What are Electrochemistry ?

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.

Lead Accumulator

  • 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.


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