Understanding Electrolysis: The conductors are of two types: Metallic Conductors and Electrolytic Conductors. Examples of metallic conductors are copper, silver, aluminum etc. The conduction through them is due to the flow of electrons. Electrolytic Conductors conduct only when they are in a molten or aqueous state. When electrolytic conductors melt to molten or aqueous state there is a formation of ions. As soon as we connect to battery their movement starts. The conduction through the solution is due to the movement of ions. Ions are the carrier of current through solutions. On passing electric current these ions move towards oppositely charged electrodes and thus carry the current. Electrolytic conductance increases with the rise in temperature. Examples are electrolytic solution of acids, bases, and salts.

Must Read: Electrolytes and its Types

Define Electrolysis

The process of chemical decomposition of an electrolyte (in the dissolved or molten state) by the passage of electric current is known as electrolysis.

For instance, when an aqueous solution of hydrogen chloride is subjected to electrolysis, hydrogen chloride gets decomposed to form hydrogen and chlorine gases.

Electrolysis Process

The process of electrolysis is carried out in a vessel known as an electrolytic tank or electrolytic cell. The tank which contains either the fused electrolyte or aqueous solution of the electrolyte. Two metallic plates or rods known as electrodes are suspended into the electrolyte. One electrode is connected to a positive terminal and is called anode. The other electrode is connected to the negative terminal and called cathode.

According to Scientist Arrhenius, when an electrolyte is fused or dissolved in water, it dissociates into positively and negatively charged particles called ions.These ions move about here and there in the solution.

Mechanism of Electrolysis

The mechanism starts on passing electric current the ions start moving towards oppositely charged electrodes. The positively charged ions move towards cathode and are called cations while negatively charged ions move towards the anode and are called anions. The movement of cations and anions is responsible for the electrical conductivity of the solutions of electrolytes.Electrolysis

In addition to electrical conductivity, the movement of the ions also causes transport of matter from one part of the system to another. As the ions reach the respective electrodes, a chemical reaction takes place. Such a reaction that takes place at the electrode between the electrolyte and electrode is called electrochemical reaction.  A few electrochemical reactions which may take place at cathode and anode are mentioned below.

(1) Reactions at Cathode. Reaction taking place at cathode is always reduction, which simply involves the gain of electrons by the cation, e.g.,


(2) Reactions at anode. Reaction taking place at anode is always oxidation which simply involves the loss of electrons by a substance e.g.,


Let us explain the electrolysis by taking examples of molten sodium chloride and aqueous solutions of sodium chloride and copper sulfate as given below.

Electrolysis of Molten Sodium Chloride

Molten sodium chloride contains free moving Na+ and Cl  ions.

NaCl ⇌ Na+ + Cl

On passing the electric current through the molten sodium chloride, Na+ ions migrate towards the cathode while Cl  ions migrate towards the anode. The following reactions take place as soon as the ions reach their respective electrodes.

  • At the cathode:  2Na+ +2e  ⇌ 2Na
  • At the anode: 2CI– → 2Cl +2e
  • 2Cl → Cl2

Therefore, the overall reaction is:  2Na+Cl–  → 2Na+ Cl2

Thus, sodium metal is deposited at the cathode while chlorine gas is liberated at the anode.

Electrolysis of Aqueous Sodium Chloride Solution

An aqueous solution of sodium chloride contains a large number of sodium ions and chloride ions along with some H+ ions and OH, ions (which come from the dissociation of water). Thus, there are two types of cations, i.e., Na+ and H+ ions along with two types of anions, i.e., CI and OH ions in the aqueous solution of sodium chloride.

On passing an electric current through the aqueous solution of sodium chloride, the following reactions take place:

At the cathode. Both Na+ and H+ ions are attracted by the cathode but neither is actually deposited. On the other hand, water molecules react with the electrons available at the cathode producing hydrogen gas as:

2H2O +2e  → H2+2OH

It is because the decomposition potential of water is lower than that of Na+ or H+ both.

The discharge potential is defined as the minimum potential that must be applied across the electrodes to bring about the electrolysis and discharge of the ion at the electrode

At the anode. Both CI and OH ions are attracted to the anode but it is only Cl ions which get deposited there as:

2Cl+  2e  →  2Cl2

Therefore, the overall reaction is:

2H2O + 2Cl →  H2 + Cl2 +2OH

If we include the Na+ ions by adding them to both sides, the overall reaction is:

2Na+2H2O →  H2 + Cl2  +2Na+ +2OH

Since both hydrogen gas and chlorine gas get liberated at their respective electrodes, the residual solution becomes rich in Na+ and OH, ions. Consequently, the residual solution becomes alkaline. Such a solution on evaporation yields solid sodium hydroxide.

Electrolysis of Copper Sulphate Solution using Copper Electrodes

When copper electrodes are used in electrolysis, the anode is attacked by anions so that it starts dissolving. The process may be represented as follows:

  • CuSO4 → Cu2+ + SO42- (Almost complete ionization)
  • H2O ⇌ H+ + OH (Weakly ionized)

At cathode. On passing electric current both Cu2+ and H+ ions move towards the cathode. However, Cu2+ ions having lower discharge potential get liberated at cathode as:

Cu2+ + 2e →Cu

At anode. Both OH– and SO42, ions move towards anode but, unlike what happens when platinum electrodes are used. None of the ions get liberated. Rather the copper electrode itself starts dissolving by losing electrons due to its lower discharge potential than that of OH– or SO42, ions. It may be shown as:

Cu → Cu2+ + 2e

It may, therefore, be noted that in this electrolysis, copper is deposited at cathode from solution and an equivalent amount of copper from the anode dissolves in solution forming Cu2+ ions.

