17.7 Electrolysis

In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell.

In electrolysis a galvanic cell produces current when an oxidation-reduction proceeds spontaneously. A similar apparatus, an electrolytic cell uses electrical energy to produce chemical change. The process of electrolysis involves:

Forcing a current through a cell to produce a chemical change for which the cell potential is negative; that is electrical work causes an otherwise non-spontaneous chemical reaction to occur.

A familiar example of electrolysis is recharging a battery, which involves use of an external power source to drive the spontaneous (discharge) cell reaction in the reverse direction, restoring to some extent the composition of the half-cells and the voltage of the battery. Perhaps less familiar is the use of electrolysis in the refinement of metallic ores, the manufacture of commodity chemicals, and the electroplating of metallic coatings on various products (e.g., jewelry, utensils, auto parts). To illustrate the essential concepts of electrolysis, a few specific processes will be considered..

The Electrolysis of Molten Sodium Chloride

Molten (liquid) sodium chloride can be electrolyzed to produce sodium metal and chlorine gas. The electrolytic cell used in the process is called a Down’s cell (see the following figure).The reactions associated with this process are:

The above Figure is  a Down’s cell which is used for the electrolysis of molten sodium chloride.

In a Down’s cell, the liquid sodium ions are reduced at the cathode to liquid sodium metal. At the anode, liquid chlorine ions are oxidized to chlorine gas. The reactions and cell potentials are shown below:

The battery must supply over 4 volts to carry out this electrolysis. This reaction is a major source of production of chlorine gas and is the only way to obtain pure sodium metal. Chlorine gas is widely used in cleaning, disinfecting, and in swimming pools.

Electrolysis of Aqueous Sodium Chloride

It may be logical to assume that the electrolysis of aqueous sodium chloride, called brine, would yield the same result through the same reactions as the process in molten NaCl. However, the reduction reaction that occurs at the cathode does not produce sodium metal, instead, the water is reduced. This is because the reduction potential for water is only −0.83V−0.83V compared to −2.71V−2.71V for the reduction of sodium ions. This makes the reduction of water preferable because its reduction potential is less negative. Chlorine gas is still produced at the anode, just as in the electrolysis of molten NaCl.

Since hydroxide ions are also a product of the net reaction, the important chemical sodium hydroxide (NaOH)(NaOH) is obtained from evaporation of the aqueous solution at the end of the hydrolysis.

The Electrolysis of Water

Water may be electrolytically decomposed in a cell similar to the one illustrated in Figure 2. To improve electrical conductivity without introducing a different redox species, the hydrogen ion concentration of the water is typically increased by addition of a strong acid. The redox processes associated with this cell are:

The cell potential as written is negative, indicating a nonspontaneous cell reaction that must be driven by imposing a cell voltage greater than +1.229 V. Keep in mind that standard electrode potentials are used to inform thermodynamic predictions here, though the cell is not operating under standard state conditions. Therefore, at best, calculated cell potentials should be considered ballpark estimates.

The above Figure shows the electrolysis of water produces stoichiometric amounts of oxygen gas at the anode and hydrogen at the cathode.

Reference: https://www.youtube.com/watch?v=7uIIq_Ofzgw

Reference: https://www.youtube.com/watch?v=dRtSjJCKkIo&t=21s

Reference: https://www.youtube.com/watch?v=GrgYXk_NCec

Reference: https://www.youtube.com/watch?v=WmtaMq36jaE

SUMMERY AND KEY CONCEPTS

Nonspontaneous redox processes may be forced to occur in electrochemical cells by the application of an appropriate potential using an external power source—a process known as electrolysis. Electrolysis is the basis for certain ore refining processes, the industrial production of many chemical commodities, and the electroplating of metal coatings on various products. Measurement of the current flow during electrolysis permits stoichiometric calculations.

Key Equation:

. Q = I x t = n x F