Electrochemistry

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

Electrochemistry — Key Facts for CUET Galvanic cell: spontaneous redox reaction → electrical energy; $E_{\text{cell}} = E_{\text{cathode}} - E_{\text{anode}}$ Nernst equation: $E = E° - \frac{0.0591}{n}\log Q$ at 298 K; $E = E° - \frac{RT}{nF}\ln Q$ (general form) Faraday’s laws: mass deposited $m = \frac{Q \times M}{n \times F} = \frac{It \times M}{n \times F}$ Standard hydrogen electrode (SHE): $E° = 0$ V; $F = 96,485$ C/mol Conductivity: $\kappa = \frac{G}{A}$; molar conductivity $\Lambda_m = \frac{\kappa \times 1000}{c}$ (S cm² mol⁻¹) ⚡ Exam tip: In a galvanic cell, oxidation occurs at anode (negative electrode) and reduction at cathode (positive electrode); electrons flow from anode to cathode externally

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🟡 Standard — Regular Study (2d–2mo)

For students who want genuine understanding of electrochemical cells and conductivity.

Electrochemistry — CUET Chemistry Study Guide

Electrochemistry deals with interconversion of chemical and electrical energy. Two main types: galvanic (voltaic) cells that generate electricity from spontaneous chemical reactions, and electrolytic cells that use electricity to drive non-spontaneous reactions.

Galvanic Cell: Consists of two half-cells, each with an electrode immersed in an electrolyte. The Zn-Cu cell (Daniell cell): Zn electrode (anode, oxidation: Zn → Zn²⁺ + 2e⁻) and Cu electrode (cathode, reduction: Cu²⁺ + 2e⁻ → Cu). Electrons flow from Zn to Cu through external wire; salt bridge maintains electrical neutrality. Cell notation: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s).

Electrode Potential: The tendency of an electrode to lose or gain electrons in solution. Absolute potentials cannot be measured directly — only differences. Standard electrode potential $E°$ is measured under standard conditions (1 M, 1 atm, 298 K) relative to SHE. More positive $E°$ = stronger oxidising agent (better at accepting electrons).

Standard Reduction Potentials (selected):

• F₂/F⁻: +2.87 V (strongest oxidising agent)
• MnO₄⁻/Mn²⁺ (acidic): +1.51 V
• Cl₂/Cl⁻: +1.36 V
• O₂/H₂O: +1.23 V
• Cu²⁺/Cu: +0.34 V
• H⁺/H₂: 0.00 V (reference)
• Fe²⁺/Fe: −0.44 V
• Zn²⁺/Zn: −0.76 V
• Na⁺/Na: −2.71 V

Cell EMF and Free Energy: $\Delta G° = -nFE°{\text{cell}}$. For a spontaneous reaction, $E°{\text{cell}} > 0$ and $\Delta G° < 0$. The relationship $\Delta G° = -RT\ln K$ gives equilibrium constant: $K = 10^{nE°/0.0591}$ at 298 K.

Nernst Equation: At non-standard conditions, $E = E° - \frac{RT}{nF}\ln Q$. At 298 K: $E = E° - \frac{0.0591}{n}\log_{10} Q$. For concentration cells, $E° = 0$, so $E = -\frac{0.0591}{n}\log\frac{[\text{reducing agent}]}{[\text{oxidising agent}]}$.

Electrolytic Cell: Uses electrical energy to drive non-spontaneous redox reactions. In electrolysis of NaCl ( Downs cell): molten NaCl → Na (l) at cathode + Cl₂ (g) at anode. Overpotential (extra voltage needed) is often required for gas evolution.

Faraday’s Laws of Electrolysis:

1. Mass deposited $m = \frac{QM}{nF} = \frac{ItM}{nF}$
2. Same current through different electrolytes → mass deposited ∝ equivalent mass

Example: Current of 3 A passed through CuSO₄ solution for 30 minutes. Find mass of Cu deposited. $Q = It = 3 \times 30 \times 60 = 5400$ C. For Cu²⁺ + 2e⁻ → Cu, $n = 2$, $M = 63.5$ g/mol. $m = \frac{5400 \times 63.5}{2 \times 96485} = \frac{342900}{192970} = 1.78$ g.

