Electrochemistry – WISDOMYSTERY

§ 01

Introduction to Electrochemistry

Electrochemistry is the branch of chemistry that explores the relationship between electrical energy and chemical reactions. It underpins technologies ranging from batteries and fuel cells to electroplating and corrosion prevention.

At its heart, electrochemistry is driven by redox reactions — processes involving simultaneous oxidation and reduction, where electrons are transferred between species.

🔺 Oxidation

Loss of electrons (LEO — Loss of Electrons is Oxidation). The oxidation number increases. Occurs at the anode.

🔻 Reduction

Gain of electrons (GER — Gain of Electrons is Reduction). The oxidation number decreases. Occurs at the cathode.

💡

Mnemonic: OIL RIG — Oxidation Is Loss, Reduction Is Gain. Or LEO the lion says GER.

General Redox Half-Reactions

Oxidation: M → Mⁿ⁺ + n e⁻ (Anode) Reduction: Mⁿ⁺ + n e⁻ → M (Cathode)

In any balanced redox equation, electrons lost = electrons gained.

Fig 1.1 — Electron transfer in a redox reaction

§ 02

Electrochemical Cells

An electrochemical cell converts chemical energy into electrical energy (or vice versa) using redox reactions. There are two primary types:

⚡ Galvanic (Voltaic) Cell

Spontaneous redox reaction produces electrical energy. ΔG < 0, E°cell > 0. Example: Daniell Cell (Zn–Cu).

🔌 Electrolytic Cell

External electrical energy drives a non-spontaneous redox reaction. ΔG > 0. Used in electroplating, electrolysis of water.

The Daniell Cell (Zn–Cu Cell)

The classic galvanic cell uses a zinc anode in ZnSO₄ solution and a copper cathode in CuSO₄ solution, connected by a salt bridge.

Fig 2.1 — Daniell Cell (Zn–Cu Galvanic Cell)

Cell Notation (IUPAC)

Zn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s)

Single bar (|) = phase boundary. Double bar (‖) = salt bridge. Anode written on left, cathode on right.

Half-Cell Reactions — Daniell Cell

Anode : Zn(s) → Zn²⁺(aq) + 2e⁻ [Oxidation] Cathode : Cu²⁺(aq) + 2e⁻ → Cu(s) [Reduction] ───────────────────────────────────────── Net : Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

E°cell = E°cathode − E°anode = 0.34 − (−0.76) = +1.10 V

Role of Salt Bridge

The salt bridge (e.g., KCl in agar gel) maintains electrical neutrality in both half-cells by allowing ions to migrate: cations (K⁺) move toward the cathode compartment and anions (Cl⁻) move toward the anode compartment. Without it, current ceases almost immediately.

§ 03

Electrode Potential & EMF

When a metal electrode is placed in a solution of its ions, a potential difference develops between the metal and the solution — called the electrode potential. It arises from the tendency of the metal to lose or gain electrons.

Standard Hydrogen Electrode (SHE) — The Reference

All standard electrode potentials (E°) are measured relative to the SHE, which is assigned a value of 0.00 V. It consists of H₂ gas at 1 atm bubbling over a platinum electrode in 1 M H⁺ solution at 25 °C.

STANDARD HYDROGEN ELECTRODE CELL DIAGRAM

Fig 3.1 — Standard Hydrogen Electrode (SHE)

Standard Reduction Potentials

The higher the reduction potential, the stronger the oxidising agent. Metals with lower (more negative) potentials are stronger reducing agents and tend to lose electrons more readily.

Half-Reaction	E° (V)	Character
Li⁺ + e⁻ → Li	−3.04	Strongest reducing agent
Zn²⁺ + 2e⁻ → Zn	−0.76	Good reducing agent
Fe²⁺ + 2e⁻ → Fe	−0.44	Moderate reducing agent
2H⁺ + 2e⁻ → H₂	0.00	Reference (SHE)
Cu²⁺ + 2e⁻ → Cu	+0.34	Mild oxidising agent
Ag⁺ + e⁻ → Ag	+0.80	Good oxidising agent
F₂ + 2e⁻ → 2F⁻	+2.87	Strongest oxidising agent

Cell EMF Formula

E°cell = E°cathode − E°anode = E°reduction(cathode) − E°reduction(anode)

A positive E°cell indicates a spontaneous reaction. The species with higher reduction potential acts as the cathode.

Gibbs Energy & EMF Relationship

ΔG° = −nFE°cell

n = moles of electrons transferred; F = Faraday constant = 96,485 C mol⁻¹ If ΔG° < 0 ↔ E°cell > 0 ↔ Spontaneous reaction

§ 04

Nernst Equation

The Nernst equation, derived by Walther Nernst in 1889, allows calculation of the electrode potential at non-standard conditions (concentrations other than 1 M, temperatures other than 25 °C). It is one of the most powerful tools in electrochemistry.

