# Water-gas shift reaction

The water-gas shift reaction (WGSR) describes the reaction of carbon monoxide and water vapor to form carbon dioxide and hydrogen:

CO + H2O ⇌ CO2 + H2

The water gas shift reaction was discovered by Italian physicist Felice Fontana in 1780. It was not until much later that the industrial value of this reaction was realized. Before the early 20th century, hydrogen was obtained by reacting steam under high pressure with iron to produce iron, iron oxide and hydrogen. With the development of industrial processes that required hydrogen, such as the Haber–Bosch ammonia synthesis, a less expensive and more efficient method of hydrogen production was needed. As a resolution to this problem, the WGSR was combined with the gasification of coal to produce a pure hydrogen product. As the idea of hydrogen economy gains popularity, the focus on hydrogen as a replacement fuel source for hydrocarbons is increasing.

## Applications

The WGSR is an important industrial reaction that is used in the manufacture of ammonia, hydrocarbons, methanol, and hydrogen. It is also often used in conjunction with steam reforming of methane and other hydrocarbons. In the Fischer–Tropsch process, the WGSR is one of the most important reactions used to balance the H2/CO ratio. It provides a source of hydrogen at the expense of carbon monoxide, which is important for the production of high purity hydrogen for use in ammonia synthesis.

The water-gas shift reaction may be an undesired side reaction in processes involving water and carbon monoxide, e.g. the rhodium-based Monsanto process. The iridium-based Cativa process uses less water, which suppresses this reaction.

### Fuel cells

The WGSR can aid in the efficiency of fuel cells by increasing hydrogen production. The WGSR is considered a critical component in the reduction of carbon monoxide concentrations in cells that are susceptible to carbon monoxide poisoning such as the proton-exchange membrane (PEM) fuel cell.[1] The benefits of this application are two-fold: not only would the water gas shift reaction effectively reduce the concentration of carbon monoxide, but it would also increase the efficiency of the fuel cells by increasing hydrogen production.[1] Unfortunately, current commercial catalysts that are used in industrial water gas shift processes are not compatible with fuel cell applications.[2] With the high demand for clean fuel and the critical role of the water gas shift reaction in hydrogen fuel cells, the development of water gas shift catalysts for the application in fuel cell technology is an area of current research interest.

Catalysts for fuel cell application would need to operate at low temperatures. Since the WGSR is slow at lower temperatures where equilibrium favors hydrogen production, WGS reactors require large amounts of catalysts, which increases their cost and size beyond practical application.[1] The commercial LTS catalyst used in large scale industrial plants is also pyrophoric in its inactive state and therefore presents safety concerns for consumer applications.[2] Developing a catalyst that can overcome these limitations is relevant to implementation of a hydrogen economy.

## Reaction conditions

The equilibrium of this reaction shows a significant temperature dependence and the equilibrium constant decreases with an increase in temperature, that is, higher hydrogen formation is observed at lower temperatures.

### Temperature dependence

A plot of the temperature dependence of Keq, the value of which approaches 1 at 1100 K.[2]

The water gas shift reaction is a moderately exothermic reversible reaction. Therefore, with increasing temperature the reaction rate increases but the carbon dioxide production becomes less favorable.[3] Due to its exothermic nature, high carbon monoxide percentage is thermodynamically favored at low temperatures. Despite the thermodynamic favorability at low temperatures, the reaction is faster at high temperatures. The water-gas shift reaction is sensitive to temperature, with a tendency to shift towards carbon monoxide as temperature increases due to Le Chatelier's principle. Over the temperature range 600–2000 K, the equilibrium constant for the WGSR has the following relationship:[2]

