Redox Reactions: NEET notes

1. Oxidation and reduction: the classical view

NCERT §7.1

• Originally, oxidation meant adding oxygen to a substance, as when magnesium burns to MgO or sulphur burns to SO₂. Oxygen is not the only reagent that counts.
• The idea was broadened to cover removal of hydrogen: in 2H₂S + O₂ → 2S + 2H₂O, hydrogen is taken away from H₂S, so H₂S is oxidised.
• It was widened again to adding an electronegative element: Mg + F₂ → MgF₂, Mg + Cl₂ → MgCl₂ and Mg + S → MgS all count as oxidation of magnesium.
• Removal of an electropositive element is oxidation too: when K₄[Fe(CN)₆] reacts with H₂O₂ to give K₃[Fe(CN)₆], one potassium is removed per complex unit.
• Reduction is the reverse in each case: removing oxygen (2HgO → 2Hg + O₂), adding hydrogen (ethene + H₂ → ethane), removing an electronegative element (2FeCl₃ + H₂ → 2FeCl₂ + 2HCl) or adding an electropositive one (2HgCl₂ + SnCl₂ → Hg₂Cl₂ + SnCl₄, where mercury(II) chloride is reduced).
• Oxidation and reduction always happen together; one substance cannot be oxidised unless another is reduced. Such paired reactions are called redox reactions.

2. Redox as electron transfer

NCERT §7.2

• In 2Na + Cl₂ → 2NaCl, each sodium atom gives up one electron and each chlorine atom accepts one. The reaction can be split into two half reactions: 2Na → 2Na⁺ + 2e⁻ and Cl₂ + 2e⁻ → 2Cl⁻.
• The half reaction in which electrons are lost is the oxidation half; the one in which electrons are gained is the reduction half.
• Electronic definitions: oxidation is loss of electrons by a species; reduction is gain of electrons by a species.
• The oxidising agent (oxidant) accepts electrons and is itself reduced. The reducing agent (reductant) donates electrons and is itself oxidised.
• In 2Na + H₂ → 2NaH, hydrogen gains electrons and becomes H⁻, so hydrogen is reduced and sodium is oxidised. Hydrogen acts here as the oxidising agent, opposite to its usual role.

3. Competitive electron transfer

NCERT §7.2.1

• A zinc strip dipped in copper nitrate solution for about an hour gets a reddish copper coating, and the blue of the solution fades: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s). Zinc loses electrons to Cu²⁺.
• The Zn²⁺ formed can be shown by passing H₂S through the colourless solution after making it ammoniacal: a white precipitate of ZnS appears.
• A copper rod in zinc sulphate solution shows no change, and H₂S gives no black CuS. So Cu does not hand electrons to Zn²⁺; the reverse transfer does not occur.
• Copper in silver nitrate solution does react: the solution turns blue as Cu²⁺ forms and silver is deposited, Cu(s) + 2Ag⁺(aq) → Cu²⁺(aq) + 2Ag(s).
• The tendency to release electrons therefore falls in the order Zn > Cu > Ag. Arranging metals by this competition for electrons gives the metal activity series (electrochemical series).
• In some pairs neither side wins completely. Cobalt in nickel sulphate solution reaches a state where both Co²⁺ and Ni²⁺ are present at moderate concentrations: the reaction reaches equilibrium.
• The same competition, when the two half reactions are kept in separate beakers joined by a wire, can drive an electric current. This is the basis of a galvanic cell.

