Redox Reactions

Welcome to HSLC Guru! In this chapter, we explore Redox Reactions — one of the most important categories of chemical reactions where electrons are transferred between species. Redox reactions are fundamental to combustion, corrosion, photosynthesis, respiration, batteries, and metallurgical processes. This ASSEB Class 11 Chemistry guide covers classical and electronic concepts of oxidation–reduction, oxidation number rules, types of redox reactions, balancing techniques (oxidation-number method and ion–electron method), electrode processes, standard electrode potentials, and the electrochemical series with worked examples and practice questions.

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Summary

Classical Concept of Oxidation and Reduction: In the classical sense, oxidation is defined as the addition of oxygen or any electronegative element to a substance, or the removal of hydrogen or any electropositive element. Reduction is the reverse — addition of hydrogen or an electropositive element, or removal of oxygen or an electronegative element. For example, when magnesium burns in air, 2Mg + O₂ → 2MgO; magnesium is oxidised. When CuO is heated with hydrogen, CuO + H₂ → Cu + H₂O; CuO is reduced. The classical definition, although useful, is limited because it cannot explain reactions that do not involve oxygen or hydrogen.

Electron-Transfer Concept: The modern view defines oxidation as loss of electrons and reduction as gain of electrons (mnemonic: OIL RIG — Oxidation Is Loss, Reduction Is Gain). The species that loses electrons is the reducing agent and gets oxidised; the species that gains electrons is the oxidising agent and gets reduced. A redox reaction always involves simultaneous oxidation and reduction. For example, Zn + Cu²⁺ → Zn²⁺ + Cu involves Zn → Zn²⁺ + 2e⁻ (oxidation) and Cu²⁺ + 2e⁻ → Cu (reduction).

Oxidation Number Concept and Rules: The oxidation number (ON) is the apparent charge an atom would have if all bonds were considered ionic. Key rules: (i) ON of an element in its free state is zero; (ii) ON of a monoatomic ion equals its charge; (iii) ON of H is +1 (–1 in metal hydrides); (iv) ON of O is –2 (–1 in peroxides, +2 in OF₂); (v) ON of F is always –1; (vi) sum of ONs in a neutral compound is zero, in an ion equals its charge. An increase in ON indicates oxidation; a decrease indicates reduction. Types of redox reactions include combination (C + O₂ → CO₂), decomposition (2KClO₃ → 2KCl + 3O₂), displacement (Fe + CuSO₄ → FeSO₄ + Cu) and disproportionation (where the same element is simultaneously oxidised and reduced, e.g., 2H₂O₂ → 2H₂O + O₂).

Balancing Redox Equations and Electrochemistry: Two methods are used. In the oxidation-number method, changes in ON are equated by suitable multiplication. In the ion–electron (half-reaction) method, the equation is split into oxidation and reduction half-reactions, balanced for atoms and charge using H⁺/OH⁻ and H₂O, then combined. In a galvanic cell (e.g., Daniell cell), oxidation occurs at the anode and reduction at the cathode, generating electrical energy. The tendency of an electrode to lose or gain electrons is measured by its standard electrode potential (E°), measured against the Standard Hydrogen Electrode (SHE) at 25 °C, 1 M, 1 atm. Arranging electrodes by E° gives the electrochemical series, which predicts spontaneity, displacement reactions, reactivity of metals, and oxidising/reducing strengths. Redox reactions find use in batteries, electroplating, metallurgy, bleaching, photography, and biological energy generation.

$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} – E^\circ_{\text{anode}}$

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Question and Answer

Very Short Answer Questions (1 Mark)

Q1. Define oxidation in terms of electron transfer.

Answer: Oxidation is the loss of one or more electrons by a species during a chemical reaction.

Q2. What is a reducing agent?

Answer: A reducing agent is a substance that donates electrons to another species and itself gets oxidised.

Q3. What is the oxidation number of S in H₂SO₄?

Answer: +6.

Q4. Give an example of a disproportionation reaction.

