Redox Reactions

In this chapter: classical and electronic ideas of oxidation and reduction, oxidising and reducing agents, competitive electron transfer, rules for oxidation numbers, combination, decomposition, displacement and disproportionation reactions, balancing by the oxidation number and half-reaction methods, redox titrations, and redox reactions in electrode processes.

Oxidation and reduction

In the classical sense, oxidation is addition of oxygen or an electronegative element, or removal of hydrogen or an electropositive element; reduction is the reverse. In 2Mg + O₂ → 2MgO, magnesium is oxidised; in CuO + H₂ → Cu + H₂O, copper oxide is reduced.

The electronic definition is more general. Oxidation is loss of electrons and reduction is gain of electrons. The two always happen together, so the whole reaction is a redox reaction.

• An oxidising agent (oxidant) accepts electrons and is itself reduced.
• A reducing agent (reductant) gives electrons and is itself oxidised.

When a zinc rod is dipped in copper sulphate solution, the blue colour fades and copper deposits on the zinc: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s). Zinc loses electrons to Cu²⁺. The reverse does not happen when copper is dipped in zinc sulphate, because zinc releases electrons more readily than copper. Comparing such pairs gives the order Zn > Cu > Ag for the tendency to release electrons.

Oxidation number

Oxidation number is the charge an atom would carry in a compound if all its bonds were treated as ionic, with the shared electrons given to the more electronegative atom. Oxidation is an increase in oxidation number and reduction is a decrease.

Rules

• An element in the free state has oxidation number 0: H₂, O₂, O₃, Cl₂, P₄, S₈, Na, Mg.
• For a monatomic ion it equals the charge: Na⁺ +1, Mg²⁺ +2, Cl⁻ −1. Alkali metals are always +1 and alkaline earth metals +2 in their compounds.
• Oxygen is −2 in most compounds, but −1 in peroxides (H₂O₂, Na₂O₂), −½ in superoxides (KO₂), +2 in OF₂ and +1 in O₂F₂.
• Hydrogen is +1, except in metal hydrides such as NaH and CaH₂, where it is −1.
• Fluorine is always −1. Cl, Br and I are −1 in halides but take positive values when combined with oxygen, as in oxoacids and oxoanions.
• The oxidation numbers of all atoms add up to zero in a neutral compound and to the charge in an ion.

Common values to know: Mn +7 in KMnO₄, Cr +6 in K₂Cr₂O₇, S +6 in H₂SO₄, N +5 in HNO₃ and −3 in NH₃, C −4 in CH₄ and +4 in CO₂.

Worked example: Find the oxidation number of (a) Cr in K₂Cr₂O₇, (b) S in Na₂S₄O₆, (c) Fe in Fe₃O₄.
Solution: (a) 2(+1) + 2x + 7(−2) = 0 gives x = +6. (b) 2(+1) + 4x + 6(−2) = 0 gives 4x = 10, x = +2.5. (c) 3x + 4(−2) = 0 gives x = +8/3. Fractional values are averages: in S₄O₆²⁻ two S atoms are +5 and two are 0, and Fe₃O₄ contains Fe²⁺ and Fe³⁺ in the ratio 1 : 2.

Types of redox reactions

For disproportionation, the element must be able to exist in at least three oxidation states, one higher and one lower than the starting one. Fluorine, the most electronegative element, cannot show a positive oxidation state and so does not disproportionate.

Balancing redox reactions

Oxidation number method

1. Write the skeletal equation and mark the oxidation numbers of the atoms that change.
2. Find the increase and decrease per formula unit and multiply so that total increase equals total decrease.
3. Balance O by adding H₂O, then H by adding H⁺ (acidic medium) or H₂O and OH⁻ (basic medium).
4. Check that atoms and charges balance on both sides.

Half-reaction (ion-electron) method

Split the reaction into oxidation and reduction half-reactions, balance each for atoms (O with H₂O, H with H⁺) and for charge (with electrons), then multiply so the electrons cancel and add.

Worked example: Balance Fe²⁺ + Cr₂O₇²⁻ → Fe³⁺ + Cr³⁺ in acidic solution.
Solution: Oxidation: Fe²⁺ → Fe³⁺ + e⁻. Reduction: Cr₂O₇²⁻ → 2Cr³⁺; add 7H₂O on the right to balance O, 14H⁺ on the left to balance H, and 6e⁻ on the left to balance charge: Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O. Multiply the oxidation half by 6 and add: Cr₂O₇²⁻ + 6Fe²⁺ + 14H⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O. Charge check: −2 + 12 + 14 = +24 on the left; 6 + 18 = +24 on the right.

In basic medium, balance as if acidic and then add as many OH⁻ ions to both sides as there are H⁺ ions; H⁺ and OH⁻ combine to form water. For permanganate and iodide in basic solution: 2MnO₄⁻ + 6I⁻ + 4H₂O → 2MnO₂ + 3I₂ + 8OH⁻.

Redox titrations

• KMnO₄ in acid is its own indicator. MnO₄⁻ is deep purple and Mn²⁺ is almost colourless, so the first excess drop gives a permanent pink colour at the end point.
• K₂Cr₂O₇ needs an external indicator such as diphenylamine, which gives an intense blue colour at the end point.
• Iodometry: I₂ liberated in a reaction is titrated with thiosulphate using starch, which gives a blue colour with iodine: I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻. The blue colour disappears at the end point.

Redox reactions and electrode processes

If the zinc and copper reactions are carried out in separate beakers, connected by a wire and a salt bridge, the electron transfer happens through the wire and produces a current. This arrangement is the Daniell cell. Each beaker with its metal is a half-cell, or redox couple (Zn²⁺/Zn and Cu²⁺/Cu).

Electrode potentials are measured against the standard hydrogen electrode, whose potential is taken as 0.00 V. Standard electrode potentials E° (1 M ions, 1 bar gas, 298 K) are written as reduction potentials. A more positive E° means the oxidised form is a stronger oxidising agent (F₂/F⁻ +2.87 V is the highest); a more negative E° means the metal is a stronger reducing agent (Li⁺/Li −3.05 V is the lowest). Zn²⁺/Zn is −0.76 V and Cu²⁺/Cu is +0.34 V, which is why zinc displaces copper and not the other way round.

Common mistakes: (1) Calling the substance that gets oxidised the oxidising agent; it is the reducing agent. (2) Taking oxygen as −2 in H₂O₂, KO₂ or OF₂. (3) Treating every decomposition as a redox reaction. (4) Forgetting to check charge balance after the atom balance. (5) Adding H⁺ to a reaction that is stated to be in basic medium.

JEE and NEET focus

• Oxidation numbers in compounds and ions, including average and fractional values (Fe₃O₄, S₄O₆²⁻, peroxides, superoxides).
• Identifying the oxidant and reductant, and which species is oxidised in a given reaction.
• Recognising disproportionation reactions and elements that cannot disproportionate.
• Balancing in acidic and basic media, and reading mole ratios from the balanced equation (MnO₄⁻ : Fe²⁺ = 1 : 5, Cr₂O₇²⁻ : Fe²⁺ = 1 : 6).
• Indicators in KMnO₄, K₂Cr₂O₇ and iodometric titrations.
• Using E° values to predict the stronger oxidising or reducing agent.

Practice questions