Electrochemistry

1. Electrode Potential

• Electrode : An assembly from which or to which electron movement is observed.
• When a metal rod is dipped into its salt solution, two tendencies are observed:
  • De-electronation: Solution pressure.
  • Electronation: Osmotic pressure.

2. Anode

• At the anode, the solution pressure tendency dominates over the osmotic pressure tendency.
• Anode characteristics:
  • Polarity: Negative (–).
  • Process: Oxidation.
  • Role: Acts as a source of electrons, supplying them to the outer circuit.
• At equilibrium, the rate of dissociation equals the rate of deposition:
  M(s) → Mⁿ⁺(aq) + ne⁻
• Due to charge separation (rod becomes negatively charged, solution becomes positively charged), a potential is developed, called electrode potential.

3. Cathode

• Cathode characteristics:
  • Polarity: Positive (+).
  • Process: Reduction.
  • Role: Attracts and gains electrons from the outer circuit.
• The process occurring at the cathode is:
  Mⁿ⁺(aq) + ne⁻ → M(s)

4. Galvanic Cell

• In a galvanic cell, electrons flow from anode to cathode, while the conventional current (i) flows from cathode to anode.
• Example: Daniel Cell
  • Anode: Zn rod in ZnSO₄(aq) solution.
  • Cathode: Cu rod in CuSO₄(aq) solution.
  • Components are connected by a wire and a KCl salt bridge.

5. Cell Functionality & Salt Bridge

• Compartment blockage: In a galvanic cell, accumulation of Zn²⁺ in the anodic compartment stops further oxidation of Zn (repulsion). In the cathodic compartment, excess SO₄²⁻ left behind stops further reduction of Cu²⁺. Without a salt bridge the cell stops working.
• Salt Bridge Structure: A U-shaped glass/polythene tube filled with a paste of an electrolyte (usually KCl) and agar-agar powder, forming a porous network through which ions can migrate.
• Criteria for Electrolyte: The speed (mobility) of the cation must be equal to the speed (mobility) of the anion, meaning their transport numbers must be equal (e.g., KCl).
• Exceptions (KCl restriction): KCl cannot be used if the compartment contains ions that form precipitates with Cl⁻ (Ag⁺, Hg₂²⁺, Hg²⁺, Tl⁺, Pb²⁺).
  • Precipitation (e.g., AgCl, TlCl, Hg₂Cl₂, HgCl₂) at the mouth of the salt bridge leads to chocking and loss of function.
  • In such cases, use KNO₃ or NH₄NO₃.
• Functions:
  • Completes the internal circuit.
  • Maintains electrical neutrality.
  • Minimizes Liquid-Liquid Junction Potential (LLJP).

6. Standard Cell Representation

• Convention: Anode on the left, cathode on the right.
• Salt bridge: Represented by (||).
• Phase boundary: Different phases are separated by (|).
• Same phase: If electron transfer occurs between species in the same phase, separate them by a comma (,).
• Conditions: Concentrations and pressures are written in front of or below the ions and gases.
• Notation: E°_cell is written on the extreme right; terminals must be metallic.

7. Gas Electrodes

• Gas-Ion Electrode Representation (Anode) :

• Examples: Pt | H₂ | H⁺ and Pt | Cl₂ | Cl^-
• Anode reactions:
  • H₂ → 2H⁺ + 2e^-
  • 2Cl^- → Cl₂ + 2e^-
• Gas-Ion Electrode Representation (Cathode) :

• Examples: || H⁺ | H₂ | Pt and || Cl^- | Cl₂ | Pt
• Cathode reactions:
  • 2H⁺ + 2e^- → H₂
  • Cl₂ + 2e^- → 2Cl^-

8. Standard Electrode Potential (E°)

• Defined as the potential difference between a metal rod and its solution when the metal rod is placed in a 1 M solution at 25°C and 1 bar pressure.
• If the concentration of the metal ion is not 1 M, the potential is not standard.
• Important points:
  • The potential of an individual electrode cannot be measured.
  • Only the difference in electrode potential between two electrodes can be measured.
  • An arbitrary electrode is chosen as a reference with an assumed standard electrode potential = 0 V.
  • Standard Hydrogen Electrode (SHE) is used to calculate the standard electrode potential of other electrodes.

