Redox Reactions

2.1 Core concepts

• Classical idea of oxidation was originally addition of oxygen (2Mg + O2 → 2MgO; S + O2 → SO2; CH4 + 2O2 → CO2 + 2H2O), later widened to addition of any electronegative element OR removal of hydrogen/electropositive element from a substance (NCERT §7.1, p. 235–236).
• Classical idea of reduction is the mirror image — removal of oxygen/electronegative element (2HgO → 2Hg + O2; 2FeCl3 + H2 → 2FeCl2 + 2HCl) or addition of hydrogen/electropositive element (CH2=CH2 + H2 → H3C-CH3; 2HgCl2 + SnCl2 → Hg2Cl2 + SnCl4). Oxidation and reduction always occur simultaneously — hence the word "redox" (NCERT §7.1, p. 236).
• Electron-transfer concept: In 2Na + Cl2 → 2NaCl, the half reactions 2Na → 2Na+ + 2e– (oxidation = loss of electrons) and Cl2 + 2e– → 2Cl– (reduction = gain of electrons) define oxidation as electron loss, reduction as electron gain. Oxidising agent = electron acceptor; reducing agent = electron donor (NCERT §7.2, p. 237).
• Competitive electron transfer is illustrated by Zn + Cu2+ → Zn2+ + Cu (equilibrium lies far to the right) and Cu + 2Ag+ → Cu2+ + 2Ag, while Co + Ni2+ ⇌ Co2+ + Ni reaches an intermediate equilibrium — giving the order of reducing tendency Zn > Cu > Ag, the embryo of the electrochemical/activity series (NCERT §7.2.1, p. 238–239).
• Oxidation number assumes complete transfer of the bonding pair to the more electronegative atom — a book-keeping device. Six rules govern its assignment: (1) ON of an element in its free state is zero (H2, O2, Cl2, P4, S8, Na, Mg, Al); (2) for monoatomic ions ON equals the charge (Na+ = +1, Mg2+ = +2, Cl– = –1); (3) ON of O is normally –2, but –1 in peroxides (H2O2, Na2O2), –½ in superoxides (KO2, RbO2), and +2 in OF2, +1 in O2F2; (4) ON of H is +1, but –1 in metal hydrides like LiH, NaH, CaH2; (5) F is always –1; other halogens are –1 with metals but positive when bonded to O; (6) algebraic sum of ON = 0 in a neutral compound and equals the charge in a polyatomic ion (NCERT §7.3, p. 239–240).
• Stock notation writes the oxidation state of a metal as a Roman numeral after its symbol — Au(I)Cl, Au(III)Cl3, Sn(II)Cl2, Sn(IV)Cl4, Mn(II)O, Mn(IV)O2, Fe(II)O, Fe2(III)O3 — useful for distinguishing reduced and oxidised forms of the same metal (NCERT §7.3, p. 241).
• Oxidation number based definitions: oxidation = increase in ON; reduction = decrease in ON; oxidant = species that increases ON of another; reductant = species that decreases ON of another; redox reaction = reaction with ON change of interacting species (NCERT §7.3, p. 241).
• Fractional oxidation numbers in C3O2 (C = +4/3), Br3O8 (Br = +16/3) and Na2S4O6 (S = +5/2) are averages — structurally the atoms exist in different whole-number states (in C3O2 two terminal C atoms are +2 and the middle C is 0; in S4O62– two terminal S are +5 and two middle S are 0). Fe3O4, Mn3O4 and Pb3O4 are mixed oxides where the same fractional-ON paradox appears (Pb3O4 is 2PbO·PbO2; Fe3O4 is FeO·Fe2O3) (NCERT §7.3 box, p. 244–245).
• Combination redox reactions A + B → C require at least one of A, B to be elemental — e.g. C + O2 → CO2; 3Mg + N2 → Mg3N2; CH4 + 2O2 → CO2 + 2H2O. All combustion of dioxygen is in this class (NCERT §7.3.1, p. 242).
• Decomposition redox reactions are the reverse — a compound breaks down with at least one product in elemental form: 2H2O → 2H2 + O2; 2NaH → 2Na + H2; 2KClO3 → 2KCl + 3O2. Not every decomposition is redox — CaCO3 → CaO + CO2 involves no ON change (NCERT §7.3.1, p. 242).
• Displacement reactions of form X + YZ → XZ + Y come in two flavours. Metal displacement — CuSO4 + Zn → ZnSO4 + Cu; V2O5 + 5Ca → 2V + 5CaO; TiCl4 + 2Mg → Ti + 2MgCl2; Cr2O3 + 2Al → Al2O3 + 2Cr — drives industrial metallurgy (thermite process etc.) (NCERT §7.3.1, p. 242–243).
