Notes of CBSE Class 12 Chemistry Chapter 2 Electrochemistry

Electrochemistry studies the interconversion of chemical and electrical energy. It encompasses spontaneous chemical reactions producing electricity (in galvanic cells ) and the use of electrical energy to drive non-spontaneous reactions (in electrolytic cells ). Key concepts include electrode potential, cell potential (emf), and their relation to Gibbs energy and equilibrium constant. The Nernst equation describes cell potential variation with concentration. Electrolytic solutions conduct electricity via ions, characterized by conductivity and molar conductivity. Kohlrausch's law explains molar conductivity at infinite dilution. Quantitative aspects of electrolysis are governed by Faraday's laws . Practical applications include batteries (primary and secondary), fuel cells , and understanding processes like corrosion .

1. Introduction to Electrochemistry

• Electrochemistry is the study of:
• Production of electricity from energy released during spontaneous chemical reactions.
• Use of electrical energy to bring about non-spontaneous chemical transformations.
• It is important for theoretical and practical reasons, including:
• Production of metals and chemicals (e.g., sodium hydroxide, chlorine, fluorine).
• Batteries and fuel cells (convert chemical energy to electrical).
• Energy-efficient and less polluting reactions.
• Understanding biological processes like nerve signal transmission.
• This unit covers important elementary aspects.

2. Electrochemical Cells

• A device where chemical energy converts to electrical energy or vice versa.
• Two main types:
• Galvanic Cell (Voltaic Cell): Converts chemical energy of a spontaneous redox reaction into electrical energy.
• Electrolytic Cell: Uses electrical energy to carry out non-spontaneous chemical reactions.
• Comparison (based on Daniell cell example):
• Galvanic Operation: Spontaneous reaction (Zn + Cu²⁺ -> Zn²⁺ + Cu) produces electrical energy. External opposite potential (Eext) < 1.1 V. Electrons flow Zn to Cu, current Cu to Zn. Zn dissolves (anode), Cu deposits (cathode).
• Equilibrium: Eext = 1.1 V. No electron or current flow. No chemical reaction.
• Electrolytic Operation: Non-spontaneous reaction (reverse of galvanic) occurs. External opposite potential (Eext) > 1.1 V. Functions as an electrolytic cell. Electrons flow Cu to Zn, current Zn to Cu. Zn is deposited (cathode), Cu dissolves (anode).

3. Components and Concepts of Galvanic Cells

Construction:

• Consists of two half-cells or redox couples. Each half-cell has a metallic electrode dipped in an electrolyte.

Half-cells:

• Oxidation half-reaction occurs at the oxidation half-cell.
• Reduction half-reaction occurs at the reduction half-cell.

Electrodes:

• Anode: Half-cell where oxidation occurs. Has a negative potential with respect to the solution in a galvanic cell. Source of electrons in a galvanic cell.
• Cathode: Half-cell where reduction occurs. Has a positive potential with respect to the solution in a galvanic cell. Receives electrons in a galvanic cell.
• Electron and Current Flow: Electrons flow from the negative electrode (anode) to the positive electrode (cathode). Current flow is opposite to electron flow.
• Salt Bridge: Connects the electrolytes of the two half-cells internally. Sometimes not needed if electrodes are in the same electrolyte.

Electrode Potential:

• Potential difference developed between the electrode and the electrolyte at their interface.
• Arises from the tendency of metal ions to deposit on the electrode or metal atoms to go into solution as ions.
• Standard Electrode Potential (E°): Electrode potential when concentrations of all species in a half-cell are unity. According to IUPAC, standard reduction potentials are called standard electrode potentials.
• Measured with respect to the Standard Hydrogen Electrode (SHE).
• SHE: Pt electrode in contact with H₂ gas (1 bar) and H⁺ ions (1 M). Assigned zero potential at all temperatures. Reaction: H⁺(aq) + e⁻ -> ½ H₂(g).
• A positive E° value means the reduced form is more stable than H₂ gas; the species is reduced more easily than H⁺ ions.
• A negative E° value means H₂ gas is more stable than the reduced form; H⁺ ions can oxidise the species (or the species can reduce H⁺ ions).
• Standard electrode potentials indicate the oxidising power (increases with increasing E°) and reducing power (increases with decreasing E°) of species.

