See Table.

`text(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 `text(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. See fig.1.


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 `text(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.

It is designed to make use of the spontaneous redox reaction between zinc and cupric ions to produce an electric current. It consists of two half-cells. The half-cells on the left contains a zinc metal electrode dipped in `ZnSO_4` solution. See fig.1.

The half-cell on the right consists of copper metal electrode in a solution `CuSO_4`. The half-cells are joined by a salt bridge that prevents the mechanical mixing of the solution.

When the zinc and copper electrodes are joined by wire, the following observations are made:

•There is a flow of electric current through the external circuit.

•The zinc rod loses its mass while the copper rod gains in mass.

•The concentration of `ZnSO_4` solution increases while the concentration of copper sulphate solution decreases.

•The solutions in both the compartments remain electrically neutral.

See fig.2.

During the passage if electric current through external circuit, electrons flow from the zinc electrode to the copper electrode. At the zinc electrode, the zinc metal is oxidized to zinc ions which go into the solution. The electrons released at the electrode travel through the external circuit to the copper electrode where they are used in the reduction of `Cu^(2+)` ions to metallic copper which is deposited on the electrode. Thus, the overall redox reaction is

`Zn(s) + Cu^(2+) → Cu(s) + Zn^(2+)(aq)`


Thus, indirect redox reaction leads to the production of electrical energy. At the zinc rod, oxidation occurs. It is the anode of the cell and is negatively charged while at copper electrode, reduction takes, place; it is the cathode of the cell and is positively charged.

Thus, the above points can be summed up as:

•Voltaic or Galvanic cell consists of two half-cells. The reactions occurring in half-cells are called half-cell reactions. The half-cell in which oxidation taking place in it is called oxidation half-cell and the reaction taking place in it is called oxidation half-cell reaction. Similarly, the half-cell occurs is called reduction half-cell and the reaction taking place in it is called reduction half-cell reaction.


•The electrode where oxidation occurs is called anode and the electrode where reduction occurs is termed cathode.

•Electrons flow from anode to cathode in the external circuit.

•Chemical energy is converted into electrical energy.

•The net reaction is the sum of two half-cell reactions. The reaction in Daniel cell can be represented as See Table 1.

`text(Electrode Signs)`

The signs of the anode and cathode in the voltaic or galvanic cells are opposite to those in the electrolytic cells. See Table 2.

See fig.3

Daniell cell has the emf value 1.09 volt. If an opposing emf exactly equal to 1.09 volt is applied to the cell, the cell reaction,

`Zn + Cu^(2+) → Cu + Zn^(2+)`

stops but if it is increased infinitesimally beyond 1.09 volt, the cell reaction is reversed.

`Cu + Zn^(2+) → Zn + Cu^(2+)`

Such a cell is termed a reversible cell. Thus, the following are the two main conditions of reversibility


•The chemical reaction of the cell stops when an exactly equal opposing emf is applied.

•The chemical reaction of the cell is reversed and the current flows in opposite direction when the opposing emf is slightly greater than that of the cell.

•Any other cell which does not obey the above two conditions is termed as irreversible.

•A cell consisting of zinc and copper electrodes dipped into the solution of sulphuric acid is irre­versible.

•Similarly, the cell `Zn|H_2SO_4(aq)|Ag` is also irreversible because when the external emf is greater than the emf of the cell, the cell reaction,

• `Zn + 2H^(+) → Zn^(2+) + H_2` is not reversed but the cell reaction becomes `2Ag + 2H^(+) → 2Ag^(+) + H_2`

See Table.

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 `H_2` gas (the `H^+` ions are present since the solution is aqueous) at the cathode, while `Cl^-` ions would be oxidised to `Cl_2` gas at the anode.

Though 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: See fig.1.

The cell consists of two beakers, one of which contains a solution of `Cu^(2+)` and a copper rod, the other a `Zn^(2+)` solution and a zinc rod. A connection is made between the two solutions by means of a `text(“salt bridge”,)` a tube containing a solution of an electrolyte, generally `NH_4NO_3` 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 occurring.

The zinc rod starts to dissolve, and copper is deposited on the copper rod. The solution of `Zn^(2+)` becomes more concentrated, and the solution of `Cu^(2+)` 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 → Zn^(2+) + 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^(–) + Cu^(2+) (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.

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.
`Zn|Zn^(2+)||Cu^(2+)|Cu`

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

See fig.1.

The galvanic cell mentioned above is represented in a short IUPAC cell notation as follows : See fig.2.
`text(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 `ZnSO_4` with concentration as `C_1` 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 `CuSO_4` with its concentration which is `C_2` moles/litre.

•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, H_2//H^(+)` or `H^(+)|H_2, 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.

`text(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 H+ ions is lower than `Na^(+)` ions. Similarly at anode, `Cl^(-)` ions are discharged in preference to `OH^(-)` ions.

`NaCl ⇋ Na^(+) + Cl^(-)`

`H_2O ⇋ H^(+) + OH^(-)`

See Table 1.

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

Electrolysis of copper sulphate solution using platinum electrodes:

`CuSO_4 ⇋ Cu^(2+) + SO_4^(2-)`

`H_2O ⇋ H^(+) + OH^(-)`

See Table 2.

Copper is discharged at cathode as `Cu^(2+)` 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:
`Na_2SO_4 ⇋ 2Na^(+) + SO_4^(2-)`

`H_2O ⇋ H^(+) + OH`

See Table 3.


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 `Na_2SO_4` 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^(+)> Ca^(2+)> Mg^(2+)> Al^(3+)>Zn^(2+)> H^(+)> Cu^(2+)> Hg^(2+)> Ag^(+)`

For anions: `SO_4^(2-)> NO_3^(-) > OH^(-) >Cl^(-) > Br^(-) > I^(-)`

Electrolysis of copper sulphate solution using copper electrodes:

`CuSO_4 ⇋ Cu^(2+) + SO_4^(2-)`

At cathode, copper is deposited.

`Cu^(2+) + 2e^(-) → Cu`

At anode, the copper of the electrode is oxidised to `Cu^(2+)` ions or ions solution dissolve equivalent amount of copper of the anode.


`Cu → Cu^(2+) + 2e^(-)`

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

Electrolysis of silver nitrate solution using silver electrodes:

`AgNO_2 ⇋ Ag^(+) + NO_4^(2-)`

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 + NO_3^(-) → AgNO_3 + e^(-)`

Some more examples of electrolysis : See Table 4.