Oxidation and Reduction


Oxidation and Reduction

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Redox (short for reduction–oxidation reaction) (pronunciation is a chemical reaction in which the oxidation states of atoms are changed. Any such reaction involves both a reduction process and a complementary oxidation process, two key concepts involved with electron transfer processes.

Redox reactions include all chemical reactions in which atoms have their oxidation state changed; in general, redox reactions involve the transfer of electrons between chemical species.

The chemical species from which the electron is stripped is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. It can be explained in simple terms:

Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion.

Reduction is the gain of electrons or a decrease in oxidation state by a molecule, atom, or ion.

As an example, during the combustion of wood, oxygen from the air is reduced, gaining electrons from carbon which is oxidized. Although oxidation reactions are commonly associated with the formation of oxides from oxygen molecules, oxygen is not necessarily included in such reactions, as other chemical species can serve the same function.

Oxidation can occur relatively slowly, as with rust, or more quickly, in the case of fire. There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide (CO2) or the reduction of carbon by hydrogen to yield methane (CH4), and more complex processes such as the oxidation of glucose (C6H12O6) in the human body.


  • Redox” is a portmanteau of the words “reduction” and “oxidation”.
  •  The word oxidation originally implied reaction with oxygen to form an oxide, since dioxygen (O2 (g)) was historically the first recognized oxidizing agent.
  • Later, the term was expanded to encompass oxygen-like substances that accomplished parallel chemical reactions.
  • Ultimately, the meaning was generalized to include all processes involving loss of electrons.

The electrochemist John Bockris has used the words electronation and deelectronation to describe reduction and oxidation processes respectively when they occur at electrodes. These words are analogous to protonation and deprotonation, but they have not been widely adopted by chemists worldwide.

The term “hydrogenation” could be used instead of reduction, since hydrogen is the reducing agent in a large number of reactions, especially in organic chemistry and biochemistry.

But, unlike oxidation, which has been generalized beyond its root element, hydrogenation has maintained its specific connection to reactions that add hydrogen to another substance (e.g., the hydrogenation of unsaturated fats into saturated fats, R−CH=CH−R + H2 → R−CH2−CH2−R).

The word “redox” was first used in 1928

Oxidizing and Reducing Agents:

  • In redox processes, the reductant transfers electrons to the oxidant.
  • Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized, and the oxidant or oxidizing agent gains electrons and is reduced.
  • The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox pair.
  • A redox couple is a reducing species and its corresponding oxidizing form, e.g., Fe2+/Fe3+.


Substances that have the ability to oxidize other substances (cause them to lose electrons) are said to be oxidative or oxidizing and are known as oxidizing agents, oxidants, or oxidizers.

That is, the oxidant (oxidizing agent) removes electrons from another substance, and is thus itself reduced. And, because it “accepts” electrons, the oxidizing agent is also called an electron acceptor. Oxygen is the quintessential oxidizer.

Oxidants are usually chemical substances with elements in high oxidation states (e.g., H2O2, MnO−4, CrO3, Cr2O2−7, OsO4), or else highly electronegative elements (O2, F2, Cl2, Br2) that can gain extra electrons by oxidizing another substance.


Substances that have the ability to reduce other substances (cause them to gain electrons) are said to be reductive or reducing and are known as reducing agents, reductants, or reducers.

The reductant (reducing agent) transfers electrons to another substance, and is thus itself oxidized. And, because it “donates” electrons, the reducing agent is also called an electron donor. Electron donors can also form charge transfer complexes with electron acceptors.

Corrosion and Rusting:

The term corrosion refers to the electrochemical oxidation of metals in reaction with an oxidant such as oxygen. Rusting, the formation of iron oxides, is a well-known example of electrochemical corrosion; it forms as a result of the oxidation of iron metal. Common rust often refers to iron(III) oxide, formed in the following chemical reaction:

4 Fe + 3 O2 → 2 Fe2O3

The oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an acid:

Fe2+ → Fe3+ + e−

H2O2 + 2 e− → 2 OH−

Overall equation:

2 Fe2+ + H2O2 + 2 H+ → 2 Fe3+ + 2 H2O.

Balancing Redox Reaction:

Describing the overall electrochemical reaction for a redox process requires a balancing of the component half-reactions for oxidation and reduction. In general, for reactions in aqueous solution, this involves adding H+, OH−, H2O, and electrons to compensate for the oxidation changes.

Acidic Medium:

In acidic media, H+ ions and water are added to half-reactions to balance the overall reaction.

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

Unbalanced reaction:  Mn2+(aq) + NaBiO3(s) → Bi3+(aq) + MnO−4 (aq)

Oxidation:       4 H2O(l) + Mn2+(aq) → MnO−4(aq) + 8 H+(aq) + 5 e−

Reduction:      2 e− + 6 H+ + BiO−3(s) → Bi3+(aq) + 3 H2O(l)

Basic Medium:

In basic media, OH− 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:  KMnO4 + Na2SO3 + H2O → MnO2 + Na2SO4 + KOH

Reduction:      3 e− + 2 H2O + MnO−4 → MnO2 + 4 OH−

Oxidation:       2 OH− + SO2−3 → SO2−4 + H2O + 2 e−