${\displaystyle K_{\mathrm {eq} }=10^{-2.4198+0.0003855T+{\frac {2180.6}{T}}}}$

## Practical concerns

In order to take advantage of both the thermodynamics and kinetics of the reaction, the industrial scale water gas shift reaction is conducted in multiple adiabatic stages consisting of a high temperature shift (HTS) followed by a low temperature shift (LTS) with intersystem cooling.[4] The initial HTS takes advantage of the high reaction rates, but is thermodynamically limited, which results in incomplete conversion of carbon monoxide and a 2-4% carbon monoxide exit composition. To shift the equilibrium toward hydrogen production, a subsequent low temperature shift reactor is employed to produce a carbon monoxide exit composition of less than 1%. The transition from the HTS to the LTS reactors necessitates intersystem cooling. Due to the different reaction conditions, different catalysts must be employed at each stage to ensure optimal activity. The commercial HTS catalyst is the iron oxide–chromium oxide catalyst and the LTS catalyst is a copper-based catalyst. The order proceeds from high to low temperature due to the susceptibility of the copper catalyst to poisoning by sulfur that may remain after the steam reformation process.[2] This necessitates the removal of the sulfur compounds prior to the LTS reactor by a guard bed in order to protect the copper catalyst. Conversely, the iron used in the HTS reaction is generally more robust and resistant toward poisoning by sulfur compounds. While both the HTS and LTS catalysts are commercially available, their specific composition varies based on vendor. An important limitation for the HTS is the H2O/CO ratio where low ratios may lead to side reactions such as the formation of metallic iron, methanation, carbon deposition, and the Fischer–Tropsch reaction.

### Low temperature shift

The typical composition of a commercial LTS catalyst has been reported as 32-33% CuO, 34-53% ZnO, 15-33% Al2O3.[2] The active catalytic species is CuO. The function of ZnO is to provide structural support as well as prevent the poisoning of copper by sulfur. The Al2O3 prevents dispersion and pellet shrinkage. The LTS shift reactor operates at a range of 200–250 °C. The upper temperature limit is due to the susceptibility of copper to thermal sintering. These lower temperatures also reduce the occurrence of side reactions that are observed in the case of the HTS. Noble metals such as platinum, supported on ceria, have also been used for LTS.[5]

### High temperature shift catalysts

The typical composition of commercial HTS catalyst has been reported as 74.2% Fe2O3, 10.0% Cr2O3, 0.2% MgO (remaining percentage attributed to volatile components).[6] The chromium acts to stabilize the iron oxide and prevents sintering. The operation of HTS catalysts occurs within the temperature range of 310 °C to 450 °C. The temperature increases along the length of the reactor due to the exothermic nature of the reaction. As such, the inlet temperature is maintained at 350 °C to prevent the exit temperature from exceeding 550 °C. Industrial reactors operate at a range from atmospheric pressure to 8375 kPa (82.7 atm).[6] The search for high performance HT WGS catalysts remains an intensive topic of research in fields of chemistry and materials science. Activation energy is a key criteria for the assessment of catalytic performance in WGS reactions. To date, some of the lowest activation energy values have been found for catalysts consisting of copper nanoparticles on ceria support materials,[7] with values as low as Ea = 34 kJ/mol reported relative to hydrogen generation.

## Mechanism

Proposed associative and redox mechanisms of the water gas shift reaction[8][9][10]

The WGSR has been extensively studied for over a hundred years. The kinetically relevant mechanism depends on the catalyst composition and the temperature.[4][11] Two mechanisms have been proposed: an associative Langmuir–Hinshelwood mechanism and a redox mechanism. The redox mechanism is generally regarded as kinetically relevant during the high-temperature WGSR (> 350 °C) over the industrial iron-chromia catalyst.[3] Historically, there has been much more controversy surrounding the mechanism at low temperatures. Recent experimental studies confirm that the associative carboxyl mechanism is the predominant low temperature pathway on metal-oxide-supported transition metal catalysts.[12][10]

### Associative mechanism

In 1920 Armstrong and Hilditch first proposed the associative mechanism. In this mechanism CO and H2O are adsorbed onto the surface of the catalyst, followed by formation of an intermediate and the desorption of H2 and CO2. In general, H2O dissociates onto the catalyst to yield adsorbed OH and H. The dissociated water reacts with CO to form a carboxyl or formate intermediate. The intermediate subsequently dehydrogenates to yield CO2 and adsorbed H. Two adsorbed H atoms recombine to form H2.

