Course Topics

Introduction to Electrochemistry

Electrochemistry is the branch of chemistry which is concerned with the inter-conversion of chemical energy and electrical energy.

Process Energy Conversion Type of Cell
Electrical to Chemical Electrical energy is converted into chemical energy Electrolytic cells
Chemical to Electrical Chemical energy is converted into electrical energy Galvanic or voltaic cells

Oxidation

  • Gain of oxygen
  • Loss of hydrogen
  • Loss of electrons (anode reactions)
  • Increase in oxidation state

Reduction

  • Loss of oxygen
  • Gain of hydrogen
  • Gain of electrons (cathode reactions)
  • Decrease in oxidation state
Memory Tip

OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons)

LEO GER: Lose Electrons Oxidation, Gain Electrons Reduction

Oxidation Number

Definition: Apparent charge on an atom of an element in a molecule or an ion is called oxidation number. It may be positive or negative or zero.

Example: K₂Cr₂O₇
(+1)×2 + 2Cr + (-2)×7 = 0
2 + 2Cr – 14 = 0
2Cr = 12
Cr = +6
Element/Compound Rules Examples
Element Molecular form = 0, Bulk/chunk = 0 H₂ = 0, S₈ = 0, Diamond (C) = 0
Compound Algebraic sum of oxidation numbers = 0 HCl: (+1) + (-1) = 0
Ion Oxidation number = Charge on ion Cl⁻ = -1, SO₄²⁻ = +6 + 4(-2) = -2
Hydrogen Usually +1 (except in metal hydrides = -1) H₂O: +1, NaH: -1
Oxygen Usually -2 (except in peroxides = -1, superoxides = -½, with F = +2) H₂O: -2, H₂O₂: -1, KO₂: -½, OF₂: +2
CRITICAL CONCEPT:
  • Oxidation number of IA group = +1
  • Oxination number of IIA group = +2
  • Oxidation number of IIIA group = +3
  • Fluorine always has oxidation number = -1

Redox Reactions

Redox Reaction: Oxidation and reduction always take place together. There can be no oxidation without reduction and vice versa.

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
Ionic: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
Oxidation (Zn) Reduction (Cu²⁺)
Zinc atom loses two electrons Copper ion gains two electrons
Oxidation state increases from 0 to +2 Oxidation state decreases from +2 to 0
Zn atom oxidizes to Zn²⁺ ion Cu²⁺ ion reduces to Cu atom
Zn acts as reducing agent (reductant) Cu²⁺ acts as oxidizing agent (oxidant)
CRITICAL CONCEPT: In a redox reaction:
  • Oxidizing agent gets reduced (gains electrons)
  • Reducing agent gets oxidized (loses electrons)
  • Electrons flow from reducing agent to oxidizing agent

Balancing Redox Equations

Two Methods for Balancing Redox Equations:

Ion-Electron Method

  • Write half-reactions for oxidation and reduction
  • Balance atoms except H and O
  • Balance O atoms by adding H₂O
  • Balance H atoms by adding H⁺
  • Balance charge by adding electrons
  • Equalize electrons in both half-reactions
  • Add half-reactions and simplify

Oxidation Number Method

  • Write skeleton equation
  • Identify elements whose oxidation number changed
  • Indicate number of electrons gained or lost
  • Equate increase/decrease in oxidation number
  • Balance equation by inspection method
Example: MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺ (acidic medium)
Balanced: MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O
Memory Tip

Acidic medium: Add H⁺ to balance H, add H₂O to balance O

Basic medium: Add OH⁻ to balance H, add H₂O to balance O

Electrolysis

Electrolysis: The process in which non-spontaneous reaction takes place at the expense of electricity.

Anode (+)
Cathode (-)
Electrolyte Cathode Product Anode Product
PbBr₂ (molten) Pb(s) Br₂(g)
NaCl (molten) Na(s) Cl₂(g)
NaCl (aq) H₂(g) Cl₂(g)
CuCl₂ (aq) Cu(s) Cl₂(g)
CuSO₄ (aq) Cu(s) O₂(g)
KNO₃ (aq) H₂(g) O₂(g)
CRITICAL CONCEPT – Order of Discharge:

