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Electrolysis and Electrochemical Cells


Table of Content


Electrochemical Cells

Electrochemical Cells Electrochemical cells of two types-

  • Galvanic cells (also known as voltaic cell) : It is a device in which a redox reaction is used to convert chemical energy into electrical energy, i.e., electricity can be obtained with the help of oxidation and reduction reaction. The chemical reaction responsible for production of electricity takes place in two separate compartments. Each compartment consists of a suitable electrolyte solution and a metallic conductor. The metallic conductor acts as an electrode. The compartments containing the electrode and the solution of the electrolyte are called half-cells. When the two compartments are connected by a salt bridge and electrodes are joined by a wire through galvanometer the electricity begins to flow. This is the simple form of voltaic cell.

  • Electrolytic cells: In this type of cells electrical energy is used to carry out a non-spontaneous reaction.

In simple words,one can say that in galvanic cells, chemical energy is converted into electrical energy, while in electrolytic cell electrical energy is converted into chemical energy. 

Difference in Electrolytic Cell and Galvanic Cell

Electrolytic cell

Galvanic cell

Electrical energy is converted into chemical energy.


Anode positive electrode. Cathode negative electrode.


 Ions are discharged on both the electrodes.


If the electrodes are inert, concentration of the electrolyte decreases when the electric current is circulated.


Both the electrodes can be fitted in the same compartment.

Chemical energy is converted into electrical energy.


Anode negative electrode. Cathode positive electrode.


Ions are discharged only on the cathode.


Concentration of the anodic half-cell increases while that of cathodic half-cell decreases when the two electrodes are joined by a wire.


The electrodes are fitted in different compartments.


Electrolytic Cell

An electrolytic cell is an arrangement in which electricity is conducted through a solution or a molten salt by the movement of ions. It can be said that electrical energy is converted to chemical energy.

The principles of electrolytic conduction are best illustrated by reference to an electrolytic cell such as that shown in figure for the electrolysis of molten NaCl between inert electrodes.

In order to pass the current through an electrolytic conductor (aqueous solution or fused electrolyte), two rods or plates (metallic conductors) are always needed which are connected with the terminals of a battery. 

These rods or plates are known as electrodes. The electrode through which the current enters the electrolytic solution is called the anode (positive electrode) with the electrode through which the current leaves the electrolytic solution is known as cathode (negative electrode).

The entire assembly except that of the external battery is known as the cell.The electrons are received from the negative end of the external battery by the negative electrode of the cell.

These are used up in the reduction reaction at this electrode.

Electrolytic Cell

The numbers of electrons received at the negative electrode are given back to the positive end of the external battery from the positive electrode of the cell where electrons are released as a result of oxidation reaction. 

Within the cell, the current is carried by the movements of ions; cations towards the negative electrode (cathode) and anions towards the positive electrode (anode)

This movement of ions gives rise to what is known as the electrolytic conduction.

Let us now take a situation where more than one type of cation is present. The ability of cation to move towards the negative electrode and get reduced depends upon the size, mass, positive charge, negative charge etc. 

It is therefore not possible to predict, qualitatively, the order of reduction of cations, as one factor might enhance it while the another factor might hamper it.

The only way we can predict this is by giving a quantitative value based on the cumulative effect of all the factors responsible for a cation ability to get reduced. 

This quantitative value is called the standard reduction potential (SRP). A cation with a higher value of SRP would get reduced in preference to a cation with a lower value of SRP.

The standard reduction potential values at 25°C are given below for some reactions.

