Revision: Std XII >> Electrochemistry MAH-MHT CET (PCM/PCB) Electrochemistry

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.

Electrolytic conductors can conduct current with the mobility of ions. This process is known as electrolytic conduction.

Metallic or electronic conductors can conduct current by the transfer of free electrons.

In a conductivity cell, the distance l between the two electrodes and the area A of the electrodes are fixed. Therefore, the quantity `l/A` is a constant for a particular A conductivity cell. This quantity is termed as cell constant.

The reciprocal of the electrical resistance is called the conductance.

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

The unit of conductance is ohm⁻¹ or mho and is denoted by Ω⁻¹. In SI, unit is S (Siemen).

The electrode at which the oxidation occur is called anode.

The electrode at which the reduction occur is called cathode.

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

Standard electrode potential is the difference of electrical potential between a metal electrode and the solution around it at equilibrium when all the substances involved are in their standard states.

The potential of an electrode assembly is referred to as the standard electrode potential when the following conditions are satisfied.

1. The temperature of the electrode assembly is 298 K (25°C).
2. The ion solution used in the assembly is of concentration 1 mol L⁻¹.
3. The pressure of the gas, if used in the assembly, is 1 atm.

The potential difference developed between the electrode and the electrolyte due to the loss or gain of electrons by the electrode is called the electrode potential.

Oxidation potential: M(s) ⇌ Mⁿ⁺(aq) + ne⁻ Reduction potential: Mⁿ⁺(aq) + ne⁻ ⇌ M(s) Reduction potential = − Oxidation potential

The cell potential of the cell is the algebraic sum of the electrode potentials. Cell potential is called the electromotive force (e.m.f.) of the cell when no current is drawn through the cell.

Standard electrode potential is the potential associated with the electrode reaction at an electrode when all solutes are 1 M, and all gases are at 1 atm and at 25°C (298 K).

It is the cell potential when the concentrations of all the species are 1 M at 25°C, and the pressure of the gas involved is 1 atm at 25°C.

It is an electrode whose potential is arbitrarily taken as zero or is exactly known. Standard Hydrogen Electrode (SHE), calomel electrode, silver-silver chloride electrode and glass electrode are some examples of reference electrode.

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).

The standard potentials of a number of electrodes have been determined using standard hydrogen electrodes. These electrodes with their half reactions are arranged according to their decreasing standard potentials; this arrangement is called an electrochemical series.

Based on the ease with which atoms of metals lose electrons to form positively charged ions, the metals are arranged in a series known as the electrochemical series.

or

The arrangement of electrodes in order of their decreasing standard reduction potentials is called electrochemical series.

At infinite dilution, each ion migrates independently of the co-ion and contributes to the total molar conductivity of an electrolyte, irrespective of the nature of the other ion to which it is associated.

Degree of dissociation = `"Molar conductance at a given concentration"/ "Molar conductance at infinite dilution"` \[=\frac{\Lambda_{m}^{c}}{\Lambda_{m}^{\infty}}\]

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

1. Gibbs Energy Relation
ΔG = −nFEcell

2. Relation with Equilibrium Constant
E°cell \[=\frac{0.0592}{n}\log_{10}K  at   25^{\circ}C\]

Reference Electrodes

Standard Hydrogen Electrode (SHE)

Types of Voltaic Cells

Important Cells

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