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Electrochemistry – Class 12 Chemistry Chapter 2 Notes & NCERT PDF

Class 12ChemistryChapter 2NCERT book: Chemistry Part I - Class 12

Notes and a simple summary of Chapter 2 of the NCERT Class 12 Chemistry book, with the official chapter PDF, flashcards, an MCQ quiz and an AI tutor for your doubts.

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Chapter 2: Electrochemistry

Introduction

Electrochemistry is the study of the production of electricity from the energy released in spontaneous chemical reactions, and the use of electrical energy to bring about non-spontaneous chemical changes. It explains how batteries, fuel cells and electroplating work, how sodium, aluminium and chlorine are manufactured, and why iron rusts. The unit connects cell potential with Gibbs energy and equilibrium, and studies how solutions conduct electricity.

Key Concepts

2.1 Electrochemical Cells

  • A galvanic (voltaic) cell converts the chemical energy of a spontaneous redox reaction into electrical energy. The Daniell cell (Zn | Zn²⁺ ‖ Cu²⁺ | Cu) has a potential of 1.1 V when both ion concentrations are 1 mol dm⁻³:
    • Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
  • If an opposing external potential is applied: when E_ext < 1.1 V, current flows from Cu to Zn (electrons from Zn to Cu) and the cell works as galvanic; at E_ext = 1.1 V no current flows; when E_ext > 1.1 V the reaction reverses and the cell works as an electrolytic cell.
  • Oxidation occurs at the anode and reduction at the cathode. In a galvanic cell the anode is negative and the cathode positive. Two half-cells are connected internally by a salt bridge.

    • 2.2 Galvanic Cells and Electrode Potential

  • The potential difference between electrode and electrolyte is the electrode potential; at unit concentration it is the standard electrode potential (E°). By IUPAC convention, standard reduction potentials are used.
  • Cell emf: E_cell = E_right − E_left = E_cathode − E_anode. The cell is written with the anode on the left, e.g. Cu(s) | Cu²⁺(aq) ‖ Ag⁺(aq) | Ag(s).
  • The standard hydrogen electrode (SHE) — platinum coated with platinum black, dipped in acid with H⁺ activity 1 and pure H₂ gas bubbled at 1 bar — is assigned E° = 0.00 V at all temperatures.
  • A negative E° means the redox couple is a stronger reducing agent than H⁺/H₂; a positive E° means a weaker reducing agent. Fluorine (E° = 2.87 V) is the strongest oxidising agent; lithium (most negative E°) is the most powerful reducing agent.

    • 2.3 Nernst Equation

    For Mⁿ⁺(aq) + ne⁻ → M(s):

    E = E° − (RT/nF) ln(1/[Mⁿ⁺]) = E° − (0.059/n) log(1/[Mⁿ⁺]) at 298 K

    For a cell reaction, E_cell = E°_cell − (0.059/n) log Q. At equilibrium E_cell = 0, giving:

    E°_cell = (0.059/n) log K_c (at 298 K)

    Gibbs energy: ΔrG = −nFE_cell and ΔrG° = −nFE°_cell. A positive E°_cell means negative ΔG° — a spontaneous reaction. F (Faraday constant) = 96487 C mol⁻¹.

    2.4 Conductance of Electrolytic Solutions

  • Resistance R = ρ(l/A); ρ = resistivity (Ω m). Conductance G = 1/R, SI unit siemens (S). Conductivity κ = 1/ρ (S m⁻¹).
  • Cell constant G* = l/A = R × κ; found using KCl solutions of known conductivity.
  • Molar conductivity: Λm = κ/c; in practical units Λm (S cm² mol⁻¹) = κ (S cm⁻¹) × 1000/molarity.
  • Electronic (metallic) conductance depends on the nature and structure of the metal and decreases with rise in temperature; electrolytic (ionic) conductance depends on the nature of the electrolyte, solvent, viscosity, concentration and temperature, and increases with temperature.
  • Variation with concentration: conductivity decreases on dilution (fewer ions per unit volume), while molar conductivity increases on dilution. Its limiting value at zero concentration is Λ°m.
  • Strong electrolytes: Λm = Λ°m − A c^½ (Debye–Hückel–Onsager) — a straight line on a Λm vs √c plot, extrapolated to get Λ°m.
  • Weak electrolytes (e.g. acetic acid) have low Λm at high concentration that rises steeply on dilution because the degree of dissociation increases; Λ°m cannot be found by extrapolation.

    • Kohlrausch law of independent migration of ions: the limiting molar conductivity of an electrolyte is the sum of the individual contributions of its anion and cation:

    Λ°m = ν₊ λ°₊ + ν₋ λ°₋

    Uses: finding Λ°m of weak electrolytes, degree of dissociation α = Λm/Λ°m, and dissociation constant K = cα²/(1 − α).

