why only few redox couple can oxidize water?
Answers
Redox (oxidation-reduction)
When molecular hydrogen (H2) is oxidized by molecular oxygen (O2) to form water (H2O), the reaction are considered as two coupled processes; the transfer of electrons from the hydrogen to the oxygen (reduction of the oxygen) and acceptance of electrons from the hydrogen by oxygen (oxidation of the hydrogen). Oxygen is the oxidizing agent and hydrogen is the reducing agent. Oxygen and hydrogen molecules do not have to be involved in redox reactions, but the movement of electrons between the reacting chemical species is central. Redox reactions in aqueous solutions are of the greatest importance in biological and environmental systems. They support and maintain life by gathering and dissipating energy to generate and propagate low-entropy living systems.
In oxidation-reduction (redox) reactions, the ability to donate or accept electrons is given by the redox potential, E. Here, E° is the standard (electrode, reduction or redox) potential at 25 °C, the measure of individual potential of the reversible electrode at standard state which, in this case, is 1 M and gases at a pressure of 101,325 Pa. Zero current is drawn when the potential is measured. E°' is this standard potential but at pH 7.0. At each electrode, the (electrode) potential is given by the Nernst equation,
Nernst equation for the cell
where F = Faraday constant (96,485 J ˣ V-1 ˣ mol-1 = 96,485 C ˣ mol-1 = 96,485 A ˣ s ˣ mol-1), R is the gas constant, n = number of electrons transferred and the Π[Aoxidized] and Π[Areduced] terms refer to all the concentration terms (multiplied) of the oxidized and reduced materials in the equation. More correctly, the activity terms should be used in place of the concentrations.
A positive redox potential indicates the ability to accept electrons (i.e., it is an oxidizing agent, oxidant) and a negative redox potential indicates the ability to donate electrons under those conditions (i.e., it is a reducing agent, reductant). The electrode potential cannot be determined on its own but only as part of a cell containing two electrodes, where the overall potential is the sum of the individual electrode potentials. Within the cell, negative electrons are passed from the electrode of lower (more negative) E (cathode) to the electrode with a more positive E (anode); and these electrons are then returned via the external circuit.
Redox reactions can be redrafted as the sum of the half-reactions for the oxidation of the reductant and the reduction of the oxidant. The half-reactions are: (the electrons (e-) are from an electrode)
(1) reduction of the oxidant half-reaction
O2 + 4H3O+ + 4e- [equilibrium arrows] 6H2O
¼O2 + H3O+ + e- [equilibrium arrows] 1½H2O equivalent reaction
easy reduction of oxygen --> back arrow hard oxidation of water
for either equivalent reactions
E° = +1.229 V
E = +1.229 V + 0.00642 ˣ 4 ˣ Ln([H3O+]) V
E = +1.229 V + 0.0257 ˣ (-2.303 ˣ pH) V
therefore at pH 7 and unit activity oxygen,
E°' = +1.229 V - 0.0257 ˣ 2.303 ˣ 7 V = +0.815 V.