Class 12 Electrochemistry - Lecture 1

Class 12 Electrochemistry

Complete Lecture Structure from Basic to Advanced Concepts

10-Lecture Course Structure

Introduction to Electrochemistry & Redox Reactions

Galvanic Cells & Electrode Potential

Standard Electrode Potential & EMF Calculations

Nernst Equation & Concentration Dependence

Electrolytic Cells & Faraday's Laws

Conductance, Resistivity & Kohlrausch's Law

Batteries (Primary & Secondary) & Fuel Cells

Corrosion: Mechanisms & Prevention

Electrochemical Series & Applications

Advanced Concepts: Concentration Cells, Overpotential

Lecture 1: Introduction to Electrochemistry & Redox Reactions

Duration: 60 minutes

Objective

Establish foundational concepts of redox reactions and their electrochemical significance.

I. Introduction to Electrochemistry (10 mins)

Electrochemistry is the branch of chemistry that deals with the relationship between electrical energy and chemical changes.

Definition

Study of interconversion between chemical energy and electrical energy.

Real-world Applications

Batteries and accumulators
Corrosion prevention
Metal extraction (electrometallurgy)
Biological processes (nerve signal transmission)
Electroplating and electrorefining

II. Redox Reactions: Core Principles (20 mins)

Key Terms

Oxidation: Loss of electrons
Zn → Zn²⁺ + 2e⁻
Reduction: Gain of electrons
Cu²⁺ + 2e⁻ → Cu
Oxidizing Agent: Accepts electrons (gets reduced)
Reducing Agent: Donates electrons (gets oxidized)

Oxidation Number Rules

Elemental state: 0 (e.g., O₂, Na)
Monatomic ions: Charge = O.N. (e.g., Fe³⁺: +3)
Oxygen: Usually -2 (exceptions: peroxides = -1)
Hydrogen: +1 with nonmetals (e.g., H₂O), -1 with metals (e.g., NaH)
Fluorine: Always -1 in compounds
Sum in compound: Must equal compound's charge

III. Balancing Redox Reactions (20 mins)

Oxidation Number Method (Example)

Reaction: Fe²⁺ + Cr₂O₇²⁻ → Fe³⁺ + Cr³⁺ (acidic medium)

Steps:

Assign O.N.: Cr⁶⁺ → Cr³⁺ (reduction), Fe²⁺ → Fe³⁺ (oxidation)
Balance atoms & charge:
Cr₂O₇²⁻ + 6Fe²⁺ + 14H⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O

Half-Reaction Method (Example)

Reaction: Zn + Ag⁺ → Zn²⁺ + Ag

Steps:

Oxidation:
Zn → Zn²⁺ + 2e⁻
Reduction:
2Ag⁺ + 2e⁻ → 2Ag
Balanced Equation:
Zn + 2Ag⁺ → Zn²⁺ + 2Ag

IV. Why Redox Matters in Electrochemistry (5 mins)

Link to Electrochemical Cells

Spontaneous redox reactions → Electric current (Galvanic cells)
Non-spontaneous redox reactions → Electrolysis (Electrolytic cells)

Redox reactions form the basis of all electrochemical processes, allowing us to harness chemical energy as electrical energy and vice versa.

V. Summary & Homework (5 mins)

Key Takeaways

Redox reactions involve electron transfer
Balancing ensures mass/charge conservation
Oxidation number rules help track electron transfer
Redox reactions are fundamental to electrochemical cells

Homework

1. Balance: MnO₄⁻ + I⁻ → MnO₂ + I₂ (basic medium)
2. Find oxidation number of Cr in K₂Cr₂O₇
3. Identify oxidizing and reducing agents in: 2H₂S + SO₂ → 3S + 2H₂O

Next Lecture Preview

Galvanic Cells - Structure, salt bridge, and electrode potential

Materials Needed: Textbook, notebook, calculator

Class 12 Electrochemistry - Lecture 2: Galvanic Cells

Class 12 Electrochemistry

Lecture 2: Galvanic Cells & Electrode Potential

10-Lecture Course Structure

Introduction to Electrochemistry & Redox Reactions

Galvanic Cells & Electrode Potential

Standard Electrode Potential & EMF Calculations

Nernst Equation & Concentration Dependence

Electrolytic Cells & Faraday's Laws

Conductance, Resistivity & Kohlrausch's Law

Batteries (Primary & Secondary) & Fuel Cells

Corrosion: Mechanisms & Prevention

Electrochemical Series & Applications

Advanced Concepts: Concentration Cells, Overpotential

Lecture 2: Galvanic Cells & Electrode Potential

Duration: 60 minutes

Objective

Understand the structure and functioning of galvanic cells, and introduce the concept of electrode potential.

I. Introduction to Galvanic Cells (15 mins)

A galvanic cell (or voltaic cell) is an electrochemical cell that derives electrical energy from spontaneous redox reactions taking place within the cell.

