Electrolytic Cell and Electrolysis

Electrolysis is the use of electrical energy from an external source to drive a non-spontaneous chemical reaction, most often to decompose a stable compound into its constituent elements. The device used to carry it out is called an electrolytic cell, and — as noted earlier — the electrochemistry occurring inside it is exactly the reverse of what occurs inside a galvanic cell.

The electrolysis of molten sodium chloride illustrates the whole process concretely. The electrolytic cell has two electrodes — a cylindrical steel cathode and a graphite anode — dipped into a bath of molten NaCl (industrially, often mixed with CaCl₂ to lower its melting point), connected through a switch to an external DC power supply. The electrode wired to the negative terminal of the supply is, by definition, the cathode, and the one wired to the positive terminal is the anode. Once the switch closes, the external supply actively drives electrons onto the cathode and simultaneously pulls electrons away from the anode — note that this is the opposite of a galvanic cell, where electrons flow spontaneously FROM the anode; here they are forced there BY the reaction that follows.

Cathode (reduction): Na⁺ ions, attracted to the negatively-charged cathode, accept the supplied electrons and are reduced to liquid sodium metal:

Anode (oxidation): Cl⁻ ions, attracted to the positively-charged anode, give up electrons and are oxidised to chlorine gas:

Combining the two half-reactions gives the overall electrolysis reaction, $2\text{Na}^{+}\text{(l)} + 2\text{Cl}^{-}\text{(l)} \rightarrow 2\text{Na(l)} + \text{Cl}_2\text{(g)}$, with an overall $E^{o} = -4.07$ V. That negative value confirms directly that the reaction is non-spontaneous exactly as expected for an electrolytic process — an external voltage greater than 4.07 V must be supplied to drive it forward. …