Answer: 2H₂O₂ → 2H₂O + O₂ (oxygen is simultaneously oxidised and reduced).

Q5. What is the oxidation number of an element in its free (elemental) state?

Answer: Zero.

Q6. Name the reference electrode used to measure standard electrode potentials.

Answer: Standard Hydrogen Electrode (SHE), with E° = 0 V.

Q7. What is a redox couple?

Answer: A redox couple is the combination of the oxidised and reduced forms of the same species, e.g., Cu²⁺/Cu.

Q8. Which electrode acts as the anode in a galvanic cell?

Answer: The electrode where oxidation occurs (negative terminal in a galvanic cell).

Q9. What is the oxidation number of Mn in KMnO₄?

Answer: +7.

Q10. Define the electrochemical series.

Answer: The electrochemical series is an arrangement of elements (electrodes) in the order of increasing standard reduction potentials.

Short Answer Questions (2–3 Marks)

Q1. Distinguish between oxidation and reduction in terms of electron transfer with an example.

Answer: Oxidation is the loss of electrons; reduction is the gain of electrons. Example: in Zn + Cu²⁺ → Zn²⁺ + Cu, Zn loses 2 electrons (Zn → Zn²⁺ + 2e⁻, oxidation) and Cu²⁺ gains 2 electrons (Cu²⁺ + 2e⁻ → Cu, reduction). Oxidation and reduction always occur together — this combined process is called a redox reaction.

Q2. Calculate the oxidation number of Cr in K₂Cr₂O₇.

Answer: Let ON of Cr be x. K = +1, O = –2. Sum = 0 for the neutral compound. Therefore, 2(+1) + 2(x) + 7(–2) = 0, i.e., 2 + 2x – 14 = 0, so 2x = 12, x = +6. Hence, the oxidation number of Cr in K₂Cr₂O₇ is +6.

Q3. Define disproportionation reaction with an example. Why is it called so?

Answer: A disproportionation reaction is a redox reaction in which the same element in one oxidation state is simultaneously oxidised and reduced. Example: 3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O. Here, Cl (0) is oxidised to ClO₃⁻ (+5) and reduced to Cl⁻ (–1). It is called so because a single element undergoes “uneven distribution” of oxidation states.

Q4. Why does fluorine never show a positive oxidation number?

Answer: Fluorine is the most electronegative element. In every compound it forms, the bonding electrons are pulled towards F, giving it a –1 oxidation number. Since no element exceeds F in electronegativity, fluorine cannot acquire a positive oxidation number.

Q5. What is the standard hydrogen electrode? Why is its potential taken as zero?

Answer: The Standard Hydrogen Electrode (SHE) consists of a platinum electrode coated with platinum black, dipped in 1 M H⁺ solution and bathed by H₂ gas at 1 atm at 25 °C. By international convention, its standard electrode potential is set to 0 V so that potentials of all other electrodes can be measured relative to it.

Q6. Identify the oxidising and reducing agents in: MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O.

Answer: Mn changes from +4 (in MnO₂) to +2 (in MnCl₂) — reduction; therefore MnO₂ is the oxidising agent. Cl changes from –1 (in HCl) to 0 (in Cl₂) — oxidation; therefore HCl is the reducing agent.

Q7. What is meant by “redox couple”? Explain with the Daniell cell example.

Answer: A redox couple consists of the oxidised and reduced forms of the same element involved in a redox process, written as Ox/Red. In the Daniell cell, the two redox couples are Zn²⁺/Zn (anode) and Cu²⁺/Cu (cathode). The cell EMF depends on the difference of the standard reduction potentials of these two couples.

Q8. Why is the activity of metals related to their position in the electrochemical series?

Answer: The lower the standard reduction potential (E°) of a metal, the stronger its tendency to lose electrons, and hence the more reactive (more electropositive) it is. Therefore, metals at the top of the electrochemical series (low/negative E°) such as K, Na, and Ca are highly reactive, while metals at the bottom (high/positive E°) such as Au and Pt are noble and unreactive.

