class 12 notes

Class 12 Chemistry Notes Chapter 3 (Electrochemistry) – Chemistry-I Book

Philoid

Philoid

— 9 min read

Chemistry-I
Alright students, let's get straight into Chapter 3: Electrochemistry. This is a very important chapter, linking chemical reactions with electrical energy, and frequently tested in various government exams. Pay close attention to the concepts and definitions.

Chapter 3: Electrochemistry - Detailed Notes

1. Introduction:

  • Electrochemistry deals with the study of the relationship between electrical energy and chemical changes.
  • It involves:
    • Production of electricity from spontaneous chemical reactions (Galvanic/Voltaic cells).
    • Use of electrical energy to carry out non-spontaneous chemical reactions (Electrolytic cells).

2. Electrochemical Cells (Galvanic or Voltaic Cells):

  • Principle: Convert chemical energy of a spontaneous redox reaction into electrical energy.
  • Example: Daniell Cell (Zn-Cu Cell):
    • Setup: A zinc rod dipped in ZnSO₄ solution and a copper rod dipped in CuSO₄ solution, connected externally by a wire and internally by a salt bridge.
    • Reactions:
      • At Anode (Oxidation): Zn(s) → Zn²⁺(aq) + 2e⁻ (Negative electrode)
      • At Cathode (Reduction): Cu²⁺(aq) + 2e⁻ → Cu(s) (Positive electrode)
    • Overall Reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
  • Salt Bridge: An inverted U-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 by providing counter-ions.
  • Cell Representation (IUPAC Convention):
    • Anode || Cathode
    • Metal | Metal Ion (conc.) || Metal Ion (conc.) | Metal
    • Example (Daniell Cell): Zn(s) | Zn²⁺(aq, C₁) || Cu²⁺(aq, C₂) | Cu(s)
    • Single vertical line (|) denotes a phase boundary.
    • Double vertical line (||) denotes the salt bridge.

3. Electrode Potential:

  • Definition: The potential difference developed between the electrode and the electrolyte surrounding it.
  • Origin: Separation of charges at the electrode-electrolyte interface due to oxidation (M → Mⁿ⁺ + ne⁻) or reduction (Mⁿ⁺ + ne⁻ → M).
  • Standard Electrode Potential (E°): The electrode potential measured under standard conditions (298 K temperature, 1 M concentration of ions, 1 bar pressure for gases).
  • Standard Hydrogen Electrode (SHE):
    • Reference electrode assigned a standard potential of 0.00 V at all temperatures.
    • Setup: Platinum electrode coated with platinum black, dipped in 1 M H⁺ ion solution, with H₂ gas bubbled at 1 bar pressure.
    • Reaction: 2H⁺(aq, 1M) + 2e⁻ ⇌ H₂(g, 1 bar)
  • Electrochemical Series (ECS): Arrangement of elements in order of their increasing (or decreasing) standard reduction potentials (SRP).
    • Applications:
      • Comparing Oxidising/Reducing Power: Higher SRP → Stronger oxidising agent; Lower SRP → Stronger reducing agent.
      • Calculating Standard Cell Potential (E°cell): E°cell = E°cathode (Reduction) - E°anode (Reduction) OR E°cell = E°cathode (Reduction) + E°anode (Oxidation)
      • Predicting Spontaneity of Redox Reactions: If E°cell is positive, the reaction is spontaneous (ΔG° < 0). If E°cell is negative, the reaction is non-spontaneous (ΔG° > 0).
      • Predicting feasibility of metal displacement: A metal can displace another metal from its salt solution if it has a lower SRP (is higher in the ECS).

