Redox Chemistry

Redox Chemistry

The chemogenesis analysis identifies five distinct types of electronic reaction mechanism, here, and one of these is redox chemistry. This page explores the different types of oxidation and reduction reaction behaviour that lie behind the easily memorised: OIL RIG, Oxidation Is [electron] Loss and Reduction Is [electron] Gain.

Introduction to Redox Chemistry

Redox chemistry is concerned with net electron flow to and from a defined centre during a chemical reaction. A defined centre may be:

A defined centre is said to be oxidised if the electron density decreases, and reduced if electron density increases, during a reaction. The rule is:

Loss of electrons equates with Oxidation

and

Gain of electrons equates with Reduction

The oxidation of a defined centre can be changed in two ways.

Firstly by Single Electron Transfer (SET) to the defined centre (reduction) or from the defined centre (oxidation). For example, the iron(III) ion, Fe³⁺, can be reduced to iron(II), Fe²⁺. The reaction can also occur in the oxidation direction.

The reduction electron can either be provided by a chemical reducing agent (often a metal) or electrochemically.

Electron flow by way of single electron transfer oxidation and reduction can be predicted using standard reduction potential data (below).

The second method of changing the oxidation number is by reversal of bond polarisation at the defined centre.

Hydrogen is electropositive and it renders the carbon of methane, CH4, electron rich and it is defined as having an oxidation number of -4. However, the carbon of carbon dioxide has an oxidation number of +4 because oxygen is more electronegative than carbon. (Each bond contributes once.)

Carbon is able to exist in several oxidation states:

The combustion of methane to carbon dioxide is an oxidation of carbon because the oxidation number of carbon increases from -4 to +4

It follows that Redox Chemistry can proceed by three types of redox reaction:

Classification of Redox Types

A considerable number of oxidising agents and reducing agents have been added to The Chemical Thesaurus reaction chemistry database. These have been classified into six general types of reducing agent and six general types of oxidising agent, although (and please note) the classification is NOT as clear-cut or rigorous as that carried out for Lewis acid and Lewis base types.

Gain of electrons, gain of hydrogen or metal, or loss of oxygen or halogen equates with reduction. This can occur with various types of chemistry:

Single Electron Transfer Electron Donor Reducing Agent, here
Hydride Complex Reducing Agent, here
Lewis Acid Hydride Donor Reducing Agent, here
Hydrogen Reducing Agent, here
Dissolved Metal Reducing Agent, here
Miscellaneous Reducing Agent, here

Loss of electrons, loss of hydrogen or metal, or gain of oxygen or halogen equates with oxidation. This can occur with various types of chemistry:

Single Electron Transfer Electron Removal Oxidising Agent, here
Hydrogen Removal Oxidising Agent, here
Per-Oxygen Oxidising Agent, here
Oxidised Main Group Element Oxidising Agent, here
Oxidised Heavy Metal Oxidising Agent, here
Miscellaneous Oxidising Agent, here

Single Electron Transfer Electron Donor Reducing Agent

Species which act as donors of electrons, including all electropositive elements including all metals and anodes.

Electron Transfer Reactions in Standard Reduction Form
Temp 298K, Pressure 1.0 arm, All solutions 1.0 molar