    2.5 Electrolytic Cells and Electrolysis

    Faraday's laws (1833–34):

  • The amount of chemical reaction at any electrode is proportional to the quantity of electricity passed.
  • The amounts of different substances liberated by the same quantity of electricity are proportional to their chemical equivalent weights (atomic mass ÷ electrons needed to reduce the cation).

    • Q = I × t (coulombs). Charge on one mole of electrons ≈ 96487 C mol⁻¹ = 1 F. Reducing 1 mol Cu²⁺ needs 2 F; 1 mol Al³⁺ needs 3 F.

    Products of electrolysis depend on the material being electrolysed and the electrodes (inert Pt/gold vs reactive), and on standard potentials plus overpotential:

  • Molten NaCl → Na metal and Cl₂ gas.
  • Aqueous NaCl → NaOH, Cl₂ and H₂. At the cathode H₂O is reduced (E° of Na⁺/Na = −2.71 V is far lower). At the anode Cl⁻ is oxidised (E° 1.36 V) rather than water (1.23 V) because of the overpotential of oxygen.
  • Dilute H₂SO₄ gives O₂ at the anode; concentrated H₂SO₄ gives S₂O₈²⁻.

    • 2.6 Batteries

  • Primary batteries (reaction occurs only once; cannot be recharged): the dry cell (Leclanché cell) — zinc container as anode, carbon (graphite) rod cathode surrounded by MnO₂ and carbon, paste of NH₄Cl and ZnCl₂; potential nearly 1.5 V. The mercury cell — zinc–mercury amalgam anode, HgO + carbon paste cathode, KOH/ZnO electrolyte; potential about 1.35 V, constant throughout life because no ion concentration changes. Used in hearing aids and watches.
  • Secondary batteries (rechargeable): the lead storage battery — lead anode, grid of lead packed with PbO₂ as cathode, 38% H₂SO₄ electrolyte:
    • Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)

    On charging the reaction reverses. The nickel–cadmium cell has a longer life than the lead storage cell but is more expensive.

    2.7 Fuel Cells

    Galvanic cells that convert the energy of combustion of fuels such as hydrogen, methane or methanol directly into electrical energy. In the H₂–O₂ fuel cell (used in the Apollo space programme, the water produced was used as drinking water for astronauts):

    Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻; Anode: 2H₂ + 4OH⁻ → 4H₂O + 4e⁻; Overall: 2H₂ + O₂ → 2H₂O

    Fuel cells are pollution-free and have an efficiency of about 70% compared with about 40% for thermal power plants.

    2.8 Corrosion

    Corrosion slowly coats metals with oxides or other salts — rusting of iron, tarnishing of silver, green coating on copper and bronze. It is essentially an electrochemical phenomenon: at a spot on iron, Fe is oxidised (anode) — 2Fe → 2Fe²⁺ + 4e⁻ — and the electrons reduce oxygen at another spot in the presence of H⁺ (cathode). Fe²⁺ is further oxidised to Fe³⁺, which forms rust, Fe₂O₃·xH₂O. Prevention: barrier protection (paint, or chemicals such as bisphenol), covering with other metals (Sn, Zn, etc.), and sacrificial protection — providing an electrode of a more active metal (Mg, Zn, etc.) that corrodes itself and saves the iron.

    Worked Examples

  • Nernst equation (Ex. 2.1): Mg | Mg²⁺(0.130 M) ‖ Ag⁺(1.0 × 10⁻⁴ M) | Ag, E°cell = 3.17 V. E = 3.17 − (0.059/2) log[0.130/(10⁻⁴)²] = 2.96 V.
  • Gibbs energy (Ex. 2.3): Daniell cell, E° = 1.1 V, n = 2. ΔrG° = −2 × 96487 × 1.1 = −212,271 J mol⁻¹ ≈ −212.27 kJ mol⁻¹ (the textbook prints −21227 J mol⁻¹, a digit slip).
  • Molar conductivity (Ex. 2.4): 0.20 M KCl, κ = 0.0248 S cm⁻¹ → Λm = 0.0248 × 1000/0.20 = 124 S cm² mol⁻¹.
  • Resistivity and conductivity (Ex. 2.5): a 0.05 M NaOH column, diameter 1 cm (A = 0.785 cm²), length 50 cm, R = 5.55 × 10³ Ω. ρ = RA/l ≈ 87.1 Ω cm, κ = 1/ρ ≈ 0.01148 S cm⁻¹, Λm = 0.01148 × 1000/0.05 ≈ 229.6 S cm² mol⁻¹.
  • Kohlrausch (Ex. 2.7): Λ°m(CaCl₂) = 119.0 + 2(76.3) = 271.6 S cm² mol⁻¹; Λ°m(MgSO₄) = 106.0 + 160.0 = 266 S cm² mol⁻¹.
  • Weak acid (Ex. 2.8–2.9): Λ°m(HAc) = 425.9 + 91.0 − 126.4 = 390.5 S cm² mol⁻¹. For 0.001028 M acetic acid with κ = 4.95 × 10⁻⁵ S cm⁻¹: Λm = 48.15, α = 48.15/390.5 = 0.1233, Ka = cα²/(1 − α) = 1.78 × 10⁻⁵ mol L⁻¹.
  • Faraday (Ex. 2.10): 1.5 A for 10 min through CuSO₄: Q = 900 C; Cu deposited = 63 × 900/(2 × 96487) = 0.2938 g.