4. Oxidation number and its rules

NCERT §7.3

• In reactions between covalent molecules, such as H₂ + Cl₂ → 2HCl, no electron is fully transferred; the shared pair only shifts toward the more electronegative atom. To treat such reactions as redox anyway, chemists assign each atom an oxidation number.
• Oxidation number is the charge an atom would carry in a compound if every bonding pair were handed entirely to the more electronegative partner.
• Rule 1: an element in the free state has oxidation number 0, whatever its form (H₂, O₂, Cl₂, O₃, S₈, P₄, Na, Mg, Al).
• Rule 2: a monatomic ion has an oxidation number equal to its charge. In their compounds alkali metals are always +1, alkaline earth metals +2 and aluminium +3.
• Rule 3: oxygen is usually −2, but −1 in peroxides (H₂O₂, Na₂O₂), −½ in superoxides (KO₂, RbO₂), +2 in OF₂ and +1 in O₂F₂. It takes positive values only when bonded to fluorine.
• Rule 4: hydrogen is usually +1, but −1 in binary hydrides of active metals such as LiH, NaH and CaH₂.
• Rule 5: fluorine is −1 in all its compounds. Chlorine, bromine and iodine are −1 in halides but take positive values when combined with oxygen, as in oxoacids and oxoanions.
• Rule 6: the oxidation numbers of all atoms in a neutral compound add to zero; in a polyatomic ion they add to the ion's charge. In CO₃²⁻, x + 3(−2) = −2 gives carbon +4.
• For main-group elements the highest oxidation number usually equals the group number for groups 1 and 2, and the group number minus 10 for groups 13 to 17. Across period 3 it rises from +1 in NaCl to +7 in HClO₄.

5. Stock notation and redox by oxidation number

NCERT §7.3

• Metals that show more than one oxidation number have it written as a Roman numeral in brackets after the symbol (Stock notation, after Alfred Stock): Au(I)Cl and Au(III)Cl₃, Sn(II)Cl₂ and Sn(IV)Cl₄, Fe(II)O and Fe₂(III)O₃, Mn(IV)O₂.
• Oxidation is an increase in the oxidation number of an element; reduction is a decrease.
• An oxidising agent contains an element whose oxidation number falls; a reducing agent contains an element whose oxidation number rises.
• A redox reaction is one in which oxidation numbers of the reacting species change. If no oxidation number changes, the reaction is not redox.
• In CH₄ + 4Cl₂ → CCl₄ + 4HCl carbon goes from −4 to +4, so methane is oxidised even though no electron is fully transferred.
• In 2Cu₂O + Cu₂S → 6Cu + SO₂, copper falls from +1 to 0 and sulphur rises from −2 to +4. Cu₂S is both the reducing agent (through S) and partly reduced (through Cu).

6. Combination, decomposition and displacement

NCERT §7.3.1

• Combination: two or more substances unite into one, A + B → C, and at least one element changes oxidation number: C + O₂ → CO₂ (0 → +4 for C) and 3Mg + N₂ → Mg₃N₂ (0 → −3 for N).
• Decomposition: one compound breaks into two or more parts, at least one of them an element: 2H₂O → 2H₂ + O₂, 2NaH → 2Na + H₂ and 2KClO₃ → 2KCl + 3O₂ (Cl +5 → −1, O −2 → 0).
• Not every decomposition is redox. CaCO₃ → CaO + CO₂ involves no change in any oxidation number.
• Metal displacement: a metal in a compound is replaced by another metal from its elemental state, as in CuSO₄ + Zn → Cu + ZnSO₄. Reducing metals such as Ca, Mg and Al are used this way to win V, Ti and Cr from their oxides or chlorides (V₂O₅ + 5Ca, TiCl₄ + 2Mg, Cr₂O₃ + 2Al).
• Hydrogen displacement: alkali metals and some alkaline earth metals (Ca, Sr, Ba) displace hydrogen from cold water; less active metals such as Mg and Fe need steam (2Fe + 3H₂O → Fe₂O₃ + 3H₂).
• Many metals, including some that ignore steam (Cd, Sn), release hydrogen from acids. Zn, Mg and Fe react with HCl at rates that show their activity, magnesium fastest and iron slowest; Ag and Au do not react with HCl at all.
• Halogen displacement: oxidising power falls from F₂ to I₂. Fluorine even displaces oxygen from water (2H₂O + 2F₂ → 4HF + O₂). Chlorine displaces Br⁻ and I⁻, and bromine displaces I⁻; the freed Br₂ or I₂ is identified by its colour in a CCl₄ layer (Layer Test).
• Since no chemical oxidant is stronger than F₂, fluoride ions can be turned into F₂ only by electrolysis.

