Private: Chapter Seventeen
Electrolysis (17.7)
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Learning Objectives
By the end of this section you will be able to:
- Describe the process of electrolysis
- Compare the operation of electrolytic cells with that of galvanic cells
- Perform stoichiometric calculations for electrolytic processes
Electrochemical cells in which spontaneous redox reactions take place (galvanic cells) have been the topic of discussion so far in this chapter. In these cells, electrical work is done by a redox system on its surroundings as electrons produced by the redox reaction are transferred through an external circuit. This final section of the chapter will address an alternative scenario in which an external circuit does work on a redox system by imposing a voltage sufficient to drive an otherwise nonspontaneous reaction, a process known as electrolysis. 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
Metallic sodium, Na, and chlorine gas, Cl2, are used in numerous applications, and their industrial production relies on the large-scale electrolysis of molten sodium chloride, NaCl(l). The industrial process typically uses a Downs cell similar to the simplified illustration shown in Figure 17.18. The reactions associated with this process are:
anode:2Cl−(l) ⟶ Cl2(g) + 2e−
cathode:Na+(l) + e− ⟶ Na(l)
cell:2Na+(l) + 2Cl−(l) ⟶ 2Na(l) + Cl2(g)
The cell potential for the above process is negative, indicating the reaction as written (decomposition of liquid NaCl) is not spontaneous. To force this reaction, a positive potential of magnitude greater than the negative cell potential must be applied to the cell.
Figure 17.18 Cells of this sort (a cell for the electrolysis of molten sodium chloride) are used in the Downs process for production of sodium and chlorine, and they typically use iron cathodes and carbon anodes.
The Electrolysis of Water
Water may be electrolytically decomposed in a cell similar to the one illustrated in Figure 17.19. 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
anode:2H2 O(l) ⟶ O2(g) + 4H+(aq) + 4e− E°anode = +1.229 V
cathode: 2H+(aq) + 2e− ⟶ H2(g) E°cathode = 0 V
cell:2H2 O(l) ⟶ 2H2(g) + O2(g)E°cell = −1.229 V
Again, the cell potential as written is negative, indicating a nons