9. Standard Hydrogen Electrode (SHE)

• Setup consists of a Pt rod immersed in a solution with [H⁺] = 1 M, with H₂ gas fed at 1 bar pressure.

• As Anode: Pt | H₂ | H⁺ || — (Reaction: H₂ → 2H⁺ + 2e^-)
• As Cathode: — || H⁺ | H₂ | Pt (Reaction: 2H⁺ + 2e^- → H₂)

10. Electrochemical Series

• The series is formed by arranging electrodes in increasing order of their Standard Reduction Potential (SRP) values, calculated with reference to SHE.

Note: Reactivity of metals decreases on moving down the series. Li⁺/Li is the strongest reducing agent; F₂/F⁻ is the strongest oxidizing agent.

11. Short Exam-friendly Series

Decreasing SRP (oxidising power): F₂ 2.87, MnO₄⁻ 1.51, Au³⁺ 1.50, Cl₂ 1.36, Cr₂O₇²⁻ 1.33, O₂ 1.23, Ag⁺ 0.80, Fe³⁺/Fe²⁺ 0.77, Cu²⁺ 0.34, H⁺ 0, Ni²⁺ −0.25, Fe²⁺ −0.44, Zn²⁺ −0.76, Al³⁺ −1.66, Mg²⁺ −2.36, Na⁺ −2.71, K⁺ −2.93, Li⁺ −3.05 V.

12. Features of Electrochemical Series

• Relation : SRP = - SOP
• Oxidizing/Reducing Strength :
  • Higher SRP value ⇒ greater tendency to be reduced ⇒ stronger Oxidizing Agent (O.A.).
  • Higher SOP value (lower SRP) ⇒ greater tendency to be oxidized ⇒ stronger Reducing Agent (R.A.).
• Cell Construction :
  • Anode: Electrode higher up in the series.
  • Cathode: Electrode lower in the series.
• Displacement : Metals higher in the series can displace metals lower in the series from their salt solutions.
  e.g., CuSO₄ + Zn → ZnSO₄ + Cu (Reaction occurs).
  e.g., ZnSO₄ + Cu → No reaction.

13. Reactivity and Electro-positive Nature

• On moving down the series, the electro-positive nature and reactivity of metals decrease.
• Reactivity with Water/Steam:
  • Alkali/Alkaline earth metals + cold water → H₂(g)↑
  • Fe, Ni, Co, Pb, Sn + steam → H₂(g)↑
  • Cu, Ag, Au + cold water or steam → No reaction.

14. Calculation of E⁰_cell

• Standard cell potential is determined by the difference between the reduction potentials of the cathode and anode.

15. Gibbs Energy and Electrode Potentials

• The Gibbs energy change (ΔG⁰) for a cell reaction or half-cell reaction is:

where n = number of electrons involved in the reaction, F = Faraday's constant (96500 C).

16. Intensive vs. Extensive Properties

• EMF (E⁰) is an intensive property: it does not depend on the amount of substance. Consequently, E⁰ values of two reactions cannot be added or subtracted directly to calculate the E⁰ of a third reaction.
• Gibbs Energy (ΔG⁰) is an extensive property: it depends on the amount of substance. Therefore, ΔG⁰ values of two reactions can be added or subtracted to determine the ΔG⁰ (and subsequently the E⁰) of another reaction.