Long Answer Questions (5–7 Marks)

Q1. Balance the following equation by the oxidation-number method: KMnO₄ + FeSO₄ + H₂SO₄ → K₂SO₄ + MnSO₄ + Fe₂(SO₄)₃ + H₂O.

Answer: Step 1 — Identify ON changes: Mn: +7 → +2 (decrease of 5, reduction). Fe: +2 → +3 (increase of 1, oxidation).

Step 2 — Equalise electron change: multiply Fe by 5 and Mn by 1. So 1 KMnO₄ and 5 FeSO₄.

Step 3 — Skeletal: 2KMnO₄ + 10FeSO₄ + 8H₂SO₄ → K₂SO₄ + 2MnSO₄ + 5Fe₂(SO₄)₃ + 8H₂O.

Step 4 — Verify atoms: K(2=2), Mn(2=2), Fe(10=10), S(18=18), O(72=72), H(16=16). Balanced equation: 2KMnO₄ + 10FeSO₄ + 8H₂SO₄ → K₂SO₄ + 2MnSO₄ + 5Fe₂(SO₄)₃ + 8H₂O.

Q2. Balance the following equation by the ion–electron (half-reaction) method in acidic medium: MnO₄⁻ + C₂O₄²⁻ → Mn²⁺ + CO₂.

Answer: Step 1 — Write half-reactions:

Reduction: MnO₄⁻ → Mn²⁺. Balance O with H₂O: MnO₄⁻ → Mn²⁺ + 4H₂O. Balance H with H⁺: MnO₄⁻ + 8H⁺ → Mn²⁺ + 4H₂O. Balance charge with e⁻: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O.

Oxidation: C₂O₄²⁻ → 2CO₂. Balance charge: C₂O₄²⁻ → 2CO₂ + 2e⁻.

Step 2 — Equalise electrons: multiply reduction by 2 and oxidation by 5. Add: 2MnO₄⁻ + 16H⁺ + 5C₂O₄²⁻ → 2Mn²⁺ + 8H₂O + 10CO₂.

Final balanced equation: 2MnO₄⁻ + 5C₂O₄²⁻ + 16H⁺ → 2Mn²⁺ + 10CO₂ + 8H₂O.

Q3. Explain the four main types of redox reactions with one example of each.

Answer: (i) Combination reactions: two substances combine to give a compound; if at least one is an element, electrons transfer. Example: 2Mg + O₂ → 2MgO.

(ii) Decomposition reactions: a compound breaks into simpler products with electron transfer. Example: 2KClO₃ → 2KCl + 3O₂.

(iii) Displacement reactions: an element displaces another from its compound; further classified as metal-displacement (Zn + CuSO₄ → ZnSO₄ + Cu) and non-metal displacement (Cl₂ + 2KBr → 2KCl + Br₂).

(iv) Disproportionation reactions: same element is simultaneously oxidised and reduced. Example: 2H₂O₂ → 2H₂O + O₂. Each type is identified by analysing changes in oxidation numbers across reactants and products.

Q4. Describe the construction and working of a Daniell cell. Write the cell reaction and calculate its standard EMF given E°(Cu²⁺/Cu) = +0.34 V and E°(Zn²⁺/Zn) = –0.76 V.

Answer: A Daniell cell consists of a Zn rod dipped in 1 M ZnSO₄ and a Cu rod dipped in 1 M CuSO₄, the two half-cells being connected by a salt bridge (saturated KCl in agar) and an external wire. At the anode (Zn), oxidation occurs: Zn → Zn²⁺ + 2e⁻. At the cathode (Cu), reduction occurs: Cu²⁺ + 2e⁻ → Cu. Net cell reaction: Zn + Cu²⁺ → Zn²⁺ + Cu. The salt bridge maintains electrical neutrality. Cell representation: Zn | Zn²⁺ (1 M) || Cu²⁺ (1 M) | Cu.

EMF: E°(cell) = E°(cathode) – E°(anode) = (+0.34) – (–0.76) = +1.10 V. Positive EMF confirms the reaction is spontaneous.

