NEET Chemistry Notes Electrochemistry – Electrochemical Cell
An electrochemical cell or simply a cell is a system or arrangement in which two electrodes are fitted in the same electrolyte or in two different electrolytes which are joined by a salt-bridge.
Electrochemical cells are of the following two types :
It is a device in which electrolysis (oxidation and reduction) is carried out by using electricity or in which conversion of electrical energy into chemical energy is done.
Electrolytic Cells and Electrolysis
- Electrolysis is a process in which electrical energy is used to bring some chemical changes. It is carried out in an electrolytic cell which involves conversion of electrical energy to chemical energy
where,= standard oxidation potential, i.e. at standard conditions (1 atm, 298 K and 1M) tendency to lose electrons.
= standard reduction potential, i.e. at standard conditions (1 atm, 25°C and 1 M) tendency to gain electrons.
- Decreasing order of oxidation potential
- Increasing order of reduction potential All metals, more reactive than hydrogen
- During electrplysis when two or more ions complete at the electrodes, the ion with higher reduction potential gets liberated at cathode while the one with lower reduction potential at the anode.
- Besides the ions of electrolyte, if some other ions (cations or anions) are present in the solution, then which of the two or more ions gets discharged at each electrode depends upon their relative discharge potential. Usually ions with lower discharge potential
are discharged in preference to those which have high discharge potential.
- For aqueous solution of salt:
(a) If metal is less reactive (like Ag, Cu) than hydrogen, metal will be deposited at cathode.
(b) If metal is more reactive than hydrogen, 2 gas will be liberated at cathode.
In aqueous solution containing any of the cation Li+, Na + , Ba2+, Ca2+, Mg2 or Al3+, it is water which is reduced at cathode and not the metal cations. In aqueous solution of the , anions are not oxidised, it is water which is oxidised.
Faraday’s Laws of Electrolysis
The quantitative relationships based on the electrochemical researches published by Faraday. These two laws given by the Faraday are given below:
Deposited mass of the substance is directly proportional to the charge passed in a voltameter.
where, w = mass,
Q = charge (in coulomb)
i = current (in amperes)
t = time (in second)
Z = electrochemical equivalent
One Faraday or 96500 C or 1 mole of electrons cause the reduction of 1 mole of monovalent cation or 1/2 mole of divalent cation or 1/3 mole of trivalent cation.
The number of equivalents of any substance produced by a given quantity of electricity during electrolysis are same
wA = deposited mass of substance A
,Ea = equivalent weight of A
wB = deposited mass of substance B
,EB = equivalent weight of B
Galvanic or Voltaic Cells
It is a device in which a redox reaction used to convert chemical energy into electrical energy, nahr
The two types are therefore the reverse of each other.
Electrode and Half-Cell
- When used in electrochemical studies, a strip of metal, M used is called electrode. The metal strip is immersed in a solution containing the metal ion Mn+. The combination of the metal electrode and sohhion is called a half-cell.
Three kinds of interactions are possible between metal atom on the electrode and metal ion in solution.
- A metal ion Mn+may collide with the electrode and undergo no change.
- A metal ion Mn+ may collide with the electrode, gain n electrons and be converted to a metal atom M. The ion is reduced.
- A metal atom M on the electrode may lose n electrons and enter the solution as ion Mn+. The metal atom is oxidised.
It is a U-shaped tube contains a gel permeated with a solution of an inert electrolyte such as Na2S04. The ions of the inert electrolyte do not react with the other ions in the solutions and they are not oxidised or reduced at the electrodes. The salt-bridge is necessary to complete the electrical circuit and to maintain electrical neutrality in both compartments (by flow of ions).