Nernst Equation — General Form

E = E° − (RT / nF) × ln Q At 25 °C (298 K): E = E° − (0.0592 / n) × log₁₀ Q

R = 8.314 J K⁻¹ mol⁻¹ (gas constant)
T = Temperature in Kelvin
n = number of electrons transferred
F = 96,485 C mol⁻¹ (Faraday's constant)
Q = Reaction quotient = [Products] / [Reactants]

Applying the Nernst Equation

For the reaction: Zn + Cu²⁺ → Zn²⁺ + Cu (n = 2):

Daniell Cell — Nernst Application

E = 1.10 − (0.0592/2) × log([Zn²⁺]/[Cu²⁺]) If [Zn²⁺] = 0.1 M and [Cu²⁺] = 1 M: Q = 0.1/1 = 0.1 E = 1.10 − (0.0296) × log(0.1) = 1.10 − (0.0296) × (−1) = 1.10 + 0.0296 = 1.1296 V ≈ 1.13 V

At Equilibrium

When a cell reaches equilibrium, E = 0 and Q = K (equilibrium constant). The Nernst equation then becomes:

Equilibrium Constant from E°cell

E°cell = (0.0592 / n) × log K (at 25 °C) ∴ log K = (n × E°cell) / 0.0592

A large positive E°cell gives a very large K, indicating the reaction goes nearly to completion.

📌

Key Insight: As a cell discharges, reactant concentration decreases, Q increases, and E decreases. When E = 0, the cell is fully discharged (equilibrium reached).

§ 05

Electrolytic Cells

An electrolytic cell uses an external source of electrical energy to drive a thermodynamically non-spontaneous redox reaction. The external DC source acts as a "pump" forcing electrons in the non-spontaneous direction.

Fig 5.1 — Electrolytic Cell (Electrolysis of HCl)

Electrolysis of Water

Dilute H₂SO₄ is used to increase conductance. Platinum electrodes are inert.

Electrolysis of Water — Half Reactions

Cathode (reduction): 4H⁺ + 4e⁻ → 2H₂ (g) Anode (oxidation): 2H₂O → O₂ (g) + 4H⁺ + 4e⁻ ──────────────────────────────────────────── Overall: 2H₂O → 2H₂ + O₂

Volume ratio H₂ : O₂ = 2 : 1 (collected at cathode : anode)

Applications of Electrolysis

🪙 Electroplating

Depositing a thin layer of one metal over another. Object to be plated = cathode. Plating metal = anode. Electrolyte = salt of plating metal.

🏭 Electrorefining

Purification of metals (e.g., copper). Impure metal = anode (dissolves). Pure metal = cathode (deposits). Impurities fall as anode mud.

§ 06

Faraday's Laws of Electrolysis

Michael Faraday (1833) established two fundamental laws quantifying the relationship between the amount of electricity passed through an electrolyte and the chemical change produced.

First Law of Electrolysis

The mass of substance deposited or dissolved at an electrode is directly proportional to the quantity of electricity (charge) passed through the electrolyte.

Faraday's First Law

m = Z × Q = Z × I × t where: Z = electrochemical equivalent = M / (n × F)

m = mass deposited (g) | Z = electrochemical equivalent (g C⁻¹)
Q = charge (Coulombs) | I = current (A) | t = time (s)

Second Law of Electrolysis

When the same quantity of charge is passed through different electrolytes, the masses of substances deposited are proportional to their equivalent weights.

Faraday's Second Law

m₁/m₂ = E₁/E₂ (where E = M/n = equivalent weight) General formula: m = (M × I × t) / (n × F)

M = molar mass (g mol⁻¹) | n = valency/electrons transferred
F = 96,485 C mol⁻¹ ≈ 96,500 C mol⁻¹ (Faraday constant)

Worked Example

Example — Copper Deposition

How much Cu deposits when 2 A flows for 30 min? (M_Cu = 63.5 g/mol, n = 2) t = 30 × 60 = 1800 s m = (63.5 × 2 × 1800) / (2 × 96500) m = 228600 / 193000 m ≈ 1.185 g

⚗️

One Faraday (96,485 C) deposits exactly one mole of a monovalent ion (e.g., Ag⁺), half a mole of a divalent ion (Cu²⁺), etc.

§ 07

Conductance of Electrolytic Solutions

The ability of an electrolytic solution to conduct electricity depends on the number of ions present, their charge, and mobility.