Cell Potential (Ecell):

Potential difference between the two electrodes of a galvanic cell.

• Calculated as the difference between the electrode potentials (reduction potentials) of the cathode and anode.
• Ecell = Eright - Eleft (where right electrode is cathode, left is anode by convention).
• Electromotive Force (emf): Cell potential when no current is drawn from the cell.

Representation of a Galvanic Cell:

• Anode on the left, cathode on the right.
• Vertical line (|) between metal and electrolyte solution.
• Double vertical line (||) between two electrolytes connected by a salt bridge.
• Example: Cu(s)|Cu²⁺(aq)||Ag⁺(aq)|Ag(s).

4. Nernst Equation

• Describes the electrode potential at any concentration, not just unity.
• For a half-reaction Mn⁺(aq) + ne⁻ ® M(s): EMⁿ ⁺/M = E °M ⁿ ⁺/M - (RT/nF) ln (1/[Mn ⁺]) (where [M] for solid is taken as unity)
• For a general cell reaction aA + bB ® cC + dD: Ecell = E°cell - (RT/nF) ln ([C]ᶜ[D] ᵈ / [A] ᵃ[B] ᵇ) or Ecell = E°cell - (RT/nF) ln Q (where Q is the reaction quotient)
• At 298 K, converting ln to log₁₀ and substituting R, F, T values: Ecell = E°cell - (0.059/n) log Q
• Ecell depends on the concentration of ions.

5. Equilibrium Constant from Nernst Equation

• As a galvanic cell operates, reactant concentrations decrease, product concentrations increase.
• Ecell decreases over time.
• At equilibrium, Ecell = 0 and Q = Kc (equilibrium constant).
• Setting Ecell = 0 in the Nernst equation gives the relationship between E°cell and Kc: E°cell = (RT/nF) ln Kc
• At 298 K: E°cell = (0.059/n) log Kc
• Kc can be calculated from E°cell values.

6. Electrochemical Cell and Gibbs Energy of Reaction

• The reversible work done by a galvanic cell equals the decrease in its Gibbs energy (ΔrG).
• Relationship between ΔrG and Ecell: ΔrG = -nFEcell
• n = number of electrons transferred in the reaction.
• F = Faraday constant (~96487 C mol⁻¹).
• ΔrG is an extensive property, depending on n. Ecell is intensive.
• Relationship between standard Gibbs energy change (ΔrG°) and standard cell potential (E°cell): ΔrG° = -nFE°cell
• From ΔrG°, the equilibrium constant K can be calculated: ΔrG° = -RT ln K

7. Conductance of Electrolytic Solutions

• Resistance (R): Measured in ohm (Ω). R ∝ l/A.
• Resistivity (ρ): Constant of proportionality in R = ρ(l/A). Units Ω m or Ω cm. Resistance of a substance 1m long with 1m² cross-section.
• Conductance (G): Inverse of resistance (G = 1/R). Measured in siemens (S).
• Conductivity (κ): Inverse of resistivity (κ = 1/ρ). Units S m⁻¹ or S cm⁻¹. Conductance of a material 1m long with 1m² cross-section.

Materials Classification by Conductivity:

• Conductors: Very large conductivity (metals, alloys, graphite, some organic polymers).
• Insulators: Very low conductivity (glass, ceramics, teflon).
• Semiconductors: Conductivity between conductors and insulators (silicon, germanium, gallium arsenide).
• Superconductors: Zero resistivity/infinite conductivity (certain metals/alloys at low T, some ceramic materials).

Types of Electrical Conductance:

• Metallic (Electronic) Conductance: Due to movement of electrons. Depends on nature/structure of metal, number of valence electrons, temperature (decreases with increasing T). Composition of conductor unchanged.
• Electrolytic (Ionic) Conductance: Due to movement of ions in solution. Depends on nature of electrolyte, size/solvation of ions, nature/viscosity of solvent, concentration of electrolyte, temperature (increases with increasing T). Can change composition due to electrochemical reactions with prolonged DC.

