Revision: Electrochemistry Chemistry Science (English Medium) Class 12 CBSE

Electrochemistry is the study of the production of electricity from energy which is released during spontaneous chemical reactions, as well as the use of electrical energy to bring about non-spontaneous chemical transformations.

The electrode at which the reduction occur is called cathode.

The electrode at which the oxidation occur is called anode.

Fuel cells are the galvanic cells in which the energy of combustion of fuels like hydrogen, methanol, etc., is directly converted into electrical energy.

Electrochemistry is the branch of chemistry that deals with the production of electricity from the energy released during spontaneous reactions and the use of electrical energy to drive non-spontaneous reactions.

Conductivity is the inverse of resistance R and may be simply defined as the speed at which current flows in a conductor.

The limiting molar conductivity of an electrolyte is defined as its molar conductivity when the concentration of the electrolyte in the solution approaches zero.

When the concentration of an electrolytic solution placed between electrodes of a conductivity cell placed at a unit distance having an area of cross-section sufficient to accommodate enough volume of solution containing one mole of electrolyte approaches zero, then the conductance of the solution is known as limiting molar conductivity.

Molar conductivity is the conductance of a volume of solution containing 1 mole of dissolved electrolyte when placed between two parallel electrodes 1 cm apart and large enough to contain between them all the solution.

A fuel cell is a galvanic cell in which the reactants are not placed within the cell, but are continuously supplied from outside, where one reactant acts as a fuel (such as hydrogen or methanol) and the other as an oxidant (such as oxygen).

Corrosion is the gradual damage of metals caused by their reaction with components of the atmosphere, such as oxygen and moisture.

The reciprocal of resistance is called conductance.

The conductance of a solution of unit length and unit cross-section is called conductivity.

The ratio of distance between electrodes to area of cross-section is called cell constant.

The resistance of a conductor of unit length and unit cross-sectional area is called resistivity.

An electrochemical cell in which electrical energy is used to bring about a non-spontaneous chemical reaction is called an electrolytic cell.

A cell in which the chemical reaction occurs only once and cannot be reversed is called a primary cell.

A cell in which the chemical reaction can be reversed by passing current in opposite direction is called a secondary cell.

A galvanic cell designed to convert the energy of combustion of fuels directly into electrical energy is called a fuel cell.

An electrochemical cell that converts chemical energy of a spontaneous redox reaction into electrical energy is called a galvanic cell.

The potential difference developed between an electrode and its electrolyte is called electrode potential.

The electrode potential measured under standard conditions (1 M, 1 bar, 298 K) is called standard electrode potential.

The reference electrode assigned zero potential at all temperatures is called the standard hydrogen electrode.

The equation which relates electrode potential with concentration of ions is called the Nernst equation.

The ratio of product concentration to reactant concentration at equilibrium is called equilibrium constant.

The thermodynamic quantity representing maximum obtainable work from a reaction is called Gibbs free energy.

\[E_{cell}=E_{cathode}-E_{anode}\]

\[E_{cell}^\circ=E_{cathode}^\circ-E_{anode}^\circ\]

\[R=\rho\frac{l}{A}\]

\[G=\frac{1}{R}\]

\[\kappa=\frac{1}{\rho}\]

\[\kappa=\frac{G^*}{R}\]

\[G^*=\frac{l}{A}\]

\[R_2=\frac{R_1R_4}{R_3}\]

For reaction:

aA + bB → cC + dD

\[E_{cell}=E_{cell}^\circ-\frac{RT}{nF}\ln\frac{[C]^c[D]^d}{[A]^a[B]^b}\]

\[\Lambda_m=\frac{\kappa}{C}\]

\[\Lambda_m=\kappa\times\frac{1000}{M}\]

Unit relation:

\[1Sm^2mol^{-1}=10^4Scm^2mol^{-1}\]

\[\Lambda_m=\Lambda_m^\circ-A\sqrt{C}\]

\[\alpha=\frac{\Lambda_m}{\Lambda_m^\circ}\]

Kohlrausch law states that at infinite dilution of the solution, each ion of electrolyte migrates independently of its co-ions and contribute independently to the total molar conductivity irrespective of the nature of other ion.