17. Nernst Equation

• For reaction: a A + b B → c C + d D

• At 298 K:

18. Derivation of Nernst Equation

• ΔG = ΔG° + RT ln Q
• ΔG = (cG_C + dG_D) - (aG_A + bG_B)
• Using ΔG = -nFE_cell and ΔG° = -nFE°_cell:
• -nFE_cell = -nFE°_cell + RT ln Q
• E_cell = E°_cell - RTnF ln Q

19. Constants and Conversions

20. Graphs of Nernst Equation

• E_cell = E°_cell - RTnF ln Q
• E_cell = E°_cell - 2.303 RTnF log Q
• Graphs of E_cell vs ln Q and E_cell vs log Q are downward sloping lines.
• Slope for ln Q: -RTnF = -0.05912.303n
• Slope for log Q: -2.303 RTnF = -0.0591n

21. Variation of Electrode Potentials with Time

• For the cell Zn | Zn²⁺ || Cu²⁺ | Cu:
• Cathode: Cu²⁺ + 2e⁻ → Cu
E_Cu²⁺/Cu = E°_Cu²⁺/Cu + 0.03 log [Cu²⁺]
• As time passes, [Cu²⁺] decreases, therefore E_Cu²⁺/Cu decreases.
• Anode: Zn²⁺ + 2e⁻ → Zn
E_Zn²⁺/Zn = E°_Zn²⁺/Zn + 0.03 log [Zn²⁺]
• As time passes, [Zn²⁺] increases, therefore E_Zn²⁺/Zn increases.
• E_cell = (E_RP)_cathode - (E_RP)_anode. As time progresses, (E_RP)_cathode decreases and (E_RP)_anode increases until E_cell = 0 at equilibrium.

22. Equilibrium and Gibbs Energy

• At equilibrium (E_cell = 0):

23. Hydrogen Electrode & pH

• For the reduction reaction: 2H⁺ + 2e⁻ → H₂
• E_H⁺|H₂ = E°_H⁺|H₂ - 0.05912 log P_H₂[H⁺]²
• Given E°_H⁺|H₂ = 0:

24. Oxidising Power Change

• For the half reaction: 5e⁻ + MnO₄⁻ + 8H⁺ → Mn²⁺ + 4H₂O

25. Types of Electrodes

• 1) Metal | Metal-ion electrode
• 2) Gas | Gas-ion electrode
• 3) Metal-insoluble Salt anion electrode
• 4) Redox electrode

26. General Representation of Salt-Anion Electrodes

• Cathodic Representation: Salt + ne⁻ → Metal + Anion
• Anodic Representation: Metal + Anion → Salt + ne⁻

Note: E_{anion/salt/metal} depends on the concentration of the anion.

27. Metal-Salt-Anion Electrode Examples

1. Ag|AgCl|Cl^⁻

• Reaction: AgCl(s) ⇌ Ag⁺(aq) + Cl^⁻(aq)
• Reduction: Ag⁺ + e^⁻ → Ag
• Overall: AgCl(s) + e^⁻ → Ag(s) + Cl^⁻(aq) (n=1)

2. Pt|Hg(l)|Hg₂Cl₂(s)|Cl^⁻(aq) (Calomel Electrode)

• Reaction: Hg₂Cl₂(s) ⇌ Hg₂²⁺(aq) + 2Cl^⁻(aq)
• Reduction: Hg₂²⁺ + 2e^⁻ → 2Hg
• Overall: Hg₂Cl₂(s) + 2e^⁻ → 2Hg(l) + 2Cl^⁻(aq) (n=2)

3. Other Examples

• Hg(l)|HgSO₄(s)|SO₄^2⁻(aq)
• Pb(s)|PbI₂(s)|I^⁻(aq)

28. Cell Potential and K_sp

• For the cell: Ag | AgCl | Cl⁻ (C₁ M) || Ag⁺ (C₂ M) | Ag
• Cell Reactions:

• E_cell calculation:

where [Ag⁺]_c = C₂ and [Cl⁻]_a = C₁.

29. Concentration Cell with sparingly soluble salt

• Reaction at electrodes:

• Standard cell potential:

E⁰_cell = (E⁰_Ag⁺/Ag + 0.0591 log (K_sp)_AgBr) − (E⁰_Ag⁺/Ag + 0.0591 log (K_sp)_AgCl)

30. Cell Potential Equation

• This relates to the concentration of silver ions:

31. Redox Electrodes

• Redox electrodes involve the oxidation or reduction of ions in solution using an inert electrode like Pt.

32. Quinhydrone Electrode

• An indicator electrode used to determine the pH of solutions.
• Quinhydrone is a 1:1 mixture of Quinone (Q) and Hydroquinone (QH₂).

