I was a young teenager when I re-meme-bered the electrolysis of salt water. I was not yet a chemist, but I was keenly interested in science. The chemistry sets of the day did little to encourage an interest in chemistry; they went so far out of their way to make sure that nothing in the set could possibly be dangerous that little remained to interest the inquisitive mind. My scientific passion was inflamed, not by the watered-down kiddie-science of the sixties and seventies, but by the boy-scientist[1] books of the twenties and thirties. Among the admittedly dangerous activities described there, the electrolysis of salt water sparked my interest. With a few common materials I could make a poison gas on the one hand and an explosive gas on the other. Is this a great world, or what?

In Chapter 21 we learned that a voltaic cell, or battery, separates the two half-reactions of a redox reaction so that they take place at two different electrodes; the oxidation takes place at the anode and the reduction at the cathode. On a process schematic the electrons are shown explicitly as if they were reactants and products and these electron "pipes" represent the wires connected to the battery. The wire on the reactant side of the schematic represents the "+" terminal of the battery and the one on the product side represents the "-" terminal. The "desire" of the electrons to cross from anode to cathode is expressed as the electromotive force (EMF), with units of volts. The rate of flow of electrons from anode to cathode is expressed as the current, with units of amps. As a convention, electrons flow from top to bottom in the process schematic of a voltaic cell, as if they were flowing downhill. Reversing the normal direction of electron flow results in an electrolytic cell, one in which electrons are "pumped" from the bottom to the top of the process schematic. Of course, "down" and "up" are merely conventional directions for these process schematics and have nothing to do with the physical geometry of the cell. Let us now explore three applications of electrolysis, the lead-acid storage battery, the Hall-Héroult aluminum process, and the Castner-Kellner chloralkali process.

Figure 25-1 shows a process schematic of the lead-acid storage battery; the anode is lead, the cathode is lead oxide, and the electrolyte is sulfuric acid. If the lead and lead oxide were to come into contact, they would react with one another, electrons flowing from one to the other. But by separating the lead from the lead oxide, we can require the electrons to pass through an external circuit to get from the anode to the cathode. Lead and lead oxide are consumed, lead sulfate is produced, and the sulfuric acid in the electrolyte is converted to water. The astute reader will be able to balance the lead-acid redox reaction after examining the schematic; reactants for the skeleton reaction are on the left of the figure and products are on the right:

? Pb(s) + ? PbO2(s) + ? H2SO4(aq) = ? PbSO4(s) + H2O(l)

Figure 25-1. The Lead-Acid Cell

In operation, the negative terminal of the battery connects to the lead anode and the positive terminal connects to the lead oxide cathode. But suppose for a moment that we have access to a stronger battery than the lead-acid cell, that is, one with a higher EMF. Further suppose that we connect the two batteries anode-to-anode and cathode-to-cathode. Which direction will electrons flow? Normally, electrons flow out of the anode, but with two anodes connected to one another, the anode of the stronger battery "wins," pushing electrons into the weaker battery and converting its anode into a cathode. Similarly the stronger cathode pulls electrons from the weaker battery, converting its cathode into an anode. This drives the redox reaction of the weaker battery backwards, turning reactants into products and vice versa. In other words, electrolysis recharges the lead-acid cell. In practice, the EMF needed to recharge a car battery comes from the car's alternator rather than from another battery. But no matter what the source of the electrical power, the process schematic of the electrolytic cell is the mirror image of the voltaic one.

A second example of an electrolytic process, the Hall-Héroult process, has an operating temperature of 950°C. The electrolyte consists of molten alumina, Al2O3, and because the melting point of alumina is 2051°C, a flux is used to lower its melting point. The cryolite flux, Na3AlF6, is not consumed and so does not appear on the process schematic, Figure 25-2. The molten aluminum produced by the process serves as the cathode. Since aluminum is more dense than alumina, it sinks to the bottom of the reactor as a separate liquid phase. The carbon anode is consumed in the process, yielding carbon dioxide gas. As with the lead-acid battery, the reactants and products of the skeleton reaction appear on the schematic and from them the redox reaction may be balanced:

? Al2O3(l) + ? C(s) = ? Al(l) + ? CO2(g)

Figure 25-2. The Hall-Héroult Process

A third example of electrolysis is provided by the chloralkali cell, shown in Figure 25-3. Similarly to the lead-acid battery, the electrolyte is aqueous, in this case a solution of sodium chloride. The anodic half-reaction is:

2 NaCl(aq) = Cl2(g) + 2 Na+(aq) + 2 e-

Figure 25-3. The Chloralkali Process

The cathodic half-reaction converts water into hydrogen gas and hydroxide ion:

2 Na+(aq) + 2 H2O(l) + 2 e- = H2(g) + 2 NaOH(aq)

A simple chloralkali cell consists simply of two inert electrodes in a salt solution. But a complication arises because the chlorine gas produced at the anode reacts with sodium hydroxide to produce sodium hypochlorite, NaOCl, the active ingredient in ordinary laundry bleach:

Cl2(g) + 2 NaOH(aq) = NaOCl(aq) + NaCl(aq) + H2O(l)

Though sodium hypochlorite is a useful product, a simple chloralkali cell produces a mixture of aqueous sodium chloride, sodium hydroxide, and sodium hypochlorite. The Castner-Kellner cell was designed to produce pure chlorine and sodium hydroxide by physically separating them from one another. The anodic half-reaction is the same as before, but the Castner-Kellner electrolysis cell employs a metallic mercury[2] cathode, at which the half-reaction is:

2 Na+(aq) + 2 e- = 2 Na(Hg)

The product of this half-reaction is metallic sodium dissolved in metallic mercury. This mercuric solution flows from the electrolysis cell to a separate tank where its dissolved sodium reacts with water:

2 Na(Hg) + 2 H2O(l) = H2(g) + 2 NaOH(aq)

The sum of these last two reactions is the same as our previous cathodic reaction; the Castner-Kellner cell simply ensures that the chlorine and sodium hydroxide are produced in separate containers. Relieved of sodium, valuable mercury is returned to the electrolysis cell and every effort is made to conserve it. Because it is not consumed in the process, mercury does not appear in Figure 25-3.

We must touch on one more topic before moving on to the construction of a simple chloralkali cell. In Chapter 15 we learned to use the formula weight to answer stoichiometric questions. In practice, electrons are measured out, not by weight, but by a combination of current and time. The unit factor, (96,500 amp·sec/mol e-) may be used to answer stoichiometric questions involving voltaic and electrolytic cells. For example:

Q: How long will it take for a chloralkali cell passing a current of 100 mA to produce 500 mL of hydrogen gas?

WarningMaterial Safety

Locate MSDS's for chlorine (CAS 7782-50-5), sodium hydroxide (CAS 1310-73-2), sodium chloride (CAS 7647-14-5), sodium hypochlorite (CAS 7681-52-9), and hydrogen (CAS 1333-74-0). Summarize the hazardous properties in your notebook, including the identity of the company which produced each MSDS and either the NFPA or HMIS ratings for each material.[3]

Your most likely exposure is to chlorine gas. If breathing becomes difficult get plenty of fresh air and call for an ambulance. Be aware that hydrogen gas is flammable and that sodium hypochlorite will bleach clothing.

You should wear safety glasses while working on this project. All activities should be performed in a fume hood or with adequate ventilation. Leftover solution may be flushed down the drain. The hydrogen produced should be carefully burned. The chlorine should be released outdoors or into a fume hood.

NoteResearch and Development

You should not remain ignorant if you are to proceed in the Work.

  • Know the meanings of those words from this chapter worthy of inclusion in the index or glossary.

  • You should have mastered the Research and Development items of Chapter 21 and Chapter 22.

  • Know one new unit factor from this chapter and be able to use it to answer stoichiometric questions involving electrons.

  • Know the hazardous properties of chlorine, sodium hypochlorite, sodium hydroxide, and hydrogen.

  • Know the balanced redox reactions for the lead-acid cell, the Hall-Héroult cell, and the Castner-Kellner cell.

  • Be able to distinguish between the hazardous properties of mercury, mercuric chloride, and methyl mercury.

  • Know the contributions of Dow, Castner, Kellner, Hall, Héroult, Faraday, and Davy to the understanding of the spark.



Such books may also be of interest to girl-scientists who are capable of overlooking the un-apologetic sexism of the era in which they were written.


Mercury has acquired a bad reputation, some of it well-deserved. While metallic mercury is essentially non-toxic and quite insoluble in water, some small fraction of the mercury used in the Castner-Kellner cell will make its way into waste water. Discharged into lakes and streams, bacteria metabolize it into very toxic compounds such as methyl mercury. These compounds make their way up the food chain, being concentrated in the carnivores at the top, such as the tuna. Like the Leblanc soda industry before it, the chloralkali industry has been pressured to address its pollution problems and has gradually replaced aging Castner-Kellner cells with other designs which do not use mercury.


The NFPA diamond was introduced in Section 15.2. You may substitute HMIS or Saf-T-Data ratings at your convenience.