Oxidation

1.0What is Oxidation?

Oxidation is a chemical reaction in which a substance loses electrons, thus increasing its oxidation state.

The process of oxidation involves the addition of oxygen atoms or the loss of hydrogen atoms. Oxidation reactions can also occur without the involvement of oxygen, such as in reactions with other oxidizing agents.

Oxidation occurs when the oxidation state of a molecule, atom, or ion increases, while reduction occurs when there is a gain of electrons or a decrease in the oxidation state.

An illustrative example is the reaction between hydrogen and fluorine gas to form hydrofluoric acid:

H₂ +F₂ →2HF

In this reaction, hydrogen is oxidized, and fluorine is reduced. We can break down the reaction into two half-reactions:

H₂   →  2H⁺ + 2e⁻

F₂ +2e⁻  → 2F⁻

Note that there is no involvement of oxygen in this reaction. It purely demonstrates the principles of oxidation and reduction without the presence of oxygen.

2.0Old Concept of Oxidation

The old definition of oxidation was defined as the addition of oxygen to a compound, primarily because oxygen gas (O₂) was the earliest recognized oxidizing agent. While the addition of oxygen typically involves electron loss and an increase in the oxidation state, the definition of oxidation has since been broadened to encompass various types of chemical reactions.

A classic example of the old definition of oxidation is the reaction where iron combines with oxygen to form iron oxide, commonly known as rust:

2 Fe  +  O₂ → Fe₂O₃​

In this reaction, iron metal oxidizes to produce iron oxide, or rust.

Electrochemical reactions provide excellent examples of oxidation reactions. When a copper wire is immersed in a solution containing silver ions, electrons are transferred from the copper metal to the silver ions, leading to the oxidation of copper. This results in the growth of silver metal whiskers onto the copper wire, while copper ions are released into the solution:

Cu(s) + 2Ag⁺ (aq) → Cu²⁺ (aq)  +  2Ag (s)

Another instance of oxidation occurs when an element combines with oxygen, such as the reaction between magnesium metal and oxygen to produce magnesium oxide:

2 Mg (s) + O₂(g) → 2 MgO (s)

Oxidation and reduction are intimately linked, occurring together in what are known as redox reactions, short for reduction-oxidation reactions. In these reactions, one species loses electrons (is oxidized), while another gains electrons (is reduced). For example, the oxidation of a metal by oxygen gas can be explained as the metal atom losing electrons to form cations, while the oxygen molecule gains electrons to form oxygen anions. In the case of magnesium, the reaction can be represented as:

2 Mg  +  O₂  →  2[Mg²⁺][O²⁻]

This reaction can be further broken down into the following half-reactions:

Mg → Mg²⁺ + 2e⁻

O₂ + 4e⁻  →  2O²⁻

3.0The Modern Concept of Oxidation

The modern concept of oxidation reaction definition involves the loss of electrons by a substance, leading to an increase in its oxidation state. This definition extends beyond oxygen transfer to include reactions where other elements or compounds gain electrons. Oxidation is thus understood as any chemical process that involves the removal of electrons from a substance, resulting in an increase in its positive oxidation state or a decrease in its negative oxidation state.

Few examples of oxidation reactions are shared below-

​(a)  Zn → Zn⁺² + 2e^–

          General Formula-       $\mathrm{M}\to {\mathrm{M}}^{\mathrm{n}+}+{\mathrm{ne}}^{-}$

(b)  Sn⁺²  → Sn⁺⁴ + (4 –2)e^–

       General Formula- ${\mathrm{M}}^{+{\mathrm{n}}_1}\to {\mathrm{M}}^{+{\mathrm{n}}_2}+\left({\mathrm{n}}_2-{\mathrm{n}}_1\right){\mathrm{e}}^{-}$

(c)  Cl^– → Cl + e^–

       General Formula-     ${\mathrm{A}}^{-\mathrm{n}}\to \mathrm{A}+\mathrm{n}{\mathrm{e}}^{-}$

(d)  MnO₄^–2 →  MnO₄^– + (2 – 1)e^–

             General Formula-  ${\mathrm{A}}^{-{\mathrm{n}}_1}\to {\mathrm{A}}^{-{\mathrm{n}}_2}+\left({\mathrm{n}}_1-{\mathrm{n}}_2\right){\mathrm{e}}^{-}$

4.0Oxidation Reaction and Oxidation State

An oxidation reaction involves the loss of electrons from a substance, resulting in an increase in its oxidation state. Oxidation states are formal charges assigned to individual atoms within a compound or ion, indicating the hypothetical charge that an atom would have if all its bonds were completely ionic. In an oxidation reaction, the oxidation state of the atom undergoing oxidation increases, while the reducing agent loses electrons.

For example, in the reaction:

Zn(s) → Zn²⁺(aq) + 2e⁻

Zinc (Zn) is oxidized from an oxidation state of 0 to +2, indicating the loss of two electrons. The oxidation state of zinc increases from 0 to +2, indicating oxidation.

5.0Oxidizing Agent and Reducing Agent

An oxidizing agent is a substance that has the ability to oxidize other substances by accepting electrons from them. In a chemical reaction, the oxidizing agent itself undergoes reduction, while the substance it oxidizes undergoes oxidation. 

Oxidizing agents are compounds capable of oxidizing others while reducing themselves during a chemical reaction. These reagents have their oxidation numbers decrease or gain electrons in a redox reaction. Examples of powerful oxidizing agents include KMnO₄, K₂Cr₂O₇, HNO₃, and concentrated H₂SO₄.

Reducing Agent (Reductant):

Reducing agents are compounds capable of reducing others while oxidizing themselves during a chemical reaction. These reagents have their oxidation numbers increase or lose electrons in a redox reaction. Examples of powerful reducing agents include KI and Na₂S₂O₃.

Note:

Some compounds can function as both oxidizing agents and reducing agents. For instance, H₂O₂.

6.0Oxidation vs Reduction