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jMThe decomposition of pure water into hydrogen and oxygen at standard temperature and pressure is not favorable in thermodynamic terms. Anode (oxidation): 2 H2O(l) → O2(g) + 4 H+(aq) + 4e− Eo = +1.23 V (for the reduction half-equation)[7] Cathode (reduction): 2 H+(aq) + 2e− → H2(g) Eo = 0.00 V Thus, the standard potential of the water electrolysis cell (Eocell = Eocathode − Eoanode) is −1.229 V at 25 °C at pH 0 ([H+] = 1.0 M). At 25 °C with pH 7 ([H+] = 1.0×10−7 M), the potential is unchanged based on the Nernst equation. The thermodynamic standard cell potential can be obtained from standard-state free energy calculations to find ΔG° and then using the equation: ΔG°= −n F E° (where E° is the cell potential and F the Faraday constant, i. e. 96,485.3321233 C/mol). For two water molecules electrolysed and hence two hydrogen molecules formed, n = 4, and ΔG° = 474.48 kJ/2 mol(water) = 237.24 kJ/mol(water). However, calculations regarding individual electrode equilibrium potentials requires some corrections taking into account the activity coefficients.[8] In practice when an electrochemical cell is "driven" toward completion by applying reasonable potential, it is kinetically controlled. Therefore, activation energy, ion mobility (diffusion) and concentration, wire resistance, surface hindrance including bubble formation (causes electrode area blockage), and entropy, require a greater applied potential to overcome these factors. The amount of increase in potential required is termed the overpotential.