CBSE Class 11 Chemistry Redox Reactions Notes Set 08

This set of Class 11 CBSE Chemistry notes on Redox Reactions picks up the topics that follow oxidation numbers and balancing: the four types of redox reactions, Stock notation for naming compounds, redox couples, the Daniell cell with its salt bridge, electrode potentials, the standard hydrogen electrode, the electrochemical series and the rusting of iron. Worked examples apply each idea.

Important Notes and Formulas

Types of Redox Reactions

Every redox reaction can be placed in one of four groups according to how the oxidation numbers change.

Not every decomposition or combination is redox: CaO + CO₂ → CaCO₃ and CaCO₃ → CaO + CO₂ involve no change in oxidation number. In a displacement reaction the element that is more reactive (a better reducing agent) displaces the less reactive one from its compound.

Stock Notation

The oxidation number of the metal is written as a Roman numeral in brackets immediately after the name (or symbol) of the metal. Zero is written as 0 and negative numbers carry a minus sign.

In the formula the same idea can be written as Au(I)Cl and Au(III)Cl₃, or Sn(II)Cl₂ and Sn(IV)Cl₄.

Redox Couple

A redox couple is the pair of the oxidised and reduced forms of a substance taking part in an oxidation or reduction half-reaction. It is written with the oxidised form first, e.g. Zn²⁺/Zn and Cu²⁺/Cu. In Zn + Cu²⁺ → Zn²⁺ + Cu there are two couples, Zn²⁺/Zn and Cu²⁺/Cu. In general a couple is written Ox/Red, and the half-reaction is Ox + ne⁻ → Red.

Electrochemical Cell and the Daniell Cell

An electrochemical cell is a device that converts the energy of a spontaneous redox reaction into electrical energy. The Daniell cell is the standard example: the redox reaction Zn + Cu²⁺ → Zn²⁺ + Cu is carried out so that the electron transfer takes place through an external wire.

1. Anode (oxidation, negative): Zn → Zn²⁺ + 2e⁻. The zinc rod dissolves.
2. Cathode (reduction, positive): Cu²⁺ + 2e⁻ → Cu. Copper deposits on the copper rod.
3. Electrons flow in the external wire from zinc to copper; conventional current flows from copper to zinc.
4. The cell is written Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s), where | is a phase boundary and || is the salt bridge.

Salt Bridge

The salt bridge is a U-tube containing a concentrated solution of an inert electrolyte such as KCl or KNO₃ (usually in agar gel). It has two jobs: it completes the electric circuit by letting ions move between the two half-cells, and it keeps each solution electrically neutral. Without it, positive charge would build up around the anode and negative charge around the cathode, and the current would stop almost at once.

Electrode Potential and the Standard Hydrogen Electrode

The potential associated with each electrode is called its electrode potential. When the concentration of every species in the electrode reaction is 1 mol per litre, any gas is at 1 bar pressure, and the temperature is 298 K, it is called the standard electrode potential, E°. Electrode potentials cannot be measured singly, so a reference electrode is used.

The standard hydrogen electrode (SHE) is the reference. It is a platinum foil coated with platinum black, dipping partly in a solution containing H⁺ ions at 1 mol per litre (unit activity), with pure hydrogen gas at 1 bar bubbled over it. Its potential is taken as 0.00 V at all temperatures. The potential of any other half-cell is measured by joining it to the SHE.

Standard cell potential: E°cell = E°cathode − E°anode. For the Daniell cell, E°cell = 0.34 − (−0.76) = 1.10 V. A positive E°cell shows the cell reaction is spontaneous.

Electrochemical Series

Arranging the metals (and hydrogen) in order of their standard reduction potentials, or in order of their reducing power, gives the electrochemical series. The more negative the E°, the stronger the reducing agent. A metal higher in the reducing order displaces one lower in it from its salt solution, and metals above hydrogen displace hydrogen from dilute acids.

Consequences: zinc displaces copper from CuSO₄, so a copper sulphate solution cannot be kept in a zinc vessel; copper does not displace zinc from ZnSO₄, so zinc sulphate can be kept in a copper vessel.

Corrosion of Iron as an Electrochemical Process

Rusting happens because a tiny electrochemical cell forms on the surface of damp iron. Iron acts as the anode and a region with dissolved oxygen acts as the cathode.

1. Oxidation (anode): Fe(s) → Fe²⁺(aq) + 2e⁻
2. Reduction (cathode): O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l)
3. Net reaction: 2Fe²⁺(aq) + 2H₂O(l) + 1/2 O₂(g) → Fe₂O₃(s) + 4H⁺(aq); the iron(III) oxide formed is rust (hydrated).

Rusting is prevented by keeping air and moisture away from the iron (painting, greasing or oiling) or by galvanisation, which is coating iron with the more reactive metal zinc so that zinc, not iron, is oxidised.

Solved Examples

Example 1. Classify each reaction as combination, decomposition or displacement, and state whether it is a redox reaction: (a) 2KClO₃ → 2KCl + 3O₂ (b) Fe + CuSO₄ → FeSO₄ + Cu (c) CaO + CO₂ → CaCO₃
Answer:
(a) Decomposition, redox. Cl goes from +5 to −1 and O from −2 to 0.
(b) Displacement, redox. Fe goes from 0 to +2 and Cu from +2 to 0.
(c) Combination, not redox. Ca stays +2, C stays +4 and O stays −2.

Example 2. Write the Stock names of MnO₂, Mn₂O₃ and Hg₂Cl₂.
Answer:
MnO₂: 2 O at −2 gives Mn +4, so manganese(IV) oxide.
Mn₂O₃: 3 O at −2 gives −6, so each Mn is +3: manganese(III) oxide.
Hg₂Cl₂: two Cl at −1 give −2, so each Hg is +1 (as Hg₂²⁺): mercury(I) chloride.

Example 3. Calculate the standard e.m.f. of the cell Zn | Zn²⁺ || Cu²⁺ | Cu given E°(Cu²⁺/Cu) = +0.34 V and E°(Zn²⁺/Zn) = −0.76 V.
Answer:
Copper has the higher reduction potential, so it is the cathode and zinc is the anode.
E°cell = E°cathode − E°anode = 0.34 − (−0.76) = 1.10 V.