Electrolytic Cells and Electrolysis: Definition, Applications, Reactions & Notes

An electrolytic cell uses external electrical energy to form a chemical reaction. It consists of two electrodes (anode and cathode). They are placed in the solution of molten salt. The power supply is given by a battery, which is connected to electrodes.

Electrolysis is the process that occurs in the electrolytic cell. It allows the direct flow of current in the electrolyte to produce a chemical change.

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Various components are required to make a functional electrolytic cell. Below is the list of component and their use.

1. Battery or DC (for power supply)
We need a battery or DC to provide the electricity to the cell. The positive terminal of the battery is connected to the anode, and the negative terminal is connected to the cathode.

2. Metallic Rods
We use two metallic rods, one behaves as an anode (+) and the other as a cathode (-). During the electrolysis process, the positive rod will accept the electron, and the negative rod will give the electron. 

3. Solutions
We can use molten salt or a liquid solution full of free ions, which can move easily during the reaction. For example: salt water (NACL) and copper sulfate solution.

4. A Vessel
Finally, we require a vessel to complete this setup. The vessel can be of glass, ceramic or plastic. 

When the circuit is completed, it triggers a non-spontaneous reaction where electrical energy is converted into chemical energy.

Now that we know what the important component of an electrolytic cell is, let's discuss its working.

A battery is connected to the metal rod. These rods are called electrodes. The positive terminal of the rod is called the anode (+), and the negative terminal is the cathode (-). The rods are dipped into a liquid solution or molten salt solution. 

When the circuit is complete, the electrolysis process takes place. The positive rod accepts the electron, and the negative rod gives off electrons. During this process, electrical energy is converted into chemical energy.

For electrolysis of aqueous NaCl using inert electrodes:

• Anode: $2{\mathrm{C}\mathrm{l}}^{-}\left(\mathrm{a}\mathrm{q}\right)\to {\mathrm{C}\mathrm{l}}_2\left(\mathrm{g}\right)+2{\mathrm{e}}^{-}$ (oxidation).
• Cathode: $2{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)+2{\mathrm{e}}^{-}\to {\mathrm{H}}_2\left(\mathrm{g}\right)+2{\mathrm{O}\mathrm{H}}^{-}\left(\mathrm{a}\mathrm{q}\right)$ (reduction).
• Overall: $2\mathrm{N}\mathrm{a}\mathrm{C}\mathrm{l}\left(\mathrm{a}\mathrm{q}\right)+2{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)\to 2\mathrm{N}\mathrm{a}\mathrm{O}\mathrm{H}\left(\mathrm{a}\mathrm{q}\right)+{\mathrm{H}}_2\left(\mathrm{g}\right)+{\mathrm{C}\mathrm{l}}_2\left(\mathrm{g}\right)$ .

The power source overcomes the positive cell potential, enabling the reaction.

Electrode materials (metallic rod) play an important role in the electrolysis process. They are responsible for the oxidation and reduction in the electrolyte cell. The nature of metallic rods used in the process can affect the flow of ions during the chemical process.

For example, the copper in copper sulphate electrolysis participates in reactions, and it dissolves in the given format.

$\mathrm{C}\mathrm{u}\left(\mathrm{s}\right)\to {\mathrm{C}\mathrm{u}}^{2+}\left(\mathrm{a}\mathrm{q}\right)+2{\mathrm{e}}^{-}$ .

During electrolysis, multiple ions may compete at the electrodes. The ion with the higher reduction potential (at the cathode) or lower oxidation potential (at the anode) reacts preferentially. For aqueous NaCl :

• Cathode: ${\mathrm{H}}_2\mathrm{O}\left({\mathrm{E}}^{\circ }=-0.83\mathrm{V}\right)$ reduces instead of ${\mathrm{N}\mathrm{a}}^{+}\left({\mathrm{E}}^{\circ }=-2.71\mathrm{V}\right)$ .
• Anode: ${\mathrm{C}\mathrm{l}}^{-}$ oxidizes instead of ${\mathrm{H}}_2\mathrm{O}$ due to a lower energy requirement.

Faraday gave the law of electrolysis in 1833-34. Two laws were introduced by Faraday. 

Faraday's First Law of Electrolysis: The chemical reaction that occurs at the electrode during electrolysis is directly proportional to the amount of current passed through the electrolyte.
$m\propto Q$ , or $m=ZQ$

Where,

• m = mass of substance (in gram)
• Z = electrochemical equivalent
• Q = total electric charge (in coulombs)

Since Q=I×t (current × time),

m = Z ⋅ I ⋅ t

Faraday's Second Law of Electrolysis: When an equal amount of electricity is passed through different electrolytes, the amount of different substances liberated is proportional to their chemical equivalent weights.

$m\propto \frac{\text{Equivalent weight}}{n}$ .

Charge $Q=It$ (current $\times$ time), and 1 Faraday ( $F=96485\mathrm{C}/\mathrm{m}\mathrm{o}\mathrm{l}$ ) deposits 1 equivalent of a substance.

Application of Faraday's Law

• Purification of metal
• Electrorefining of metals
• Calculating the amount of substance deposited during electrolysis.

The equivalent weight in electrolysis is the molar mass divided by the number of electrons transferred per ion:

Equation:

Equivalent weight $=\frac{\text{Molar mass}}{n}$ .

Example:

For ${\mathrm{C}\mathrm{u}}^{2+},n=2$ ,

So Equivalent weight $=\frac{63.5}{2}=31.75$ g. This is critical for Faraday's law calculations.