Measurement of Conductivity:

• Difficulties with ionic solutions: DC changes composition; solutions cannot be connected directly to bridge.
• Resolved by using AC power source and a conductivity cell .
• Conductivity cell has two platinized platinum electrodes of area A separated by distance l. Solution between electrodes has resistance R = ρ(l/A) = (1/κ)(l/A).
• Cell Constant (G): * Quantity l/A. Units length⁻¹ (m⁻¹ or cm⁻¹). Determined by measuring resistance of a solution with known conductivity (e.g., KCl solutions). G* = Rκ.
• Resistance measured using a Wheatstone bridge with an AC source and detector. Unknown resistance R₂ = (R₁R₄)/R₃.
• Conductivity of solution κ = G*/R.

Molar Conductivity (Λm):

• Defined as Λm = κ/c.
• c is concentration in mol m⁻³ (if κ in S m⁻¹) or mol L⁻¹ (if κ in S cm⁻¹).
• Units: S m² mol⁻¹ or S cm² mol⁻¹. Relationship: 1 S m²mol⁻¹ = 10⁴ S cm²mol⁻¹.
• Represents conductance of a volume containing 1 mole of electrolyte between electrodes 1 unit distance apart.

Variation with Concentration:

• Conductivity (κ): Always decreases with decrease in concentration (dilution) for both strong and weak electrolytes. Explained by the decreasing number of ions per unit volume.
• Molar Conductivity (Λm): Always increases with decrease in concentration (dilution). Explained by the volume containing one mole of electrolyte increasing significantly on dilution, more than compensating the decrease in κ.
• Limiting Molar Conductivity (Λ°m): Molar conductivity when concentration approaches zero (infinite dilution).

Strong Electrolytes:

• Λm increases slowly with dilution.
• Follows approximately the equation: Λm = Λ°m - A c¹/² .
• Λ°m can be obtained by extrapolation of the Λm vs. c¹/² plot to c=0.
• Constant 'A' depends on the type of electrolyte (e.g., 1-1, 2-1, 2-2).

Weak Electrolytes:

• Λm increases steeply with dilution, especially at low concentrations.
• Increase is primarily due to the increase in the degree of dissociation (α) with dilution.
• Λ°m cannot be obtained by extrapolation.
• Λ°m is obtained using Kohlrausch's law.

Kohlrausch Law of Independent Migration of Ions:

• Statement: Limiting molar conductivity (Λ°m) of an electrolyte is the sum of the individual contributions of the anion and cation at infinite dilution.
• Equation: Λ°m = n ₊ λ° ₊ + n ₋ λ° ₋
• n₊, n₋ are the number of cations and anions per formula unit.
• λ°₊, λ°₋ are the limiting molar conductivities of the cation and anion, respectively.
• Applications: Calculating Λ°m for any electrolyte (including weak electrolytes), determining the dissociation constant of weak electrolytes.
• Degree of Dissociation (α) for a weak electrolyte: Approximated as the ratio of molar conductivity at concentration c (Λm) to limiting molar conductivity (Λ°m). α = Λm / Λ°m
• Dissociation Constant (Kc) for a weak electrolyte (like acetic acid, HAc ⇌ H⁺ + Ac⁻): Kc = (cα²) / (1-α) Substituting α = Λm/Λ°m: Kc = (c(Λm/Λ°m)²) / (1 - Λm/Λ°m)

8. Electrolytic Cells and Electrolysis

• Electrolysis uses an external voltage source to drive a non-spontaneous chemical reaction.
• Example: Electrolysis of aqueous CuSO₄ with copper electrodes.
• Cathode (negative electrode, attracts Cu²⁺): Cu²⁺(aq) + 2e⁻ ® Cu(s) (Copper deposition).
• Anode (positive electrode, attracts SO₄²⁻ but Cu is oxidised): Cu(s) ® Cu²⁺(aq) + 2e⁻ (Copper dissolution).
• Used for purifying copper.
• Electrolytic reduction is used for large-scale production of metals where chemical reducing agents are unsuitable (e.g., Na, Mg, Al).
• Na, Mg from fused chlorides. Al from electrolysis of aluminium oxide in cryolite.

