Class 12 Chemistry Notes Chapter 3 (Electrochemistry) – Examplar Problems Book

Alright, let's dive straight into Chapter 3: Electrochemistry. This is a crucial chapter, linking chemical reactions with electrical energy, and frequently tested in various government exams. We'll cover the core concepts based on the NCERT syllabus and Exemplar problems.

Electrochemistry: Detailed Notes for Exam Preparation

1. Introduction:

• Definition: Electrochemistry deals with the relationship between electrical energy and chemical changes. It studies chemical reactions that produce electricity (in galvanic/voltaic cells) and the use of electrical energy to drive non-spontaneous chemical reactions (in electrolytic cells).

2. Electrochemical Cells (Galvanic/Voltaic Cells):

• Principle: Converts chemical energy of a spontaneous redox reaction into electrical energy.
• Setup: Consists of two half-cells.
  • Anode: Electrode where oxidation occurs (negative pole in galvanic cell).
  • Cathode: Electrode where reduction occurs (positive pole in galvanic cell).
  • Salt Bridge: A U-shaped tube containing an inert electrolyte (like KCl, KNO3, NH4NO3 in agar-agar gel). Functions: Completes the electrical circuit by allowing ion flow; maintains electrical neutrality in the half-cells.
• Example: Daniell Cell (Zn-Cu cell)
  • Anode (Oxidation): Zn(s) → Zn²⁺(aq) + 2e⁻
  • Cathode (Reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)
  • Overall Reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
• Cell Representation: Anode || Cathode (Salt bridge indicated by ||)
  • Daniell Cell: Zn(s) | Zn²⁺(aq, C1) || Cu²⁺(aq, C2) | Cu(s)
• Electrode Potential (E): The potential difference developed between an electrode and its electrolyte. It's a measure of the tendency of an electrode to lose or gain electrons.
• Standard Electrode Potential (E°): Electrode potential measured under standard conditions (298 K, 1 atm pressure for gases, 1 M concentration for ions).
• Standard Hydrogen Electrode (SHE): Reference electrode with an assigned E° value of 0.00 V at all temperatures.
  • Reaction: 2H⁺(aq, 1M) + 2e⁻ ⇌ H₂(g, 1 atm)
  • Representation: Pt(s) | H₂(g, 1 atm) | H⁺(aq, 1M)
• Electrochemical Series: Arrangement of elements in order of their increasing (or decreasing) standard reduction potentials.
  • Applications:
    • Predicting relative oxidizing/reducing power (Higher E° = stronger oxidizing agent; Lower E° = stronger reducing agent).
    • Calculating Standard EMF of a cell (E°cell).
    • Predicting spontaneity of a redox reaction.
    • Predicting whether a metal can displace another from its salt solution or liberate H₂ from acid.
• EMF (Electromotive Force) of a Cell (Ecell): Potential difference between the two electrodes when no current is drawn.
  • Ecell = E_reduction (Cathode) - E_reduction (Anode)
  • Alternatively, Ecell = E_oxidation (Anode) + E_reduction (Cathode)
  • Standard EMF (E°cell): E°cell = E°cathode - E°anode (using standard reduction potentials)
  • A positive E°cell indicates a spontaneous reaction.

3. Nernst Equation:

• Gives the relationship between electrode potential (or cell potential) and the concentration of species involved.
• For an electrode reaction: Mⁿ⁺(aq) + ne⁻ → M(s)
  • E_(Mⁿ⁺/M) = E°_(Mⁿ⁺/M) - (RT/nF) ln([M]/[Mⁿ⁺])
  • Since [M(s)] = 1, E_(Mⁿ⁺/M) = E°_(Mⁿ⁺/M) - (RT/nF) ln(1/[Mⁿ⁺])
• For a cell reaction: aA + bB → cC + dD
  • Ecell = E°cell - (RT/nF) ln([C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ)
  • At 298 K: Ecell = E°cell - (0.0591/n) log₁₀([C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ)
  • Where: R = Gas constant (8.314 J K⁻¹ mol⁻¹), T = Temperature (K), n = number of electrons transferred in the balanced redox reaction, F = Faraday constant (≈ 96485 C mol⁻¹).