• Non-metal displacement is dominated by hydrogen displacement: alkali metals and Ca/Sr/Ba displace H2 from cold water (2Na + 2H2O → 2NaOH + H2); Mg and Fe need steam (Mg + 2H2O → Mg(OH)2 + H2; 2Fe + 3H2O → Fe2O3 + 3H2); less active metals (Zn, Mg, Fe, Cd, Sn) liberate H2 from acids; Ag and Au don't react with HCl. Halogen displacement runs F > Cl > Br > I (Cl2 + 2KBr → 2KCl + Br2; Cl2 + 2KI → 2KCl + I2; Br2 + 2I– → 2Br– + I2) — basis of the "layer test" (NCERT §7.3.1, p. 243).
• Disproportionation — one species in an intermediate oxidation state is simultaneously oxidised and reduced. The element must have at least three accessible oxidation states. Examples: 2H2O2 → 2H2O + O2 (O at –1 → –2 and 0); P4 + 3OH– + 3H2O → PH3 + 3H2PO2– (P: 0 → –3 and +1); S8 + 12OH– → 4S2– + 2S2O32– + 6H2O (S: 0 → –2 and +2); Cl2 + 2OH– → ClO– + Cl– + H2O — basis of household bleach (Cl: 0 → +1 and –1) (NCERT §7.3.1, p. 244).
• Fluorine deviation: F2 + 2OH– → 2F– + OF2 + H2O — F has no positive oxidation state available, so it does not disproportionate. Among ClO–, ClO2–, ClO3–, ClO4–, only ClO4– does not disproportionate because Cl is already in its highest state +7 (NCERT §7.3.1, p. 244).
• Balancing — oxidation number method: (1) write skeletal equation; (2) assign ON to identify oxidant and reductant; (3) make ON-increase equal to ON-decrease by suitable multiplication; (4) balance ionic charges with H+ (acidic medium) or OH– (basic medium); (5) balance H atoms with H2O — final cross-check that O atoms balance (NCERT §7.3.2, p. 246).
• Balancing — half-reaction (ion-electron) method: (1) write unbalanced ionic equation; (2) split into oxidation and reduction halves; (3) balance atoms other than O and H first; (4) in acidic medium balance O with H2O and H with H+; (5) balance charge with electrons; (6) equalise electrons across the two halves and add; (7) verify. For basic medium, balance as if acidic, then add equal OH– to both sides to neutralise H+; combine H+ + OH– → H2O (NCERT §7.3.2, p. 246–249). Worked examples in chapter: Cr2O72– + 3SO32– + 8H+ → 2Cr3+ + 3SO42– + 4H2O; 2MnO4– + Br– + H2O → 2MnO2 + BrO3– + 2OH–; 6Fe2+ + Cr2O72– + 14H+ → 6Fe3+ + 2Cr3+ + 7H2O; 6I– + 2MnO4– + 4H2O → 3I2 + 2MnO2 + 8OH–.
• Redox titrations rely on indicator-controlled colour change. (i) KMnO4 acts as its own indicator — pink persists at MnO4– ~10–6 mol L–1 just past equivalence. (ii) K2Cr2O7 is not a self-indicator; diphenylamine is oxidised immediately past the equivalence point giving intense blue. (iii) Iodometric titrations use the starch–iodine deep-blue with Cu2+ liberating I2 from KI, the I2 then consumed by thiosulphate via I2 + 2S2O32– → 2I– + S4O62– (NCERT §7.3.3, p. 249).
• Redox in electrochemical cells: the same Zn + CuSO4 reaction, when zinc and copper rods are separated into two beakers connected by a salt bridge (U-tube of KCl/NH4NO3 in agar) and joined by an external wire, becomes the Daniell cell. Electrons travel externally from Zn anode (oxidation) to Cu cathode (reduction); ions migrate via the salt bridge (NCERT §7.4, p. 250).
• Redox couple = oxidised/reduced form of the same species (Zn2+/Zn, Cu2+/Cu) — oxidised form is written first. The electrode potential measures the tendency of the couple to remain in oxidised or reduced form. At unit activity and 298 K it is the standard electrode potential E°, with E°(H+/H2) = 0.00 V by convention. Negative E° means stronger reducing agent than H+/H2; positive E° means weaker reducing agent than H+/H2 (NCERT §7.4, p. 250–251).
• Electrochemical series (Table 7.1): in descending oxidising strength — F2 (+2.87), Co3+ (+1.81), H2O2 (+1.78), MnO4– (+1.51 acidic), Au3+ (+1.40), Cl2 (+1.36), Cr2O72– (+1.33), O2 (+1.23), MnO2 (+1.23), Br2 (+1.09), Ag+ (+0.80), Fe3+/Fe2+ (+0.77), I2 (+0.54), Cu2+ (+0.34), H+/H2 (0.00), Pb2+ (–0.13), Sn2+ (–0.14), Ni2+ (–0.25), Fe2+ (–0.44), Cr3+ (–0.74), Zn2+ (–0.76), Al3+ (–1.66), Mg2+ (–2.36), Na+ (–2.71), Ca2+ (–2.87), K+ (–2.93), Li+ (–3.05). Reducing strength increases down the table; oxidising strength increases up (NCERT Table 7.1, p. 251).