Q5. What is the electrochemical series? List any five important applications.

Answer: The electrochemical series is the arrangement of metals (and other electrodes) in the increasing order of their standard reduction potentials. Lithium has the lowest E° (most easily oxidised; strongest reducing agent), and fluorine has the highest E° (strongest oxidising agent).

Applications: (i) Predicting feasibility of redox reactions — a positive E°(cell) means spontaneous. (ii) Reactivity of metals — metals higher in the series displace those below from their salt solutions. (iii) Reaction of metals with acids — metals above hydrogen liberate H₂ from dilute acids. (iv) Comparing oxidising/reducing strength — higher E° means stronger oxidising agent. (v) Selecting electrode materials for galvanic cells, batteries, and corrosion-protective metals like sacrificial anodes (e.g., Zn for Fe protection).

Q6. Balance the following half-reaction in basic medium: MnO₄⁻ → MnO₂.

Answer: Step 1 — Balance Mn (already balanced). Step 2 — Balance O by adding H₂O on the side deficient in O: MnO₄⁻ → MnO₂ + 2H₂O (left has 4 O, right needs 4, so add 2 H₂O). Step 3 — Balance H by adding H⁺: MnO₄⁻ + 4H⁺ → MnO₂ + 2H₂O. Step 4 — Balance charge with electrons: MnO₄⁻ + 4H⁺ + 3e⁻ → MnO₂ + 2H₂O. Step 5 — Convert to basic medium by adding 4 OH⁻ to both sides: MnO₄⁻ + 4H₂O + 3e⁻ → MnO₂ + 2H₂O + 4OH⁻. Simplify: MnO₄⁻ + 2H₂O + 3e⁻ → MnO₂ + 4OH⁻.

Q7. List five everyday and industrial applications of redox reactions.

Answer: (i) Combustion of fuels for energy (CH₄ + 2O₂ → CO₂ + 2H₂O). (ii) Respiration and photosynthesis in living systems — oxidation of glucose generates ATP. (iii) Batteries and dry cells — chemical to electrical energy conversion in cell phones, vehicles, and torches. (iv) Metallurgy and electroplating — extraction of metals (e.g., reduction of Fe₂O₃ in blast furnace) and protective coating with Cr/Zn. (v) Bleaching, photography, and sanitation — chlorine bleaches fabrics by oxidation; AgBr in photographic films is reduced to metallic Ag on exposure to light; chlorination of water disinfects pathogens.

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Multiple Choice Questions (MCQs)

Q1. Oxidation involves:
(a) gain of electrons (b) loss of electrons (c) gain of protons (d) loss of protons
Answer: (b) loss of electrons.

Q2. The oxidation number of N in NH₄NO₃ on the average is:
(a) +1 (b) –3 (c) +5 (d) Both –3 and +5 are present
Answer: (d) –3 in NH₄⁺ and +5 in NO₃⁻.

Q3. Which of the following is a disproportionation reaction?
(a) Zn + Cu²⁺ → Zn²⁺ + Cu (b) 2KClO₃ → 2KCl + 3O₂ (c) 3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O (d) C + O₂ → CO₂
Answer: (c).

Q4. Standard electrode potential of Zn²⁺/Zn is:
(a) +0.34 V (b) –0.76 V (c) +0.80 V (d) 0 V
Answer: (b) –0.76 V.

Q5. The oxidation number of O in H₂O₂ is:
(a) –2 (b) –1 (c) 0 (d) +1
Answer: (b) –1.

Q6. The strongest oxidising agent in the electrochemical series is:
(a) Li (b) Na (c) F₂ (d) Zn
Answer: (c) F₂.

Q7. In the reaction 2Na + Cl₂ → 2NaCl, sodium is:
(a) reduced (b) oxidised (c) unchanged (d) precipitated
Answer: (b) oxidised.

Q8. The oxidation number of S in Na₂S₂O₃ is:
(a) +2 (b) +4 (c) +6 (d) –2
Answer: (a) +2 (average).

