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 (electrolytic cells).

1. Redox Fundamentals

Oxidation and reduction always occur together:

Term Definition Electrode
Oxidation Loss of electrons, increase in oxidation number Anode
Reduction Gain of electrons, decrease in oxidation number Cathode

Mnemonic: OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons). Another: RED CAT, AN OX — REDuction at CAThode, OXidation at ANode.

[Zn+2]

[Cu+2]

2. Two Types of Electrochemical Cells

Feature Galvanic (Voltaic) Cell Electrolytic Cell
Energy conversion Chemical → Electrical Electrical → Chemical
Spontaneity Spontaneous ($\Delta G < 0$) Non-spontaneous ($\Delta G > 0$)
Anode charge Negative (−) Positive (+)
Cathode charge Positive (+) Negative (−)
Anode process Oxidation Oxidation
Cathode process Reduction Reduction
$E°_{cell}$ $> 0$ $< 0$
Example Battery, Daniell cell Electroplating, electrolysis

[!tip] Galvanic vs Electrolytic Polarity GAEP — Galvanic: A-node (−), E-node (+), Produces electricity ECPO — Electrolytic: C-athode (−), P-ositive anode, Outsource needed (external power)

Or: GAL-AN — in a GALvanic cell, ANode is negative. Since electrolytic is the opposite: anode = positive.

Key Insight

In both cell types, oxidation always occurs at the anode and reduction at the cathode. The difference is which electrode is positive/negative — determined by whether the cell produces or consumes electrical energy.

3. Galvanic Cells in Detail

Cell Notation (Cell Diagram)

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

  • | = phase boundary
  • || = salt bridge
  • Inert electrode (Pt, C) used when no solid metal participates

Example — Daniell Cell: $$\text{Zn}(s) \vert \text{Zn}^{2+}(aq) \parallel \text{Cu}^{2+}(aq) \vert \text{Cu}(s)$$

Example — Gas Electrodes: $$\text{Pt}(s) \vert \text{Br}^-(aq), \text{Br}_2(l) \parallel \text{Cl}_2(g), \text{Cl}^-(aq) \vert \text{Pt}(s)$$

BrBr

ClCl

How It Works

  1. Zn electrode erodes as Zn oxidises to Zn²⁺
  2. Cu²⁺ deposits on Cu electrode as it reduces to Cu
  3. Anode solution becomes positively charged (Zn²⁺ accumulation)
  4. Cathode solution becomes negatively charged (Cu²⁺ depletion)
  5. Salt bridge neutralises charge buildup (anions → anode, cations → cathode)

Standard Hydrogen Electrode (SHE)

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

Standard Cell Potential

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

  • $E°_{cell} > 0$ → spontaneous (galvanic)
  • $E°_{cell} < 0$ → non-spontaneous (electrolytic)
  • $E°_{cell} = 0$ → equilibrium (dead battery)

Relative Strength of Agents

  • Stronger oxidising agent → more positive $E°_{reduction}$
  • Stronger reducing agent → more negative $E°_{reduction}$
  • Anode = more negative $E°$ (better reducing agent)
  • Cathode = more positive $E°$ (better oxidising agent)

4. Electrolytic Cells in Detail

An electrolytic cell uses electrical energy to force a non-spontaneous reaction.

Electrolysis of CuSO₄ with Copper Electrodes

  • Anode (impure Cu): $\text{Cu}(s) \rightarrow \text{Cu}^{2+}(aq) + 2e^-$ — active electrode dissolves
  • Cathode (pure Cu): $\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)$ — pure Cu deposited
  • Impurity falls as anode mud; pure copper collected at cathode
[Cu+2]

5. Selective Discharge of Ions

When multiple ions are present in aqueous solution, only one cation and one anion are discharged. Three factors determine priority:

Factor 1: Electrochemical Series (E° Values)

Cations (discharged at cathode) — ease of discharge ↓ down the series: $$\text{K}^+ < \text{Na}^+ < \text{Ca}^{2+} < \text{Mg}^{2+} < \text{Al}^{3+} < \text{Zn}^{2+} < \text{Fe}^{2+} < \text{Sn}^{2+} < \text{Pb}^{2+} < \text{H}^+ < \text{Cu}^{2+} < \text{Ag}^+$$