Kohlrausch’s law states that the molar conductivity of an electrolyte at infinite dilution is the same as the sum of the anions' and cations' limited molar conductivities.

`∧_m^° = v_+  λ_+^° + v_-  λ_-^°`

Here `λ_+^°` and `λ_-^°` are limiting molar conductivities of cations and anions.

Kohlrausch’s Law states that at infinite dilution, each ion contributes independently to the total molar conductivity of an electrolyte, and the limiting molar conductivity is equal to the sum of individual ionic conductivities.

Mathematically,

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

where λ^∘₊ and λ^∘₋​ are limiting molar conductivities of cation and anion respectively.

Electrode potential varies with concentration and temperature.

\[E=E^\circ-\frac{RT}{nF}\ln Q\]

At 298 K:

\[E=E^\circ-\frac{0.059}{n}\log Q\]

Faraday’s Second Law of Electrolysis states that when the same quantity of electricity is passed through different electrolytes, the masses of substances deposited are proportional to their chemical equivalent weights.

Mathematically,

\[\frac{m_1}{m_2}=\frac{E_1}{E_2}\]

where m is mass deposited and E is equivalent weight.

Faraday’s First Law of Electrolysis states that the mass of a substance deposited or liberated at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte.

Mathematically,

m ∝ Q

m = ZQ

where m is mass deposited, Q is charge passed, and Z is electrochemical equivalent.

Electrolysis of NaCl

1. Molten NaCl:

• Oxidation: Cl⁻ → Cl₂ (gas)

• Reduction: Na⁺ → Na (metal)

• Products: Na (cathode), Cl₂ (anode)

2. Aqueous NaCl:

• Oxidation: Cl⁻ → Cl₂

• Reduction: H₂O → H₂ + OH⁻

• Products: H₂ (cathode), Cl₂ (anode), NaOH formed

Components of a Galvanic Cell

6. Cell Notation 

• Anode written on the left, cathode on the right

• Example:

  Cu(s) | Cu²⁺(aq) || Ag⁺(aq) | Ag(s)

• Single line (|) → phase boundary

• Double line (||) → salt bridge

The Nernst equation is used to calculate the electrode or cell potential under non-standard conditions.

\[E_{cell}=E_{cell}^\circ-\frac{RT}{nF}\ln Q\]

At 25°C, it becomes:

\[E_{cell}=E_{cell}^\circ-\frac{0.0591}{n}\log Q\]

Where E°cell is the standard cell potential, n is the number of electrons transferred, and Q is the reaction quotient.

The equation helps determine the direction and spontaneity of a reaction:

• Ecell > 0 → spontaneous
• Ecell = 0 → equilibrium (Q = K)

It also relates to Gibbs energy:

ΔG = −nFE_cell

Thus, the Nernst equation is important for electrochemical calculations and equilibrium analysis.

Electrical conductance and resistance:

\[\mathrm{K}=\mathrm{G}\frac{l}{A}\]

K = Conductivity

G = Conductance

\[\mathrm{G}=\mathrm{}\frac{1}{R}\]

R = Resistance

\[\mathrm{K}=\mathrm{}\frac{l}{RA}\]

Reactions

• Anode:
  2H₂ + 4OH⁻ → 4H₂O + 4e⁻
• Cathode:
  O₂ + 4H₂O + 4e⁻ → 4OH⁻
• Overall reaction:
  2H₂ + O₂ → 2H₂O

Applications

• Spacecraft (electric power)
• Power generators (homes, hospitals)
• Automobiles (experimental)
• Clean energy for industries

Drawbacks

• Hydrogen gas is hazardous
• High cost of hydrogen preparation