4. Nernst Equation:

  • Gives the relationship between electrode potential/cell potential and the concentration of ions and temperature.
  • For an Electrode Reaction: Mⁿ⁺(aq) + ne⁻ → M(s)
    • E_(Mⁿ⁺/M) = E°_(Mⁿ⁺/M) - (RT / nF) ln([M(s)] / [Mⁿ⁺(aq)])
    • E_(Mⁿ⁺/M) = E°_(Mⁿ⁺/M) - (RT / nF) ln(1 / [Mⁿ⁺]) (Since activity of pure solid [M(s)] = 1)
    • At 298 K: E = E° - (0.0591 / n) log(1 / [Mⁿ⁺])
  • For a Cell Reaction: aA + bB → cC + dD
    • Ecell = E°cell - (RT / nF) ln([C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ) (Qc = Reaction Quotient)
    • At 298 K: Ecell = E°cell - (0.0591 / n) log(Qc)
    • 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).

5. Equilibrium Constant (Kc) from Nernst Equation:

  • At equilibrium, Ecell = 0 and Qc = Kc.
  • 0 = E°cell - (RT / nF) ln(Kc)
  • E°cell = (RT / nF) ln(Kc)
  • At 298 K: E°cell = (0.0591 / n) log(Kc)

6. Gibbs Energy and Cell Potential:

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

7. Conductance of Electrolytic Solutions:

  • Resistance (R): Obstruction to the flow of current. 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).
  • Conductance (G): Ease with which current flows. 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 area of cross-section (or conductance of 1 m³ of solution). κ = G (l/A) = (1/R) * (l/A) = 1/ρ. Unit: S m⁻¹ (or S cm⁻¹). (l/A) is the cell constant (G*). κ = G × G*.
  • Molar Conductivity (Λm): Conductivity of a volume V of solution containing 1 mole of electrolyte kept between two electrodes with area A and distance l.
    • Λm = κ / C (where C is molar concentration in mol m⁻³)
    • Λm (S m² mol⁻¹) = κ (S m⁻¹) / [Concentration (mol m⁻³)]
    • Commonly used: Λm (S cm² mol⁻¹) = [κ (S cm⁻¹) × 1000 (cm³/L)] / [Molarity (mol/L)]
  • Variation of Conductivity (κ) and Molar Conductivity (Λm) with Concentration:
    • Conductivity (κ): Decreases with decrease in concentration (dilution) for both strong and weak electrolytes, because the number of ions per unit volume decreases.
    • Molar Conductivity (Λm): Increases with decrease in concentration (dilution) for both types.
      • Strong Electrolytes: Increase is gradual. Due to increased ionic mobility as inter-ionic attractions decrease. Λm = Λ°m - A√C (Debye-Hückel-Onsager equation). Λ°m is molar conductivity at infinite dilution (limiting molar conductivity).
      • Weak Electrolytes: Increase is steep at high dilutions. Due to an increase in the degree of dissociation (α) as well as ionic mobility. Λ°m cannot be obtained by extrapolation.

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

  • Statement: The limiting molar conductivity of an electrolyte (Λ°m) can be represented as the sum of the individual contributions of the anion and cation of the electrolyte.
  • Λ°m = ν₊ λ°₊ + ν₋ λ°₋
    • Where λ°₊ and λ°₋ are the limiting molar conductivities of the cation and anion, respectively.
    • ν₊ and ν₋ are the number of cations and anions per formula unit of the electrolyte.
  • 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) of weak electrolytes: α = Λm / Λ°m.
    • Calculation of dissociation constant (Ka) for weak electrolytes: Ka = Cα² / (1-α).
    • Determination of solubility (S) and solubility product (Ksp) of sparingly soluble salts: Λm ≈ Λ°m for saturated solution. κ = Λ°m × S (where S is solubility in mol/L or mol/m³ depending on units of Λ°m). Ksp can be calculated from S.

