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Reduction of oxyggen and oxidation of hydrogen

Reduction of oxyggen and oxidation of hydrogen to form water

Water Redox processes

The driving forces of life's processes are redox (oxidation-reduction) reactions involving the transfer of electrons.

link; Thermodynamics

V Redox (oxidation-reduction)
V The redox potential of water

link; Redox reactions of O2


OIL RIGoxidation is loss of electrons, reduction is gain of electrons

Redox (oxidation-reduction)

When molecular hydrogen (H2) is oxidized by molecular oxygen (O2) to form water (H2O), the reaction may be 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.


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 [Aoxidized] and [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) and a negative redox potential indicates the ability to donate electrons under those conditions (i.e. it is a reducing agent). 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.


The half reactions are: (the electrons (e-) are from an electrode)


O2 + 4H3O+ + 4e- [equilibrium arrows] 6H2O

    easy reduction of oxygen -->              back arrow hard oxidation of water


E° = +1.229 V


therefore at pH 7 and unit activity oxygen, E°' = +1.229 V + 2.303 ˣ (-7) ˣ 0.0257 V = +0.815 V.



2H3O+ + 2e- [equilibrium arrows] H2 + 2H2O

hard reduction of hydrogen ions -->              back arrow easy oxidation of hydrogen


E° = 0.00 V (the standard hydrogen electrode; exactly zero by definition)



therefore at pH 7 and unit activity hydrogen, E°' = 0.00 V + 2.303 ˣ (-7) ˣ 0.0257 V = -0.414 V.

Noting that a compound with the more positive potential will oxidize the reduced form of a substance of lower (more negative) potential., then electrons flow from negative potential to positive potential. In this example, the electrons flow from the hydrogen (E°' = -0.414 V) to the oxygen (E°' = +0.815 V).


O2 + 4H3O++ 4e- -> 6H2O                E°' = +0.815 V
           2H2 + 4H2O -> 4H3O++ 4e-      E°' = -(-0.414 V)


                                               2H2 + O2 -> 2H2O                 E°' = +0.815 + 0.414V = 1.229 V     


The redox potentials are related to the Gibbs free energy (ΔG) by:


ΔG = -nFE


and E is the net redox potential (working cell potential in volts), related to the conditions by the appropriate Nernst equation,

   ΔG°' = -nFE°'
    ΔG°' = -4 ˣ 96.49 ˣ 1.229 kJ ˣ mol-1
    ΔG°' = -474.3 kJ ˣ mol-1


Consider burning hydrogen gas                      2H2 + O2 -> 2H2O  + energy                                            (a)


where ΔG°' (the standard free energy change at pH 7.0) = -474.3 kJ ˣ mol-1. The enthalpy of this reaction (ΔH°') is -571.66 kJ ˣ mol-1 (heat is given out) and the entropy (ΔS°') is -326.8 J K-1 (fewer molecules produced, therefore more order and less entropy). ΔG°' = ΔH°' - TΔS°' = -571.66 + 0.3268 ˣ 298 = -474.3 kJ ˣ mol-1. To reverse this reaction (a) requires a large input of energy from the electrolysis process in order to achieve both the oxidation of water and the reduction of hydrogen ions. If heat is not supplied from the environment, the heat required for the entropy change will cool the electrolytic process.

Redox potential of water

Redox sensor


Schematic of a typical ORP electrode
The redox potential of liquid water varies over a range of more than two volts according to the solutes it contains. This potential can be determined using an Oxidation-Reduction Potential (ORP) electrode (see right) and is a measure of the collective redox potential of everything in the water, including dissolved gasses such as oxygen. The potential of the solution is determined relative to the standard potential generated by the reference electrode a and then corrected for that potential.


The key redox half-reaction in liquid water is,

O2 + 4H3O++ 4e- -> 6H2O                 E°' = +0.815 V


The redox potentials of aqueous solutions clearly depend on both the dissolved oxygen and hydrogen ion concentrations.


              As the concentration of molecular oxygen increases, the redox potential increases

              As the concentration of molecular oxygen decreases, the redox potential decreases

              As the concentration of hydrogen ions increases (and pH decreases), the redox potential increases

              As the concentration of hydrogen ions decreases (and pH increases), the redox potential decreases


A decrease of one pH unit is accompanied by an increase in redox potential of 58 mV. Anoxic waters may have negative redox potentials.


The approximate redox potential of some solutions

Aqueous material Redox potential, mV Aqueous material Redox potential, mV
Water saturated with H2 -600 Distilled water +250
Human cells -170 ~ -290 Mineral water +300 ~ +400
Anaerobic yeast fermentation -180 Tap water +300 ~ +400
Anaerobic soil -100 Surface seawater ~+400
Green tea -100 Deep seawater (~2000 m) ~+430
Mother's milk -70 Swimming pool +400 ~ +475
Degassed pure water +200 Water saturated with O2 +600


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a The electrode reaction for the Ag/AgCl/saturated KCl reference electrode is

AgCl + e- -> Ag0 + Cl-                E = +0.197 V (in saturated KCl)


The potential is relative to the Standard Hydrogen Electrode (SHE)

2 H+ + 2 e- -> H2              E0 = 0 V at all temperatures

where the electrode is platinum foil covered in platinum black (finely divided platinum). The H+ is unit activity (~ 1 m HCl) and the gas pressure is 1 atm H2 through a bubbler. For practical purposes, if a hydrogen electrode is needed, the Normal Hydrogen Electrode (NHE) is used; where the potential is calculated from the acid strength used and the ambient pressure [3380].




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This page was established in 2015 and last updated by Martin Chaplin on 1 August, 2018

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