Editing Electrolysis (section)

==Competing half-reactions in solution electrolysis==
Using a cell containing inert platinum electrodes, electrolysis of aqueous solutions of some salts leads to the reduction of the cations (such as metal deposition with, for example, zinc salts) and oxidation of the anions (such as the evolution of bromine with bromides). However, with salts of some metals (such as sodium) hydrogen is evolved at the cathode, and for salts containing some anions (such as sulfate {{chem|SO|4|2−}}) oxygen is evolved at the anode. In both cases, this is due to water being reduced to form hydrogen or oxidized to form oxygen.
In principle, the voltage required to electrolyze a salt solution can be derived from the [[standard electrode potential]] for the reactions at the anode and cathode. The standard electrode potential is directly related to the [[Gibbs free energy]], Δ''G'', for the reactions at each electrode and refers to an electrode with no current flowing. An extract from the [[table of standard electrode potentials]] is shown below.

:{| class="wikitable"
|-
! [[Half-reaction]]
! ''E''° (V)
! Ref.
|-
| [[Sodium|Na]]<sup>+</sup> + e<sup>−</sup> {{eqm}} Na{{abbr|(s)|solid}}
| −2.71 ||<ref name="Atk">{{cite book|last1=Atkins|first1=Peter|year=1997|title=Physical Chemistry|edition=6|publisher=W.H. Freeman and Company|location=New York}}</ref>
|-
| [[Zinc|Zn]]<sup>2+</sup> + 2 e<sup>−</sup> {{eqm}} Zn{{abbr|(s)|solid}}
| −0.7618 ||<ref name="van88">{{cite book|last=Vanýsek|first=Petr|year=2007|url=http://www.hbcpnetbase.com/articles/08_08_88.pdf|chapter=Electrochemical Series|archive-url=https://web.archive.org/web/20170724011402/http://www.hbcpnetbase.com/|archive-date=24 July 2017|title=Handbook of Chemistry and Physics|edition=88|publisher=Chemical Rubber Company}}</ref>
|-
| '''2 H<sup>+</sup> + 2 e<sup>−</sup> {{eqm}} H<sub>2</sub>{{abbr|(g)|gaseous}}'''
| '''≡ 0'''||<ref name="van88"/>
|-
| Br<sub>2</sub>{{abbr|(aq)|aqueous}} + 2 e<sup>−</sup> {{eqm}} 2 Br<sup>−</sup>
| +1.0873 ||<ref name=van88/>
|-
| O<sub>2</sub>{{abbr|(g)|gaseous}} + 4 H<sup>+</sup> + 4 e<sup>−</sup> {{eqm}} 2 H<sub>2</sub>O
| +1.23 ||<ref name="Atk"/>
|-
| Cl<sub>2</sub>{{abbr|(g)|gaseous}} + 2 e<sup>−</sup> {{eqm}} 2 Cl<sup>−</sup>
| +1.36 ||<ref name="Atk"/>
|-
| {{chem|S|2|O|8|<sup>2−</sup>}} + 2 e<sup>−</sup> {{eqm}} 2 {{chem|SO|4|2−}}
| +2.07 ||<ref name="Atk"/>
|}
In terms of electrolysis, this table should be interpreted as follows:

* Moving ''down'' the table, ''E''° becomes more positive, and species on the ''left'' are more likely to be ''reduced'': for example, zinc ions are more likely to be reduced to zinc metal than sodium ions are to be reduced to sodium metal.
* Moving ''up'' the table, ''E''° becomes more negative, and species on the ''right'' are more likely to be ''oxidized'': for example, sodium metal is more likely to be oxidized to sodium ions than zinc metal is to be oxidized to zinc ions.

Using the [[Nernst equation]] the [[electrode potential]] can be calculated for a specific concentration of ions, temperature and the number of electrons involved. For pure water ([[pH]]&nbsp;7):
* the electrode potential for the reduction producing hydrogen is −0.41&nbsp;V,
* the electrode potential for the oxidation producing oxygen is +0.82&nbsp;V.
Comparable figures calculated in a similar way, for 1&nbsp;M [[zinc bromide]], ZnBr<sub>2</sub>, are −0.76&nbsp;V for the reduction to Zn metal and +1.10&nbsp;V for the oxidation producing bromine.
The conclusion from these figures is that hydrogen should be produced at the cathode and oxygen at the anode from the electrolysis of water—which is at variance with the experimental observation that zinc metal is deposited and bromine is produced.<ref>{{cite journal|doi=10.1021/ed029p319.1|title=A Textbook of Quantitative Inorganic Analysis (Vogel, Arthur I.) |year=1952 |last1=Hall |first1=Norris F. |journal=Journal of Chemical Education |volume=29 |issue=6 |page=319 |bibcode=1952JChEd..29..319H |doi-access=free }}</ref>
The explanation is that these calculated potentials only indicate the thermodynamically preferred reaction. In practice, many other factors have to be taken into account such as the kinetics of some of the reaction steps involved. These factors together mean that a higher potential is required for the reduction and oxidation of water than predicted, and these are termed [[overpotential]]s. Experimentally it is known that overpotentials depend on the design of the cell and the nature of the electrodes.

For the electrolysis of a neutral (pH&nbsp;7) sodium chloride solution, the reduction of sodium ion is thermodynamically very difficult and water is reduced evolving hydrogen leaving hydroxide ions in solution. At the anode the oxidation of chlorine is observed rather than the oxidation of water since the overpotential for the oxidation of [[chloride]] to [[chlorine]] is lower than the overpotential for the oxidation of [[water]] to [[oxygen]]. The [[hydroxide ion]]s and dissolved chlorine gas react further to form [[hypochlorous acid]]. The aqueous solutions resulting from this process is called [[electrolyzed water]] and is used as a disinfectant and cleaning agent.