7. Disproportionation and fractional oxidation numbers

NCERT §7.3.1

• In a disproportionation reaction one element in one oxidation state is oxidised and reduced at the same time. The reacting substance must contain that element in a state from which it can both rise and fall.
• Hydrogen peroxide decomposes as 2H₂O₂ → 2H₂O + O₂: oxygen at −1 goes both to −2 (in water) and to 0 (in O₂).
• White phosphorus in alkali gives phosphine and hypophosphite (P 0 → −3 and 0 → +1); sulphur in alkali gives sulphide and thiosulphate (S 0 → −2 and 0 → +2).
• Chlorine in cold dilute alkali: Cl₂ + 2OH⁻ → ClO⁻ + Cl⁻ + H₂O, with Cl going from 0 to +1 and to −1. The hypochlorite formed is the active part of household bleach.
• Fluorine cannot disproportionate. It is the most electronegative element, cannot take a positive state, and in alkali gives F⁻ and OF₂ in a reaction where oxygen, not fluorine, is oxidised.
• Among the chlorine oxoanions, ClO⁻, ClO₂⁻ and ClO₃⁻ can disproportionate. ClO₄⁻ cannot, because chlorine is already at its highest state, +7, and can only be reduced.
• An averaged oxidation number can be fractional when atoms of one element sit in different environments: C₃O₂ gives +4/3, Br₃O₈ gives +16/3 and S₄O₆²⁻ gives +2.5. The real structures hold whole-number states, for example +5, 0, 0, +5 for the four S in tetrathionate.
• Mixed oxides such as Fe₃O₄, Mn₃O₄ and Pb₃O₄ also show fractional averages. Pb₃O₄ behaves as 2PbO + PbO₂, and the O₂⁺ and O₂⁻ ions give oxygen +½ and −½.

8. Balancing redox equations

NCERT §7.3.2

• Oxidation number method: write the skeletal equation, find the atoms whose oxidation numbers change, and multiply so that the total increase equals the total decrease.
• Then balance the charge (with H⁺ in acid or OH⁻ in base), then hydrogen with water, and finally check oxygen.
• Example: dichromate and sulphite in acid give Cr₂O₇²⁻ + 3SO₃²⁻ + 8H⁺ → 2Cr³⁺ + 3SO₄²⁻ + 4H₂O. Each Cr falls by 3 (6 for the pair) and each S rises by 2, so three sulphites are needed.
• Half-reaction method: split the reaction into oxidation and reduction halves, balance the atoms other than O and H, then O with H₂O and H with H⁺ (in acid). Add electrons to balance charge, scale the halves to equal electron counts and add them.
• Example: Fe²⁺ with dichromate in acid gives 6Fe²⁺ + Cr₂O₇²⁻ + 14H⁺ → 6Fe³⁺ + 2Cr³⁺ + 7H₂O, since each Fe²⁺ gives one electron and one Cr₂O₇²⁻ takes six.
• In basic medium, balance as if acidic and then add as many OH⁻ to both sides as there are H⁺, combining H⁺ + OH⁻ into water. Permanganate and iodide in base give 2MnO₄⁻ + 6I⁻ + 4H₂O → 2MnO₂ + 3I₂ + 8OH⁻.
• A balanced redox equation must balance atoms of every element and the total charge on both sides.

9. Redox titrations and limits of the concept

NCERT §7.3.3, §7.3.4

• In a redox titration the end point is found by a colour change, just as in an acid-base titration.
• Permanganate is its own indicator. The purple MnO₄⁻ is decolourised as it reacts with a reductant such as Fe²⁺ or C₂O₄²⁻; the first slight excess leaves a faint pink, visible at a concentration as low as 10⁻⁶ mol L⁻¹.
• Dichromate gives no sharp colour change of its own, so an indicator such as diphenylamine is added. Just after the equivalence point it is oxidised to an intense blue.
• Iodometric titration: Cu²⁺ oxidises iodide, 2Cu²⁺ + 4I⁻ → Cu₂I₂ + I₂, and the freed iodine is titrated with thiosulphate, I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻. Starch forms an intense blue with iodine, and the blue vanishes when the iodine is used up.
• Iodine is only slightly soluble in water, but it stays in solution in the presence of KI as KI₃.
• The electron-transfer picture is an idealisation. In covalent compounds it is more accurate to describe oxidation as a decrease in electron density around an atom and reduction as an increase.