    • Formula Sheet

  • E_cell = E_cathode − E_anode (reduction potentials)
  • E = E° − (0.059/n) log Q at 298 K; E°cell = (0.059/n) log Kc
  • ΔrG° = −nFE°cell; ΔrG° = −RT ln K
  • G* = l/A = Rκ; Λm = κ × 1000/c
  • Λm = Λ°m − A√c (strong electrolyte); Λ°m = ν₊λ°₊ + ν₋λ°₋
  • α = Λm/Λ°m; Ka = cα²/(1 − α)
  • Q = It; mass deposited = (M × Q)/(n × F)

    • Board Exam Focus

  • Nernst numericals — write the cell reaction first, find n, then substitute. Also E°cell → K and ΔG°.
  • Conductivity numericals — cell constant, κ, Λm, α and Ka for weak acids using Kohlrausch law.
  • Faraday's law numericals — mass deposited, charge needed (e.g. 3 F for 1 mol Al).
  • Theory — variation of κ and Λm with dilution (with graph for strong and weak electrolytes), products of electrolysis of aqueous NaCl, lead storage battery reactions, H₂–O₂ fuel cell, mechanism of rusting.
  • Answer outline — rusting: anode reaction at iron spot, cathode reduction of O₂ in presence of H⁺, further oxidation to Fe³⁺, formation of Fe₂O₃·xH₂O, and two methods of prevention.

    • Key Terms

  • Galvanic cell (गैल्वेनी सेल), Electrolytic cell (वैद्युत अपघटनी सेल)
  • Anode (ऐनोड), Cathode (कैथोड)
  • Salt bridge (लवण सेतु)
  • Standard electrode potential (मानक इलेक्ट्रोड विभव)
  • Nernst equation (नेर्न्स्ट समीकरण)
  • Conductivity (चालकता), Molar conductivity (मोलर चालकता)
  • Cell constant (सेल स्थिरांक)
  • Kohlrausch law (कोलराउश नियम)
  • Electrolysis (वैद्युत अपघटन)
  • Fuel cell (ईंधन सेल)
  • Corrosion (संक्षारण), Rusting (जंग लगना)

    • Common Mistakes

  • Subtracting oxidation potentials instead of using reduction potentials for both electrodes.
  • Using the wrong n in the Nernst equation — n is the electrons transferred in the balanced cell reaction.
  • Thinking conductivity and molar conductivity both increase on dilution — only Λm increases; κ decreases.
  • Trying to extrapolate Λm of a weak electrolyte to get Λ°m — use Kohlrausch law instead.
  • Predicting O₂ at the anode for aqueous NaCl — Cl₂ is formed due to the overpotential of oxygen.
  • Forgetting to convert minutes to seconds in Q = It.

    • 💡 Key Learning Points

    • ✓Describe galvanic and electrolytic cells, the Daniell cell and the SHE.
    • ✓Apply the Nernst equation and relate E°cell to ΔG° and the equilibrium constant.
    • ✓Define conductivity, molar conductivity and cell constant and explain their variation with concentration.
    • ✓Use Kohlrausch law to find Λ°m, degree of dissociation and dissociation constant of weak electrolytes.
    • ✓Apply Faraday's laws, predict electrolysis products, and explain batteries, fuel cells and corrosion.

    👨‍🏫 Teaching Tips

    • →Build a lemon or potato cell with Zn and Cu strips to introduce galvanic cells.
    • →Draw the Daniell cell with arrows for electron flow, ion flow and salt bridge; ask students to label signs.
    • →Use the electrochemical series table to let students predict which metal displaces another.
    • →Teach Nernst numericals with a fixed four-step routine: reaction, n, Q, substitute.
    • →Plot Λm vs √c for KCl and acetic acid side by side to contrast strong and weak electrolytes.
    • →Show an old car battery or dry cell cut-away image and link each part to its reaction.
    • →Leave iron nails in water, salt water and oil for a week to observe rusting conditions.
    • →Connect fuel cells and batteries to electric vehicles and clean-energy discussions.

    📋 Assessment Questions

    1. Calculate emf of a cell using the Nernst equation at given concentrations.
    2. Find ΔrG° and Kc from E°cell.
    3. Calculate κ, Λm, α and Ka from conductance data.
    4. Compute mass deposited or charge required using Faraday's laws.
    5. Write electrode reactions of the lead storage battery, fuel cell and rusting.

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    Keep learning

    All chapters in Chemistry Part I - Class 12

    1. 1. Solutions
    2. 2. Electrochemistry
    3. 3. Chemical Kinetics
    4. 4. The d- and f-Block Elements
    5. 5. Coordination Compounds