Cations: Higher reduction potential ions discharge first at cathode

Anions: Lower oxidation potential ions discharge first at anode

Ease of discharging anions: I⁻ > Br⁻ > Cl⁻ > OH⁻ > NO₃⁻ > SO₄²⁻

Electrochemical Cells

Voltaic/Galvanic Cell Electrolytic Cell
Chemical energy → Electrical energy Electrical energy → Chemical energy
Anode is negative, Cathode is positive Anode is positive, Cathode is negative
Spontaneous redox reaction Non-spontaneous redox reaction
Salt bridge is used No salt bridge needed
Generates heat (exothermic) Consumes heat (endothermic)
Examples: Daniel’s cell, Ni-Cd cell Examples: Nelson’s Cell, Down’s Cell
CRITICAL CONCEPT – Daniel’s Cell:
  • Zn(s) | Zn²⁺(1M) || Cu²⁺(1M) | Cu(s)
  • E°cell = 1.10 Volts
  • Anode: Zn(s) → Zn²⁺(aq) + 2e⁻ (Oxidation)
  • Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) (Reduction)
  • Salt bridge: Aqueous KCl in gel

Standard Electrode Potential

Standard Electrode Potential (E°): The potential setup when an electrode is in contact with one molar solution of its own ions at 298K.

E°cell = E°cathode – E°anode
E°cell = E°reduction + E°oxidation
CRITICAL CONCEPT – Standard Hydrogen Electrode (SHE):
  • Reference electrode with E° = 0.00 V
  • Consists of Pt foil coated with Pt black
  • H₂ gas at 1 atm pressure
  • 1M HCl solution
  • Reactions: H₂(g) ⇌ 2H⁺(aq) + 2e⁻

Calculating Standard Cell Potential:

  1. Electrode with higher reduction potential acts as cathode
  2. Electrode with lower reduction potential acts as anode
  3. E°cell = E°cathode – E°anode
  4. If E°cell is positive, reaction is spontaneous

Electrochemical Series (ECS)

Electrochemical Series: When elements are arranged in the order of their increasing standard reduction potentials on the hydrogen scale.

Memory Mnemonic

Little (Li), Kinza (K), Can (Ca), Send (Na), Monkey (Mg), Elephant (Al), Zebras (Zn), Crows (Cr), In (Fe³⁺), Cubic (Co), Nickel (Ni), Tin (Sn) Lead (Pb), Hut (H₂), Cute (Cu²⁺) Cute (Cu⁺), Imported (I₂), Iron (Fe²⁺), Silver (Ag), Hyderabadi (Hg₂²⁺), Bangles (Br₂), Collected (Cl₂), All over (Au), Finland (F₂)

Element Reduction Reaction E° (Volts)
Li Li⁺ + e⁻ → Li -3.05
K K⁺ + e⁻ → K -2.93
Zn Zn²⁺ + 2e⁻ → Zn -0.76
H₂ 2H⁺ + 2e⁻ → H₂ 0.00
Cu Cu²⁺ + 2e⁻ → Cu +0.34
Ag Ag⁺ + e⁻ → Ag +0.80
Au Au³⁺ + 3e⁻ → Au +1.50
CRITICAL CONCEPT – Applications of ECS:
  • Predict feasibility of redox reactions
  • Determine oxidizing/reducing strength
  • Predict displacement reactions
  • Determine corrosion tendencies
  • Calculate EMF of cells

Cell Potential Calculations

Example: Calculate E°cell for Zn-Cu cell

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

E°red(Cu²⁺/Cu) = +0.34 V
E°red(Zn²⁺/Zn) = -0.76 V

E°cell = E°cathode – E°anode
E°cell = 0.34 – (-0.76)
E°cell = 1.10 V

Nernst Equation: For non-standard conditions

E = E° – (RT/nF) ln Q
At 298K: E = E° – (0.059/n) log Q
Memory Tip

Positive E°cell = Spontaneous reaction

Negative E°cell = Non-spontaneous reaction

Zero E°cell = Equilibrium condition

Applications of Electrochemistry

Practical Applications:

  • Batteries: Lead-acid, Li-ion, Ni-Cd, Fuel cells
  • Electroplating: Chromium plating, silver plating, gold plating
  • Corrosion Prevention: Galvanization, sacrificial anodes
  • Electrorefining: Purification of metals (Cu, Al, Zn)
  • Electrosynthesis: Production of chemicals (NaOH, Cl₂, Al)
  • Sensors: pH meters, glucose sensors, gas detectors
  • Energy Storage: Supercapacitors, redox flow batteries
CRITICAL CONCEPT – Electroplating:
  • Anode: Made of metal to be plated
  • Cathode: Object to be plated
  • Electrolyte: Salt solution of plating metal
  • Metal dissolves from anode and deposits on cathode
  • Examples: Silver plating, chromium plating