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Reduction of oxygen 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


'...there is one very important class of chemical reactions that deserves special study. These reactions are oxidation-reduction reactions...'

Linus Pauling, 1957



Oxidation 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 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 [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, 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.


(2) oxidation of the reductant half-reaction

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

                  H3O+ + e- [equilibrium arrows] ½H2 + H2O                equivalent reaction


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


for either equivalent reactions

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


At unit activity H2

E = 0.01284 ˣ 2 ˣ Ln([H3O+]) V

E = 0.0257 ˣ (-2.303 ˣ pH) V


therefore at pH 7 and unit activity hydrogen,

E°' = -0.0257 ˣ 2.303 ˣ 7 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


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


At pH = 7, under standard conditions,

   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 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.

The 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 (pH). More acidic solutions favor aerobic conditions and more positive redox potential, and more alkaline solutions favor anaerobic conditions.


              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
Electrolytic catholyte (H2) -600 Deep well water 0
Water associated with oil deposits -500 Degassed pure water +200
Organic-rich saline -400 Distilled water +250
Euxinic water (H2S) -250 Groundwater +250
Healthy human cells -170 ~ -290 Mineral water +200 ~ +400
Anaerobic yeast fermentation -180 Tap water +220 ~ +380
Anaerobic water-logged soil -100 Surface seawater ≈ +400
Green tea -100 Deep seawater (≈ 2000 m) ≈ +430
Vegetable juice -70 Swimming pool +400 ~ +475
Mother's milk -70 Rainwater +600
Human internal environment -70 Electrolytic anolyte (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 at unit activity (≈ 1 m HCl), and the gas pressure is at 1 atm (101.325 kPa) 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 4 November, 2018

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