Key Components

Anode: Electrode where oxidation occurs (electrons flow out)
Cathode: Electrode where reduction occurs (electrons flow in)
Salt Bridge: Maintains electrical neutrality by ion migration
Electrolytes: Conducting solutions in which electrodes are immersed

Working Principle

Spontaneous redox reaction produces electrical energy:

Daniell Cell Example

Zn | Zn²⁺(1M) || Cu²⁺(1M) | Cu

Anode: Zn → Zn²⁺ + 2e⁻ (Oxidation)

Cathode: Cu²⁺ + 2e⁻ → Cu (Reduction)

Overall: Zn + Cu²⁺ → Zn²⁺ + Cu

II. Cell Notation & Representation (15 mins)

Standard Cell Notation

Anode on left, cathode on right
Single vertical line (|) represents phase boundary
Double vertical line (||) represents salt bridge
Concentrations specified in parentheses

Example: Daniell Cell

Zn(s) | Zn²⁺(aq, 1M) || Cu²⁺(aq, 1M) | Cu(s)

Electrode Representation

Metal-metal ion electrode: M | Mⁿ⁺
Gas electrode: Pt, Cl₂(g) | Cl⁻
Ion-ion electrode: Pt | Fe²⁺, Fe³⁺

III. Electrode Potential (20 mins)

Concept of Electrode Potential

The tendency of an electrode to lose or gain electrons when in contact with its own ions.

Mⁿ⁺ + ne⁻ ⇌ M

Measured in volts (V) relative to Standard Hydrogen Electrode (SHE)

Standard Hydrogen Electrode (SHE)

Reference electrode with zero potential
Pt electrode coated with platinum black
H₂ gas at 1 atm pressure
H⁺ concentration of 1M at 25°C

2H⁺(aq) + 2e⁻ ⇌ H₂(g) (E° = 0.00 V)

Measuring Electrode Potential

The electrode potential of a half-cell is measured by connecting it to SHE:

E°_cell = E°_cathode - E°_anode

Example: Zn²⁺/Zn Electrode

Zn | Zn²⁺ || H⁺ | H₂, Pt → E°_cell = +0.76 V

Since Zn is anode: E°_Zn = -0.76 V

IV. Calculating EMF (10 mins)

EMF of a Cell

Electromotive force (EMF) is the maximum potential difference between two electrodes.

E°_cell = E°_reduction (cathode) - E°_reduction (anode)

Example Calculation

For Daniell cell: Zn | Zn²⁺(1M) || Cu²⁺(1M) | Cu

E°_Zn²⁺/Zn = -0.76 V, E°_Cu²⁺/Cu = +0.34 V

E°_cell = E°_Cu²⁺/Cu - E°_Zn²⁺/Zn = 0.34 - (-0.76) = 1.10 V

V. Summary & Homework (10 mins)

Key Takeaways

Galvanic cells convert chemical energy to electrical energy
Anode = oxidation, Cathode = reduction
Salt bridge maintains electrical neutrality
Electrode potential measured relative to SHE
EMF = E°_cathode - E°_anode

Homework

1. Represent the galvanic cell where Ni and Ag electrodes are dipped in Ni²⁺ and Ag⁺ solutions
2. Calculate EMF for Mg|Mg²⁺||Al³⁺|Al cell (E°_Mg = -2.37V, E°_Al = -1.66V)
3. Why is platinum used in standard hydrogen electrode?
4. Draw a labeled diagram of Daniell cell

Next Lecture Preview

Standard Electrode Potential & EMF Calculations - Detailed calculations and applications

Materials Needed: Electrochemical series chart, calculator

Class 12 Electrochemistry - Lecture 3: Standard Electrode Potential & EMF

Class 12 Electrochemistry

Lecture 3: Standard Electrode Potential & EMF Calculations

10-Lecture Course Structure

Introduction to Electrochemistry & Redox Reactions

Galvanic Cells & Electrode Potential

Standard Electrode Potential & EMF Calculations

Nernst Equation & Concentration Dependence

Electrolytic Cells & Faraday's Laws

Conductance, Resistivity & Kohlrausch's Law

Batteries (Primary & Secondary) & Fuel Cells

Corrosion: Mechanisms & Prevention

Electrochemical Series & Applications

Advanced Concepts: Concentration Cells, Overpotential

Lecture 3: Standard Electrode Potential & EMF Calculations

Duration: 60 minutes

Objective

Understand standard electrode potential, the electrochemical series, and perform EMF calculations for electrochemical cells.

I. Standard Electrode Potential (15 mins)

The standard electrode potential (E°) is the electrode potential measured under standard conditions (1M concentration, 1 atm pressure, 25°C) relative to the Standard Hydrogen Electrode (SHE).

Measurement Process

Connect the electrode to SHE through a voltmeter
Measure the potential difference when no current flows
Sign indicates the tendency to gain/lose electrons

Example: Zn Electrode

Zn²⁺ + 2e⁻ → Zn E° = -0.76 V

Negative value indicates Zn is a stronger reducing agent than H₂

II. Electrochemical Series (20 mins)

The electrochemical series arranges elements in order of their standard reduction potentials.

Electrochemical Series (Partial)

Electrode Reaction	E° (V)	Tendency
F₂ + 2e⁻ → 2F⁻	+2.87	Strong oxidizing agent
Au³⁺ + 3e⁻ → Au	+1.50
Ag⁺ + e⁻ → Ag	+0.80
Cu²⁺ + 2e⁻ → Cu	+0.34
2H⁺ + 2e⁻ → H₂	0.00	Reference point
Pb²⁺ + 2e⁻ → Pb	-0.13
Sn²⁺ + 2e⁻ → Sn	-0.14
Ni²⁺ + 2e⁻ → Ni	-0.25
Fe²⁺ + 2e⁻ → Fe	-0.44
Zn²⁺ + 2e⁻ → Zn	-0.76	Strong reducing agent
Al³⁺ + 3e⁻ → Al	-1.66
Mg²⁺ + 2e⁻ → Mg	-2.37

Applications of Electrochemical Series

Predicting feasibility of redox reactions
Determining relative strengths of oxidizing/reducing agents
Predicting displacement reactions
Calculating EMF of electrochemical cells

III. EMF Calculations (20 mins)

EMF Calculation Formula

The EMF of a cell is calculated as:

E°_cell = E°_cathode - E°_anode

Where E°_cathode is the reduction potential of the cathode and E°_anode is the reduction potential of the anode.

