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# Electrolytic Cell

The electronic device which converts the electrical energy to chemical energy or chemical energy to electrical energy is called as cell.  Cell has two ends; positive and negative. It pushes the electric current from positive terminal to negative terminal which is conventional flow of current. Remember electrons always flow in opposite direction to electrons or flow of current. The combination of two or more cells in series forms a battery.

Voltaic cells are good examples of electrochemical cell which produce electricity due to spontaneous chemical reaction. In electrolytic cell, the chemical reaction induces by flowing electric current in cell as their chemical reactions are non-spontaneous in nature therefore requires some induction. The electrolytic cells are based on the electrolysis process of electrolyte due to flow of electric current. Since these chemical reactions are non-spontaneous therefore such reactions require some induction.

## Electrolytic Cell Definition

In a cell, the general reaction can be written as given below.

 Spontaneous$\rightarrow$ Reactants $\rightleftharpoons$ Products + Electrical Energy $\leftarrow$Non spontaneous

So we can define an electrolytic cell as the device that does work on a chemical system which drives by electric current. The chemical reaction that occurs in the cell is called as electrolysis that is decomposition of substance due to flow of electric current. Electrolysis is an oxidation-reduction and non-spontaneous reaction.

If a cell has a voltage of 1.5 volts and are connected in series then the resultant battery will have 3 V or 6 V. On the basis of reaction and flow of current, cell can be classified as electrochemical cell and electrolytic cell.

Electrolytic cell is composed of two half-cells—one. Out of these two half cells, one is reduction half-cell and another is oxidation half-cell. If we reverse the direction of electron flow in cells, the direction of reaction also reversed. Here reduction takes place at the cathode and oxidation occurs at the anode.

So we can say that an electrolytic cell converts the electrical energy to chemical energy. The supply of electricity initiates the chemical reaction. In such cells both electrodes are placed in a same container and container is filled with molten electrolyte. The negative electrode is called as cathode and positive electrode is called as anode.

The Faraday's law of electrolysis helps us to understand the concept of electrolysis. This law states that the amount of a substance consumed or produced at one of the electrodes in an electrolytic cell is directly proportional to the amount of electricity that passes through the cell.

## Electrolytic Cell Example

An electrolytic cell with molten sodium chloride is one of the most common examples of electrolytic cell. The chemical reaction that initiates with the flow of electric current is as given below.

 $\rightarrow$ Non spontaneous ( electrolytic cell ) 2 Nacl (I) $\rightleftharpoons$ 2 Na (s) + Cl2 (g)

In this electrolytic cell, the molten NaCl(l) is placed into the container with carbon electrodes.

Both terminals are attached with +ve and –ve terminals of a battery. As we pass electricity, electrons flow from the negative terminal to the cathode.
At cathode sodium ions reduce to sodium atoms. Similarly negatively charged chloride ions move to the anode and oxidize to chlorine atoms. It releases chlorine gas in the form of bubbles. In an electrolytic cell, the strongest reducing substance will undergo oxidation whereas the strongest oxidizing agent will be reduced.

For example in aqueous solution of sodium chloride hydrogen would undergo reduction instead of sodium. In this electrolytic cell, the 1.36 V of potential is required to oxidize (Cl-) ions to Cl2 and -2.71 V is required to reduce (Na+) ions to sodium metal. So overall potential difference would be [1.36 - (- 2.71)] = 4.07V.

## Electrolytic Cell Diagram

An electrolytic cell with aqueous solution of sodium chloride involves following reactions.

 cathode 2 H2O + 2 e– $\rightarrow$ H2(g) + 2 OH– E = +0.41 v ([OH–] = 10-7 M) anode Cl–  $\rightarrow$ $\frac{1}{2}$ Cl2(g) + e– E° = - 1.36 v net Cl- + H2O $\rightarrow$ 2 H2(g) +  $\frac{1}{2}$ Cl2(g) + 2 OH– E = - 0.95 v

Similarly electrolysis of water is also a good example of such device. Here electrolysis occurs in the presence small amount of acid or salt that makes water ionic and decomposition will be easy.

 cathode 2 H2O + 2 e– $\rightarrow$ H2(g) + 2 OH– E° = –0.83 v anode H2O $\rightarrow$ $\frac{1}{2}$ O2(g) + 2 H+ + 2 e– E° = –1.23 v net 3 H2O(l)  $\rightarrow$ H2(g) + $\frac{1}{2}$ O2(g) E = –2.06 v