Resistance & Conductance

R = ρ × (l / A) G = 1/R = κ × (A/l)

ρ = resistivity; κ = conductivity (S m⁻¹)
l = length; A = cross-section area

Molar Conductivity

Λm = κ / c

c = concentration (mol m⁻³)
Units: S m² mol⁻¹

Kohlrausch's Law

For strong electrolytes, molar conductivity varies with concentration:

Kohlrausch's Law (Strong Electrolytes)

Λm = Λ°m − K√c

Λ°m = limiting molar conductivity (at infinite dilution)
K = Kohlrausch constant; c = concentration (mol L⁻¹)
Plot of Λm vs √c is a straight line with intercept Λ°m

Law of Independent Migration of Ions

Λ°m = ν₊λ°₊ + ν₋λ°₋

λ°₊ and λ°₋ = limiting molar conductivity of cation and anion
ν₊, ν₋ = stoichiometric numbers
Used to find Λ°m for weak electrolytes (e.g., CH₃COOH) indirectly.

Degree of Dissociation from Conductance

For Weak Electrolytes

α = Λm / Λ°m Ka = cα² / (1 − α) ≈ cα² (when α « 1)

α = degree of dissociation; Ka = acid dissociation constant

§ 08

Batteries & Fuel Cells

A battery is a self-contained galvanic cell (or set of cells) that stores chemical energy and converts it to electrical energy on demand.

Primary vs Secondary Batteries

🔋 Primary Cells (Non-rechargeable)

Reactions are irreversible. Once discharged, cannot be recharged. Examples: Leclanche dry cell (1.5 V), Mercury cell (1.35 V), Alkaline cell.

⚡ Secondary Cells (Rechargeable)

Reactions are reversible. Can be recharged by applying external current. Examples: Lead–Acid battery, Ni–Cd cell, Li-ion battery.

Lead–Acid Storage Battery

Lead–Acid Battery — Reactions

Anode (discharge): Pb + SO₄²⁻ → PbSO₄ + 2e⁻ Cathode (discharge): PbO₂ + 4H⁺ + SO₄²⁻ + 2e⁻ → PbSO₄ + 2H₂O Overall: Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O EMF ≈ 2 V per cell (6 cells → 12 V battery)

On recharging, the reactions are reversed. H₂SO₄ density rises on charging, falls on discharging — used to monitor battery state.

Fuel Cells

A fuel cell continuously converts the chemical energy of a fuel (H₂) and oxidant (O₂) directly into electrical energy, with water as the only by-product. Much more efficient (~70%) than internal combustion engines (~35%).

Hydrogen–Oxygen Fuel Cell

Anode : H₂ + 2OH⁻ → 2H₂O + 2e⁻ [Alkaline] Cathode : O₂ + 2H₂O + 4e⁻ → 4OH⁻ ──────────────────────────────────────────── Overall : 2H₂ + O₂ → 2H₂O (E°cell ≈ 1.23 V)

Platinum/palladium acts as a catalyst at both electrodes. Used in spacecraft, buses, and stationary power plants.

Fig 8.1 — Battery & Fuel Cell Comparison

§ 09

Corrosion

Corrosion is the oxidative deterioration of a metal due to its reaction with environmental agents (O₂, H₂O, CO₂, acids). It is essentially an electrochemical process involving the formation of tiny galvanic cells on the metal surface.

Rusting of Iron (Fe)

Rusting is the most economically significant form of corrosion. It requires both oxygen and moisture. Impurities in iron create anodic and cathodic regions, forming a microgalvanic cell.

Iron Corrosion — Mechanism

Anode : Fe → Fe²⁺ + 2e⁻ (Oxidation) Cathode : O₂ + 2H₂O + 4e⁻ → 4OH⁻ (Reduction) ──────────────────────────────────────────── Fe²⁺ + 2OH⁻ → Fe(OH)₂ 2Fe(OH)₂ + ½O₂ → Fe₂O₃·H₂O (Rust)

Rust (Fe₂O₃·xH₂O) is porous — it does not protect the underlying metal, unlike aluminium oxide which forms a protective passivating layer.

Factors Affecting Corrosion Rate

⬆️ Factors that Accelerate

High humidity • Presence of electrolytes (NaCl) • Low pH (acid rain) • High temperature • Galvanic contact with noble metals • Scratches exposing fresh metal

🛡️ Prevention Methods

Galvanisation (Zn coat) • Tinning (Sn coat) • Painting/oiling • Alloying (stainless steel) • Cathodic protection • Sacrificial anode (Mg, Zn)

Cathodic Protection

Cathodic Protection — Sacrificial Anode

Sacrificial metal (Mg or Zn) attached to steel pipeline: Anode : Mg → Mg²⁺ + 2e⁻ (Mg corrodes) Cathode : Fe²⁺ + 2e⁻ → Fe (Fe protected) ∴ Steel pipe remains as cathode — protected from corrosion.

Mg and Zn have lower reduction potentials than Fe, so they preferentially oxidize. Must be periodically replaced.

🔬

Aluminium Paradox: Although Al has a very negative E° (−1.66 V), it resists corrosion due to a thin, adherent, impermeable Al₂O₃ layer that forms instantly on exposure to air — a process called passivation.