2.2 Definitions to memorise

2.3 Diagrams / processes to remember

• Fig. 7.1 — Zn strip in aqueous Cu(NO3)2: zinc strip becomes coated with reddish copper; blue colour fades as Cu2+ is reduced (NCERT §7.2.1, p. 238).
• Fig. 7.2 — Cu strip in aqueous AgNO3: solution turns blue (Cu → Cu2+) and silver deposits on the strip (NCERT §7.2.1, p. 239).
• Fig. 7.3 — Daniell cell: Zn rod in ZnSO4 + Cu rod in CuSO4 connected by a salt bridge and external wire with ammeter. Electrons flow Zn → Cu through wire; conventional current is opposite (NCERT §7.4, p. 250).
• Highest-oxidation-number table (p. 241): Na(+1), Mg(+2), Al(+3), Si(+4), P(+5), S(+6), Cl(+7) — highest ON of a representative element generally equals its group number (or group number – 10 for groups 13–17).
• Table 7.1 — Standard electrode potentials at 298 K: memorise that F2 sits at the top (+2.87 V, strongest oxidant) and Li at the bottom (–3.05 V, strongest reductant); H+/H2 is the zero (NCERT Table 7.1, p. 251).
• Structural pictures of "fractional ON" species (p. 245): C3O2 = O=C=C*=C=O (terminal C +2, middle C 0); S4O62– (terminal S +5, middle S 0); Br3O8 (terminal Br +6, middle Br +4). These show fractional ON is an average artefact.

2.4 Common confusions / NTA trap points

• "Every decomposition reaction is a redox reaction" — false. CaCO3 → CaO + CO2 has no change in oxidation number; only decompositions where at least one product is elemental are redox (e.g. 2KClO3 → 2KCl + 3O2 is redox, 2NaH → 2Na + H2 is redox, but CaCO3 decomposition is not) (NCERT §7.3.1, p. 242).
• ON of oxygen is not always –2: in peroxides it is –1 (H2O2, Na2O2), in superoxides it is –½ (KO2, RbO2), in OF2 it is +2 and in O2F2 it is +1 — NTA loves to put H2O2 or OF2 in the stem (NCERT §7.3 rule 3, p. 240).
• ON of hydrogen is +1 except in metal hydrides (LiH, NaH, CaH2, NaBH4 — where it's –1). Students who blindly write +1 for H in NaH will miss the reduction step in 2NaH → 2Na + H2 (NCERT §7.3 rule 4, p. 240).
• Fluorine never shows a positive oxidation state and therefore never disproportionates. Among ClO–, ClO2–, ClO3–, ClO4–, the one that does not disproportionate is ClO4– (Cl already in +7, no higher state available) — NTA frequently sets this as the trap (NCERT §7.3.1 + Problem 7.5, p. 244).
• KMnO4 acts as its own indicator because of its intense purple colour; K2Cr2O7 does not — it needs diphenylamine. Mixing these up is a classic trap (NCERT §7.3.3, p. 249).
• The half-reaction method in basic medium does not add OH– directly to balance O — you balance as in acid first with H+/H2O, then neutralise each H+ with an OH– on both sides, combining the new H+ + OH– into H2O. Skipping this routine gives the wrong stoichiometry.