Example 1: Daniell Cell

Zn | Zn²⁺(1M) || Cu²⁺(1M) | Cu

E°_Zn²⁺/Zn = -0.76 V, E°_Cu²⁺/Cu = +0.34 V

E°_cell = E°_Cu²⁺/Cu - E°_Zn²⁺/Zn = 0.34 - (-0.76) = 1.10 V

Example 2: Silver-Zinc Cell

Zn | Zn²⁺(1M) || Ag⁺(1M) | Ag

E°_Zn²⁺/Zn = -0.76 V, E°_Ag⁺/Ag = +0.80 V

E°_cell = E°_Ag⁺/Ag - E°_Zn²⁺/Zn = 0.80 - (-0.76) = 1.56 V

IV. Predicting Spontaneity (10 mins)

Spontaneity Criteria

A redox reaction is spontaneous if:

E°_cell > 0

This corresponds to a negative ΔG (Gibbs free energy change).

Example: Will Zinc Reduce Copper Ions?

Reaction: Zn + Cu²⁺ → Zn²⁺ + Cu

E°_cell = E°_Cu²⁺/Cu - E°_Zn²⁺/Zn = 0.34 - (-0.76) = 1.10 V > 0

Since E°_cell > 0, the reaction is spontaneous.

V. Summary & Homework (15 mins)

Key Takeaways

Standard electrode potential (E°) measured vs SHE
Electrochemical series arranges elements by E°
EMF = E°_cathode - E°_anode
Positive EMF indicates spontaneous reaction
Higher E° = stronger oxidizing agent

Homework

1. Calculate EMF for: Al | Al³⁺(1M) || Ag⁺(1M) | Ag (E°_Al = -1.66V, E°_Ag = +0.80V)
2. Predict if Fe can reduce Sn²⁺ to Sn (E°_Fe = -0.44V, E°_Sn = -0.14V)
3. Arrange in increasing oxidizing power: Cl₂, F₂, Br₂, I₂
4. Calculate EMF for: Pt | Fe²⁺, Fe³⁺ || Cl⁻ | Cl₂, Pt

Next Lecture Preview

Nernst Equation & Concentration Dependence - Effect of concentration on cell potential

Materials Needed: Calculator, concentration problems worksheet

Class 12 Electrochemistry - Lecture 4: Nernst Equation

Class 12 Electrochemistry

Lecture 4: Nernst Equation & Concentration Dependence

10-Lecture Course Structure

Introduction to Electrochemistry & Redox Reactions

Galvanic Cells & Electrode Potential

Standard Electrode Potential & EMF Calculations

Nernst Equation & Concentration Dependence

Electrolytic Cells & Faraday's Laws

Conductance, Resistivity & Kohlrausch's Law

Batteries (Primary & Secondary) & Fuel Cells

Corrosion: Mechanisms & Prevention

Electrochemical Series & Applications

Advanced Concepts: Concentration Cells, Overpotential

Lecture 4: Nernst Equation & Concentration Dependence

Duration: 60 minutes

Objective

Understand the Nernst equation and how concentration affects electrode potential and cell EMF.

I. Introduction to Nernst Equation (15 mins)

The Nernst equation relates the electrode potential to the concentration of reactants and products in an electrochemical cell.

Derivation

For a general redox reaction:

aA + bB ⇌ cC + dD

The Nernst equation is:

E = E° - \frac{RT}{nF} \ln Q

Where:

E = Electrode potential
E° = Standard electrode potential
R = Gas constant (8.314 J/mol·K)
T = Temperature in Kelvin
n = Number of electrons transferred
F = Faraday's constant (96485 C/mol)
Q = Reaction quotient

Simplified Form at 25°C

E = E° - \frac{0.059}{n} \log Q

II. Application to Half-Cells (15 mins)

General Form for Half-Cell

For a reduction reaction:

M^{n+} + ne^{-} ⇌ M

The Nernst equation becomes:

E = E° - \frac{0.059}{n} \log \frac{1}{[M^{n+}]}

Example: Zinc Electrode

Zn²⁺ + 2e⁻ ⇌ Zn (E° = -0.76 V)

At [Zn²⁺] = 0.1 M, T = 25°C

E = -0.76 - \frac{0.059}{2} \log \frac{1}{[0.1]} = -0.76 - 0.0295 \times 1 = -0.7895 V

III. Application to Full Cells (20 mins)

For Complete Electrochemical Cells

The Nernst equation for a cell reaction:

E_{\text{cell}} = E°_{\text{cell}} - \frac{0.059}{n} \log Q

Example: Daniell Cell

Zn + Cu²⁺ ⇌ Zn²⁺ + Cu (E° = 1.10 V)

At [Zn²⁺] = 0.1 M, [Cu²⁺] = 0.01 M, T = 25°C

Q = \frac{[Zn^{2+}]}{[Cu^{2+}]} = \frac{0.1}{0.01} = 10

E_{\text{cell}} = 1.10 - \frac{0.059}{2} \log 10 = 1.10 - 0.0295 \times 1 = 1.0705 V

IV. Concentration Dependence (10 mins)

Effect on Electrode Potential

For reduction potential: E increases as [ion] increases
For oxidation potential: E increases as [ion] decreases

Concentration vs. Electrode Potential

Concentration (M)

Electrode Potential (V)

As concentration increases, electrode potential increases for reduction reactions.