Examples of Electrolysis

Electrolyte Electrode Cathode Reaction Anode Reaction
Silver nitrate Pt Ag+ + e → Ag 2OH– → 1/2 O2 + H2O + 2e
Sodium nitrate Pt 2H+ + 2e → H2 2OH → 1/2 O2 + H2O + 2e
 Molten PbBr2 Pt Pb2+ + 2e → Pb  2Br  → Br2 + 2e
CuCl2 solution Pt Cu2+ + 2e→ Cu 2Cl → Cl2 + 2e

Faraday’s Laws of electrolysis

In early nineteenth century, Michael Faraday carried out a number of experiments on electrolysis and discovered two important laws of electrolysis as:

Faraday’s First Law of Electrolysis

It states that the amount of any substance deposited at an electrode is directly proportional to the quantity of electricity passed through the electrolyte solution.

Mathematically: W ∝ Q

where W is the weight (in grams) of the substance deposited and Q is the quantity of electricity (in coulombs) which is passed through the electrolyte.

Since the quantity of electricity Q = I x t, where I is the current strength in amperes and t is the time in seconds, therefore, the above expression can be written as:

W ∝ I x t  Or   W=Z×I×t

where z is a constant of proportionality and is called the electrochemical equivalent of the substance. If I = 1 ampere and t = 1 second, then: W = Z

Thus, an electrochemical equivalent is defined as the weight of a substance deposited by the passage of one ampere current for one second (or one coulomb of electricity).

Importance of First Law of Electrolysis

This law is used to calculate:

  • The values of electrochemical equivalents of different ions.
  • The weights of different ions deposited by passing different quantities of electricity through their electrolytes.

It has been found out experimentally that by passing one Faraday, i.e., 96500 coulombs of electricity through an electrolyte, it results in the decomposition of one gram equivalent (i.e., equivalent weight expressed in grams) of the substance being deposited on the electrode. Therefore:

On the other hand, one coulomb of electricity deposits electrochemical equivalent (z) of the substance Hence, we may conclude that:

One Faraday or 96500 coulombs deposit one gram equivalent of the substance.

Electrochemical equivalent (Z) x 96500 = Gram equivalent

Example of First Law of Electrolysis

A current of 3 amperes strength on passing through silver nitrate solution for 20 minutes deposits 4 grams of silver. What is the electrochemical equivalent of silver?


  • Strength of the current (I) = 3 amperes
  • Time (t) = 20 minutes = 20 x 60 seconds
  • Weight of silver deposited (W) = 4 grams

According to Faraday’s Law of Electrolysis:

  • W = Z x I x t
  • 4= Z×3×20×60
  • Z = 4/3600 = 0.00111 g

Faraday’s Second Law of Electrolysis

It states that when the same quantity of electricity is passed through different electrolytes,  the amounts of different substances produced at the electrodes are directly proportional to their equivalent weights.

Let us explain the above law. Consider three cells, one contains water and a little HCl, the second cell contains copper sulfate solution and the third cell contains silver nitrate solution. Let these cells be connected in series with an electric battery.Electrolysis

Since all the three cells have been connected in series, they will receive the same quantity of electricity during the same period of time. On passing the electric current through these cells, hydrogen, copper, and silver will be deposited at their respective cathodes in the ratio of their chemical equivalent, i.e., 31.75: 108. In other words,


Electrolysis of Water Equation

Electrolysis of Water is used to generate oxygen for the International Space Station. The overall Equation for the electrolysis of water is given below.

  • Ions are attracted to the oppositely charged electrode.
  • Reduction at cathode: 2 H+(aq) + 2e → H2(g)
  • Oxidation at  Anode:  2 H2O(l) → O2(g) + 4 H+(aq) + 4e.

 Uses of Electrolysis

  • A large number of chemicals used in industry are obtained by electrolysis. For example, hydrogen and oxygen are manufactured by electrolysis of acidulated water.
  • Caustic soda is obtained by the electrolysis of aqueous sodium chloride in suitable cells. Chlorine is obtained as a byproduct.
  • Sodium metal is obtained by the electrolysis of either fused caustic soda or fused sodium chloride. Similarly, aluminum is extracted by the electrolysis of bauxite (i.e., Al2O3) in fused cryolite.
  • For preservation. Metals like iron are electroplated with tin, nickel or chromium to protect them from rusting.
  • For decoration. Cheap jewelry and other fancy articles made of copper are electroplated with gold or silver to enhance their beauty.
  • For repairs. Broken parts of machinery are sometimes repaired by electrodeposition of the metal between the broken parts.

Application of Electrolysis

The phenomenon of electrolysis finds numerous applications in industry and theoretical chemistry. Let us explain some of the important applications of electrolysis.

  1. Electroplating. We may define electroplating as the process of depositing a superior metal like gold, silver or nickel over a baser metal like copper or iron with the help of electricity. 
  2. Electrorefining of Metals. We may define electro-refining as the process of removal of impurities from an impure metal by electrolysis. Metals like copper, silver, gold, tin, aluminum, etc. are nowadays refined electrolytically.
  3. Electrotyping. We may define electrotyping as the process of obtaining the impressions of letters with the help of electrolysis. This process is employed in large-scale printing.
  4. Electrometallurgy.We may define Electrometallurgy as the process of extraction of metals from their ores by electrolysis.
  5. Determination of Equivalent Weights. Equivalent weights of elements can be determined from Faraday’s second law of electrolysis.