Electrolytic Cells and Electrolysis

Electrolysis is a fundamental part of chemistry, which involves breaking of electrolyte(the solution) and formation of positive and negative ions. This is the principle behind the electrolytic cells. The process may look complicated but it's just like a cake walk.

Electrolysis is defined as a process of decomposition of an electrolyte by the passage of electricity through its aqueous solution or molten (fused) state.

Electrolytic cell

The apparatus used to bring about electrolysis is electrolytic cell.It consists of a vessel containing electrolytes either as a aqueous solution or in molten state.Two metal rods are dipped into it and are connected to two terminal of a electricity source.These rods are called electrodes;the one connected to negative pole of the battery is called cathode and other connected to the positive pole is called anode.

Mechanism of electrolysis

During electrolysis a electrolyte is decomposed into positive and negative ions.Positive ions are attracted to cathode(negative electrode),accept electron from cathode and become neutral in charge.While negative ions migrate towards anode(positive electrode),lose electron and become neutral .The conversion of ions into neutral species at their respective electrode is called the primary change.The product formed as a result of primary change may be collected as such or may undergo secondary change to form the final product.

In case where are there are more than one positive and negative ions,who will preferably gain or lose electrons depends on their respective discharge potential.For example,in case of electrolysis of aqueous solution of NaCl there are two positive ions \(\text{Na}^+\),\(\text{H}^+\) and two negative ions \(\text{Cl}^-\),\(\text{OH}^-\) .These extra ions \(\text{H}^+\) and \(\text{OH}^-\) are produced from decomposition of water since it is also a weak electrolyte.On cathode, \(\text{H}^+\) get discharged since it has lower electrode potential than \(\text{Na}^+\) and \(\text{Na}^+\) will remain in the solution .Similarly on anode, \(\text{Cl}^-\) has lower discharge potential than \(\text{OH}^-\).Hence \(\text{Cl}^-\) will get discharged in preference to \(\text{OH}^-\) and \(\text{OH}^-\)will remain in the solution.

\[\text{At cathode:}\quad \quad \text{H}^++e^{-} \to H \quad \quad\text{(Reduction, primary change)}\] \[\quad \quad \space \text{H}^.+\text{H}^. \to \text{H}_2 \quad \quad \text{(Secondary change)}\]

\[\text{At Anode:}\quad \quad \text{Cl}^- \to \text{Cl}^.+e^{-} \quad \quad \text{(Oxidation, primary change)}\] \[\quad \quad \space \text{Cl}^.+\text{Cl}^. \to \text{Cl}_2 \quad \quad \text{(Secondary change)}\]

An easy way to remember which electrode is positive is that positive electrode attracts cations (thus cathode) and vice versa. Another mnemonic to remembering whether there is oxidation or reduction happening at the electrode is "RCOA" Reduction at Cathode and Oxidation at Anode.

First law of electrolysis

First law of electrolysis states that the quantity of elements liberated or deposited in electrolysis is proportional to the quantity of electric charge passed through the circuit.

\[m \propto Q\]

\[m=Z\times Q\]

Since \(Q=I \times t\)

\[m=Z\times I\times t\]

Where,

\( \text{Q } \) - Quantity of electricity
\( \text{I } \) - Current supplied to electrolyte
\( \text{t } \) - Time for which electricity is passed
\(\text{Z } \) -Proportionality constant and is called electrochemical equivalent of the substance deposited.

If 1 ampere of current is passed for one second through electrolyte ,quantity of substance deposited or liberated is \( m=Z \times 1 \times 1=Z\)

Electrochemical equivalent of a substance may be defined as the mass of the substance deposited when a current of one ampere is passed for one second i.i a quantity of electricity equal to one coulomb is passed.

Second law of electrolysis

Second law of electrolysis states that when the same amount of electricity is passed through different electrolytes/elements connected in series, the mass of substance liberated/deposited at the electrodes is directly proportional to their equivalent weights.

For example, for AgNO\(_3\) solution and CuSO\(_4\) solution connected in series, if the same quantity of electricity is passed.

\[\frac{\text{Mass of Ag deposited}}{\text{Mass of Cu deposited}}=\frac{\text{Eq.Wt. of Ag}}{\text{Eq.Wt. of Cu}}\]

Qualitative aspect of Electrolysis

Consider a reaction

\[\text{Na}^++e^- \to \text{Na}\]

What will be amount of electricity required to deposit one mole of Sodium?

We see that one electron is required to deposit one sodium atom.To deposit one mole of sodium that is \(6.022\times 10^{23}\) sodium atom ,we need \(6.022\times 10^{23}\) electrons. Charge carried by one electron is \(1.6021 \times 10^{-19}\) coulomb. Total charge carried by \(6.022\times 10^{23}\) electron is \(6.022\times 10^{23}\times 1.6021 \times 10^{-19}=96500\) coulomb.This amount of electricity deposit one mole of sodium and is called one faraday.

\[1 \space \text{Faraday}=96500 \space \text{C mol}^{-1}\]

Thus it may be concluded that If n electrons are involved in the electrode reaction, the passage of n faraday of electricity will liberate one mole of the substance.

In terms of equivalent weight,it may be remembered that One faraday of electricity deposits one gram equivalent of the substance.

Two important results

(1) As one faraday (96500 coulomb) deposits one gram equivalent of the substance, hence electrochemical equivalent can be calculated from the equivalent weight.

\[Z=\frac{\text{Eq.wt.of the substance}}{96500}\]

(2) Equivalent weight can be calculated from the weight of substance deposited W and amount of electricity Q

\[\text{Eq.wt.of the substance}=\frac{W}{Q}\times 96500\]

Electrolysis is very important industrial process which include mass production of variety of compound

• Production of chlorine by electrolysis of aqueous NaCl solution
• Production of heavy water (\(D_2O\))
• Production of oxygen for spacecraft and nuclear submarines
• Electroplating and electrolytic refining of metals
• Manufacture of hydrogen by electrolysis of water