Electrochemistry

Unit 2: Electrochemistry

Introduction

Electrochemistry deals with the relationship between chemical energy and electrical energy — specifically how chemical reactions produce electricity and how electricity drives chemical reactions. It underpins batteries, fuel cells, electroplating, and corrosion.

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Electrochemical Cells

• An electrochemical cell converts chemical energy into electrical energy (or vice versa). Two main types:
• Galvanic (Voltaic) cell: Spontaneous redox reaction generates electricity (e.g., Daniel cell)
• Electrolytic cell: Electrical energy drives a non-spontaneous chemical reaction (e.g., electrolysis of water)

• Daniel Cell: Zn rod in ZnSO4 solution (anode) and Cu rod in CuSO4 solution (cathode), connected by a salt bridge.
• At anode (oxidation): Zn → Zn2+ + 2e-
• At cathode (reduction): Cu2+ + 2e- → Cu
• Salt bridge maintains electrical neutrality by allowing ion migration.

Cell notation (cell representation):
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
(Single line = phase boundary; Double line = salt bridge)

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Standard Electrode Potential

The Standard Electrode Potential (E°) is the potential of a half-cell under standard conditions (1 M concentration, 298 K, 1 bar pressure) relative to the Standard Hydrogen Electrode (SHE), which is assigned E° = 0.00 V.

• Reduction potential: tendency of a species to be reduced. Higher value = stronger oxidising agent.
• EMF of cell = E°_cathode - E°_anode (both written as reduction potentials)
• Spontaneous reaction: E°_cell > 0 (corresponds to negative delta G)

Relationship: delta G° = -nFE°_cell (F = Faraday constant = 96500 C/mol)
delta G° = -RT ln K (equilibrium constant)
So: E°_cell = (RT/nF) ln K = (0.0591/n) log K at 298 K

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Nernst Equation

For a cell not at standard conditions:
E_cell = E°_cell - (RT/nF) ln Q
At 298 K: E_cell = E°_cell - (0.0591/n) log Q
where Q is the reaction quotient.

At equilibrium, E_cell = 0, and Q = K.

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Conductance and Resistivity

Resistance (R): Measured in ohm (Omega). R = rho x l/A (rho = resistivity)
Conductance (G): G = 1/R, unit = siemens (S)
Conductivity (kappa): kappa = 1/rho = G x (l/A). Unit: S m^-1
Cell constant: l/A

Molar conductivity (Lambda_m): Conductance of solution containing 1 mol of electrolyte, between electrodes 1 unit apart.
Lambda_m = kappa x 1000 / M (when kappa in S/cm and M in mol/L)

Kohlrausch's Law: At infinite dilution, the molar conductivity of an electrolyte is the sum of contributions of cation and anion:
Lambda_m° = v+ x lambda°_+ + v- x lambda°_-

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Electrolysis and Faraday's Laws

Electrolysis is the decomposition of an electrolyte by passing electric current.

Faraday's First Law: Mass of substance deposited/liberated is directly proportional to the quantity of electricity passed.
m = Z x Q = Z x I x t (Z = electrochemical equivalent)

Faraday's Second Law: When the same quantity of electricity is passed through different electrolytes, the masses deposited are proportional to their equivalent masses.
m proportional to E (equivalent weight)

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Batteries and Fuel Cells

• Primary batteries: Non-rechargeable (e.g., dry cell — Leclanche cell, mercury cell)
• Secondary batteries: Rechargeable (e.g., lead-acid battery, nickel-cadmium, lithium-ion)
• Lead-acid battery: E = 2V per cell, 6 cells = 12V. Anode: Pb; Cathode: PbO2; Electrolyte: H2SO4
• Fuel cells: Produce electricity from continuous supply of fuel (H2) and oxidant (O2). Used in space shuttles. Reactants not stored inside; they are fed continuously.

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Corrosion

Corrosion is the deterioration of a metal due to attack by environmental agents (water, oxygen, acids). Rusting of iron is an electrochemical process:
Anode (oxidation): Fe → Fe2+ + 2e-
Cathode (reduction): O2 + H2O + 4e- → 4OH-
Fe2+ + 2OH- → Fe(OH)2 → Fe2O3.xH2O (rust)
Prevention: Galvanisation (zinc coating), alloying, cathodic protection, painting.

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Worked Examples

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