V. Summary & Homework (10 mins)

Key Takeaways

Nernst equation: $ E = E° - \frac{0.059}{n} \log Q $
Electrode potential depends on concentration
EMF decreases as reaction quotient Q increases
At equilibrium, E_cell = 0 and Q = K

Homework

1. Calculate E for Ag⁺/Ag electrode when [Ag⁺] = 0.01 M (E° = 0.80 V)
2. Find EMF of cell: Zn|Zn²⁺(0.001M)||Cu²⁺(0.1M)|Cu (E° = 1.10 V)
3. Calculate equilibrium constant for Daniell cell at 25°C
4. Predict if reaction is spontaneous: [Zn²⁺]=2M, [Cu²⁺]=0.5M

Next Lecture Preview

Electrolytic Cells & Faraday's Laws - Electrolysis and quantitative aspects

Materials Needed: Calculator, electrolysis problems worksheet

Class 12 Electrochemistry - Lecture 5: Electrolytic Cells & Faraday's Laws

Class 12 Electrochemistry

Lecture 5: Electrolytic Cells & Faraday's Laws

10-Lecture Course Structure

Introduction to Electrochemistry & Redox Reactions

Galvanic Cells & Electrode Potential

Standard Electrode Potential & EMF Calculations

Nernst Equation & Concentration Dependence

Electrolytic Cells & Faraday's Laws

Conductance, Resistivity & Kohlrausch's Law

Batteries (Primary & Secondary) & Fuel Cells

Corrosion: Mechanisms & Prevention

Electrochemical Series & Applications

Advanced Concepts: Concentration Cells, Overpotential

Lecture 5: Electrolytic Cells & Faraday's Laws

Duration: 60 minutes

Objective

Understand the working of electrolytic cells, electrolysis, and Faraday's laws for quantitative calculations.

I. Electrolytic Cells (15 mins)

Electrolytic cells use electrical energy to drive non-spontaneous chemical reactions (electrolysis).

Comparison with Galvanic Cells

Feature	Galvanic Cell	Electrolytic Cell
Energy Conversion	Chemical → Electrical	Electrical → Chemical
Spontaneity	Spontaneous	Non-spontaneous
Anode	Oxidation (negative)	Oxidation (positive)
Cathode	Reduction (positive)	Reduction (negative)
Salt Bridge	Required	Not required

Electrolysis Process

Decomposition of electrolyte by electric current:

Power Source

→

Anode (+)

→

Electrolyte

→

Cathode (-)

II. Electrolysis Examples (15 mins)

Electrolysis of Molten NaCl

Anode: 2Cl⁻ → Cl₂ + 2e⁻
Cathode: 2Na⁺ + 2e⁻ → 2Na
Overall: 2NaCl → 2Na + Cl₂

Electrolysis of Aqueous NaCl

Anode: 2Cl⁻ → Cl₂ + 2e⁻
Cathode: 2H₂O + 2e⁻ → H₂ + 2OH⁻
Products: Cl₂, H₂, NaOH

Factors Affecting Electrolysis

Nature of electrodes
Concentration of electrolyte
Standard electrode potential
Overpotential

III. Faraday's Laws (20 mins)

Faraday's First Law

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

m = Z \times Q = Z \times I \times t

Where:

m = Mass deposited (grams)
Z = Electrochemical equivalent
Q = Charge (coulombs)
I = Current (amperes)
t = Time (seconds)

Faraday's Second Law

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

\frac{m_1}{E_1} = \frac{m_2}{E_2} = \text{constant}

Where E = Equivalent weight = Molar mass / n

Faraday Constant (F)

Charge of 1 mole of electrons (96,485 C/mol)

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

IV. Calculations (10 mins)

Example 1: Copper Deposition

Calculate mass of Cu deposited when 2A current passes through CuSO₄ solution for 30 minutes.

Cu²⁺ + 2e⁻ → Cu (M = 63.5 g/mol, n=2)

Q = I × t = 2 × 30 × 60 = 3600 C
m = (Q × M) / (F × n) = (3600 × 63.5) / (96485 × 2) = 1.18 g

Example 2: Comparing Substances

Same current deposits 2.16g Ag (E=108) in 30 min. What mass of Al (E=9) would be deposited?

m₁/E₁ = m₂/E₂
2.16/108 = m₂/9
m₂ = (2.16 × 9) / 108 = 0.18 g

V. Summary & Homework (10 mins)

Key Takeaways

Electrolytic cells use electricity for non-spontaneous reactions
Anode = oxidation (positive), Cathode = reduction (negative)
Faraday's 1st Law: m ∝ Q
Faraday's 2nd Law: m ∝ E
F = 96,485 C/mol (charge of 1 mole electrons)

Homework

1. Calculate time to deposit 5g Ag from AgNO₃ using 3A current (M=108, n=1)
2. Find current needed to produce 2L Cl₂ at STP in 1 hour from NaCl
3. Why is sodium metal not obtained from aqueous NaCl electrolysis?
4. Calculate mass ratio of Ag, Cu, Al deposited by same charge

Next Lecture Preview

Conductance, Resistivity & Kohlrausch's Law - Electrolytic conduction and its applications

