Editing Redox (section)

== Redox reactions in soils ==
[[Electron transfer]] reactions are central to myriad processes and properties in soils, and [[redox potential]], quantified as Eh (platinum electrode potential ([[voltage]]) relative to the standard hydrogen electrode) or pe (analogous to pH as −log electron activity), is a master variable, along with pH, that controls and is governed by chemical reactions and biological processes. Early theoretical research with applications to flooded soils and [[paddy rice]] production was seminal for subsequent work on thermodynamic aspects of redox and plant root growth in soils.<ref>{{cite book |doi=10.1016/S0065-2113(08)60633-1 |title=The Chemistry of Submerged Soils |series=Advances in Agronomy |date=1972 |last1=Ponnamperuma |first1=F.N. |volume=24 |pages=29–96 |isbn=978-0-12-000724-0 }}</ref> Later work built on this foundation, and expanded it for understanding redox reactions related to heavy metal oxidation state changes, [[pedogenesis]] and morphology, organic compound degradation and formation, [[free radical]] chemistry, [[wetland]] delineation, [[soil remediation]], and various methodological approaches for characterizing the redox status of soils.<ref>{{cite journal |last1=Bartlett |first1=Richmond J. |last2=James |first2=Bruce R. |date=1991 |title=Redox chemistry of soils |journal=Advances in Agronomy |volume=39 |pages=151–208}}</ref><ref>{{cite book|last1=James |first1=Bruce R. |title=Handbook of soil sciences: properties and processes |edition=second |last2=Brose |first2=Dominic A. |publisher=[[CRC Press]] |year=2012 |isbn=978-1-4398-0305-9 |editor-last1=Huang |editor-first1=Pan Ming |location=Boca Raton, Florida |pages=14-1 -- 14-24 |chapter=Oxidation-reduction phenomena |editor-last2=Li |editor-first2=Yuncong |editor-last3=Sumner |editor-first3=Malcolm E.}}</ref>

<!-- wikipedia is not a textbook==Balancing redox reactions==
Describing the overall electrochemical reaction for a redox process requires a ''balancing'' of the component [[half-reaction]]s for oxidation and reduction. In general, for reactions in aqueous solution, this involves adding [[Proton|H<sup>+</sup>]], [[Hydroxide|OH<sup>−</sup>]], [[Water (molecule)|H<sub>2</sub>O]], and electrons to compensate for the oxidation changes.

===Acidic media===
In acidic aqueous media, H<sup>+</sup> ions and water are added to half-reactions to balance the overall reaction.

For instance, when [[manganese]](II) reacts with [[sodium bismuthate]]:

:{|
|Unbalanced reaction:
|Mn<sup>2+</sup>(aq) + NaBiO<sub>3</sub>(s) → Bi<sup>3+</sup>(aq) + {{chem|MnO|4|−}} (aq)
|-
|Oxidation:
|4&nbsp;H<sub>2</sub>O(l) + Mn<sup>2+</sup>(aq) → {{chem|MnO|4|-}}(aq) + 8&nbsp;H<sup>+</sup>(aq) + 5&nbsp;e<sup>−</sup>
|-
|Reduction:
|2&nbsp;e<sup>−</sup> + 6&nbsp;H<sup>+</sup> + {{chem|BiO|3|-}}(s) → Bi<sup>3+</sup>(aq) + 3&nbsp;H<sub>2</sub>O(l)
|}

The reaction is balanced by scaling the two half-cell reactions to involve the same number of electrons (multiplying the oxidation reaction by the number of electrons in the reduction step and vice versa):

:8&nbsp;H<sub>2</sub>O(l) + 2&nbsp;Mn<sup>2+</sup>(aq) → 2&nbsp;{{chem|MnO|4|-}}(aq) + 16&nbsp;H<sup>+</sup>(aq) + 10&nbsp;e<sup>−</sup>
:10&nbsp;e<sup>−</sup> + 30&nbsp;H<sup>+</sup> + 5&nbsp;{{chem|BiO|3|-}}(s) → 5&nbsp;Bi<sup>3+</sup>(aq) + 15&nbsp;H<sub>2</sub>O(l)

Adding these two reactions eliminates the electrons' terms and yields a balanced reaction:

:14&nbsp;H<sup>+</sup>(aq) + 2&nbsp;Mn<sup>2+</sup>(aq) + 5&nbsp;NaBiO<sub>3</sub>(s) → 7&nbsp;H<sub>2</sub>O(l) + 2&nbsp;{{chem|MnO|4|-}}(aq) + 5&nbsp;Bi<sup>3+</sup>(aq) + 5&nbsp;{{chem|Na|+}}(aq)

===Basic media===
In primary aqueous media, [[Hydroxide|OH<sup>−</sup>]] ions and water are added to half-reactions to balance the overall reaction.

For example, in the reaction between [[potassium permanganate]] and [[sodium sulfite]]:

:{|
|Unbalanced reaction:
|KMnO<sub>4</sub> + Na<sub>2</sub>SO<sub>3</sub> + H<sub>2</sub>O → MnO<sub>2</sub> + Na<sub>2</sub>SO<sub>4</sub> + KOH
|-
|Reduction:
|3&nbsp;e<sup>−</sup> + 2&nbsp;H<sub>2</sub>O + {{chem|MnO|4|−}} → MnO<sub>2</sub> + 4&nbsp;OH<sup>−</sup>
|-
|Oxidation:
|2&nbsp;OH<sup>−</sup> + {{chem|SO|3|2−}} → {{chem|SO|4|2−}} + H<sub>2</sub>O + 2&nbsp;e<sup>−</sup>
|}

Balancing the number of electrons in the two half-cell reactions gives:

:6&nbsp;e<sup>−</sup> + 4&nbsp;H<sub>2</sub>O + 2&nbsp;{{chem|MnO|4|−}} → 2&nbsp;MnO<sub>2</sub> + 8&nbsp;OH<sup>−</sup>
:6&nbsp;OH<sup>−</sup> + 3&nbsp;{{chem|SO|3|2−}} → 3&nbsp;{{chem|SO|4|2−}} + 3&nbsp;H<sub>2</sub>O + 6&nbsp;e<sup>−</sup>

Adding these two half-cell reactions together gives the balanced equation:

:2&nbsp;KMnO<sub>4</sub> + 3&nbsp;Na<sub>2</sub>SO<sub>3</sub> + H<sub>2</sub>O → 2&nbsp;MnO<sub>2</sub> + 3&nbsp;Na<sub>2</sub>SO<sub>4</sub> + 2&nbsp;KOH-->