Anions (discharged at anode) — ease of discharge ↓ down: $$\text{F}^- < \text{SO}_4^{2-} < \text{NO}_3^- < \text{Cl}^- < \text{Br}^- < \text{I}^- < \text{OH}^-$$

[!tip] Electrochemical Series Mnemonic Cation ease-of-discharge (discharged at cathode):
King Napoleon Can't Mgnify All Zinc's Fent Snail Pbone Healing Cures Again
— K⁺ → Na⁺ → Ca²⁺ → Mg²⁺ → Al³⁺ → Zn²⁺ → Fe²⁺ → Sn²⁺ → Pb²⁺ → H⁺ → Cu²⁺ → Ag⁺

Anion ease-of-discharge (discharged at anode):
Fat SOda Never Cleans Brown Ink On Hands
— F⁻ → SO₄²⁻ → NO₃⁻ → Cl⁻ → Br⁻ → I⁻ → OH⁻

Note: In aqueous solutions, $\text{H}^+$ and $\text{OH}^-$ from water autoprotolysis compete with solute ions.

Factor 2: Concentration

Higher concentration favours discharge even if the ion is lower in the electrochemical series.

Factor 3: Type of Electrode

  • Active electrode (e.g., Cu, Ag): Participates in the reaction; can dissolve
  • Inert electrode (e.g., Pt, graphite, C): Only provides surface; does not react

6. Overpotential (Overvoltage)

Standard electrode potentials alone do not fully predict electrolysis products — overpotential is the extra voltage required beyond the theoretical value.

Oxygen Overvoltage

  • Producing $\text{O}_2$ at the anode requires an additional ~0.4–0.6 V above $E°$
  • Causes: slow reaction kinetics and bubble formation at electrode surface
  • Effective potential for $\text{H}_2\text{O} \rightarrow \text{O}_2$: ~1.22–1.42 V

Practical Consequence

Enables $\text{Cl}^-$, $\text{Br}^-$, $\text{I}^-$ to be oxidised preferentially over water:

Process Standard $E°$ Overpotential Effective
$2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^-$ −1.36 V minimal ≈ −1.36 V
$2\text{H}_2\text{O} \rightarrow \text{O}_2 + 4\text{H}^+ + 4e^-$ −1.23 V ~0.5 V ≈ −1.73 V

Example: Electrolysis of dilute NaCl — $\text{Cl}_2$ gas is produced at anode instead of $\text{O}_2$ despite unfavourable standard potentials.

7. Industrial Electrolysis Applications

The Downs Cell — Sodium Production

  • Process: Electrolysis of molten NaCl (not aqueous)
  • Mixture: NaCl + CaCl₂ (2:3 ratio) lowers melting point from 801°C to ~580°C
  • Products: Liquid Na metal at cathode, $\text{Cl}_2$ gas at anode
  • $\text{Na}^+ + e^- \rightarrow \text{Na}(l)$ at cathode
  • $2\text{Cl}^- \rightarrow \text{Cl}_2(g) + 2e^-$ at anode

The Hall Process — Aluminium Extraction

  • Process: Electrolysis of $\text{Al}_2\text{O}_3$ (alumina) dissolved in molten cryolite ($\text{Na}_3\text{AlF}_6$)
  • Temperature: ~940–980°C
  • Anode: Carbon (consumed as $\text{C} + 2\text{O}^{2-} \rightarrow \text{CO}_2 + 4e^-$)
  • Cathode: Iron vessel (molten Al collects at bottom)
  • Overall: $2\text{Al}_2\text{O}_3 + 3\text{C} \rightarrow 4\text{Al} + 3\text{CO}_2$
[Na+]

[Al+3]

8. Faraday's Laws of Electrolysis

First Law

$$m \propto Q \quad\Rightarrow\quad m = ZQ = ZIt$$

The mass of substance produced at an electrode is directly proportional to the quantity of electricity passed.

Second Law

$$\frac{m_1}{m_2} = \frac{E_1}{E_2}$$

Masses of different substances produced by the same quantity of electricity are proportional to their equivalent weights $E = M/n$.