9. Electrolytic Cells and Electrolysis:

  • Principle: Use electrical energy to drive a non-spontaneous chemical reaction.
  • Process: Passing direct current through an electrolyte (molten state or aqueous solution) causes chemical decomposition.
  • Mechanism: Cations move to the cathode (negative electrode) and undergo reduction. Anions move to the anode (positive electrode) and undergo oxidation.
  • Products of Electrolysis depend on:
    • Nature of the electrolyte (molten vs. aqueous).
    • Concentration of the solution.
    • Nature of the electrodes (inert like Pt, graphite vs. active like Cu, Ag).
    • Standard electrode potentials of the species present. Overpotential may also play a role.
  • Examples:
    • Molten NaCl: Cathode: Na⁺ + e⁻ → Na(l); Anode: 2Cl⁻ → Cl₂(g) + 2e⁻
    • Aqueous NaCl (concentrated):
      • Cathode: Competition between Na⁺ + e⁻ → Na (E° = -2.71 V) and 2H₂O + 2e⁻ → H₂(g) + 2OH⁻ (E° = -0.83 V at pH 7). H₂ is preferentially discharged.
      • Anode: Competition between 2Cl⁻ → Cl₂(g) + 2e⁻ (E° = +1.36 V) and 2H₂O → O₂(g) + 4H⁺ + 4e⁻ (E° = +1.23 V). Due to overpotential of oxygen, Cl₂ is preferentially discharged.
      • Overall: 2NaCl(aq) + 2H₂O(l) → 2NaOH(aq) + H₂(g) + Cl₂(g)
    • Aqueous CuSO₄ (using Cu electrodes):
      • Cathode: Cu²⁺ + 2e⁻ → Cu(s)
      • Anode: Cu(s) → Cu²⁺ + 2e⁻ (Active electrode participates - used in refining of copper).

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
    • Since Q = I × t (I = current in Amperes, t = time in seconds), w ∝ I × t
    • w = Z × I × t
    • Z = Electrochemical Equivalent (mass deposited by 1 C of charge). Unit: kg C⁻¹ or g C⁻¹.
  • Second Law: When the same quantity of electricity is passed through different electrolytes connected in series, the masses of the substances liberated/deposited at the electrodes are directly proportional to their chemical equivalent weights (E).
    • w ∝ E
    • w₁ / w₂ = E₁ / E₂
  • Relationship between Z, E, and F:
    • Equivalent weight (E) = Molar Mass (M) / n-factor (valency/electrons transferred)
    • Charge required to deposit 1 mole of substance = nF Coulombs.
    • Charge required to deposit 1 equivalent of substance = 1F Coulomb.
    • Mass deposited by nF charge = M grams.
    • Mass deposited by 1F charge = M/n = E grams.
    • Mass deposited by 1 Coulomb charge (Z) = E / F.
    • Z = E / F = M / (nF)
    • Substituting Z in First Law: w = (E / F) × I × t = (M / nF) × I × t

11. Batteries:

  • Galvanic cells arranged in series to produce higher voltage.
  • Primary Batteries: Non-rechargeable. Reaction occurs only once.
    • Dry Cell (Leclanché Cell): Anode: Zn container; Cathode: Carbon (graphite) rod surrounded by MnO₂ powder; Electrolyte: Paste of NH₄Cl and ZnCl₂. Voltage ≈ 1.5 V.
    • Mercury Cell: Anode: Zn-Hg amalgam; Cathode: Paste of HgO and Carbon; Electrolyte: Paste of KOH and ZnO. Constant voltage ≈ 1.35 V. Used in watches, hearing aids.
  • Secondary Batteries: Rechargeable. Cell reaction can be reversed by passing external current.
    • Lead Storage Battery: Commonly used in automobiles. Anode: Spongy lead; Cathode: Grid of lead packed with PbO₂; Electrolyte: 38% H₂SO₄ solution. Voltage ≈ 2 V per cell (usually 6 cells = 12 V).
      • Discharging: Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)
      • Charging: Reverse reaction.
    • Nickel-Cadmium (Ni-Cd) Cell: Anode: Cadmium; Cathode: Nickel(III) oxide (NiO(OH)); Electrolyte: KOH solution. Longer life than lead storage but more expensive. Voltage ≈ 1.2 V.