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Electrochemistry — Complete CUET Chemistry Notes

Fuel Cells: Convert chemical energy of a fuel (H₂) directly to electricity. In PEM fuel cell: at anode, H₂ → 2H⁺ + 2e⁻; at cathode, ½O₂ + 2H⁺ + 2e⁻ → H₂O. Overall: H₂ + ½O₂ → H₂O. Efficiency ~60% (vs ~40% for internal combustion engines). Byproducts are water and heat. Used in spacecraft and increasingly in vehicles.

Corrosion: Electrochemical process of metal deterioration. For iron: anodic areas (Fe → Fe²⁺ + 2e⁻) and cathodic areas (O₂ + 2H₂O + 4e⁻ → 4OH⁻) form on the surface. Rust = Fe₂O₃·H₂O. Prevention: galvanising (Zn coating), cathodic protection (sacrificial anode), painting, or alloying (stainless steel).

Batteries:

• Primary (non-rechargeable): Leclanche cell (dry cell) — Zn/MnO₂, $E = 1.5$ V; alkaline cell — better performance
• Secondary (rechargeable): Lead-acid battery — Pb/PbO₂ in H₂SO₄, each cell = 2 V; 6 cells in series = 12 V. Nicke -cadmium (NiCd), Nickel-Metal Hydride (NiMH), Lithium-ion (Li-ion) — used in phones and laptops

Conductivity of Electrolytes: Molar conductivity $\Lambda_m = \frac{\kappa \times 1000}{c}$ where $c$ is molarity in mol/L. For strong electrolytes, $\Lambda_m$ increases slowly with dilution (decreased interionic attractions). For weak electrolytes, $\Lambda_m$ increases sharply near infinite dilution (degree of dissociation increases). Kohlrausch’s law: $\Lambda_m° = \lambda°{\text{cation}} + \lambda°{\text{anion}}$.

Battery Capacity: Expressed in ampere-hours (Ah). A 100 Ah battery can deliver 5 A for 20 hours. For Li-ion, capacity depends on current — higher current → lower usable capacity (P = IV, heating losses).

Electroplating: Electrolytic deposition of metal coating. To plate silver on an object: object is cathode, silver anode, silver nitrate electrolyte. Faraday’s law determines thickness: $t = \frac{m}{\rho A}$. Factors affecting quality: current density, temperature, electrolyte concentration.

Concentration Cells: Both electrodes are the same material but in different concentrations. Example: Cu(s) | Cu²⁺(0.01 M) || Cu²⁺(1 M) | Cu(s). $E° = 0$, so $E = -\frac{0.0591}{2}\log\frac{0.01}{1} = +0.0295$ V. The cell with more concentrated solution becomes the cathode (standard convention).

CUET Exam Patterns (2022–2024):

• Nernst equation calculations are most frequent (2–3 marks)
• Cell notation and identification of anode/cathode is very common
• Faraday’s law problems (mass deposited, current, time relationships) appear every year
• Standard reduction potentials and reactivity series questions are common
• Common mistakes: getting the sign wrong in Nernst equation; confusing galvanic and electrolytic cell direction

⚡ Key insight: In Nernst equation, $Q$ is the reaction quotient in the same form as the balanced cell reaction. For the cell Zn | Zn²⁺(1 M) || Cu²⁺(1 M) | Cu, the reaction is Zn + Cu²⁺ → Zn²⁺ + Cu, so $Q = \frac{[Zn^{2+}]}{[Cu^{2+}]}$. Remember to use $n$ = total electrons transferred (2 for this cell). For $\log_{10}$, use 0.0591; for $\ln$, use $\frac{RT}{F} = \frac{8.314 \times 298}{96485} = 0.0257$ V.

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