Materials Needed: Conductivity meter diagram, calculator

Class 12 Electrochemistry - Lecture 6: Conductance & Kohlrausch's Law

Class 12 Electrochemistry

Lecture 6: Conductance, Resistivity & Kohlrausch's Law

10-Lecture Course Structure

Introduction to Electrochemistry & Redox Reactions

Galvanic Cells & Electrode Potential

Standard Electrode Potential & EMF Calculations

Nernst Equation & Concentration Dependence

Electrolytic Cells & Faraday's Laws

Conductance, Resistivity & Kohlrausch's Law

Batteries (Primary & Secondary) & Fuel Cells

Corrosion: Mechanisms & Prevention

Electrochemical Series & Applications

Advanced Concepts: Concentration Cells, Overpotential

Lecture 6: Conductance, Resistivity & Kohlrausch's Law

Duration: 60 minutes

Objective

Understand electrolytic conductance, conductivity, resistivity, and Kohlrausch's law for calculating molar conductivity of electrolytes.

I. Basic Concepts (15 mins)

Resistivity (ρ)

Resistance offered by a conductor of unit length and unit cross-sectional area.

\rho = R \times \frac{A}{l}

Where R = resistance, l = length, A = cross-sectional area

Conductivity (κ)

Reciprocal of resistivity. Measure of a material's ability to conduct electric current.

\kappa = \frac{1}{\rho} = \frac{l}{R \times A}

Units: S m⁻¹ (Siemens per meter)

Molar Conductivity (Λₘ)

Conductivity of all ions produced by one mole of electrolyte in solution.

\Lambda_m = \frac{\kappa}{c}

Where c = concentration (mol m⁻³)

Units: S m² mol⁻¹

II. Concentration Dependence (15 mins)

Strong vs Weak Electrolytes

Strong electrolytes: Λₘ decreases slightly with concentration increase

Weak electrolytes: Λₘ decreases sharply with concentration increase

Molar Conductivity vs √c

Λₘ (S m² mol⁻¹)

Dilute

Strong Electrolyte

Medium

Concentrated

Dilute

Weak Electrolyte

Medium

Concentrated

III. Kohlrausch's Law (20 mins)

Statement of Law

The molar conductivity of an electrolyte at infinite dilution can be expressed as the sum of the contributions from its individual ions.

\Lambda_m^\circ = \nu_+ \lambda_+^\circ + \nu_- \lambda_-^\circ

Where:

Λₘ° = molar conductivity at infinite dilution
ν₊, ν₋ = number of cations and anions per formula unit
λ₊°, λ₋° = molar ionic conductivities at infinite dilution

Applications

Calculate Λₘ° for weak electrolytes
Determine dissociation constant of weak electrolytes
Calculate ionic product of water
Determine solubility of sparingly soluble salts

Example: Weak Electrolyte

Calculate Λₘ° for acetic acid (CH₃COOH) given:

λ°(H⁺) = 349.6 S cm² mol⁻¹, λ°(CH₃COO⁻) = 40.9 S cm² mol⁻¹

\Lambda_m^\circ = \lambda_{\text{H}^+}^\circ + \lambda_{\text{CH}_3\text{COO}^-}^\circ

\Lambda_m^\circ = 349.6 + 40.9 = 390.5 \text{S cm}^2 \text{mol}^{-1}

IV. Calculations (10 mins)

Example 1: Strong Electrolyte

Calculate Λₘ° for MgCl₂ given:

λ°(Mg²⁺) = 106.0 S cm² mol⁻¹, λ°(Cl⁻) = 76.3 S cm² mol⁻¹

\Lambda_m^\circ = \lambda_{\text{Mg}^{2+}}^\circ + 2 \times \lambda_{\text{Cl}^-}^\circ

\Lambda_m^\circ = 106.0 + 2 \times 76.3 = 258.6 \text{S cm}^2 \text{mol}^{-1}

Example 2: Weak Electrolyte

Calculate dissociation constant of acetic acid (0.1 M) with Λₘ = 5.2 S cm² mol⁻¹

Λₘ° = 390.5 S cm² mol⁻¹

\alpha = \frac{\Lambda_m}{\Lambda_m^\circ} = \frac{5.2}{390.5} = 0.0133

K_a = \frac{c \alpha^2}{1 - \alpha} \approx c \alpha^2 = 0.1 \times (0.0133)^2 = 1.77 \times 10^{-5}

V. Summary & Homework (10 mins)

Key Takeaways

Conductivity (κ) = 1/resistivity
Molar conductivity (Λₘ) = κ/c
Λₘ decreases with concentration increase
Kohlrausch's Law: Λₘ° = Σνλ°
Applications: weak electrolytes, solubility, dissociation constants

Homework

1. Calculate Λₘ° for Al₂(SO₄)₃ given λ°(Al³⁺)=189 S cm² mol⁻¹, λ°(SO₄²⁻)=160 S cm² mol⁻¹
2. Find Λₘ° for HF if λ°(H⁺)=349.6, λ°(F⁻)=55.4 S cm² mol⁻¹
3. Why does Λₘ for weak electrolytes show sharp decrease with concentration?
4. Calculate K_a for 0.01M HA with Λₘ=15.0 S cm² mol⁻¹ and Λₘ°=380 S cm² mol⁻¹

Next Lecture Preview

Batteries (Primary & Secondary) & Fuel Cells - Working principles of common batteries

Materials Needed: Battery samples, fuel cell diagram

Class 12 Electrochemistry - Lecture 7: Batteries & Fuel Cells

Class 12 Electrochemistry

Lecture 7: Batteries (Primary & Secondary) & Fuel Cells

10-Lecture Course Structure

Introduction to Electrochemistry & Redox Reactions

Galvanic Cells & Electrode Potential

Standard Electrode Potential & EMF Calculations

Nernst Equation & Concentration Dependence

Electrolytic Cells & Faraday's Laws

Conductance, Resistivity & Kohlrausch's Law

Batteries (Primary & Secondary) & Fuel Cells

Corrosion: Mechanisms & Prevention

Electrochemical Series & Applications

Advanced Concepts: Concentration Cells, Overpotential

Lecture 7: Batteries (Primary & Secondary) & Fuel Cells

Duration: 60 minutes

Objective

Understand the working principles, applications, and chemistry of different types of batteries and fuel cells.

