Heterogeneous water oxidation

Water oxidation is one of the half reactions of water splitting:

2H2O → O2 + 4H+ + 4e−   Oxidation (generation of dioxygen)

4H+ + 4e− → 2H2  Reduction (generation of dihydrogen)

2H2O → 2H2 + O2  Total reaction

Of the two half-reactions, the oxidation step is the most demanding because it requires the coupling of four electron and proton transfers and the formation of an oxygen-oxygen bond. This process occurs naturally in plants photosystem II to provide protons and electrons for the photosynthesis process and release oxygen to the atmosphere. Since hydrogen can be used as an alternative clean burning fuel, there has been a need to split water efficiently. However, there are known materials that can mediate the reduction step efficiently therefore much of the current research is aimed at the oxidation half reaction also known as the Oxygen Evolution Reaction (OER). Current research focuses on understanding the mechanism of OER and development of new materials that catalyze the process.[1]

Thermodynamics

Both the oxidation and reduction steps are pH dependent. Figure 1 shows the standard potentials at pH 0 (strongly acidic) as referenced to the normal hydrogen electrode (NHE).

2 half reactions (at pH = 0)
Oxidation 2H2O → 4H+ + 4e− + O2 E° = 1.23 V vs. NHE

Reduction 4H+ + 4e− → 2H2 E° = 0.00 V vs. NHE

Overall 2H2O → 2H2 + O2 E°cell = 1.23 V; ΔG = -475 kJ/mol


Water splitting can be done at higher pH values as well however the standard potentials will vary according to the Nernst equation and therefore shift by -59 mV for each pH unit increase. However, the total cell potential (difference between oxidation and reduction half cell potentials) will remain 1.23 V. This potential can be related to Gibbs free energy (ΔG) by:

ΔG°cell = −nFE°cell

Where n is the number of electrons per mole products and F is the Faraday constant. Therefore, it takes 475 kJ of energy to make one mole of O2 as calculated by thermodynamics. However, in reality no process can be this efficient. Systems always suffer from an overpotential that arise from activation barriers, concentration effects and voltage drops due to resistance. The activation barriers or activation energy is associated with high energy transition states that are reached during the electrochemical process of OER. The lowering of these barriers would allow for OER to occur at lower overpotentials and faster rates.

Mechanism

Heterogeneous OER is sensitive to the surface which the reaction takes place and is also affected by the pH of the solution. The general mechanism for acidic and alkaline solutions is shown below. Under acidic conditions water binds to the surface with the irreversible removal of one electron and one proton to form a platinum hydroxide.[2] In an alkaline solution a reversible binding of hydroxide ion coupled to a one electron oxidation is thought to precede a turnover-limiting electrochemical step involving the removal of one proton and one electron to form a surface oxide species.[3] The shift in mechanism between the pH extremes has been attributed to the kinetic facility of oxidizing hydroxide ion relative to water. Using the Tafel equation, one can obtain kinetic information about the kinetics of the electrode material such as the exchange current density and the Tafel slope.[4] OER is presumed to not take place on clean metal surfaces such as Platinum, but instead an oxide surface is formed prior to oxygen evolution.[5]

OER under acidic conditions.
OER under alkaline conditions.

Catalyst Materials

OER has been studied on a variety of materials including:

Preparation of the surface and electrolysis conditions have a large effect on reactivity (defects, steps, kinks, low coordinate sites) therefore it is difficult to predict an OER material's properties by its bulk structure. Surface effects have a large influence on the kinetics and thermodynamics of OER.

Platinum

Platinum has been a widely studied material for OER because it is the catalytically most active element for this reaction.[13] It exhibits exchange current density values on the order of 10−9 A/cm2. Much of the mechanistic knowledge of OER was gathered from studies on platinum and its oxides.[2] It was observed that there was a lag in the evolution of oxygen during electrolysis. Therefore, an oxide film must first form at the surface before OER begins.[5] The Tafel slope, which is related to the kinetics of the electrocatalytic reaction, was shown to be independent of the oxide layer thickness at low current densities but becomes dependant on oxide thickness at high current densities [14]

Ruthenium oxide

Ruthenium oxide (RuO2) shows some of the best performance as an OER material in acidic environments. It has been studied since the early 1970s as a water oxidation catalyst with one of the lowest reported overpotentials for OER at the time.[15] It has since been investigated for OER in Ru(110) single crystal oxide surfaces,[16] compact films,[17] Titanium supported films.[18] RuO2 films can be prepared by thermal decomposition of ruthenium chloride on inert substrates.[17]

References

  1. ↑ "Anode - Lewis Research Group". Nsl.caltech.edu. Retrieved 2012-08-05.
  2. 1 2 3 Conway, B. E. and Liu, T. C. Langmuir 1990, 6, 268– 276
  3. ↑ Birss, V. I., Damjanovic, A., and Hudson, P. G. J. Electrochem. Soc. 1986, 133, 1621
  4. ↑ K. Zeng, D. Zhang / Progress in Energy and Combustion Science 36 (2010) 307–326
  5. 1 2 1980 Damjanovic J. electrochem. Soc. 127, 1945
  6. ↑ Matsumoto, Y. and Sato, E. Mater. Chem. Phys. 1986, 14, 397– 426
  7. ↑ Reaction Kinetics and Catalysis Letters Volume 21, Number 3, 195
  8. ↑ Bockris, J. O’M. and Otagawa, T. J. Phys. Chem. 1983, 87, 2960
  9. ↑ Das et al. Angew. Chem. Int. Ed., 2013, 52, 7224-7227
  10. ↑ Terebilenko K. V., Bychkov K.L. Dalton Trans,2016,45,3895-3904
  11. ↑ http://onlinelibrary.wiley.com/doi/10.1002/anie.201301327/abstract
  12. ↑ Hansen and Das, Energy & Environ Sci.(http://pubs.rsc.org/en/content/articlelanding/2013/ee/c3ee43040e#!divAbstract)
  13. ↑ H. Dau, C. Limberg, T. Reier, M. Risch, S. Roggan, P. Strasser, ChemCatChem 2010, 2, 724–761.
  14. ↑ Birss, V. I., Damjanovic, A. J. Electrochem. Soc. 134, 113.
  15. ↑ S. Trasatti and G. Buzzanca, J. Electroanal. Chem. 29 (1971), p. 1.
  16. ↑ P. Castelli, S. Trasatti, F. H. Pollak, W. E. O’Grady, J. Electroanal. Chem. 1986, 210, 189.
  17. 1 2 G. Lodi, E. Sivieri, A. D. Battisti, S. Trasatti, J. Appl. Electrochem. 1978, 8, 135.
  18. ↑ S. Trasatti, Electrochim. Acta 2000, 45, 2377.
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