Unit 2: Electrochemistry

The interplay between electrical energy and chemical reactions — from batteries to corrosion

2.1 Electrochemical Cells

An electrochemical cell converts chemical energy into electrical energy (galvanic) or vice versa (electrolytic). A spontaneous redox reaction drives a galvanic cell, producing an electric current.

        Daniel Cell: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

        Anode (oxidation): Zn(s) → Zn²⁺(aq) + 2e⁻

        Cathode (reduction): Cu²⁺(aq) + 2e⁻ → Cu(s)

        Overall: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)   E°cell = +1.1 V

Convention: Anode is written on the left (negative electrode in galvanic cell), cathode on the right. Electrons flow from anode to cathode through the external circuit. The salt bridge maintains electrical neutrality by allowing ion migration.

Electrode Potentials

The tendency of a electrode to lose or gain electrons is measured as electrode potential. By convention, the standard hydrogen electrode (SHE) is assigned E° = 0 V.

Half-Reaction	E° (V)
Li⁺ + e⁻ → Li	−3.04
Zn²⁺ + 2e⁻ → Zn	−0.76
2H⁺ + 2e⁻ → H₂	0.00 (SHE)
Cu²⁺ + 2e⁻ → Cu	+0.34
Ag⁺ + e⁻ → Ag	+0.80
F₂ + 2e⁻ → 2F⁻	+2.87

        E°cell = E°cathode − E°anode

        More negative E° → stronger reducing agent (tendency to lose electrons)

        More positive E° → stronger oxidizing agent (tendency to gain electrons)

2.2 Nernst Equation

The Nernst equation relates electrode potential to the concentration of reacting species under non-standard conditions.

        Nernst Equation (at 298 K):

        E = E° − (0.0592 / n) log Q

        For the Daniel Cell:

        E = E° − (0.0592/2) log([Zn²⁺] / [Cu²⁺])

Gibbs Energy and Equilibrium

Gibbs Energy

ΔG = −nFE
ΔG° = −nFE°
where F = 96487 C mol⁻¹

At Equilibrium

E = 0, Q = K
E° = (0.0592/n) log K
K = 10^(nE°/0.0592)

If E°cell is positive, ΔG° is negative → reaction is spontaneous. If E°cell is negative, ΔG° is positive → reaction is non-spontaneous.

2.3 Conductance of Electrolytic Solutions

Unlike metallic conduction (electrons), electrolytic solutions conduct via movement of ions. The conductance depends on the nature of the electrolyte, ion size, solvent viscosity, concentration, and temperature.

Conductivity (κ)

G = 1/R = κ · A/l
SI unit: S m⁻¹
1 S cm⁻¹ = 100 S m⁻¹

Molar Conductivity (Λm)

Λm = κ / c
Units: S m² mol⁻¹
or S cm² mol⁻¹

Variation with Concentration

κ decreases on dilution
Λm increases on dilution
Λm = Λ°m − A√c (strong electrolytes)

Measurement of Conductivity

Conductivity is measured using a Wheatstone bridge with an AC source (to avoid electrolysis) and a conductivity cell (platinized Pt electrodes).

        Cell constant: G* = l / A (unit: m⁻¹ or cm⁻¹)

        Conductivity: κ = G* / R

        Known G* is determined using KCl solutions of known conductivity.

2.4 Kohlrausch Law

Kohlrausch's law of independent migration of ions states that the limiting molar conductivity of an electrolyte is the sum of the individual contributions of the anion and cation.

Kohlrausch Law of Independent Migration of Ions

        Λ°m = n₊ λ°₊ + n₋ λ°₋

        Applications:
        Calculate Λ°m for any electrolyte from ionic values
Determine Λ°m for weak electrolytes (cannot extrapolate directly)
Calculate dissociation constant: α = Λm / Λ°m

Dissociation Constant of Weak Electrolytes

For a weak electrolyte like acetic acid:

Kₐ = (c · α²) / (1 − α) where α = Λm / Λ°m

At infinite dilution, even weak electrolytes dissociate completely (α → 1), but conductivity becomes too low to measure. Kohlrausch law provides the alternative route to Λ°m.

2.5 Applications of Electrochemistry

Batteries

Battery	Anode	Cathode	EMF
Dry cell (Leclanché)	Zn	MnO₂ + NH₄⁺	~1.5 V
Lead storage	Pb	PbO₂	~2.0 V
Mercury cell	Zn(Hg)	HgO	~1.35 V

Fuel Cells

In a H₂-O₂ fuel cell, hydrogen and oxygen react to produce electricity. The electrolyte is aqueous KOH.

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

Fuel cells are highly efficient (70–80%) and are used in space shuttles and submarines.

Corrosion

Corrosion is an electrochemical process where a metal is oxidized by its environment. Iron rusting involves:

Anode: 2Fe → 2Fe²⁺ + 4e⁻
Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻
Overall: 2Fe + O₂ + 2H₂O → 2Fe(OH)₂ → Fe₂O₃·xH₂O

        Prevention of Corrosion:
        Galvanization: Coating iron with zinc
Painting: Barrier against moisture and air
Sacrificial anodes: Attaching more reactive metal (Mg/Zn) which corrodes instead
Alloying: Stainless steel (Cr + Ni added to Fe)

    

Summary

Electrochemistry studies the interconversion of chemical and electrical energy. Galvanic cells produce electricity from spontaneous redox reactions; electrolytic cells use electricity to drive non-spontaneous reactions. The Nernst equation relates electrode potential to concentration. Conductivity and molar conductivity measure how well electrolytic solutions carry current. Kohlrausch's law allows calculation of limiting molar conductivity from individual ionic contributions. Applications include batteries, fuel cells, and corrosion prevention.