# Electrochemistry Class 12 Notes Chemistry Chapter 3 - CBSE

## Standard Electrode Potential

The standard electrode potential (E0 half-cell) is the potential of a given half-reaction when all components are in their standard states, i.e., solutions are at 1M concentrations, gases at 1 atm pressure and solids and liquids are in pure form with all at 25°C.
By convention, all standard electrode potentials refer to the half-reaction written as a reduction. The standard cell potential equals to the difference between the abilities of the two electrodes to act as reducing agents.
E0 cell = E0 cathode (reduction) – E0 anode (oxidation)

## Relation Between Gibbs Energy Change And Emf Of A Cell

The maximum cell potential is directly related to the free energy difference between the reactants and the products in the cell.

𐤃G° = –nFε°cell

• Volts = work (J)/charge (C)
• Number of moles transferred per mole of rectant and product
• Faraday = 96,485 coulombs of charge per mole of electrons
 ε° ΔG° Spontaneity Positive Negative spontaneous reaction Negative Positive non-spontaneous reaction Zero Zero reaction is at equilibrium

## Emf Of A Cell

• The cell potential or emf of cell is the difference between electrode (reduction potentials) of the cathode and anode. When there is no current in cell.
• Ecell = Ecathode – Eanode = Eright – Eleft

## Nernst Equation

Ecell = E0cell RT/nF – ln(Q) = E0cell(-0.0592/n)log10 (Q) (at 25°C) because ln(Q) = 2.303 log10 (Q) R = 8.31 J/(mol.K) T = Temperature in K, Q = Reaction Quotient n = moles of electrons in balanced redox equation F = Faraday constant = 96,485 coulombs/mol e

## Molar Conductance

The molar conductance is defined as the conductance of all the ions produced by ionization of 1 g mole of an electrolyte when present in V mL of solution. It is denoted by λm. λm =λ × 1000/c Where λm is molar conductance in S cm2 mol–1, λ is specific conductance in S cm–1 and c is concentration in mol L–1.

## Redox Reactions ## Kohlrausch’s Law

Statement : “At time infinite dilution, the molar conductance of an electrolyte can be expressed as the sum of the contributions from its individual ions.” i.e., 𐤃m = v+λ + + vλ where, v+ and v are the number of cation and anion per formula unit of electrolyte respectively and, λ + and λ are the molar conductivities of the cation and anion at infinite dilution respectively.

## Application Of Nernst Equation To Chemical Cell

Nernst equation can be used to calculate:
• Single electrode reduction or oxidation potential at any conditions.
• Standard electrode potentials.
• EMF of an electrochemical cell.
• Unknown ionic Concentration.
• pH of solutions and solubility of sparingly soluble salts.

## Conductance In Electrolytic Solutions

The ability of electolytic solutions ot allow the passage of electric current through them is known as electrolytic conductance. Note : Electrolytes are those substances that dissolve in a solvent and dissociate into charged ions (cations) (+ve), anions (–ve). The electolytes can conduct electricity only in molten or aqueous state and not in any solid form. eg-KNO3, NaCl, KCl etc.

## Factors Affecting Electrolytic Conductance

• Concentration of ion : electrolytic conductance ∞ concentration of ions.
• Temperature : Electrolytic conductance ∞ Temperature

## Variations Of Conductivity With Concentration

Conductivity of an electrolyte decreases with decrease in concentration (both for weak and strong electrolyte), whereas molar conductivity increases with decrease in concentration.

## Electrolysis And Law Of Electr Olysis

It is a type of electric battery, used for home and portable elecrtronic devices because it consists of low moisture inmobilized electrolytes in form of a paste.

• It is an electrochemical cell devoloped by Cast Gassner in 1886.
• It is an electrochemical cell devoloped by Cast Gassner in 1886.

Depending on nature of dry cell, it can be classified as a primay cell and secondary cell.

• Primary Cell: These cells are neither reusable nor rechargeable. E.g. Zinc Carbon Cell, Mercury cell.
• Secondary Cell: These cells are rechargeable and reusable. E.g. Ni-cd Cell, lithium-ion cell.

## Dry Cell

The process of carrying out non-spontaneous reactions under the influence of electric engergy is termed as electrolysis.

• Faraday’s First law : During electrolysis, the amount of chemical reaction (which occures at any electrode under the influence of electrical energy) is proportional to Quantity of electricity of electricity passed through electrolyte.
• Faraday’s Second law : During electrolysis, a number of different substances liberated (at some quantity of electricity) are proportional to their chemical equivalent weights. Equivalent weight = atomic mass of metal / number of e– required for reducing the cation.

## Electrolytic Cell

It is a device through which conversion of electrical energy in to chemical energy and vice versa takes place. The cell reaction is non-spontaneous.

## Galvanic Cell Or Voltaic Cell

• The chemical energy is converted to electrical energy.
• The cell reaction is spontaneous.
• Oxidation and reduction reactions takes place simultaneously.

Note: When a secondary cell is discharging, it acts as a galvanic cell and when a secondary cell is recharging, it acts as an electrolytic cell.

• It is a secondary cell because electrical energy is not generated within the cell itself but it is previously stored in it from external source.
• Energy stored in the form of chemical energy so the cell is a storage cell or accumulator or storage battery.
• It consists of lead plates as negative electrode and the lead plates are impregnated with lead oxide which act as positive electrode.
• Emf of cell = 2.041V

## Fuel Cell

An electrochemical cell that converts chemical energy of a fuel (often hydrogen) with oxidizing agent (often oxygen) into electrical energy through a pair of redox reactions, is known as fuel cell.

## Corrosion

It is an electrochemical process that causes transformation of pure metals into undesirable substances when they react with substances like water or air.

E.g. Tarnishing of silver : 2Ag(s) + H2S(g) → Ag2S(s) + H2(g)

Rusting of iron : Fe2+ + 3O2 → 2Fe2O3

Fe2O3 + H2O → Fe2O3 . × H2O

(rust)