12. Fuel Cells:

  • Galvanic cells that convert the energy of combustion of fuels (like H₂, CH₄, CO, CH₃OH) directly into electrical energy.
  • Example: H₂-O₂ Fuel Cell:
    • Reactants (H₂, O₂) supplied continuously. Products (water) removed continuously.
    • 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 (≈ 70%), continuous operation, pollution-free (water is the only product in H₂-O₂ cell).

13. Corrosion:

  • Slow deterioration of metals due to reaction with their environment (air, water, chemicals). It's essentially an electrochemical phenomenon.
  • Example: Rusting of Iron:
    • Requires oxygen and water.
    • Impure iron surface acts as a miniature electrochemical cell.
    • Anodic regions (Oxidation): Fe(s) → Fe²⁺(aq) + 2e⁻ (E° = -0.44 V)
    • Electrons flow through the metal to cathodic regions.
    • Cathodic regions (Reduction): O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l) (E° = 1.23 V) (H⁺ comes from H₂CO₃ formed from CO₂ in air and water).
    • Overall (simplified): 2Fe(s) + O₂(g) + 4H⁺(aq) → 2Fe²⁺(aq) + 2H₂O(l)
    • Fe²⁺ is further oxidised by atmospheric oxygen to Fe³⁺, which forms hydrated ferric oxide (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-Cu cell), the flow of electrons is from:
    (a) Copper electrode to Zinc electrode externally
    (b) Zinc electrode to Copper electrode externally
    (c) Copper electrode to Zinc electrode through the salt bridge
    (d) Zinc electrode to Copper electrode through the salt bridge

  2. The standard electrode potential (E°) of the SHE (Standard Hydrogen Electrode) is:
    (a) 1.00 V
    (b) -1.00 V
    (c) 0.00 V
    (d) 3.14 V

  3. According to the Nernst equation, the electrode potential of the Mⁿ⁺/M electrode increases when:
    (a) The concentration of Mⁿ⁺ ions increases
    (b) The concentration of Mⁿ⁺ ions decreases
    (c) The temperature decreases
    (d) The pressure increases

  4. Kohlrausch's law is used to determine:
    (a) Standard cell potential of a galvanic cell
    (b) Limiting molar conductivity of a weak electrolyte
    (c) Rate constant of a reaction
    (d) Equilibrium constant using E°cell

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

  6. Faraday's first law of electrolysis states that the mass of substance deposited is proportional to:
    (a) Current (I) only
    (b) Time (t) only
    (c) Quantity of electricity passed (Q)
    (d) Equivalent weight (E)

  7. Which of the following is a secondary battery?
    (a) Dry cell
    (b) Mercury cell
    (c) Lead storage battery
    (d) H₂-O₂ fuel cell

  8. The unit of molar conductivity (Λm) is:
    (a) S cm⁻¹
    (b) S cm² mol⁻¹
    (c) S⁻¹ cm mol⁻¹
    (d) S cm mol

  9. For a spontaneous cell reaction, the standard cell potential (E°cell) and standard Gibbs energy change (ΔG°) are respectively:
    (a) Positive, Positive
    (b) Negative, Negative
    (c) Positive, Negative
    (d) Negative, Positive

  10. Rusting of iron is an electrochemical process involving the formation of:
    (a) Ferrous oxide (FeO)
    (b) Ferric oxide (Fe₂O₃)
    (c) Hydrated ferric oxide (Fe₂O₃·xH₂O)
    (d) Hydrated ferrous oxide (FeO·xH₂O)

Answer Key for MCQs:

  1. (b)
  2. (c)
  3. (a)
  4. (b)
  5. (d)
  6. (c)
  7. (c)
  8. (b)
  9. (c)
  10. (c)

Study these notes thoroughly. Focus on understanding the principles behind each concept, the equations (especially Nernst and Faraday's laws), and the applications of the electrochemical series and Kohlrausch's law. Practice numerical problems based on these concepts as well. Good luck!

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