Electrochemistry

The study of chemical processes that involve the transfer of electrons, encompassing both spontaneous redox reactions (galvanic cells) and non-spontaneous reactions driven by electrical energy (electrolysis).

Redox Reactions

Oxidation and reduction always occur together in redox reactions:

Oxidation: Loss of electrons, increase in oxidation number
Reduction: Gain of electrons, decrease in oxidation number
Oxidising agent: Species that is reduced (gains electrons)
Reducing agent: Species that is oxidised (loses electrons)

Electrochemical Cells

Galvanic (Voltaic) Cells

Spontaneous cells that produce electrical energy from chemical reactions:

Anode: Site of oxidation (negative electrode, electrons flow out)
Cathode: Site of reduction (positive electrode, electrons flow in)
Electrons flow from anode to cathode through external circuit
Current flows from cathode to anode (opposite direction)
Salt bridge: Maintains electrical neutrality by allowing ion flow between half-cells

Electrolytic Cells

Non-spontaneous cells that use electrical energy to drive chemical reactions:

External power source forces electron flow
Anode: positive terminal (oxidation)
Cathode: negative terminal (reduction)

Cell Notation (Cell Diagram)

Standard notation for representing electrochemical cells:

$$ \text{Anode} \mid \text{Anode electrolyte} \parallel \text{Cathode electrolyte} \mid \text{Cathode} $$

Single vertical line (|) represents phase boundary
Double vertical line (||) represents salt bridge
Inert electrodes (Pt, graphite) used when no solid metal participates

Examples:

$\text{Zn}(s) \mid \text{Zn}^{2+}(aq) \parallel \text{Cu}^{2+}(aq) \mid \text{Cu}(s)$
$\text{Pt}(s) \mid \text{Br}^-(aq), \text{Br}_2(l) \parallel \text{Cl}_2(g), \text{Cl}^-(aq) \mid \text{Pt}(s)$

BrBr

ClCl

Standard Electrode Potential (E°)

The potential of a half-cell under standard conditions (1 M solutions, 1 atm gases, 25°C).

Standard Hydrogen Electrode (SHE)

Reference electrode with $E° = 0.00$ V: $2\text{H}^+(aq) + 2e^- \rightarrow \text{H}_2(g)$

[H][H]

Standard Cell Potential (E°cell)

$$ E°_{cell} = E°_{cathode} - E°_{anode} $$

Positive $E°_{cell}$ indicates spontaneous reaction
More positive cathode: stronger oxidising agent
More negative anode: stronger reducing agent

Electromotive Force (EMF) and Gibbs Free Energy

$$ \Delta G° = -nFE°_{cell} $$

Where:

$n$ = number of electrons transferred
$F$ = Faraday constant (96,485 C/mol)
Negative $\Delta G°$ = spontaneous reaction

Nernst Equation

Relates cell potential to concentration:

$$ E_{cell} = E°_{cell} - \frac{RT}{nF} \ln Q $$

At 25°C:

$$ E_{cell} = E°_{cell} - \frac{0.0592}{n} \log Q $$

Electrolysis and Faraday's Laws

Faraday's First Law: Mass of substance produced is proportional to quantity of electricity passed.

Faraday's Second Law: Masses of different substances produced by the same quantity of electricity are proportional to their equivalent weights.

$$ m = \frac{Q \times M}{n \times F} $$