There has been significant controversy surrounding the kinetically relevant intermediate during the associative mechanism. Experimental studies indicate that both intermediates contribute to the reaction rate over metal oxide supported transition metal catalysts.[12][10] However, the carboxyl pathway accounts for about 90% of the total rate owing to the thermodynamic stability of adsorbed formate on the oxide support. The active site for carboxyl formation consists of a metal atom adjacent to an adsorbed hydroxyl. This ensemble is readily formed at the metal-oxide interface and explains the much higher activity of oxide-supported transition metals relative to extended metal surfaces.[10] The turn-over-frequency for the WGSR is proportional to the equilibrium constant of hydroxyl formation, which rationalizes why reducible oxide supports (e.g. CeO2) are more active than irreducible supports (e.g. SiO2) and extended metal surfaces (e.g. Pt). In contrast to the active site for carboxyl formation, formate formation occurs on extended metal surfaces. The formate intermediate can be eliminated during the WGSR by using oxide-supported atomically dispersed transition metal catalysts, further confirming the kinetic dominance of the carboxyl pathway.[13]

### Redox mechanism

The redox mechanism involves a change in the oxidation state of the catalytic material. In this mechanism, CO is oxidized by an O-atom intrinsically belonging to the catalytic material to form CO2. A water molecule undergoes dissociative adsorption at the newly formed O-vacancy to yield two hydroxyls. The hydroxyls disproportionate to yield H2 and return the catalytic surface back to its pre-reaction state.

### Homogeneous models

The mechanism entails nucleophilic attack of water or hydroxide on a M-CO center, generating a metallacarboxylic acid.[1][14]

### Thermodynamics

The WGSR is exergonic, with the following thermodynamic parameters at room temperature (298 K):

Free energy ΔG⊖ = –6.82 kcal ΔH⊖ = –9.84 kcal ΔS⊖ = –10.1 cal/deg

In aqueous solution, the reaction is less exergonic.[15]

## Reverse water-gas shift

Since the water-gas shift reaction is an equilibrium reaction, there isn't a ‘reverse’ water-gas shift reaction. Water gas is defined as a fuel gas consisting mainly of carbon monoxide (CO) and hydrogen (H2). The term ‘shift’ in water-gas shift means changing the water gas composition (CO:H2) ratio. The ratio can be increased by adding CO2 or reduced by adding steam to the reactor.