4. Equilibrium Constant (Kc) from Nernst Equation:

• At equilibrium, Ecell = 0. The reaction quotient Q equals Kc.
• 0 = E°cell - (RT/nF) ln Kc
• E°cell = (RT/nF) ln Kc
• At 298 K: E°cell = (0.0591/n) log₁₀ Kc

5. Gibbs Energy (ΔG) and Cell Potential:

• The electrical work done by a galvanic cell is related to the Gibbs energy change of the reaction.
• ΔrG = -nFEcell
• Under standard conditions: ΔrG° = -nFE°cell
• Relationship with Kc: ΔrG° = -RT ln Kc
• Spontaneity:
  • If E°cell > 0, ΔrG° < 0 (Spontaneous reaction)
  • If E°cell < 0, ΔrG° > 0 (Non-spontaneous reaction)
  • If E°cell = 0, ΔrG° = 0 (Equilibrium)

6. Conductance in Electrolytic Solutions:

• Electrolytic Conduction: Flow of current due to the movement of ions in a solution or molten electrolyte.
• Resistance (R): Opposition to current flow. Unit: Ohm (Ω). R = ρ (l/A)
• Resistivity (ρ): Resistance of a conductor of 1 m length and 1 m² cross-sectional area. Unit: Ohm-meter (Ω m) or Ohm-cm (Ω cm).
• Conductance (G): Ease of current flow. Reciprocal of resistance. G = 1/R. Unit: Siemens (S) or ohm⁻¹ or mho.
• Conductivity (κ, kappa): Conductance of a solution of 1 m length and 1 m² cross-sectional area (or 1 cm length and 1 cm² area). Reciprocal of resistivity. κ = 1/ρ = G (l/A). Unit: S m⁻¹ or S cm⁻¹.
  • l/A is the Cell Constant (G)*. Unit: m⁻¹ or cm⁻¹.
  • κ = G × G*
• Molar Conductivity (Λm): Conductivity of a volume of solution containing one mole of electrolyte placed between two electrodes 1 m (or 1 cm) apart with a large enough area to contain the entire solution.
  • Λm = κ / C (where C is molar concentration in mol m⁻³)
  • If κ is in S cm⁻¹ and C is in mol L⁻¹ (Molarity): Λm (S cm² mol⁻¹) = (κ × 1000) / C
  • Unit: S m² mol⁻¹ or S cm² mol⁻¹.

7. Variation of Conductivity and Molar Conductivity with Concentration:

• Conductivity (κ): Decreases with decreasing concentration (dilution) for both strong and weak electrolytes. Reason: The number of ions per unit volume carrying the current decreases upon dilution.
• Molar Conductivity (Λm): Increases with decreasing concentration (dilution) for both strong and weak electrolytes.
  • Strong Electrolytes: Λm increases slowly with dilution. Due to increased ionic mobility as inter-ionic attractions decrease. Can be extrapolated to zero concentration (infinite dilution) using the Debye-Hückel-Onsager equation: Λm = Λ°m - A√C, where Λ°m is the limiting molar conductivity.
  • Weak Electrolytes: Λm increases steeply at very low concentrations. Due to a significant increase in the degree of dissociation (α) upon dilution. Cannot be easily extrapolated to find Λ°m.
  • Limiting Molar Conductivity (Λ°m): Molar conductivity when the concentration approaches zero (infinite dilution).

8. Kohlrausch's Law of Independent Migration of Ions:

• Statement: The limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation of the electrolyte.
• Λ°m (AxBy) = x λ°(Aʸ⁺) + y λ°(Bˣ⁻)
  • Where λ° is the limiting molar ionic conductivity.
• Applications:
  • Calculation of Λ°m for weak electrolytes (e.g., Λ°m(CH₃COOH) = Λ°m(CH₃COONa) + Λ°m(HCl) - Λ°m(NaCl)).
  • Calculation of degree of dissociation (α) for weak electrolytes: α = Λm / Λ°m.
  • Calculation of dissociation constant (Ka) for weak electrolytes: Ka = Cα² / (1-α) = C(Λm/Λ°m)² / (1 - Λm/Λ°m).
  • Determination of solubility (S) and solubility product (Ksp) of sparingly soluble salts (e.g., AgCl, BaSO₄). For these, the saturated solution is very dilute, so Λm ≈ Λ°m. κ = Λ°m × S (where S is molar solubility).