Standard Reduction Potentials at 25° C

Reducation half reaction

E°, V

Reducation half reaction

E°, V

F2 + 2e → 2F


AgCl + e → Ag + Cl


S2O82-+2e- →  2SO42-


PdI42- + 2e → Pd + 4I


Co3+ + e →  Co+2


Cu2+ + e → Cu+


H2O2 + 2H+ + 2e → 2H2O


Sn4+ + 2e → Sn2+


MnO4- + 4H+ + 3e → MnO2 + 2H2O


Ag(S2O3)23-+e→  Ag+ 2S2O32-


PbO2 + 4H+ + SO42+ 2e → PbSO4 + 2H2O


2H+ + 2e → H2


Ce4+ + e → Ce3+


Pb2+ + 2e → Pb

– 0.126

MnO4- + 8H+ + 5e– →  Mn+2 + 4H2O


Sn2+ + 2e → Sn

– 0.14

Au3+ + 3e → Au


2CuO + H2O + 2e →  Cu2O + 2OH

– 0.15

Cl2 + 2e → 2Cl


AgI + e → Ag + I

– 0.151

Cr2O72- + 14H+ + 6e → 2Cr+3 + 7H2O


CuI + e → Cu + I

– 0.17

Ti3+ + 2e → Ti+


Ni2+ + 2e → Ni

– 0.25

MnO2 + 4H+ + 2e  → Mn2+ + 2H2O


Co2+ + 2e → Co

– 0.28

O2 + 4H+ + 4e–  → 2H­2O


PbSO4 + 2e → Pb + SO42-

– 0.31

2IO3- + 12H+ + 10e  → I2 + 6H2O


Ti+ + e → Ti

– 0.336

Br2 + 2e → 2Br


Cu2O + H2O + 2e → 2Cu + 2OH

– 0.34

AuCl4- + 3e → Au + 4Cl


Cd2+ + 2e → Cd

– 0.403

OCl + H2O + 2e → Cl + 2OH


Fe2+ + 2e → Fe

– 0.44

Pd2+ + 2e → Pd


Cr3+ + 3e → Cr

– 0.74

2Hg2+ + 2e → Hg22+


Zn2+ + 2e → Zn

– 0.7628

Cu2+ + I  + e → CuI


2H2O + 2e → H2 + 2OH

– 0.828

Ag+ + e → Ag


Mn2+ + 2e → Mn

– 1.18

Fe3+ + e → Fe2+


H2 + 2e → 2H

– 2.25

O2 + 2H+ + 2e → H2O2


Mg2+ + 2e → Mg

– 2.37

Cu2+ + Cl + e → CuCl


Ce3+ + 3e → Ce

– 2.48

I2 + 2e–- → 2I


Na+ + e → Na

– 2.713

Cu+ + e →  Cu


Ca2+ + 2e → Ca

– 2.87

Cu2+ + 2e → Cu


K+ + e → K

– 2.93

Hg2Cl2 + 2e → 2Hg + 2Cl


Li+ + e → Li


Hg2Cl2 + 2e → 2Hg + 2Cl (satd KCl)




Ge2+ + 2e → Ge




For anions the ability to get oxidized is given by the standard oxidation potential which is the reverse of the standard reduction potential of a molecule to form the anion.

According to the table, if we take an aqueous solution of NaCl and do its electrolysis, H+ would be reduced to H2 gas (the H+ ions are present since the solution is aqueous) at the cathode, while Cl ions would be oxidised to Cl2 gas at the anode.

 Galvanic cellThough what we have stated just now is used in solving problem, it is not always valid. This is because the ability of a cation to be reduced or an anion to be oxidized not only depends on their SRP’s, but also depends on their concentrations. This means that it is possible to reduce a cation in preference to another cation even though the SRP of the former may be less than that of the latter, just by adjusting concentrations. 

A most remarkable feature of oxidation - reduction reactions is that they can be carried out with the reactants separated in space and linked only by an electrical connection. That is to say, chemical energy is converted to electrical energy. Consider figure ,a representation of a galvanic cell which involves the reaction between metallic
zinc and cupric ion:

The cell consists of two beakers, one of which contains a solution of Cu2+ and a copper rod, the other a Zn2+ solution and a zinc rod. A connection is made between the two solutions by means of a “salt bridge”, a tube containing a solution of an electrolyte, generally NH4NO3 or KCl.

Flow of the solution from the salt bridge is prevented either by plugging the ends of the bridge with glass wool, or by using a salt dissolved in a gelatinous material as the bridge electrolyte.

When the two metallic rods are connected through an ammeter, a deflection is  observed in  ammeter which is an evidence that a chemical reaction is Salt bridgeoccurring.

The zinc rod starts to dissolve, and copper is deposited on the copper rod.  The solution of Zn2+ becomes more concentrated, and the solution of Cu2+ becomes more dilute.

The ammeter indicates that electrons are flowing from the Zinc rod to the copper rod. This activity is continuous as long as the electrical connection and the salt bridge are maintained, and visible amounts of reactants remain.

Now let us analyze what happens in each beaker more carefully. We note that electrons flow from the Zinc rod through the external circuit, and that Zinc ions are produced as the Zinc rod dissolves. We can summarize these observations by writing,
Zn →  Zn2+ + 2e (at the zinc rod).

