Electrochemistry Key Concepts

Overview

This lecture covers the key principles and concepts of Electrochemistry, including energy conversion, electrochemical and electrolytic cells, electrode potentials, conductivity, electrolysis, and laws of electrolysis, emphasizing board exam-relevant topics.

Introduction to Electrochemistry

• Electrochemistry studies conversion between chemical energy and electrical energy.
• Devices converting chemical energy to electricity are called electrochemical/voltaic/galvanic cells.
• Devices converting electrical energy to chemical change are called electrolytic cells.

Electrochemical Cell Construction & Operation

• Electrochemical cell uses two half-cells (e.g., zinc and copper rods in respective salt solutions) linked by a salt bridge.
• Salt bridge (inverted U-tube with KCl/KNO₃ in agar/gelatin) completes the circuit and maintains neutrality.
• Zinc electrode undergoes oxidation (Zn → Zn²⁺ + 2e⁻), and copper electrode undergoes reduction (Cu²⁺ + 2e⁻ → Cu).
• Electrons flow from zinc to copper; conventional current flows from copper to zinc.
• In electrochemical cells: anode is negative (oxidation), cathode is positive (reduction).

Electrode Potential

• Electrode potential arises from charge separation at interface of metal electrode and solution.
• Standard electrode potential (E⁰) is measured with 1M concentration.

Cell Potential Calculations

• Cell potential (E⁰cell) = E⁰(right/reduction) − E⁰(left/oxidation).
• If both sides are reduction or oxidation, add the values.

Standard Hydrogen Electrode (SHE)

• SHE is the reference electrode (E⁰ = 0 V) for measuring other half-cell potentials.
• Conditions: 1 bar H₂ gas, 1M HCl, platinum electrode; can act as anode or cathode.
• Disadvantages: Hard to maintain 1M HCl and 1 bar H₂; easily poisoned by impurities.

Nernst Equation

• Used when concentrations are not standard (not 1M).
• Nernst equation: Ecell = E⁰cell − 0.0591/n × log Qc (at 25°C/298K).
• Qc is the reaction quotient (product ion concentrations over reactant ion concentrations).
• n = number of moles of electrons exchanged (from balanced reaction).

Equilibrium Constant from Nernst Equation

• At equilibrium, Ecell = 0; thus, E⁰cell = 0.0591/n × log Kc (Kc = equilibrium constant).

Cell Representation

• Cell is represented as: Zn/Zn²⁺ || Cu²⁺/Cu (|| indicates salt bridge).

Gibbs Free Energy and Cell Potential

• ΔG = −nFEcell; ΔG° = −nFE⁰cell.
• F (Faraday constant) = 96500 C/mol e⁻.
• Use Nernst equation value for Ecell; for standard, use E⁰cell.

Electrolytes: Types and Properties

• Electrolytes dissociate into ions in solution.
• Strong electrolytes (e.g., NaCl, KCl) completely dissociate; weak electrolytes partially dissociate.
• Reciprocal of resistance is conductance (G = 1/R); reciprocal of resistivity is conductivity (κ = 1/ρ).

Relation Between Conductance and Conductivity

• Cell constant (G*) = L/A (length/area between electrodes).
• κ = conductance × cell constant.

Conductivity and Molar Conductivity

• Conductivity (κ) = conductance of 1 unit cube of solution.
• On dilution, conductivity decreases because number of ions per volume decreases.
• Molar conductivity (λm) = κ/molarity; increases with dilution due to increased volume.

Variation of Molar Conductivity with Dilution

• For strong electrolytes, λm increases slightly due to increased ion movement.
• For weak electrolytes, λm increases sharply with dilution due to increased dissociation/ion production.

Kohlrausch’s Law

• At zero concentration/infinite dilution, molar conductivity is sum of conductivities of cation and anion.
• λm⁰(electrolyte) = λm⁰(cation) + λm⁰(anion).

Degree of Dissociation (α)

• α = λm/λm⁰; represents fraction of electrolyte dissociated.

Electrolysis and Electrolytic Cells

• Electrolysis uses electrical energy to drive chemical changes (e.g., decomposing NaCl).
• In electrolytic cells, anode is positive (oxidation), cathode is negative (reduction).
• Products of electrolysis depend on the electrolyte state (molten vs. aqueous).

Faraday’s Laws of Electrolysis

• First Law: Mass of substance produced is proportional to charge passed (m = ZIt).
• Z = electrochemical equivalent = atomic weight/(n × 96500).
• Second Law: If same charge passes through different electrolytes, mass produced is proportional to equivalent weight.

Corrosion and Prevention

• Rusting (corrosion) of metals occurs due to reaction with air and water.
• Prevention methods include painting, greasing, and galvanization (zinc coating as sacrificial metal).

Key Terms & Definitions

• Electrochemical Cell — Device converting chemical energy to electrical energy.
• Electrolytic Cell — Device using electrical energy for chemical change.
• Salt Bridge — U-tube with electrolyte maintaining circuit and neutrality.
• Electrode Potential — Potential developed due to charge separation at electrode-solution interface.
• Standard Hydrogen Electrode (SHE) — Reference electrode with E⁰ = 0 V.
• Nernst Equation — Formula for cell potential under non-standard conditions.
• Conductance (G) — Inverse of resistance.
• Conductivity (κ) — Inverse of resistivity; ability to conduct electric current.
• Molar Conductivity (λm) — Conductivity per unit molarity.
• Degree of Dissociation (α) — Fraction of electrolyte molecules ionized.
• Faraday Constant (F) — 96500 C/mol e⁻, charge per mole of electrons.
• Equivalent Weight (E) — Atomic weight/n, used in Faraday’s laws.

Action Items / Next Steps

• Review and memorize electrode reactions and products from electrolysis in NCERT.
• Practice cell notation, cell potential calculations, and use of Nernst equation.
• Study types and prevention of corrosion.
• Complete homework and assigned textbook readings as directed by instructor.