I. Primary Batteries (15 mins)

Primary batteries are non-rechargeable batteries that convert chemical energy to electrical energy until reactants are exhausted.

Dry Cell (Leclanché Cell)

Cathode

Anode: Zn → Zn²⁺ + 2e⁻
Cathode: 2MnO₂ + 2NH₄⁺ + 2e⁻ → Mn₂O₃ + 2NH₃ + H₂O
Voltage: 1.5 V

Applications: Flashlights, remote controls, portable devices

Alkaline Battery

Improved version of dry cell with alkaline electrolyte (KOH)

Anode: Zn + 2OH⁻ → ZnO + H₂O + 2e⁻
Cathode: 2MnO₂ + H₂O + 2e⁻ → Mn₂O₃ + 2OH⁻
Voltage: 1.5 V, Longer shelf life

II. Secondary Batteries (15 mins)

Secondary batteries are rechargeable batteries that can undergo multiple discharge-charge cycles.

Lead-Acid Battery

PbO₂

Discharge:
Anode: Pb + SO₄²⁻ → PbSO₄ + 2e⁻
Cathode: PbO₂ + 4H⁺ + SO₄²⁻ + 2e⁻ → PbSO₄ + 2H₂O
Voltage: 2.0 V per cell (6 cells = 12V battery)

Applications: Automobiles, UPS systems

Lithium-Ion Battery

Anode: LiC₆ → Li⁺ + C₆ + e⁻
Cathode: Li₁₋ₓCoO₂ + xLi⁺ + xe⁻ → LiCoO₂
Voltage: 3.7 V, High energy density

Applications: Smartphones, laptops, electric vehicles

III. Fuel Cells (20 mins)

Fuel cells convert chemical energy of fuel directly into electricity through electrochemical reactions.

Fuel (H₂)

→

Anode

→

Electricity

→

Cathode

→

H₂O

Hydrogen-Oxygen Fuel Cell

Anode: 2H₂ + 4OH⁻ → 4H₂O + 4e⁻
Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻
Overall: 2H₂ + O₂ → 2H₂O
Voltage: 0.6-0.8 V per cell

Applications: Space missions, electric vehicles, power plants

Advantages of Fuel Cells

High efficiency (40-60%)
Zero emissions (only water produced)
Continuous operation with fuel supply
Quiet operation

IV. Comparison (10 mins)

Feature	Primary Batteries	Secondary Batteries	Fuel Cells
Rechargeable	No	Yes	N/A (continuous fuel supply)
Energy Density	Moderate	High	Very High
Lifetime	Single use	500-1000 cycles	Continuous with fuel
Applications	Portable devices	EVs, electronics	Spacecraft, power plants
Environmental Impact	High (disposal)	Medium	Low (only water)

V. Summary & Homework (10 mins)

Key Takeaways

Primary batteries: non-rechargeable (dry cell, alkaline)
Secondary batteries: rechargeable (lead-acid, Li-ion)
Fuel cells: continuous electricity from fuel (H₂/O₂)
Each has specific applications based on energy density, cost, and environmental impact

Homework

1. Write electrode reactions for alkaline battery
2. Calculate energy density ratio of Li-ion (150 Wh/kg) vs lead-acid (30 Wh/kg)
3. Why are fuel cells more efficient than heat engines?
4. Compare environmental impact of different battery types

Next Lecture Preview

Corrosion: Mechanisms & Prevention - Understanding rusting and protection methods

Materials Needed: Rust samples, corrosion prevention charts

Class 12 Electrochemistry - Lecture 8: Corrosion

Class 12 Electrochemistry

Lecture 8: Corrosion (Mechanisms & Prevention)

10-Lecture Course Structure

Introduction to Electrochemistry & Redox Reactions

Galvanic Cells & Electrode Potential

Standard Electrode Potential & EMF Calculations

Nernst Equation & Concentration Dependence

Electrolytic Cells & Faraday's Laws

Conductance, Resistivity & Kohlrausch's Law

Batteries (Primary & Secondary) & Fuel Cells

Corrosion: Mechanisms & Prevention

Electrochemical Series & Applications

Advanced Concepts: Concentration Cells, Overpotential

Lecture 8: Corrosion (Mechanisms & Prevention)

Duration: 60 minutes

Objective

Understand the electrochemical mechanisms of corrosion and various prevention techniques.

I. Introduction to Corrosion (10 mins)

Corrosion is the gradual destruction of materials (usually metals) by chemical or electrochemical reaction with their environment.

