commercialoperation. Ceria and titania- based platinum catalysts are the front runners aspotential water gas catalysts in fuel cell applications. Apart from their highcosts, some of their major drawbacks include their low activity below 250uCand deactivation in long – term operations, especially at lower temperaturesand high pressures. Formation of hydrocarbons at low temperatures and highpressures (Fischer – Tropsch activity) is yet another drawback of thesecatalysts. The present copper – based catalysts do not form significantamounts of hydrocarbons (like methane) under LTS conditions. Attempts toimprove the long-term life of noble metal catalysts by incorporating acid orbasic additives in the oxide support have been described above (see Section 9).Modifying the electronic properties of the noble metal, by alloying surface Ptatoms 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. Recenttheoretical calculations, by Knudsen et al. (224) indicate that this may indeedbe a promising approach. These authors find that a Cu-Pt surface-alloy bindsCO more weakly than pure Pt (Figure 52) and, hence, CO poisoning at lowtemperatures is less likely with the alloy than with the pure metal; in atemperature programmed desorption experiment, adsorbed CO desorbs atlower 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 H2Omore 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).



several metal surfaces. The higher WGS activity of Pt-Cu compared to Pt-alone catalysts had already been claimed by researchers from the HondaMotor Company (152). Similarly, the higher activity of Pt-Re and Pt-Agcompared to Pt catalysts has also been claimed by Sud-Chemie Inc. (225).Inaddition to the above challenges, development of a single catalyst formulationfor both the hydrocarbon reforming and water gas shift reactions will go a longway 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 thehigh temperature reforming/ partial oxidation as well as the low temperaturewater gas shift reaction, will have to be discovered and developed to meet thischallenge. Major efforts in this direction are in progress worldwide.

Challenges in Fundamental studies: There are also more basicfundamental problems and key issues that remain to be addressed andclarified on the mechanistic aspects of the water gas shift, especially the LTSreaction; (a) The role of oxygen mobility in the oxide component of the noblemetal- partially reducible oxide catalyst needs to be investigated further; inthe redox mechanism, the role of oxygen mobility is very clear and obvious.However, the role of surface oxygen mobility is also crucial in the associativemechanism because intermediates (e.g., formates and carbonates) are boundto the oxide by their surface oxygen atoms and presumably move across theoxide 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 andassociative mechanisms; (b) The need to confirm, unambiguously, that H2Ocan indeed reoxidize partially reduced cerium oxide at low temperatures (150–250uC). Such a reoxidation is a fundamental assumption in the redoxmechanism and its occurrence at high temperatures is well established inthe case of the Fe2O3-Cr2O3 catalysts. There has been no confirmatoryevidence of such a reoxidation process at low temperatures in the case ofpartially reduced cerium oxide. Does the presence of a metal (like platinum)facilitate the reoxidation of partially reduced ceria by water molecules at lowtemperatures? While it is known that Pt enhances the reduction of ceria byhydrogen and creation of surface oxygen vacancies and OH groups, it is notconfirmed that Pt also facilitates the reoxidation of the reduced ceria by H2Oat low temperatures. These experiments are crucial to confirm the redoxmechanism at low temperatures; (c) The need to establish a standard protocolto estimate the noble metal dispersion when they are supported on partiallyreducible oxides like ceria, titania, ceria-zirconia, chromia etc.; while manylabs(especially in the industry) have evolved in-house empirical methods forcatalyst screening purposes, a more scientific foundation is desirable;(d) Theneed to model, more accurately, the transition state of formate decompositionin the presence of co-adsorbed water. In the absence of water, formate is quite



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 (asduring the WGS reaction), the formate decomposes very rapidly even below150uC 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 thetheoretical models developed so far have not explicitly considered this majorinfluence of co-adsorbed water molecules in enhancing the decomposition offormate ions to CO2 and H2 at temperatures typical of the WGS reaction andhave treated the decomposition only as a thermal decomposition; hence,theoretical calculations taking into account the original transition statepicture 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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