Unified Formula

$$m = \frac{QM}{nF} = \frac{ItM}{nF}$$

[!tip] Faraday's Formula Mnemonic: QuIM NoF $m = \dfrac{Q \cdot M}{n \cdot F} = \dfrac{\color{blue}{I} \color{black}{\cdot t \cdot M}}{\color{blue}{n} \color{black}{\cdot F}}$ Qu = Q (charge), I = I (current), M = M (molar mass), NoF = n · F (electrons × Faraday) — or: "It's My NoF formula" (I × t × M) / (n × F)

Where:

  • $m$ = mass produced (g)
  • $Q$ = charge passed (C) = $I \times t$
  • $I$ = current (A)
  • $t$ = time (s)
  • $M$ = molar mass (g·mol⁻¹)
  • $n$ = number of electrons per ion
  • $F$ = Faraday constant = $96{,}485\ \text{C·mol}^{-1}$

9. Thermodynamic Relationships

Gibbs Free Energy and Cell Potential

$$\Delta G° = -nFE°_{cell}$$

[!tip] ΔG-E Mnemonic: "Negative Nancy Eats Fries" $\Delta G° = \mathbf{-}n\mathbf{F}E°_{cell}$ — The minus sign links $\Delta G$ and $E_{cell}$ oppositely: When $E_{cell}$ is positive → $\Delta G$ is negative → spontaneous. Think: "Nancy Fries Eggs" → $\mathbf{n}\mathbf{F}\mathbf{E}$ with a minus in front.

$\Delta G°$ $E°_{cell}$ Spontaneity
$< 0$ $> 0$ Spontaneous (galvanic)
$> 0$ $< 0$ Non-spontaneous (electrolytic)
$= 0$ $= 0$ Equilibrium

Equilibrium Constant

$$E°_{cell} = \frac{RT}{nF} \ln K$$

At 25°C: $E°_{cell} = \dfrac{0.0592}{n} \log K$

10. The Nernst Equation

Relates cell potential to concentration under non-standard conditions:

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

At 25°C:

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

[!tip] Nernst Mnemonic: "Standard Minus 0.0592 over n log Q" The Nernst equation subtracts from $E°$ — higher product concentration (larger $Q$) lowers $E_{cell}$. At equilibrium ($Q = K$): $E_{cell} = 0$, so $E°_{cell} = \dfrac{0.0592}{n} \log K$. Think: "E is E-standard minus Reaction Progress Penalty" — as reaction proceeds, voltage drops.

The Dead Battery Analogy

A dead battery isn't out of chemicals — the forward and reverse reactions have reached equilibrium ($Q = K$). At equilibrium, $E_{cell} = 0$ because both directions proceed at equal rates, creating no net electron push.

Concentration Cell

Two identical half-cells with different ion concentrations create a potential difference: $$E_{cell} = -\frac{0.0592}{n} \log \frac{[\text{dilute}]}{[\text{concentrated}]}$$

Electrons flow from the more dilute side (higher potential to give electrons) to the more concentrated side — the cell equalises concentrations.

11. Common Applications

Application Type Key Chemistry
Lead-acid battery Galvanic $\text{Pb} + \text{PbO}_2 + 2\text{H}_2\text{SO}_4 \rightleftharpoons 2\text{PbSO}_4 + 2\text{H}_2\text{O}$
Lithium-ion battery Galvanic Li⁺ intercalation in graphite and metal oxides
H₂-O₂ Fuel Cell Galvanic $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$
Electroplating Electrolytic Metal cation reduction at cathode
Chlor-alkali process Electrolytic $\text{NaCl}(aq) \rightarrow \text{NaOH} + \text{Cl}_2 + \text{H}_2$
Corrosion prevention Galvanic Sacrificial anode (more reactive metal oxidises instead)

Related Pages

  • FAD1018 - Basic Chemistry II — course page
  • Redox Reactions — oxidation-reduction fundamentals
  • Gibbs Free Energy — thermodynamic spontaneity
  • Nernst Equation — concentration effects on cell potential
  • Electrolytic Cell — non-spontaneous cells, selective discharge
  • Faraday's Laws — quantitative electrolysis
  • Electrochemical Series — ion discharge order
  • Overpotential — overvoltage in electrolysis

Sources