Faraday's Laws of Electrolysis:

Describe the quantitative aspects.

• First Law: Amount of chemical reaction at an electrode is proportional to the quantity of electricity passed.
• Second Law: Amounts of different substances liberated by the same quantity of electricity are proportional to their chemical equivalent weights (Atomic Mass / number of electrons for reduction).
• Quantity of Electricity (Q): Q = It (current × time). Measured in Coulombs (C).
• Faraday (F): The charge on one mole of electrons. 1 F ≈ 96487 C mol ⁻ ¹ (approx. 96500 C mol⁻¹).
• Reduction of 1 mol Ag⁺ requires 1 mol e⁻ (1F).
• Reduction of 1 mol Mg²⁺ requires 2 mol e⁻ (2F).
• Reduction of 1 mol Al³⁺ requires 3 mol e⁻ (3F).

Products of Electrolysis:

• Depend on the material being electrolysed and the electrodes used (inert vs. reactive).
• Depend on the oxidising and reducing species present and their standard electrode potentials (reactions with more positive E° are preferred for reduction at cathode; reactions with less positive E° are preferred for oxidation at anode).
• Overpotential: Extra potential needed for some kinetically slow processes, which can alter expected products based solely on E° values.
• Example: Electrolysis of molten NaCl yields Na and Cl₂.
• Example: Electrolysis of aqueous NaCl yields NaOH, H₂, Cl₂.
• Cathode: Competition between Na⁺ + e⁻ ® Na (E° = -2.71V) and H⁺ + e⁻ ® ½ H₂ (E° = 0.00V). H⁺ reduction is preferred based on E°. Net cathode reaction: H₂O + e⁻ ® ½ H₂ + OH⁻.
• Anode: Competition between Cl⁻ ® ½ Cl₂ + e⁻ (E° = 1.36V) and 2H₂O ® O₂ + 4H⁺ + 4e⁻ (E° = 1.23V). Water oxidation has lower E°, but Cl₂ is preferred due to overpotential of oxygen.
• Example: Electrolysis of H₂SO₄. Dilute gives O₂ at anode (2H₂O ® O₂ + 4H⁺ + 4e⁻, E°=1.23V). Concentrated can give S₂O₈²⁻ at anode (2SO₄²⁻ ® S₂O₈²⁻ + 2e⁻, E°=1.96V).

9. Batteries

• Basically galvanic cells that convert chemical energy to electrical energy.
• Desirable properties: light, compact, stable voltage.
• Primary Batteries: Non-rechargeable; reaction occurs once.
• Dry Cell (Leclanche Cell): Anode: Zinc container (Zn ® Zn²⁺ + 2e⁻). Cathode: Graphite rod surrounded by MnO₂ and carbon paste (MnO₂ + NH₄⁺ + e⁻ ® MnO(OH) + NH₃). Electrolyte: Moist NH₄Cl and ZnCl₂ paste. Potential ~1.5V.
• Mercury Cell: Anode: Zinc-mercury amalgam (Zn(Hg) + 2OH⁻ ® ZnO(s) + H₂O + 2e⁻). Cathode: Paste of HgO and carbon (HgO + H₂O + 2e⁻ ® Hg(l) + 2OH⁻). Electrolyte: KOH and ZnO paste. Overall: Zn(Hg) + HgO(s) ® ZnO(s) + Hg(l). Potential ~1.35V, remains constant.
• Secondary Batteries: Rechargeable by passing current in the opposite direction.
• Lead Storage Battery: Used in automobiles/invertors. Anode: Lead (Pb(s) + SO₄²⁻(aq) ® PbSO₄(s) + 2e⁻). Cathode: Lead grid with PbO₂ (PbO₂(s) + SO₄²⁻(aq) + 4H⁺(aq) + 2e⁻ ® PbSO₄(s) + 2H₂O(l)). Electrolyte: 38% H₂SO₄ solution. Overall discharge: Pb(s) + PbO₂(s) + 2H₂SO₄(aq) ® 2PbSO₄(s) + 2H₂O(l). Recharging reverses this reaction.
• Nickel-Cadmium Cell: Longer life than lead storage, more expensive. Overall discharge: Cd(s) + 2Ni(OH)₃(s) ® CdO(s) + 2Ni(OH)₂(s) + H₂O(l).