Economic Impact

Corrosion costs 3-4% of GDP in industrialized nations
Annual global cost: ~$2.5 trillion
Significant safety risks in infrastructure

Electrochemical Nature

Corrosion is an electrochemical process where:

Metal acts as an anode and undergoes oxidation
Oxygen or other species act as cathodes and undergo reduction
Electrolyte (moisture) completes the circuit

II. Mechanism of Rusting (20 mins)

Rusting of iron is the most common form of corrosion:

Anodic Reaction

Fe(s) → Fe²⁺(aq) + 2e⁻ (Oxidation)

Cathodic Reaction

1/2 O₂(g) + H₂O(l) + 2e⁻ → 2OH⁻(aq) (Reduction)

Rust Formation

Fe²⁺

Fe²⁺ ions

→

Combine with OH⁻

→

Fe(OH)₂

Ferrous hydroxide

→

Oxidation by O₂

→

Fe₂O₃

Hydrated ferric oxide (rust)

4Fe(OH)₂ + O₂ → 2Fe₂O₃·H₂O + 2H₂O

III. Factors Affecting Corrosion (15 mins)

Nature of Metal

Position in electrochemical series (more active metals corrode faster)
Purity (impurities create galvanic cells)
Physical state (stressed areas corrode faster)

Environmental Factors

Humidity (critical humidity for iron: ~70%)
Temperature (rate doubles with every 10°C rise)
Presence of electrolytes (salt water accelerates corrosion)

pH Effect

Low pH (acidic): H⁺ reduction accelerates corrosion
High pH (alkaline): Protective oxide layer forms

IV. Prevention Methods (15 mins)

Barrier Protection

Painting, oiling, greasing, or plastic coating

Sacrificial Protection

Using more active metals (Zn, Mg) as anodes

Alloying

Stainless steel (Fe + Cr, Ni)

Cathodic Protection

Impressed current technique

Galvanization

Coating iron with a layer of zinc:

Zn → Zn²⁺ + 2e⁻ (Sacrificial anode)
Fe protected as cathode

Impressed Current Technique

Using an external DC source to make metal a cathode:

DC Power Source

→

Metal becomes cathode

→

Corrosion prevented

V. Summary & Homework (10 mins)

Key Takeaways

Corrosion is an electrochemical process with anodic oxidation and cathodic reduction
Rust formation: Fe → Fe²⁺ → Fe(OH)₂ → Fe₂O₃·H₂O
Factors: Metal activity, impurities, humidity, temperature, pH
Prevention: Barrier methods, sacrificial protection, alloying, cathodic protection

Homework

1. Write the complete reaction sequence for rust formation
2. Calculate corrosion rate if 2.8g iron rusts in 10 days
3. Why does corrosion accelerate in coastal areas?
4. Compare galvanization and tinning as corrosion prevention methods

Next Lecture Preview

Electrochemical Series & Applications - Predicting redox reactions and practical applications

Materials Needed: Electrochemical series chart, metal samples

Class 12 Electrochemistry - Lecture 9: Electrochemical Series & Applications

Class 12 Electrochemistry

Lecture 9: Electrochemical Series & Applications

10-Lecture Course Structure

Introduction to Electrochemistry & Redox Reactions

Galvanic Cells & Electrode Potential

Standard Electrode Potential & EMF Calculations

Nernst Equation & Concentration Dependence

Electrolytic Cells & Faraday's Laws

Conductance, Resistivity & Kohlrausch's Law

Batteries (Primary & Secondary) & Fuel Cells

Corrosion: Mechanisms & Prevention

Electrochemical Series & Applications

Advanced Concepts: Concentration Cells, Overpotential

Lecture 9: Electrochemical Series & Applications

Duration: 60 minutes

Objective

Understand the electrochemical series and its practical applications in predicting redox reactions and corrosion behavior.

I. Introduction to Electrochemical Series (15 mins)

The electrochemical series is a list of elements arranged in order of their standard electrode potentials (E°).

Key Features

Arranged from most negative E° (strong reducing agents) to most positive E° (strong oxidizing agents)
Measured under standard conditions (1M, 25°C, 1 atm)
Based on the standard hydrogen electrode (SHE) as reference (E° = 0V)

Electrode Reaction Standard Potential (E° in volts) Tendency Li⁺ + e⁻ ⇌ Li -3.04 Strong reducing agent K⁺ + e⁻ ⇌ K -2.93 Strong reducing agent Ca²⁺ + 2e⁻ ⇌ Ca -2.87 Strong reducing agent Zn²⁺ + 2e⁻ ⇌ Zn -0.76 Moderate reducing agent Fe²⁺ + 2e⁻ ⇌ Fe -0.44 Moderate reducing agent 2H⁺ + 2e⁻ ⇌ H₂ 0.00 Reference point Cu²⁺ + 2e⁻ ⇌ Cu +0.34 Moderate oxidizing agent Ag⁺ + e⁻ ⇌ Ag +0.80 Strong oxidizing agent Au³⁺ + 3e⁻ ⇌ Au +1.50 Very strong oxidizing agent

The tendency to lose electrons: Li > K > Ca > Na > Mg > Al > Zn > Fe > Ni > Sn > Pb > H > Cu > Ag > Au

II. Predicting Redox Reactions (20 mins)

A metal higher in the series can displace metals lower in the series from their salt solutions.

Displacement Principle

Any reducing agent will reduce the ions of oxidizing agents below it in the series.