## References

1. ^ a b c d Vielstich, Wolf; Lamm, Arnold; Gasteiger, Hubert A., eds. (2003). Handbook of fuel cells: fundamentals, technology, applications. New York: Wiley. ISBN 978-0-471-49926-8.
2. Callaghan, Caitlin (2006). Kinetics and catalysis of the water-gas-shift reaction: A Microkinetic and Graph Theoretic Approach (PDF) (PhD). Worcester Polytechnic Institute.
3. ^ a b Ratnasamy, Chandra; Wagner, Jon P. (September 2009). "Water Gas Shift Catalysis". Catalysis Reviews. 51 (3): 325–440. doi:10.1080/01614940903048661. S2CID 98530242.
4. ^ a b Smith R J, Byron; Muruganandam Loganthan; Murthy Shekhar Shantha (2010). "A Review of the Water Gas Shift Reaction". International Journal of Chemical Reactor Engineering. 8: 1–32. doi:10.2202/1542-6580.2238. S2CID 96769998.
5. ^ Jain, Rishabh; Maric, Radenka (April 2014). "Synthesis of nano-Pt onto ceria support as catalyst for water–gas shift reaction by Reactive Spray Deposition Technology". Applied Catalysis A: General. 475: 461–468. doi:10.1016/j.apcata.2014.01.053.
6. ^ a b Newsome, David S. (1980). "The Water-Gas Shift Reaction". Catalysis Reviews: Science and Engineering. 21 (2): 275–318. doi:10.1080/03602458008067535.
7. ^ Rodriguez, J.A.; Liu, P.; Wang, X.; Wen, W.; Hanson, J.; Hrbek, J.; Pérez, M.; Evans, J. (15 May 2009). "Water-gas shift activity of Cu surfaces and Cu nanoparticles supported on metal oxides". Catalysis Today. 143 (1–2): 45–50. doi:10.1016/j.cattod.2008.08.022.
8. ^ Gokhale, Amit A.; Dumesic, James A.; Mavrikakis, Manos (2008-01-01). "On the Mechanism of Low-Temperature Water Gas Shift Reaction on Copper". Journal of the American Chemical Society. 130 (4): 1402–1414. doi:10.1021/ja0768237. ISSN 0002-7863. PMID 18181624.
9. ^ Grabow, Lars C.; Gokhale, Amit A.; Evans, Steven T.; Dumesic, James A.; Mavrikakis, Manos (2008-03-01). "Mechanism of the Water Gas Shift Reaction on Pt: First Principles, Experiments, and Microkinetic Modeling". The Journal of Physical Chemistry C. 112 (12): 4608–4617. doi:10.1021/jp7099702. ISSN 1932-7447.
10. ^ a b c d Nelson, Nicholas C.; Szanyi, János (2020-05-15). "Heterolytic Hydrogen Activation: Understanding Support Effects in Water–Gas Shift, Hydrodeoxygenation, and CO Oxidation Catalysis". ACS Catalysis. 10 (10): 5663–5671. doi:10.1021/acscatal.0c01059.
11. ^ Yao, Siyu; Zhang, Xiao; Zhou, Wu; Gao, Rui; Xu, Wenqian; Ye, Yifan; Lin, Lili; Wen, Xiaodong; Liu, Ping; Chen, Bingbing; Crumlin, Ethan (2017-06-22). "Atomic-layered Au clusters on α-MoC as catalysts for the low-temperature water-gas shift reaction". Science. 357 (6349): 389–393. doi:10.1126/science.aah4321. ISSN 0036-8075. PMID 28642235. S2CID 206651887.
12. ^ a b Nelson, Nicholas C.; Nguyen, Manh-Thuong; Glezakou, Vassiliki-Alexandra; Rousseau, Roger; Szanyi, János (October 2019). "Carboxyl intermediate formation via an in situ-generated metastable active site during water-gas shift catalysis". Nature Catalysis. 2 (10): 916–924. doi:10.1038/s41929-019-0343-2. ISSN 2520-1158. S2CID 202729116.
13. ^ Nelson, Nicholas C.; Chen, Linxiao; Meira, Debora; Kovarik, Libor; Szanyi, János (2020). "In Situ Dispersion of Palladium on TiO2 During Reverse Water–Gas Shift Reaction: Formation of Atomically Dispersed Palladium". Angewandte Chemie International Edition. n/a (n/a). doi:10.1002/anie.202007576. ISSN 1521-3773. PMID 32589820.
14. ^ Barakat, Tarek; Rooke, Joanna C.; Genty, Eric; Cousin, Renaud; Siffert, Stéphane; Su, Bao-Lian (1 January 2013). "Gold catalysts in environmental remediation and water-gas shift technologies". Energy & Environmental Science. 6 (2): 371. doi:10.1039/c2ee22859a.
15. ^ King, A. D.; King, R. B.; Yang, D. B., "Homogeneous catalysis of the water gas shift reaction using iron pentacarbonyl", J. Am. Chem. Soc. 1980, vol. 102, pp. 1028-1032. doi:10.1021/ja00523a020