Q9. In a galvanic cell, the cathode is:
(a) positive electrode where reduction occurs (b) negative electrode where reduction occurs (c) positive electrode where oxidation occurs (d) negative electrode where oxidation occurs
Answer: (a).

Q10. Which of the following metals will displace copper from CuSO₄ solution?
(a) Au (b) Hg (c) Fe (d) Ag
Answer: (c) Fe.

Fill in the Blanks

Q1. A reducing agent ____________ electrons during a redox reaction.
Answer: donates (loses).

Q2. The oxidation number of hydrogen in metal hydrides is ____________.
Answer: –1.

Q3. The electrode potential of SHE at 25 °C is ____________ volt.
Answer: 0 (zero).

Q4. The reaction in which the same element is simultaneously oxidised and reduced is called a ____________ reaction.
Answer: disproportionation.

Q5. In the electrochemical series, metals above hydrogen liberate ____________ from dilute acids.
Answer: hydrogen gas (H₂).

Q6. The oxidation number of carbon in CO₂ is ____________.
Answer: +4.

Q7. The salt bridge in a galvanic cell maintains ____________ neutrality.
Answer: electrical (charge).

True or False

Q1. Oxidation and reduction always occur simultaneously. (True)

Q2. The oxidation number of fluorine in any compound is always +1. (False — it is always –1.)

Q3. In a Daniell cell, copper acts as the cathode. (True)

Q4. H₂O₂ can act both as an oxidising and a reducing agent. (True)

Q5. A negative E°(cell) indicates a spontaneous redox reaction. (False — positive E°(cell) indicates spontaneity.)

Q6. Lithium occupies the topmost position in the electrochemical series and is the strongest reducing agent in solution. (True)

Q7. The oxidation-number method and the half-reaction method give different balanced equations for the same redox reaction. (False — both yield the same balanced equation.)

Glossary

Term	Meaning
Redox Reaction	A chemical reaction in which oxidation and reduction occur together.
Oxidation	Loss of electrons; increase in oxidation number.
Reduction	Gain of electrons; decrease in oxidation number.
Oxidising Agent	Species that gains electrons (gets reduced) and oxidises another species.
Reducing Agent	Species that loses electrons (gets oxidised) and reduces another species.
Oxidation Number	Apparent charge on an atom assuming complete electron transfer in bonds.
Disproportionation	Redox reaction where the same element is both oxidised and reduced.
Half-Reaction	Either the oxidation or reduction part of a redox equation written separately.
Galvanic Cell	Electrochemical cell that converts chemical energy into electrical energy.
Salt Bridge	Tube containing electrolyte (e.g., KCl in agar) that maintains charge balance between half-cells.
Standard Electrode Potential (E°)	Potential of an electrode relative to SHE under standard conditions (1 M, 1 atm, 25 °C).
Standard Hydrogen Electrode (SHE)	Reference electrode with E° = 0 V; Pt | H₂ (1 atm) | H⁺ (1 M).
Electrochemical Series	Arrangement of electrodes in increasing order of standard reduction potentials.
Anode	Electrode where oxidation occurs (negative in galvanic, positive in electrolytic cell).
Cathode	Electrode where reduction occurs (positive in galvanic, negative in electrolytic cell).
EMF	Electromotive force; the potential difference between cathode and anode under standard conditions.
Redox Couple	The pair of oxidised and reduced forms of the same species (e.g., Cu²⁺/Cu).
Combination Reaction	Two or more substances combine to form a single product.
Decomposition Reaction	A single compound breaks down into two or more simpler products.
Displacement Reaction	An element displaces another from its compound.
Combustion	Rapid reaction with oxygen producing heat and light; an oxidation process.
Corrosion	Slow oxidation of metals (e.g., rusting of iron) by atmospheric agents.
Oxidation-Number Method	Balancing technique that equalises increase and decrease in oxidation numbers.
Ion–Electron Method	Balancing technique using separate oxidation and reduction half-reactions.