33. Concentration Cell

• An electrochemical cell where both electrodes are identical, but the electrolyte concentrations are different.
• Produces electrical energy due to the concentration difference between two half-cells.

Condition for Spontaneity :

• For the cell to work (produce energy), the cell reaction must be spontaneous, i.e., E_cell > 0.
• Anode : Lower ion concentration or lower gas pressure.
• Cathode : Higher ion concentration or higher gas pressure.

34. Thermodynamics of Galvanic Cell

Entropy (ΔS) and Enthalpy (ΔH) :

• From G = H - TS and dG = VdP - SdT (at constant P, dP=0):
• S = -(dGdT)_P
• Using ΔG = -nFE_cell, we get:

(dE_celldT) is the temperature coefficient.

• Since ΔH = ΔG + TΔS:

35. Efficiency of a Galvanic Cell

• The efficiency (η) of a cell represents the ratio of useful electrical work to the total enthalpy change.

36. Reference Electrodes & Electrode Types

• Calomel Electrode: Categorized as a Metal - Insoluble Salt - Anion electrode.
• Reference Electrode Selection: A half-cell with a very low temperature coefficient (∂E∂T) is preferred as a reference electrode (e.g., a cell with (∂E∂T) = 0.1 × 10⁻⁴ is preferred over 0.2 × 10⁻⁴, 1 × 10⁻⁴, or 2 × 10⁻⁴).

37. Electrochemical vs Electrolytic Cells

• Electrochemical Cell (Discharging) : Uses spontaneous reactions to produce electrical energy.
  • Anode (Zn): Zn(s) → Zn²⁺(aq) + 2e⁻
  • Cathode (Cu): Cu²⁺(aq) + 2e⁻ → Cu(s)
  • Net Reaction: Cu²⁺(aq) + Zn(s) → Cu(s) + Zn²⁺(aq)
  • E°_cell = E°_Cu²⁺/Cu − E°_Zn²⁺/Zn = 0.34 V − (−0.76 V) = 1.10 V
• Electrolytic Cell (Charging) : Uses electrical energy (E_ext > E_cell) to drive a non-spontaneous reaction.
  • Anode (Cu): Cu(s) → Cu²⁺(aq) + 2e⁻
  • Cathode (Zn): Zn²⁺(aq) + 2e⁻ → Zn(s)
  • Net Reaction: Cu(s) + Zn²⁺(aq) → Cu²⁺(aq) + Zn(s)

38. Electrolysis Principles

• Electrolysis: Decomposition of an electrolyte by passing electricity through its aqueous solution or molten state.
• Electrolytic Cell: A device that converts electrical energy into chemical energy.
• At electrodes (general):
  • Cathode: A⁺ + e⁻ → A
  • Anode: B⁻ → B + e⁻

39. Comparison Table

40. Factors Affecting Electrolysis Products

• Electrolyte used.
• Electrode used:
  • Inert electrode: Does not participate in reaction (e.g., Pt, C, Pd).
  • Active electrode: Participates in reaction (e.g., Cu, Ag, Zn, Ni).
• Overvoltage (bubble voltage).

41. Oxidation Tendency at Anode

• Anions with a higher Standard Oxidation Potential (SOP) than water (−1.23 V) are preferentially oxidized.
• F⁻, SO₄²⁻, MnO₄⁻, NO₃⁻, CO₃²⁻ are never oxidized before water; oxygen gas is liberated in these cases.
• Oxidation Tendency Order: I⁻ > Br⁻ > Cl⁻ and CH₃COO⁻ > OH⁻ > F⁻ > SO₄²⁻ | CO₃²⁻ | MnO₄⁻ | NO₃⁻.

42. Electrolytic Reduction of Cations

• For s-block cations (Group 1 and Group 2): They cannot be reduced before H₂O because the SRP of H₂O (−0.83 V) is greater than their reduction potential. Consequently, H₂ gas is liberated at the cathode.
• Cations with an SRP greater than that of water can be reduced before H₂O:

43. Order of Deposition

Cations (Increasing order of deposition at cathode):

Anions (Increasing order of deposition at anode):

• Stronger reducing agents (lower SRP or higher SOP) are liberated first.