Also, we observe that electrons flow to the copper rod as cupric ions leave the solution and metallic copper is deposited. We can represent these occurrences by
2e + Cu2+ (aq) →  Cu (at the copper rod).

In addition, we must examine the purpose of the salt bridge. Since Zinc ions are produced as electrons leave the zinc electrode, we have a process which tends to produce a net positive charge in the left beaker.

The purpose of the salt bridge is to prevent any net charge accumulation in either beaker, diffuse through the bridge, and enter the left beaker.

At the same time, there can be a diffusion of positive ions from left to right.

If this diffusional exchange of ions did not occur, the net charge accumulating in the beakers would immediately stop the electron flow through the external circuit, and the oxidation reduction reaction would stop.

Thus, while the salt bridge does not participate chemically in the cell reaction, it is necessary if the cell is to operate.

Significance of Salt Bridge

The following are the func­tions of the salt bridge:

  • It connects the solutions of two half-cells and completes the cell circuit.

  • It prevents transference or diffusion of the solutions from one half-cell to the other.

  • It keeps the solutions in two half-cells electrically neutral. In anodic half cell, positive ions pass into the solution and there shall be accumulation of extra positive charge in the solution around the anode which will prevent the flow of electrons from anode. This does not happen because negative ions are provided by salt bridge. Similarly, in cathodic half-cell negative ions will accumulate around cathode due to deposition of positive ions by reduction. To neutralize these negative ions, sufficient number of positive ions is provided by salt bridge. Thus, salt bridge maintains electrical neutrality.

  • It prevents liquid-liquid junction-potential, i.e., the potential difference which arises between two solutions when in contact with each other.

  • A broken vertical line or two parallel vertical lines in a cell reaction indicates the salt bridge.

  • Salt bridge can be replaced by a porous partition which allows the migration of ions without allowing the solutions to intermix.

IUPAC Cell Representation 

  Galvanic cell

The galvanic cell mentioned above is represented in a short IUPAC cell notation as follows:

It is important to note that:

  • First of all the anode (electrode of the anode half cell) is written. In the above case, it is Zn.

  • After the anode, the electrolyte of the anode should be written with concentration. In this case it is ZnSO4 with concentration as C1 moles/litre.

  • A slash (|) is put in between the Zn rod and the electrolyte. This slash denotes a surface barrier between the two as they exist in different phases.

  • Then we indicate the presence of a salt bridge by a double slash (||).

  • Now, we write the electrolyte of the cathode half-cell which is CuSO4 with its concentration which is C2   moles/ l.

  • Finally we write the cathode electrode of the cathode half – cell.

  • A  slash (/) between the electrolyte and the electrode in the cathode half – cell.

  • In case of a gas, the gas to be indicated after the electrode in case of anode and before the electrode in case of cathode. Example: Pt, H2/H+ or H+|H2, Pt.


Electrolysis can be defined as the process of process of separating any compound into its constituent elements by passing an electric current through its aqueous solution.

Preferential Discharge Theory

If an electrolytic solution consists of more than two ions and the electrolysis is done, it is observed that all the ions are not discharged at the electrodes simultaneously but certain ions are liberated at the electrodes in preference to others. This is explained by preferential discharge theory.

It states that if more than one type of ions are attracted towards a particular electrode, then the one discharged is the ion which requires least energy.

The potential at which the ion is discharge or deposition potential.

The values of discharge potential are different for different ions. For example, the discharge potential of H+ ions is lower than Na+ ions when platinum or most of the other metals  are used as cathodes. Similarly, discharge potential of Cl- ions is lower than that of OH- ions. This can be explained by some examples given below:

Electrolysis of Sodium Chloride Solution

The solution of sodium chloride besides Na+ and Cl- ions possesses H+ and OH-ions due to ionization of water. However, the number is small as water is a weak electrolyte. When potential difference is established across the two electrodes, Na+ and H+ ions move towards cathode and Cl- and OH- ions move towards anode. At cathode H+ ions are discharged in preference to Na+ ions as the discharge potential of Hions is lower than Na+ ions. Similarly at anode, Cl- ions are discharged in preference to OH- ions.

