Electrochemistry

1. Introduction

Electrochemistry studies the relationship between chemical energy and electrical energy: how spontaneous redox reactions can produce electricity (galvanic cells) and how electricity can drive non-spontaneous reactions (electrolysis). Every electrochemical process involves oxidation (loss of electrons) at one electrode and reduction (gain of electrons) at the other.

2. Electrode Potential and Standard Electrode Potential

When a metal rod is dipped in a solution of its own ions, a potential difference develops at the metal-solution interface because the metal tends either to lose ions into solution (oxidation) or to deposit ions on itself (reduction). This is the electrode potential.

• The standard electrode potential ($E^\circ$) is the electrode potential measured under standard conditions: $1\,\text{M}$ ion concentration, $298\,\text{K}$, $1\,\text{atm}$ for any gas, measured relative to the standard hydrogen electrode (assigned $0.00\,\text{V}$).
• By convention, tabulated values are standard reduction potentials. A more positive $E^\circ$ means a greater tendency to be reduced (a stronger oxidising agent).

The half reaction for a metal electrode:

$\ce{M^n+ (aq) + n e- <=> M(s)}$

3. Reference Electrodes

An electrode's potential cannot be measured alone; it is measured against a reference electrode of known, constant potential.

3.1 Standard Hydrogen Electrode (SHE)

The primary reference, assigned $E^\circ = 0.00\,\text{V}$. Hydrogen gas at $1\,\text{atm}$ bubbles over a platinised platinum foil dipped in $1\,\text{M}$ $\ce{H+}$:

$\ce{2H+ (aq) + 2e- <=> H2(g)}\qquad E^\circ = 0.00\,\text{V}$

It is inconvenient (needs pure $\ce{H2}$, clean Pt), so secondary references are used in practice.