44. Electrolysis Examples

Molten NaCl :

• Species present: Na⁺, Cl⁻
• Cathode: Na⁺ + e⁻ → Na
• Anode: Cl⁻ → 12Cl₂ + e⁻

Molten PbBr₂ :

• Species present: Pb²⁺, Br⁻
• Cathode: Pb²⁺ + 2e⁻ → Pb(s)
• Anode: Br⁻ → 12Br₂ + e⁻

Aqueous NaBr :

• Species present: Na⁺, Br⁻, H⁺, OH⁻, H₂O
• Cathodic competition: H₂O + e⁻ → H₂ (preferred over Na⁺).
• Anodic competition: Br⁻ → 12Br₂ + e⁻ (E°_OP = −1.08 V) is preferred over 2H₂O → O₂ + 4H⁺ + 4e⁻ (E°_OP = −1.23 V).

45. Electrolysis of Aqueous NaCl Solution

• Ions present: Na⁺, Cl⁻, H⁺, OH⁻, H₂O.
• Cathode (C): H₂O → H₂
• Anode (A):
  • 2H₂O → O₂ + 4H⁺ + 4e⁻     E⁰_op = −1.23 V
  • Cl⁻ → 1/2 Cl₂ + e⁻     E⁰_op = −1.36 V
• Note: Based on E⁰_op values, water should be oxidised at the anode liberating O₂. However, due to low rates of O₂ liberation, extra voltage (overvoltage or bubble voltage) is applied. At this overvoltage, Cl₂ liberation dominates. Thus, Cl₂ gas is obtained at the anode.

46. Electrolysis of Dilute Aqueous NaCl

• Using inert electrodes.
• Cathode (C): H₂O + 2e⁻ → H₂ + 2OH⁻
• Anode (A): 2H₂O → O₂ + 4H⁺ + 4e⁻
• Since the solution is dilute, [Cl⁻] is very low, leading to O₂ liberation.

47. Electrolysis of Aqueous NaCl using Mercury Electrodes

• Cathode (C): Na⁺ + e⁻ → Na ; then Na + Hg → Na-Hg (amalgam).
• Anode (A): Cl⁻ → 1/2 Cl₂ + e⁻

48. Electrolysis of Concentrated H₂SO₄

• At cathode: 2H⁺ + 2e⁻ → H₂(g)
• At anode: 2SO₄^2⁻ → S₂O₈^2⁻ + 2e⁻

49. Electrolysis of Aqueous CH₃COONa

• Cathode (C): H₂O + e⁻ → H₂
• Anode (A):
  • CH₃COO⁻ → CH₃COO + e⁻
  • CH₃COO → CH₃ + CO₂
  • CH₃ + CH₃ → CH₃-CH₃ (ethane)

50. Faraday's First Law of Electrolysis

• The weight or amount of any substance produced during electrolysis is proportional to the amount of charge or electricity passed into the solution.

• For constant current:

where z = electrochemical equivalent ; z = E96500 ; E = equivalent weight.

51. Faraday's Second Law of Electrolysis

• When two or more electrolytic cells are connected in series, the number of gram equivalents deposited at the electrodes is the same.

52. Electrolytic Calculation Concept

• Using the weight deposited at the cathode (W₁), one can determine the moles, weight, or volume of gas liberated at the anode (W₂).

This relationship allows for stoichiometric calculations across different substances in series.

53. Electroplating

• For electroplating of Ag on an iron spoon:

• The relation between deposit weight and thickness:

where A = surface area, h = thickness, d = density, E = equivalent weight, i = current (A), t = time (s).

54. Electrolytic Conductance

• Resistance (R): The ratio of applied potential difference (V) to the current (i). R = V / i. (Unit: ohm, Ω).
• Conductance (G): Reciprocal of resistance; measures the ease of current flow.

• Units of G: ohm⁻¹, mho, or Siemens (S).