NaCl \rightleftharpoons  Na+ + Cl-

H2O \rightleftharpoons  H+ + OH-

At cathode        

     At Anode

H+ + e- →  H   

2H→  H2  

Cl- →  Cl + e-

2Cl →  Cl2

Thus, Na+ and OH- ions remain in solution and the solution when evaporated yields crystals of sodium hydroxide.

Refer to the following video for electrolysis of sodium chloride

Electrolysis of Copper Sulphate Solution using Platinum Electrodes

CuSO4 \rightleftharpoons Cu2+ + SO42-

H2O \rightleftharpoons  H+ + OH-

At cathode        

     At Anode

 Cu2+ + 2e- → Cu

2OH- → H2O + O + 2e-

O + O → O2

Copper is discharged at cathode as Cu2+ ions have lower discharge potential than H+ ions. OH- ions are discharged at anode as these have lower discharge potential than  ions. Thus, copper is deposited at cathode and oxygen gas is evolved at anode.

Electrolysis of Sodium Sulphate Solution using Inert Electrodes

Na2SO4 \rightleftharpoons 2Na+ + SO42-

H2O \rightleftharpoons  H+ + OH-

At cathode        

     At Anode

H+ + e- →  H   

2H→  H2  

OH- →  H2O + O + 2e-

O + O →  O2

Hydrogen is discharged at cathode as H+ ions have lower discharge potential than Na+ ions. OH- ions are discharged at anode as these have lower discharge potential than  ions. Thus, hydrogen is evolved at cathode and oxygen is evolved at anode, i.e., the net reaction describes the electrolysis of water. The ions of Na2SO4 conduct the current through the solution and take no part in the overall chemical reaction.

The decreasing order of discharge potential or the increasing order of deposition of some of the ions is given below:

For cations:  K+> Na+> Ca2+> Mg2+> Al3+>Zn2+> H+> Cu2+> Hg2+> Ag+

For anions:  SO42-> NO3-> OH->Cl-> Br-> I-

Electrolysis of Copper Sulphate Solution using Copper Electrodes

CuSO4 \rightleftharpoons Cu2+ +  SO42- 

At cathode, copper is deposited.

Cu2+ + 2e- →  Cu

At anode, the copper of the electrode is oxidised to Cu2+ ions or ions solution dissolve equivalent amount of copper of the anode.

Cu  → Cu2+ + 2e-

Thus, during electrolysis, copper is transferred from anode to cathode.

Electrolysis of Silver Nitrate Solution using Silver Electrodes

AgNO2 \rightleftharpoons Ag+ +  NO42-

At cathode, silver is deposited.

Ag+ + e- → Ag

At anode, the silver of the electrode is oxidised to Ag+ ions which go into the solution or ions dissolve equivalent amount of silver of the electrode.

Ag →  Ag+ + e-

Ag +  NO3-  → AgNO3 + e-

Some more Examples of Electrolysis



Cathodic reaction

Anodic reaction


Aqueous acidified CuCl2 solution


Molten PbBr2


Sodium chloride solution


Silver nitrate solution


Sodium nitrate solution














  Cu2+ + 2e-→ Cu


  Pb2+ + 2e- → Pb


  2Na+ + 2e- → 2Na


  Ag+ + e- --> Ag


  2H+ + 2e- → H2



   2Cl- → Cl2 + 2e-

   2br- --> Br2 + 2e-


  2Cl- → Cl2 + 2e-


1/2 O2 + H2O + 2e-


  2OH- → 1/2 O2 + H2O + 2e-

Question 1: Electrolytic cell converts..

a. electrical energy into chemical energy.

b. chemical energy into electrons.

c. chemical energy into electrical energy.

d. electrical energy into ions.

Question 2:  Which of the following objects is not part of electrochemical cell

a. Anode

b. Battery 

c. Cathode

d. Electrolyte

Question 3: Which of the following statements is incorrect regarding salt bridge?

a. It connects the solutions of two half-cells and completes the cell circuit.

b. It prevents transference or diffusion of the solutions from one half-cell to the other.

c. It keeps the solutions in two half-cells electrically neutral.

d. It causes liquid-liquid junction-potential

Question 4: Which of the following gases is produced at anode during electrolysis of sodium chloride?

a. Oxygen

b. Hydrogen

c. Chlorine

d. Nitrogen

Question 5: Which of the following ions has maximum discharge potential?

a. K+

b. Na+

c. Ca2+

d. Mg2











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