Electrochemistry Class 12 Notes: Electrochemical cells, Nernst Equation, and Electrolysis

As per NCERT, “Electrochemistry is the study of the production of electricity from energy released during spontaneous chemical reactions and the use of electrical energy to bring about non-spontaneous chemical transformations”.

Simply explaining this:

Electrochemistry is the study of how chemical reactions can produce electricity or electrical energy, and this energy can drive chemical reactions that wouldn’t happen on their own.

For example, consider your phone's battery. When you use your phone, chemical reactions inside convert stored chemical energy into electricity. When you charge it, electricity is being used to reverse those chemical reactions and store energy again. 

This conversion of chemical reactions into electricity and electricity to chemical energy explains what electrochemistry is all about.

An electrochemical cell is a device that either produces electricity from chemical reactions or uses electricity to make chemical reactions. 

Every electrochemical cell has these essential parts:

• Two conducting surfaces where reactions occur called electrode
• A solution containing ions that can carry an electrical current called electrolyte
• Wires that allow electrons to flow between electrodes called external circuit.

Types of Electrochemical Cells

Electrochemical cells are of two types: a Galvanic cell and an Electrolytic cell.

Galvanic Cells

Galvanic cells convert chemical energy into electrical energy through spontaneous reactions.

Example: Your car battery, flashlight batteries

Electrolytic Cells

Electrolytic cells use external electrical energy to force non-spontaneous chemical reactions.

Example: Metal plating, water splitting for hydrogen production

As stated earlier, Galvanic cells are the generators of electricity. They convert chemical energy into electrical energy.

The best example of a Galvanic cell is a Daniell cell, which covert chemical energy into electrical energy through a redox reaction.

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

This reaction occurs in two half-reactions:

At the anode, Zn(s) → Zn²⁺(aq) + 2e⁻ (oxidation), and 

at the cathode, Cu²⁺(aq) + 2e⁻ → Cu(s) (reduction).

Electrode Potential:

Every electrode develops a potential or voltage when placed in an electrolyte solution. Because

• Metal atoms may dissolve as ions, leaving electrons on the electrode
• Metal ions from the solution may deposit on the electrode
• An equilibrium forms, creating a charge separation

Anode vs Cathode:

Anode is the part of the cell where oxidation happens or electrons are lost.

• Anode has a negative potential in galvanic cells
• Electrons flow out of the anode

 The cathode is the other half of the cell where reduction happens or electrons are gained.

• The cathode has a positive potential in galvanic cells
• Electrons flow into the cathode

Standard Electrode Potentials (E°)

Since we can’t measure the potential or voltage of a single electrode, scientists created a reference system using the Standard Hydrogen Electrode (SHE).

• Reference potential: 0.00 V (by definition)
• Conditions: 1 M H⁺ ions, 1 bar H₂ gas, 25°C
• Reaction: 2H⁺(aq) + 2e⁻ → H₂(g)

Using the Standard Hydrogen Electrode (SHE) as reference:

• Positive E°: Species gets reduced more easily than H⁺
• Negative E°: Species gets reduced less easily than H⁺

For any galvanic cell:

E°cell = E°cathode - E°anode

Example: For the Daniell cell

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

The Nernst equation explains how cell potential (voltage) changes when conditions aren’t standard. This equation was developed by Walther Nernst, responsible for the concentration effects on cell potential.

Nernst Equation for a general reaction:

E(cell) = E°(cell) - (RT/nF) ln Q

Where:

• R: Gas constant
• T: Temperature in Kelvin
• n: Number of electrons transferred
• F: Faraday constant 
• Q: Reaction quotient

Simplified Form of Nernst equation at 25°C:

At 298 K or 25°C, the equation becomes:

E(cell) = E°(cell) - (0.059/n) log Q

Equilibrium Constant from Nernst Equation

When a cell reaches equilibrium, E(cell) becomes 0 means no net current flows, and Q = K (reaction quotient becomes equal to the equilibrium constant.

This gives us: E°(cell) = (0.059/n) log K

This helps calculate equilibrium constants from standard potentials.

Connection to Gibbs Energy

The Nernst equation connects electrochemistry to thermodynamics:

ΔG = -nFE(cell)

For standard conditions:

ΔG° = -nFE°(cell)

This means:

• Positive E°: Negative ΔG° 
• Negative E°: Positive ΔG°

Electricity flows through solutions through ions. In metals, electrons carry current. But in solutions, the ions carry the current.

Resistance and Conductance

• Resistance (R): Opposition to current flow (measured in ohms, Ω)
• Conductance (G): Ability to conduct current = 1/R (measured in siemens, S)

Resistivity and Conductivity

Just like how different materials have different heat resistance, they also vary in electrical resistance.

Resistance is directly proportional to the length and inversely proportional to the area of cross-section.

