Electrochemistry

Introduction

• Electrochemistry involves the conversion of chemical energy into electrical energy and vice versa.
• Important for both theoretical and practical purposes in producing chemicals like sodium hydroxide, chlorine, etc.
• Applications in batteries and fuel cells for energy conversion.
• Less polluting and energy-efficient electrochemical reactions.
• Electrochemical processes are involved in biological sensory signal transmission and communication between cells.

Objectives

• Understand electrochemical cells: differentiate galvanic and electrolytic cells.
• Use the Nernst equation to calculate the emf of galvanic cells.
• Relate standard potential, Gibbs energy, and equilibrium constant.
• Define resistivity, conductivity, and molar conductivity for ionic solutions.
• Differentiate between ionic and electronic conductivity.
• Measure and calculate conductivity and molar conductivity.
• Kohlrausch law and its applications.
• Quantitative aspects of electrolysis.
• Construction of primary and secondary batteries, and fuel cells.
• Corrosion as an electrochemical process.

Electrochemical Cells

Daniell Cell

• Converts chemical energy from redox reactions into electrical energy.
• Reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s).
• Functioning stops when the external voltage equals cell potential (1.1 V).

Galvanic Cells

• Converts spontaneous redox reaction energy into electrical work.
• Consists of two half-cells: oxidation at anode (negative), reduction at cathode (positive).
• Electrode potential develops at the electrode-electrolyte interface.
• Standard electrode potential is measured against the standard hydrogen electrode.

Measurement of Electrode Potential

• Standard hydrogen electrode (SHE) used as a reference (0 V).
• Electrode potential can be calculated using measured emf and standard conditions.

Nernst Equation

• Accounts for concentration effects on electrode potential.
• Given by: E = E° - (RT/nF) ln(Q), where Q is the reaction quotient.
• Useful in calculating equilibrium constants from emf measurements.

Conductivity of Electrolytic Solutions

• Conductivity (κ) is the inverse of resistivity.
• Dependence on concentration, solvent nature, and temperature.
• Molar conductivity (Λm) defined as Λm = κ/c.
• Conductivity decreases with dilution, but molar conductivity increases.

Kohlrausch Law

• Molar conductivity at infinite dilution is the sum of individual ionic contributions.
• Useful for determining dissociation constants and limiting conductivities.

Electrolysis

• Electrical energy drives non-spontaneous reactions.
• Faraday’s Laws: Amount of substance liberated is proportional to the current and molar mass/equivalent weight.

Batteries

Primary Batteries

• Single-use; reaction occurs once (e.g., dry cell, mercury cell).

Secondary Batteries

• Rechargeable (e.g., lead-acid battery, nickel-cadmium cell).
• Chemical reactions are reversible.

Fuel Cells

• Convert chemical energy directly into electrical energy using continuous reactant supply.
• Example: Hydrogen-oxygen fuel cell with high efficiency and low pollution.

Corrosion

• Electrochemical oxidation of metals in presence of water and oxygen.
• Preventative methods: coatings, sacrificial anodes.

The Hydrogen Economy

• Aims to replace carbon-based fuels with hydrogen for a cleaner energy source.

Summary

• Electrochemical cells convert chemical energy to electrical energy and vice versa.
• Galvanic cells use spontaneous reactions; electrolytic cells drive non-spontaneous reactions.
• Important concepts: electrode potential, Nernst equation, conductivity, and applications in technology and industry.