+2.87 F₂ → F^–

+2.07 O₃ → O₂

+1.81 Co³⁺ → Co²⁺

+1.77 NO₂ → N₂

+1.77 H₂O₂ → H₂O

+1.70 MnO₄^– → MnO₂

+1.69 PbO₂ → Pb⁴⁺

+1.68 Au⁺ → Au

+1.66 Pb⁴⁺ → Pb²⁺

+1.63 ClOH → Cl₂

+1.59 BrOH → Br₂

+1.51 MnO₄^– → Mn²⁺

+1.46 PbO₂ → Pb²⁺

+1.45 IO^– → I₂

+1.41 Au³⁺ → Au⁺

+1.36 Cl₂ → Cl^–

+1.33 Cr₂O₇^2– → Cr³⁺

+1.28 Ce⁴⁺ → Ce³⁺

+1.26 Tl³⁺ → Tl⁺

+1.23 MnO₂ → Mn²⁺

+1.23 O₂ → H₂O

+1.20 Pt²⁺ → Pt

+1.20 IO₃^– → I₂

+1.17 Ag₂O → Ag

+1.09 Br₂ → Br^–

+1.00 HNO₂ → NO

+0.99 Pd²⁺ → Pd

+0.96 NO₃^– → NO

+0.94 NO₃^– → HNO₂

+0.91 Hg²⁺ → Hg⁺

+0.88 NO₃^– → [NH₄]⁺

+0.85 Hg²⁺ → Hg

+0.80 NO₃^– → N₂O₄

+0.80 Ag⁺ → Ag

+0.79 Hg⁺ → Hg

+0.77 Fe³⁺ → Fe²⁺

+0.68 O₂ → H₂O₂

+0.62 I₂ → I^–

+0.54 Cu²⁺ → CuCl

+0.52 Cu⁺ → Cu

+0.45 H₂SO₃ → S

+0.40 Cd²⁺ → Cd

+0.40 O₂ → HO^–

+0.34 Cu²⁺ → Cu

+0.27 Hg₂Cl₂ → Hg

+0.25 PbO → Pb

+0.23 Ge²⁺ → Ge

+0.22 AgCl → Ag

+0.17 S → H₂S

+0.17 SO₄^– → H₂SO₃

+0.16 Cu²⁺ → Cu⁺

+0.15 Sn⁴⁺ → Sn²⁺

+0.12 CuCl → Cu

+0.10 Si → SiH₄

+0.06 P → PH₃

0.00 H⁺ → H₂

-0.06 Fe³⁺ → Fe

-0.12 CO₂ → CO

-0.13 Pb²⁺ → Pb

-0.14 Sn²⁺ → Sn

-0.18 V²⁺ → V

-0.20 CO₂ → HCOOH

-0.23 In⁺ → In

-0.26 V³⁺ → V²⁺

-0.27 Ni²⁺ → Ni

-0.28 H₃PO₄ → P(OH)₃

-0.28 Co²⁺ → Co

-0.34 Tl⁺ → Tl

-0.36 PbSO₄ → Pb

-0.37 Se → SeH₂

-0.37 Ti³⁺ → Ti²⁺

-0.40 Ga³⁺ → Ga⁺

-0.40 In³⁺ → In⁺

-0.41 Cr³⁺ → Cr²⁺

-0.44 Fe²⁺ → Fe

-0.48 S → S^2–

-0.49 CO₂ → (COOH)₂

-0.50 P(OH)₃ → HP(OH)₂

-0.51 HP(OH)₂ → P

-0.51 Sb → SbH₃

-0.56 Ga³⁺ → Ga

-0.60 As → AsH₃

-0.61 U⁴⁺ → U³⁺

-0.74 Cr³⁺ → Cr

-0.76 Zn²⁺ → Zn

-0.86 SiO₂ → Si

-0.91 Cr²⁺ → Cr

-0.92 Se → Se^2–

-1.14 Te → Te^2–

-1.18 Mn²⁺ → Mn

-1.63 Ti²⁺ → Ti

-1.66 Al³⁺ → Al

-1.80 U³⁺ → U

-1.85 Be²⁺ → Be

-2.36 Mg²⁺ → Mg

-2.48 Ce³⁺ → Ce

-2.52 La³⁺ → La

-2.71 Na⁺ → Na

-2.87 Ca²⁺ → Ca

-2.89 Sr²⁺ → Sr

-2.90 Ba²⁺ → Ba

-2.92 Cs⁺ → Cs

-2.92 Rb⁺ → Rb

-2.93 K⁺ → K

-3.03 Li⁺ → Li