9. Electrolytic Cells and Electrolysis:

• Principle: Uses external electrical energy to drive a non-spontaneous chemical reaction.
• Setup: Electrodes dipped in an electrolytic solution or molten electrolyte, connected to a DC power source.
  • Anode: Connected to the positive terminal of the battery; Oxidation occurs here.
  • Cathode: Connected to the negative terminal of the battery; Reduction occurs here.
• Mechanism: Cations move towards the cathode (negative electrode) and get reduced. Anions move towards the anode (positive electrode) and get oxidized.
• Products of Electrolysis: Depend on:
  • Nature of the electrolyte (molten vs. aqueous).
  • Standard electrode potentials of the species present.
  • Concentration of the solution.
  • Nature of the electrodes (inert like Pt, graphite vs. active like Cu, Ag).
  • Overpotential: Sometimes, a reaction with higher E° might occur preferentially if the kinetically favoured reaction requires extra voltage (overpotential) to proceed at a reasonable rate (e.g., liberation of O₂ often requires overpotential compared to oxidation of Cl⁻).
• Examples:
  • Electrolysis of molten NaCl: Na⁺ + e⁻ → Na (at cathode); 2Cl⁻ → Cl₂(g) + 2e⁻ (at anode).
  • Electrolysis of aqueous NaCl: 2H₂O + 2e⁻ → H₂(g) + 2OH⁻ (at cathode, E° = -0.83 V vs Na⁺ + e⁻ → Na, E° = -2.71 V); 2Cl⁻ → Cl₂(g) + 2e⁻ (at anode, preferred over 2H₂O → O₂(g) + 4H⁺ + 4e⁻, E° = 1.23 V, due to overpotential of O₂).
  • Electrolysis of aqueous CuSO₄ using inert electrodes: Cu²⁺ + 2e⁻ → Cu (at cathode); 2H₂O → O₂(g) + 4H⁺ + 4e⁻ (at anode).
  • Electrolysis of aqueous CuSO₄ using Cu electrodes: Cu²⁺ + 2e⁻ → Cu (at cathode); Cu(s) → Cu²⁺ + 2e⁻ (at anode - active electrode participates).

10. Faraday's Laws of Electrolysis:

• First Law: The mass (w) of a substance deposited or liberated at any electrode during electrolysis is directly proportional to the quantity of electricity (Q) passed through the electrolyte.
  • w ∝ Q
  • w = Z Q = Z I t (where Q = I × t; I = current in Amperes, t = time in seconds)
  • Z is the Electrochemical Equivalent. Z = Equivalent weight (E) / F = (Molar mass / n-factor) / 96485.
• Second Law: When the same quantity of electricity is passed through different electrolytic solutions connected in series, the masses of the substances liberated at the electrodes are directly proportional to their chemical equivalent weights (E).
  • w₁ / w₂ = E₁ / E₂

11. Batteries:

• Arrangement of one or more electrochemical cells used as a source of direct electric current.
• Primary Batteries: Reaction occurs only once; become dead after use and cannot be recharged.
  • Example: Dry Cell (Leclanché cell), Mercury Cell.
• Secondary Batteries: Can be recharged by passing current through them in the opposite direction; reaction can be reversed.
  • Example: Lead Storage Battery (common car battery), Nickel-Cadmium (Ni-Cd) cell.
  • Lead Storage Battery:
    • Anode: Pb(s) + SO₄²⁻(aq) → PbSO₄(s) + 2e⁻
    • Cathode: PbO₂(s) + SO₄²⁻(aq) + 4H⁺(aq) + 2e⁻ → PbSO₄(s) + 2H₂O(l)
    • Overall (Discharging): Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)
    • Recharging reverses these reactions.

12. Fuel Cells:

• Galvanic cells designed to convert the chemical energy from the combustion of fuels (like H₂, CH₄, CO, CH₃OH) directly into electrical energy.
• Example: H₂-O₂ Fuel Cell:
  • Anode: 2H₂(g) + 4OH⁻(aq) → 4H₂O(l) + 4e⁻
  • Cathode: O₂(g) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)
  • Overall: 2H₂(g) + O₂(g) → 2H₂O(l)
• Advantages: High efficiency (up to 70%), continuous operation (as long as fuel is supplied), pollution-free (if H₂/O₂ used).

