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Oxidation and Reduction Reactions

Oxidation Reaction

Oxidation reactions can be defined in terms of addition of oxygen or the removal of hydrogen. There are a number of reactions in which oxygen is added, or in which hydrogen is removed.

Examples: Addition of Oxygen:

(a). CuO(s) + H₂(g) → Cu(s) + H₂O(l)

CuO adds oxygen to H₂;  CuO is therefore the oxidizing agent or oxidant, it oxidizes the H₂ to H₂O.

(b). C(s) + H₂O(g) → H₂(g) + CO(g)

H₂O is the oxidizing agent; it gives oxygen to carbon, which is oxidized to CO.

Removal of Hydrogen:

(a). H₂S(g) + Cl₂(g) → 2HCl(g) + S(s)

Chlorine is the oxidant, it removes hydrogen from H₂S, thereby oxidizing it to sulphur.

(b). 2H₂O(l) + 2F₂(g) → 4HF(aq) + O₂(g)

Fluorine removes hydrogen from the water, therefore, fluorine is the oxidant. The water is oxidized to oxygen.

Reduction Reaction

Reduction reactions can be defined as reactions that involve the removal of oxygen or the addition of hydrogen.

Example, Removal of Oxygen:

(a). CuO(s) + H₂(g) → Cu(s) + H₂O(l)

Hydrogen removes oxygen from CuO, thereby reducing it to Cu. Hydrogen is therefore the reducing agent or reductant.

(b). H₂O(g) + C(s) → H₂(g) + CO(g)

Carbon is the reductant, it removes oxygen from steam, thereby reducing it to hydrogen.

Addition of Hydrogen:

(a). H₂S(g) + Cl₂(g) → 2HCl(g) + S(s)

H₂S adds hydrogen to Cl₂(g). H₂S is the reducing agent. It reduces Cl₂(g) to HCl(g).

(b). 2H₂O(l) + 2F₂(g) → 4HF(aq) + O₂(g)

H₂O here adds hydrogen to F₂, thereby reducing it to HF. H₂O is the reductant in this reaction.

Redox Reactions

- It is not possible to have only oxidation or reduction reaction taking place alone. Both reactions must occur together. Oxidation - reduction reactions are simply called redox reactions.

Example, .g. in the reaction: CuO(s) + H₂(g) → Cu(s) + H₂O(l)

CuO is the oxidant, while H₂ is the reductant . The oxidant always gets reduced, while the reductant always gets oxidized. I.e. CuO is reduced to Cu, while H₂ is oxidized to H₂O.

- Do not regard any particular substance as either an oxidant or reductant until you have examined the reaction, to see how it acts. This is because some substances can act as both oxidizing and reducing agent, depending on the reaction.

Example, in the reactions: a. H₂O(g) + C(s) → H₂(g) + CO(g)

b. 2H₂O(l) + 2F₂(g) → 4HF(aq) + O₂(g)

Water acts as an oxidant in reaction (a), but as a reductant in reaction (b). Also, SO₂ can act as an oxidant, i.e. in the reaction: 2H₂S(g) + SO₂(g) → 3S(s) + 2H₂O(l) and as a reductant in the presence of water.

Example, in the reaction: 5SO₂(g) + 2MnO₄^- (aq) + 2H₂O(l) → 5SO₄^2-(aq) + 2Mn²⁺(aq) + 4H⁺(aq).

When a substance which can act both as an oxidant and reductant is reacted with a stronger reductant, it behaves as the oxidant. And when reacted with a stronger oxidant, it behaves as a reductant.

Oxidation and Reduction Reaction in Terms of Electron Transfer

This is the most recent definition of oxidation and reduction reactions. Here, oxidation is defined as the process of losing electrons; while reduction is defined as the process of electron gain.

Note:

In a redox reaction, an oxidizing agent is the substance which accepts electrons and become reduced.

A reducing agent is the substance which donates electrons and become oxidized.

Oxidation leads to increase in oxidation number.

Reduction leads to decrease in oxidation number.

The reducing agent (reductant), which is oxidized increases in oxidation number (i.e., loss of electrons increases oxidation number).

The oxidizing agent (oxidant), which is reduced decreases in oxidation number (i.e., acceptance of electrons reduces oxidation number).

Metals are usually reducing agents – they have high electropositivity. They reduce their counterparts in redox reactions by donating electrons to them, while they themselves become oxidized (i.e. increase in oxidation number). Example, K, Na, Ca and Mg.

Non-metals have high electronegativity (i.e. they accept electrons readily). They therefore act as oxidizing agents, and become reduced. Examples, F₂, O₂, Cl₂.

Examples:

a. Mg(s) + ½ O₂(g) → MgO(s)

Mg losses two electrons: Mg → Mg²⁺ + 2e- (oxidation reaction).

An atom of oxygen gains the two electrons: ½ O₂ + 2e- → O^2- (reduction).

Combining the two: Mg + ½ O₂ → MgO

Mg is the reducing agent, while oxygen is the oxidizing agent.

b. H₂S(g) + Cl₂(g) → 2HCl(g) + S(s)

In H₂S (we have 2H⁺ + S^2-) . The sulphide ion S^2- gives its two electrons to Cl₂ and then gets oxidized. H₂S is the reducing agent. Cl₂ is the oxidizing agent, it accepts the two electrons and become reduced.

c. CuO(s) + H₂(g) → Cu(s) + H₂O(g)

Hydrogen is more electropositive than copper, it gives to Cu²⁺ the two electrons that copper metal initially gave to oxygen to form CuO (CuO consists of Cu²⁺ + O^2-).

I.e. Cu²⁺(aq) + 2e^-(aq) → Cu(s), then 2H⁺(aq) + O^2-(aq) → H₂O(l). Therefore, CuO is the oxidizing agent, and it becomes reduced to Cu; while H₂ is the reducing agent, which gets oxidized to H₂O.