R = ρ(l/A)

Where, 

• ρ: Resistivity 
• l: Length of conductor
• A: Cross-sectional area
• Units: Ω⋅m

Conductance is the opposite of resistance

G= κ(A/l)

Where,

• κ: Conductivity

The inverse of the resistivity is called conductivity.

Conductivity (κ): κ = 1/ρ

• Units: S⋅m⁻¹ or S⋅cm⁻¹

Measurement of Conductivity of Ionic Solutions

We can measure the conductivity of ionic solutions through a conductivity cell.

Conductivity Cell

The conductivity cell contains two platinum electrodes with area A, separated by a distance l.

So the resistance will be:

R = ρ(l/A)

R = l/κA

Here,

• Cell constant: G* = l/A = Rκ
• Measurement: κ = G*/R

Molar Conductivity

Molar conductivity (Λm) is used to identify how well one mole of an electrolyte conducts electricity.

Λm = κ/c

Where.

• κ: Conductivity
• c: Molar concentration
• Units: S⋅cm²⋅mol⁻¹

Variation of Conductivity and Molar Conductivity with Concentration

Strong Electrolytes

In strong electrolytes, molar conductivity (Λm) decreases slowly with increasing concentration because of complete dissociation, but ions interfere with each other at higher concentrations

Kohlrausch equation: Λm = Λ°m - A√c

Weak Electrolytes

In weak electrolytes, molar conductivity (Λm) decreases rapidly with increasing concentration because the degree of dissociation decreases with concentration.

Need Kohlrausch law for Λ°m

Kohlrausch Law of Independent Migration

This law states that the limiting molar conductivity of an electrolyte equals the sum of individual ionic contributions.

Λ°m = n₊λ°₊ + n₋λ°₋

Where:

• λ°₊, λ°₋: Individual ionic conductivities
• n₊, n₋: Number of cations and anions

Applications of the Kohlrausch Law

• Calculate Λ°m for weak electrolytes
• Find degree of dissociation: α = Λm/Λ°m
• Calculate dissociation constants: K = α²c/(1-α)

While galvanic cells generate electricity from chemical reactions, electrolytic cells do the opposite; they use electricity to drive chemical reactions that wouldn’t happen on their own.

In electrolytic cells, we force electrons to flow in the opposite direction.

Faraday’s Laws of Electrolysis

Michael Faraday discovered quantitative relationships governing electrolysis. Faraday’s two laws of electrolysis are:

1. The amount of chemical change is directly proportional to the quantity of electricity passed.

Chemical change ∝ Charge passed

2. Different substances liberated by the same quantity of electricity are proportional to their chemical equivalent weights.

Amount liberated ∝ (Atomic mass)/(Charge on ion)

Quantitative Relationships:

• Faraday constant (F): 96,487 C/mol (charge on 1 mole of electrons)
• For practical calculations: F ≈ 96,500 C/mol

Example: To deposit 1 mole of Cu from Cu²⁺:

• Electrons needed: 2 moles
• Charge required: 2F = 2 × 96,500 = 193,000 C

Products of Electrolysis

What gets produced during electrolysis depends on several factors:

Competing Reactions: When multiple species can be oxidized or reduced, the one with the more favorable potential wins.

At cathode (reduction):

• Most positive E° gets reduced first

Example: In aqueous NaCl, H₂O gets reduced instead of Na⁺ because H₂O reduction has a higher (less negative) potential

At anode (oxidation):

• The most negative E° gets oxidized first

Some reactions need extra voltage to occur at reasonable rates

For Example,

1. Electrolysis of molten NaCl:

• Cathode: Na⁺ + e⁻ → Na (sodium metal)
• Anode: Cl⁻ → ½Cl₂ + e⁻ (chlorine gas)

2. Electrolysis of aqueous NaCl:

• Cathode: 2H₂O + 2e⁻ → H₂ + 2OH⁻ (hydrogen gas)
• Anode: 2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas)

Net result: NaOH solution, H₂ gas, and Cl₂ gas

Read more: Electrolytic Cell and Electrolysis

Batteries are portable galvanic cells that store chemical energy and convert it to electrical energy when needed. They’re everywhere in our world, from digital watches to car batteries.

There are two kinds of batteries: Primary Batteries and Secondary Batteries.

Primary Batteries

These are like single-use items; once the chemical reactions are complete, they’re done, you can’t recharge them.

Characteristics of primary batteries:

• Single-use: Cannot be recharged
• Convenient: Ready to use immediately
• Limited life: Reactions eventually stop
• Examples: Alkaline batteries, dry cells

Secondary Batteries

These are the rechargeable batteries that can be used multiple times.

Characteristics of secondary batteries:

• Rechargeable: Can be reused with external electricity
• Reusable: Multiple charge-discharge cycles
• Higher initial cost, but economical, as it has a long life
• Examples: Car batteries, phone batteries