10. Fuel Cells

• Galvanic cells that directly convert the energy of combustion of fuels (like H₂, methane, methanol) into electrical energy.
• High efficiency (~70%) compared to thermal plants (~40%).
• Reactants fed continuously, products removed continuously.

Hydrogen-Oxygen Fuel Cell:

• Used in Apollo space program.
• Fuel: Hydrogen (H₂). Oxidiser: Oxygen (O₂). Product: Water (H₂O).
• Electrodes: Porous carbon with catalysts (e.g., Pt, Pd).
• Electrolyte: Concentrated aqueous NaOH solution.
• Cathode: O₂(g) + 2H₂O(l) + 4e⁻ ® 4OH⁻(aq).
• Anode: 2H₂(g) + 4OH⁻(aq) ® 4H₂O(l) + 4e⁻.
• Overall: 2H₂(g) + O₂(g) ® 2H₂O(l).
• Pollution free.

11. Corrosion

• Process where metals are oxidised by loss of electrons, forming oxides or salts.
• Causes damage to metal objects.
• Examples: Rusting of iron, tarnishing of silver, green coating on copper.
• Essentially an electrochemical phenomenon .
• Rusting of Iron:
• Occurs in the presence of water and air.
• A spot on the iron surface acts as the anode where oxidation occurs: Fe(s) ® Fe²⁺(aq) + 2e⁻ (E° = -0.44V).
• Electrons travel through the metal to another spot, which acts as the cathode . Here, oxygen is reduced in the presence of H⁺ ions (from dissolution of CO₂ in water, or other acidic oxides). Cathode reaction: O₂(g) + 4H⁺(aq) + 4e⁻ ® 2H₂O(l) (E° ≈ 1.23V).
• Overall electrochemical reaction: 2Fe(s) + O₂(g) + 4H⁺(aq) ® 2Fe²⁺(aq) + 2H₂O(l) (E°cell ≈ 1.67V).
• The Fe²⁺ ions are further oxidised by atmospheric oxygen to Fe³⁺, which forms rust (hydrated ferric oxide, Fe₂O₃ . xH₂O).
• Prevention:
• Prevent contact with atmosphere (e.g., painting, chemicals).
• Covering with other metals (e.g., Sn, Zn) that are inert or react preferentially.
• Using a sacrificial electrode (a more reactive metal like Mg or Zn) which corrodes instead of the object.

12. Hydrogen Economy

• Vision for using hydrogen as a renewable, non-polluting energy source.
• Electrochemical principles are relevant for:
• Production of hydrogen by electrolysis of water.
• Combustion of hydrogen in fuel cells.

Frequently Asked Questions:

1. What is the difference between a galvanic cell and an electrolytic cell?
   • A galvanic cell converts spontaneous chemical energy into electrical energy.
   • An electrolytic cell uses electrical energy to drive a non-spontaneous chemical reaction.
   • Electrodes have opposite polarities (anode is negative in galvanic, positive in electrolytic when reversing the same reaction).
2. How does conductivity change with dilution for strong and weak electrolytes?
   • Conductivity (κ) always decreases with dilution for both, as the number of ions per unit volume decreases.
   • Molar conductivity (Λm) always increases with dilution.
   • For strong electrolytes, Λm increases slowly; for weak electrolytes, Λm increases steeply due to increased dissociation.
3. What are Faraday's laws of electrolysis?
   • First law states that the amount of substance reacted at an electrode is proportional to the quantity of electricity passed.
   • Second law states that the amounts of different substances liberated by the same quantity of electricity are proportional to their chemical equivalent weights.
   • The quantity of electricity is measured in Coulombs (Q=It), where 1 Faraday (F) is the charge of one mole of electrons (~96487 C/mol).