55. Resistivity (Specific Resistance, ρ)

• Resistance is directly proportional to length (l) and inversely proportional to area (A).

• If l = 1 cm and A = 1 cm², then ρ = R.
• Resistivity is the resistance offered by 1 cm³ of electrolytic solution.
• Unit of ρ: ohm cm (Ω cm).

56. Specific Conductance (Conductivity, κ)

• It is the reciprocal of specific resistance (ρ).
• Relationship with resistance (R) : R = ρ lA.

where G = conductance (1R), G* = cell constant (lA).

• Specific conductivity is defined as the conductance offered by 1 cm³ of electrolytic solution.
• Unit : ohm⁻¹ cm⁻¹ (or S cm⁻¹ or mho cm⁻¹).

57. Molar Conductivity (Λ_m, λ_m or μ)

• It is defined as the conductance of all the ions produced by one mole of electrolyte present in a given volume of solution.
• Relationship: Λ_m = κ × V, where V is the volume of solution containing 1 mole of electrolyte.
• If concentration of the solution is M (mol L⁻¹):

Unit: Ω⁻¹ cm² mol⁻¹ (or S cm² mol⁻¹). Note: 1 cm³ = 1 mL.

58. Units and Conversions

• If κ is in S cm⁻¹, then Λ_m = κ × 1000M S cm² mol⁻¹.
• If κ is in S m⁻¹, then Λ_m = κ1000 × M S m² mol⁻¹.
• 1 m³ = 1000 L.

59. Relation between Λ_m and Λ_eq

• Molar conductivity: Λ_m = κ × 1000M
• Equivalent conductivity: Λ_eq = κ × 1000N
• Since Normality (N) = Valency factor (V.F.) × Molarity (M):

• V.F. (n-factor) = total positive (or negative) charge per formula unit.
• Examples: NaCl = 1, CaCl₂ = 2, Al₂(SO₄)₃ = 6.

60. Cell Constant

• The cell constant (G*) is defined by the relationship between specific conductance (κ) and conductance (C) or resistance (R).

where κ = specific conductance, R = resistance, C = conductance (Ω⁻¹).

61. Factors Affecting Electrolytic Conductance

• Inter-ionic attraction: Higher attraction between solute ions results in lower conductance. (Attraction ↑ ⇒ Conductance ↓).
• Polarity of solvent: Solvents with high dielectric constants increase ionization and thus increase conductance.
• Viscosity: Increasing the viscosity of the medium decreases conductance.
• Temperature: Increasing temperature increases conductance because kinetic energy of ions increases, attraction forces decrease, and viscosity of the medium decreases.
• Hydrated size: Conductance decreases as the hydrated size of ions increases.

62. Effect of Dilution

• Conductance (G): Increases with dilution.
  • Strong electrolytes: Inter-ionic attraction decreases with dilution, causing conductance to increase.
  • Weak electrolytes: Degree of dissociation (α) increases with dilution, causing conductance to increase.
• Specific conductance (K): Decreases with dilution because the number of ions per mL of solution decreases.
• Equivalent and Molar conductance (Λ_eq, Λ_m): Increase with dilution because normality or molarity decreases.

63. Variation of Molar Conductivity (Λ_m) with Concentration

• Strong Electrolytes: Λ_m increases slowly with dilution as concentration decreases, following the Debye-Hückel-Onsager equation:
  Λ_m = Λ°_m − b√c
  where Λ°_m is limiting molar conductivity and b is a constant depending on the nature of the solvent and temperature. The graph of Λ_m vs √c is a straight line.
• Weak Electrolytes: Λ_m increases sharply with dilution because the degree of dissociation (α) increases. At infinite dilution, Λ_m approaches a maximum value (Λ°_m).

64. Important Concepts Summary

• Conductivity (κ) always decreases with a decrease in concentration for both strong and weak electrolytes (due to fewer ions per unit volume).
• Molar conductivity (Λ_m) increases with a decrease in concentration.
• Water types: Pure/distilled water has the lowest ionic conductance due to the lack of dissolved electrolytes compared to sea water or saline water.
• For organic acids (0.1 M), conductivity follows the order of their acid dissociation constants (K_a). Formic acid (A) > Benzoic acid (C) > Acetic acid (B).