10. Redox couples and electrode processes

NCERT §7.4

• A redox couple is the oxidised and reduced forms of one substance taking part in a half reaction, written with the oxidised form first: Zn²⁺/Zn, Cu²⁺/Cu.
• In the Daniell cell a zinc rod in ZnSO₄ solution and a copper rod in CuSO₄ solution are kept in separate beakers. A salt bridge, a U-tube of KCl or NH₄NO₃ solution set in agar, joins them and completes the circuit without letting the solutions mix.
• When the two rods are joined by a wire through an ammeter and the switch is closed, zinc is oxidised (Zn → Zn²⁺ + 2e⁻) and copper ions are reduced (Cu²⁺ + 2e⁻ → Cu). Electrons flow from zinc to copper through the wire, while conventional current flows from copper to zinc.
• The potential of an electrode in contact with its ions is its electrode potential. When all species are at unit concentration (gases at 1 atm) and 298 K it is the standard electrode potential, E°.
• The hydrogen electrode is assigned E° = 0.00 V by convention, and other couples are measured against it.
• A negative E° means the couple is a stronger reducing agent than the H⁺/H₂ couple; a positive E° means a weaker reducing agent, that is, a stronger oxidising agent than H⁺.
• In the series fluorine sits at the top (E° = +2.87 V for F₂/F⁻), so F₂ is the strongest oxidant and F⁻ the weakest reductant. Lithium sits at the bottom (E° = −3.05 V for Li⁺/Li), so Li is the strongest reductant listed.
• Other values worth recalling: MnO₄⁻/Mn²⁺ +1.51 V, Cl₂/Cl⁻ +1.36 V, Cr₂O₇²⁻/Cr³⁺ +1.33 V, Ag⁺/Ag +0.80 V, Cu²⁺/Cu +0.34 V, Zn²⁺/Zn −0.76 V, Na⁺/Na −2.71 V.

Must-know facts

1. Oxidation: loss of electrons or rise in oxidation number. Reduction: gain of electrons or fall in oxidation number.
2. The oxidising agent is itself reduced; the reducing agent is itself oxidised.
3. Oxidation and reduction always occur together in a redox reaction.
4. A free element has oxidation number 0 in every allotrope.
5. Oxygen: −2 normally, −1 in peroxides, −½ in superoxides, +2 in OF₂, +1 in O₂F₂.
6. Hydrogen: +1 normally, −1 in active metal hydrides such as NaH and CaH₂.
7. Fluorine is −1 in every compound and never disproportionates.
8. Oxidation numbers add to zero in a molecule and to the charge in an ion.
9. Zn displaces Cu from Cu²⁺ solution and Cu displaces Ag from Ag⁺ solution, but Cu does not displace Zn.
10. Four types of redox reaction: combination, decomposition, displacement and disproportionation.
11. CaCO₃ → CaO + CO₂ is a decomposition that is not redox.
12. In disproportionation the same element in one state is both oxidised and reduced.
13. ClO₄⁻ cannot disproportionate because Cl is at its highest state, +7.
14. Halogen oxidising power: F₂ > Cl₂ > Br₂ > I₂.
15. Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O; each dichromate takes six electrons.
16. KMnO₄ acts as a self-indicator; dichromate titrations use diphenylamine.
17. Iodometry: I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻ with starch as indicator.
18. In the Daniell cell electrons flow from Zn to Cu in the external wire; the salt bridge completes the circuit.
19. E° of the standard hydrogen electrode is 0.00 V by convention.
20. More negative E° means a stronger reducing agent; more positive E° means a stronger oxidising agent.

Common traps

Formulas

Key terms

Redox reaction

A reaction in which oxidation and reduction take place together.

Half reaction

The oxidation part or the reduction part of a redox reaction, written with electrons shown.

Oxidising agent

A species that accepts electrons and is reduced.

Reducing agent

A species that donates electrons and is oxidised.

Oxidation number

The charge an atom would carry if all its bonding pairs went to the more electronegative atom.

Stock notation

A Roman numeral in brackets giving a metal's oxidation number, as in Fe(III).

Disproportionation

A reaction in which one element in one state is both oxidised and reduced.

Redox couple

The oxidised and reduced forms of a species in a half reaction, oxidised form first.

Salt bridge

A tube of electrolyte in agar that joins two half cells without mixing them.

Standard electrode potential

The electrode potential with all species at unit concentration or 1 atm and 298 K.

Iodometric titration

A titration in which liberated iodine is measured with thiosulphate using starch.