Zn (E° = -0.76V)

→

Can displace

→

Cu²⁺ (E° = +0.34V)

Zn + Cu²⁺ → Zn²⁺ + Cu (Spontaneous reaction, E°cell = 0.34 - (-0.76) = 1.10V

Cu (E° = +0.34V)

→

Cannot displace

→

Zn²⁺ (E° = -0.76V)

Predicting Cell EMF

EMF of a cell = E°cathode - E°anode

Example Calculation

Calculate EMF for Zn|Zn²⁺||Cu²⁺|Cu cell:

E°cell = E°Cu²⁺/Cu - E°Zn²⁺/Zn = 0.34 - (-0.76) = 1.10V

III. Practical Applications (15 mins)

Corrosion Prediction

Metals above hydrogen corrode easily; noble metals below hydrogen resist corrosion

Battery Design

Selection of electrode pairs with large potential difference

Electroplating

Coating with metals higher in series provides sacrificial protection

Displacement Reactions

Predicting metal displacement in metallurgy

Environmental Applications

Predicting heavy metal mobility in soil
Wastewater treatment processes
Predicting redox reactions in natural waters

IV. Limitations & Special Cases (10 mins)

Concentration Effects

Actual behavior may differ at non-standard concentrations

Temperature Effects

E° values change with temperature

Passivation

Some metals form protective layers (Al, Cr)

Anomalous Cases

Aluminum (reactive but protected by oxide layer)
Chromium (used in stainless steel despite negative E°)
Hydrogen displacement by some metals (Na reacts violently with water)

V. Summary & Homework (10 mins)

Key Takeaways

Electrochemical series orders elements by standard electrode potential
Higher metals are stronger reducing agents
Applications: corrosion prediction, battery design, electroplating
Limitations: concentration effects, passivation, kinetics

Homework

1. Arrange these metals in order of increasing reducing power: Ag, Mg, Cu, Zn, Na
2. Calculate EMF for Mg|Mg²⁺||Ag⁺|Ag cell (E°Mg = -2.37V, E°Ag = +0.80V)
3. Why can't copper displace zinc from zinc sulfate solution?
4. Explain why aluminum doesn't corrode easily despite its position in the series

Next Lecture Preview

Advanced Concepts: Concentration Cells, Overpotential - Exploring specialized electrochemical systems

Materials Needed: Concentration cell apparatus, voltage measurement tools

Class 12 Electrochemistry - Lecture 10: Concentration Cells & Overpotential

Class 12 Electrochemistry

Lecture 10: Advanced Concepts - Concentration Cells & Overpotential

10-Lecture Course Structure

Introduction to Electrochemistry & Redox Reactions

Galvanic Cells & Electrode Potential

Standard Electrode Potential & EMF Calculations

Nernst Equation & Concentration Dependence

Electrolytic Cells & Faraday's Laws

Conductance, Resistivity & Kohlrausch's Law

Batteries (Primary & Secondary) & Fuel Cells

Corrosion: Mechanisms & Prevention

Electrochemical Series & Applications

Advanced Concepts: Concentration Cells, Overpotential

Lecture 10: Concentration Cells & Overpotential

Duration: 60 minutes

Objective

Understand advanced electrochemical concepts including concentration cells and overpotential, and their practical applications.

I. Concentration Cells (25 mins)

Concentration cells are electrochemical cells where the EMF arises from a difference in concentration of the same electrolyte.

Low Concentration

High Concentration

Electrode Concentration Cells

Anode: M → Mⁿ⁺(dilute) + ne⁻
Cathode: Mⁿ⁺(concentrated) + ne⁻ → M
EMF: E = (RT/nF) ln(C₂/C₁)

Example: Two hydrogen electrodes at different pressures

Electrolyte Concentration Cells

Anode: M → Mⁿ⁺ + ne⁻ (in dilute solution)
Cathode: Mⁿ⁺ + ne⁻ → M (in concentrated solution)

Example: Two silver electrodes in AgNO₃ solutions of different concentrations

Example Calculation

A concentration cell has 0.1M Ag⁺ at anode and 1.0M Ag⁺ at cathode at 25°C. Calculate EMF.

E = (0.059/1) log(1.0/0.1) = 0.059 × 1 = 0.059 V

II. Overpotential (25 mins)

Overpotential (η) is the difference between the actual potential and the theoretical potential required for an electrochemical reaction.

Current-Potential Relationship

Current Density

Overpotential (η)

Tafel Region

Activation Control

Activation Overpotential

Energy barrier for charge transfer at electrode surface

Concentration Overpotential

Mass transport limitations near electrode

Resistance Overpotential

Ohmic losses in electrolyte and connections

Tafel Equation

η = a + b log(i)
Where:
η = Overpotential
a, b = Tafel constants
i = Current density

III. Industrial Applications (10 mins)

pH Measurement

Concentration cells used in pH meters with glass electrodes

E = E° - (0.059) pH

Electrolytic Refining

Overpotential considerations in copper refining:

Prevents hydrogen evolution at cathode
Ensures pure metal deposition

Chlor-Alkali Industry

Overpotential affects electrode reactions in brine electrolysis:

Anode: 2Cl⁻ → Cl₂ + 2e⁻ (instead of O₂ evolution due to overpotential)

Course Completion!

Congratulations on completing the 10-lecture Electrochemistry course

IV. Summary & Homework (10 mins)

Key Takeaways

Concentration cells generate EMF from concentration differences
Overpotential: difference between theoretical and actual potential
Types: Activation, concentration, and resistance overpotential
Applications: pH meters, electrorefining, industrial electrolysis

Homework

1. Calculate EMF for a Cu²⁺ concentration cell with 0.01M and 0.1M solutions
2. Using Tafel equation: η = 0.3 + 0.1 log(i), find η at i=100 mA/cm²
3. Why is overpotential important in chlor-alkali industry?
4. How do concentration cells help in measuring ionic concentrations?