65. Kohlrausch Law of Independent Migration of Ions

Examples of Limiting Conductivity :

66. Applications of Kohlrausch Law

• Determination of Λ°_m of weak electrolytes:

Explanation: λ°_m(CH₃COO^-) + λ°_m(H⁺) ← [λ°_m(H⁺) + λ°_m(Cl^-)] + [λ°_m(CH₃COO^-) + λ°_m(Na⁺)] − [λ°_m(Na⁺) + λ°_m(Cl^-)]

67. Degree of Dissociation (α)

• Λ_m = molar conductivity at concentration C.
• Λ°_m = limiting molar conductivity.
• At limiting molar conductance, degree of dissociation = 1.

68. Ionisation Constant of Weak Electrolytes

• For a weak acid (HA), the dissociation constant K_a is related to molar conductivity (Λ_m) and molar conductivity at infinite dilution (Λ°_m):

Where C = concentration.

69. Solubility of Sparingly Soluble Salt

• A saturated solution of a sparingly soluble salt can be considered as an infinitely diluted solution.
• Therefore, Λ_m ≈ Λ⁰_m.

where S is the solubility (mol L⁻¹) and K is the conductivity.

70. Limiting Molar Conductivity Trends

• Conductance of H⁺ and OH⁻ ions is greater than any other ion.
• Jump Mechanism: H⁺ ions exhibit anomalously high conductivity in water due to a "jump" or "Grotthuss" mechanism, where the proton effectively shifts between water molecules by breaking and forming O-H bonds along a chain of hydrogen-bonded molecules.
• Order of limiting molar conductivity for cations in water at 298 K: H⁺ > Ca²⁺ > Mg²⁺ > K⁺ > Na⁺.

71. Conductometric Titrations

1. Strong Acid vs Strong Base (HCl + NaOH)

• Reaction: HCl + NaOH → NaCl + H₂O
• The graph of Λ_m vs volume of NaOH starts high, decreases linearly to an equivalence point, then increases linearly due to excess NaOH.

2. Weak Acid vs Strong Base (CH₃COOH + NaOH)

• Reaction: CH₃COOH + NaOH → CH₃COONa + H₂O
• Initial conductivity is low due to weak dissociation of CH₃COOH.
• Graph shows Λ_m decreasing slightly (due to common ion effect) before rising linearly upon addition of excess NaOH.

3. Strong Acid vs Weak Base (HCl + NH₄OH)

• Reaction: HCl + NH₄OH → NH₄Cl + H₂O
• The conductivity decreases linearly as H⁺ ions are replaced by NH₄⁺ ions.
• After the equivalence point, the conductivity remains nearly constant because excess weak base (NH₄OH) is not dissociated.

4. Weak Acid vs Weak Base (CH₃COOH + NH₄OH)

• Reaction: CH₃COOH + NH₄OH → CH₃COONH₄ + H₂O
• Λ_m increases linearly from the start as salt is formed, then remains constant due to the common ion effect and suppression of dissociation of the weak base.

5. Mixture of Strong Acid and Weak Acid vs Strong Base (HCl + CH₃COOH + NaOH)

• Step 1: HCl + NaOH → NaCl + H₂O
• Step 2: CH₃COOH + NaOH → CH₃COONa + H₂O
• Graph decreases to the first equivalence point (HCl neutralized), then increases to the second equivalence point (CH₃COOH neutralized).

72. Types of Electrochemical Cells

• Primary Cells: Cells which cannot be recharged again and again.
• Secondary Cells: Cells which can be recharged again and again.

73. Dry Cell (Leclanché Cell)

• Construction :
• Anode : Zinc container (Zn).
• Cathode : Carbon rod (graphite) coated with MnO₂ and carbon black.
• Electrolyte : Paste of NH₄Cl and ZnCl₂.
• Separator : Porous paper or starch paste.