13. Corrosion:

• Slow deterioration of a metal due to unwanted chemical or electrochemical reaction with its environment.
• Electrochemical Theory of Rusting (Iron): Impure iron surface behaves like a small electrochemical cell in the presence of water containing dissolved O₂ or CO₂.
  • Anode (Pure Iron): Fe(s) → Fe²⁺(aq) + 2e⁻
  • Cathode (Impure sites/presence of O₂): O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l) (H⁺ from H₂CO₃ formed by CO₂ + H₂O)
  • Atmospheric oxygen further oxidizes Fe²⁺ to Fe³⁺, which forms hydrated ferric oxide (rust): 2Fe²⁺ + 2H₂O + ½O₂ → Fe₂O₃ + 4H⁺; Fe₂O₃ + xH₂O → Fe₂O₃·xH₂O (Rust).
• Prevention: Barrier protection (painting, oiling), sacrificial protection (coating with more reactive metal like Zn - galvanization), using anti-rust solutions, cathodic protection.

Multiple Choice Questions (MCQs)

1. In the Daniell cell, Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s), the negatively charged electrode is:
   a) Cu electrode
   b) Zn electrode
   c) The salt bridge
   d) Both electrodes

2. The standard electrode potential E° for the half-cell Zn²⁺(aq) + 2e⁻ → Zn(s) is -0.76 V. What does this indicate?
   a) Zn²⁺ ions are easily reduced.
   b) Zn metal is easily oxidized.
   c) Zn metal is an oxidizing agent.
   d) Zn²⁺ ions are easily oxidized.

3. The units of molar conductivity (Λm) are:
   a) S cm⁻¹
   b) S cm² mol⁻¹
   c) S⁻¹ cm² mol⁻¹
   d) S cm mol⁻¹

4. According to Kohlrausch's law, the limiting molar conductivity of CaCl₂ can be represented as:
   a) λ°(Ca²⁺) + λ°(Cl⁻)
   b) λ°(Ca²⁺) + 2λ°(Cl⁻)
   c) 2λ°(Ca²⁺) + λ°(Cl⁻)
   d) ½λ°(Ca²⁺) + λ°(Cl⁻)

5. How many Coulombs of charge are required to deposit 27 g of Aluminium (Molar mass = 27 g/mol) from a molten AlCl₃ solution? (F = 96500 C/mol)
   a) 96500 C
   b) 3 × 96500 C
   c) 1 × 96500 / 3 C
   d) 27 × 96500 C

6. Which of the following relationships is correct for a spontaneous reaction in a galvanic cell?
   a) E°cell = positive, ΔG° = positive
   b) E°cell = negative, ΔG° = positive
   c) E°cell = positive, ΔG° = negative
   d) E°cell = negative, ΔG° = negative

7. During the electrolysis of aqueous NaCl solution using inert electrodes, the product obtained at the cathode is:
   a) Na metal
   b) Cl₂ gas
   c) O₂ gas
   d) H₂ gas

8. A lead storage battery is an example of a:
   a) Primary cell
   b) Secondary cell
   c) Fuel cell
   d) Concentration cell

9. In the electrochemical corrosion of iron (rusting), the iron metal acts as the:
   a) Cathode
   b) Anode
   c) Electrolyte
   d) Salt bridge

10. If the concentration of Zn²⁺ ions in the cell Zn | Zn²⁺ || Cu²⁺ | Cu is increased, according to the Nernst equation, the Ecell will:
    a) Increase
    b) Decrease
    c) Remain unchanged
    d) Become zero

Answers to MCQs:

1. b) Zn electrode (Anode is negative in galvanic cells)
2. b) Zn metal is easily oxidized (Negative E° means low tendency for reduction, high tendency for oxidation)
3. b) S cm² mol⁻¹ (or S m² mol⁻¹)
4. b) λ°(Ca²⁺) + 2λ°(Cl⁻) (Based on stoichiometry)
5. b) 3 × 96500 C (Al³⁺ + 3e⁻ → Al; 1 mole Al requires 3 moles of electrons, i.e., 3F charge)
6. c) E°cell = positive, ΔG° = negative (Condition for spontaneity)
7. d) H₂ gas (Reduction of water occurs preferentially over reduction of Na⁺)
8. b) Secondary cell (It is rechargeable)
9. b) Anode (Iron gets oxidized: Fe → Fe²⁺ + 2e⁻)
10. b) Decrease (Ecell = E°cell - (0.0591/n) log([Zn²⁺]/[Cu²⁺]). Increasing [Zn²⁺] increases the log term, making Ecell smaller)

Study these notes thoroughly. Focus on understanding the concepts behind the formulas and definitions, especially the Nernst equation, conductivity variations, Faraday's laws, and the principles of different cell types. Good luck with your preparation!