• Key Observations :
• Manganese is reduced from +4 to +3 oxidation state at the cathode.
• The NH₃ produced forms a complex with Zn²⁺ : Zn²⁺ + 4NH₃ → [Zn(NH₃)₄]²⁺.
• Overall reaction : Zn + 2MnO₂ + 2NH₄⁺ → Zn²⁺ + Mn₂O₃ + 2NH₃ + H₂O.
• Features : Potential ≈ 1.5 V, low cost, widely available, prone to leakage due to zinc corrosion.

74. Mercury Cell (Button Cell)

• Also known as the Ruben–Mallory cell.
• Construction :
• Anode : Zinc amalgam (Zn(Hg)).
• Cathode : Mercury(II) oxide (HgO).
• Electrolyte : Alkaline paste of KOH or NaOH.
• Container : Steel sealed design.

• Overall reaction : Zn(Hg) + HgO(s) → ZnO(s) + Hg(l).
• Features :
• Potential ≈ 1.35 V.
• The potential remains constant throughout its life because the overall reaction involves no ions in solution whose concentration changes.

75. Lead Storage Battery: Construction

• Anode: Spongy lead (Pb).
• Cathode: Lead dioxide (PbO₂).
• Electrolyte: Dilute H₂SO₄ (30-40%).
• Container: Hard rubber/plastic case.
• Plates: Multiple positive and negative plates arranged alternatively.
• Separators: Porous insulators between plates.

76. Lead Storage Battery: Discharging Reactions

77. Lead Storage Battery: Features

• Rechargeable battery (secondary cell).
• High voltage per cell ~ 2.0 V.
• Heavy weight due to lead plates.
• High current output.
• Maintenance: Battery needs recharging when the density of the H₂SO₄ solution falls below 1.2 g/mL.

78. Lead Storage Battery

• Components: Grid of lead packed with PbO₂ as cathode; electrolyte is ~38% H₂SO₄ solution.
• Charging Process:
  • PbSO₄ on anode is converted to Pb.
  • PbSO₄ on cathode is converted to PbO₂.
  • Reaction: 2PbSO₄(s) + 2H₂O(l) → Pb(s) + PbO₂(s) + 2H₂SO₄(aq).
• Note: PbO₂ acts as a strong oxidizing agent and is amphoteric in nature.

79. Nickel Cadmium (NiCd) Cell

[DIAGRAM: Cross-sectional illustration of a cylindrical NiCd battery. Features include a spring-loaded cap, vent hole, nickel-plated steel case, and insulating washer. Internal layers consist of a nickelic hydroxide electrode, a cadmium hydroxide electrode, and a porous separator.]

Construction :

• Anode: Cadmium (Cd).
• Cathode: Nickel oxyhydroxide (NiO(OH)).
• Electrolyte: Concentrated KOH.
• Separator: Nylon/porous polymer sheet.
• Container: Sealed metal cylinder (e.g., AA, AAA rechargeable).
• Usage: Emergency lights, toys, medical equipment.

80. NiCd Cell: Working Reactions

Discharging :

Charging :

• Cd(OH)₂ → Cd
• Ni(OH)₂ → NiO(OH)
• Since the electrolyte is not consumed, the cell has a long life.

81. H₂ – O₂ Fuel Cell

[DIAGRAM: Flow schematic showing H₂ and O₂ input; Anode, Electrolyte, Cathode, and Catalyst stack; and outputs of excess H₂, electric power, and water (H₂O).]

• An electrochemical device that converts the chemical energy of fuels (H₂ and O₂) directly into electrical energy.
• Eco-friendly and used in space missions (e.g., Apollo space ships).

82. Fuel Cell

• Converts the energy of combustion directly into electrical energy.
• Reaction: 2H₂(g) + O₂(g) → 2H₂O(l).

83. Corrosion of Iron

• An electrochemical process where iron acts as an electrochemical cell.

• Ferrous ions (Fe²⁺) are further oxidized by atmospheric oxygen to ferric ions (Fe³⁺).
• These form rust as hydrated ferric oxide (Fe₂O₃·xH₂O), with further production of hydrogen ions (H⁺).

84. Quick Formula Revision