Electrolysis

Electrolysis

Electrolysis of water

Production of sodium and chlorine

Production of magnesium

Production of sodium hydroxide, chlorine and hydrogen

Production of aluminum

Refining of copper

Electroplating

Resources

Electrolysis is the process of causing a chemical reaction to occur by passing an electric current through a substance or mixture of substances, most often in liquid form. Electrolysis frequently results in the decomposition of a compound into its elements. To carry out an electrolysis, two electrodes, a positive electrode (anode) and a negative electrode (cathode), are immersed into the material to be electrolyzed and connected to a source of direct (DC) electric current.

The apparatus in which electrolysis is carried out is called an electrolytic cell. The roots -lys and -lyt come from the Greek lysis and lytos, meaning to cut or decompose; electrolysis in an electrolytic cell is a process that can decompose a substance.

The substance being electrolyzed must be an electrolyte, a liquid that contains positive and negative ions and therefore is able to conduct electricity. There are two kinds of electrolytes. One kind is an ion compound solution of any compound that produces ions when it dissolves in water, such as an inorganic acid, base, or salt. The other kind is a liquefied ionic compound such as a molten salt.

In either kind of electrolyte, the liquid conducts electricity because its positive and negative ions are free to move toward the electrodes of opposite charge—the positive ions toward the cathode and the negative ions toward the anode. This transfer of positive charge in one direction and negative charge in the opposite direction constitutes an electric current because an electric current is, after all, only a flow of charge, and it does not matter whether the carriers of the charge are ions or electrons. In an ionic solid such as sodium chloride, for example, the normally fixed-in-place ions become free to move as soon as the solid is dissolved in water or as soon as it is melted.

During electrolysis, the ions move toward the electrodes of opposite charge. When they reach their respective electrodes, they undergo chemical oxidation-reduction reactions. At the cathode, which is pumping electrons into the electrolyte, chemical reduction takes place—a taking on of electrons by the positive ions. At the anode, which is removing electrons out of the electrolyte, chemical oxidation takes place— a loss of electrons by the negative ions.

In electrolysis, there is a direct relationship between the amount of electricity that flows through the cell and the amount of chemical reaction that takes place. The more electrons are pumped through the electrolyte by the battery, the more ions will be forced to give up or take on electrons, thereby being oxidized or reduced. To produce one mole’s worth of chemical reaction, one mole of electrons must pass through the cell. A mole of electrons, that is, 6.02←× 10²³ of electrons, is called a faraday. The unit is named after Michael Faraday (1791–1867), the English chemist and physicist who discovered this relationship between electricity and chemical change. He is also credited with first using the words anode, cathode, electrode, electrolyte, and electrolysis.

Various kinds of electrolytic cells can be devised to accomplish specific chemical objectives.

Electrolysis of water

Perhaps the best known example of electrolysis is the electrolytic decomposition of water to produce hydrogen and oxygen:

Because water is such a stable compound, scientists can only make this reaction go by pumping energy into it—in this case, in the form of an electric current. Pure water, which does not conduct electricity very well, must, first, be made into an electrolyte by dissolving an acid, base, or salt in it. Then, an anode and a cathode, usually made of graphite or some non-reacting metal such as platinum, can be inserted and connected to a battery or other source of direct current.

At the cathode, where electrons are being pumped into the water by the battery, they are taken up by water molecules to form hydrogen gas:

At the anode, electrons are being removed from water molecules:

The net result of these two electrode reactions added together is

(Note that when these two equations are added together, the four H⁺ ions and four OH^- ions on the right-hand side are combined to form four H₂ O molecules, which then cancel four of the H₂ O molecules on the left-hand side.) Thus, every two molecules of water have been decomposed into two molecules of hydrogen and one molecule of oxygen.

The acid, base, or salt that made the water into an electrolyte was chosen so that its particular ions cannot be oxidized or reduced (at least at the voltage of the battery), so they do not react chemically and serve only to conduct the current through the water. Sulfuric acid, H₂ SO₄, is commonly used.

Production of sodium and chlorine

By electrolysis, common salt, sodium chloride, NaCl, can be broken down into its elements, sodium and chlorine. This is an important method for the production of sodium. It is used also for producing other alkali metals and alkaline earth metals from their salts.

To obtain sodium by electrolysis, scientists will first melt some sodium chloride by heating it above its melting point of 1, 474°F (801°C). Then they will insert two inert (non-reacting) electrodes into the melted salt. The sodium chloride must be molten in order to permit the Na⁺ and Cl^- ions to move freely between the electrodes; in solid sodium chloride, the ions are frozen in place. Finally, scientists will pass a direct electric current (DC) through the molten salt.

The negative electrode (the cathode) will attract Na⁺ ions and the positive electrode (the anode) will attract Cl^- ions, whereupon the following chemical reactions take place.

At the cathode, where electrons are being pumped in, they are being grabbed by the positive sodium ions:

At the anode, where electrons are being pumped out, they are being ripped off the chloride ions:

(The chlorine atoms immediately combine into diatomic molecules, Cl₂ .) The result is that common salt has been broken down into its elements by electricity.

Production of magnesium

Another important use of electrolysis is in the production of magnesium from seawater. Seawater is a major source of that metal, since it contains more ions of magnesium than of any other metal except sodium. First, magnesium chloride, MgCl₂, is obtained by precipitating magnesium hydroxide from seawater and dissolving it in hydrochloric acid. The magnesium chloride is, then, melted and electrolyzed. Similar to the production of sodium from molten sodium chloride, above, the molten magnesium is deposited at the cathode, while the chlorine gas is released at the anode. The overall reaction is MgCl₂ → Mg + Cl₂.

Production of sodium hydroxide, chlorine and hydrogen

Sodium hydroxide, NaOH, also known as lye and caustic soda, is one of the most important of all industrial chemicals. As of 2004, it is produced at the rate of over 25 billion pounds (11 billion kilograms) each year in the United States alone. World production, in that same year, is over 100 billion pounds (44 billion kilograms). The major method for producing it is the electrolysis of brine or salt water, a solution of common salt, sodium chloride in water. Chlorine and hydrogen gases are produced as valuable byproducts.

When an electric current is passed through salt water, the negative chloride ions, Cl^-, migrate to the positive anode and lose their electrons to become chlorine gas.

(The chlorine atoms then pair up to form Cl₂ molecules.) Meanwhile, sodium ions, Na⁺, are drawn to the negative cathode. However, they do not pick up electrons to become sodium metal atoms as they do in molten salt. This is because in a water solution the water molecules themselves pick up electrons more easily than sodium ions do. What happens at the cathode, then, is

The hydroxide ions, together with the sodium ions that are already in the solution, constitute sodium hydroxide, which can be recovered by evaporation.

This so-called chloralkali process is the basis of an industry that has existed for well over one hundred years. By electricity, it converts cheap salt into valuable chlorine, hydrogen, and sodium hydroxide. Among other uses, the chlorine is used in the purification of water, the hydrogen is used in the hydrogenation of oils, and the lye is used in making soap, industrial drain and oven cleaner, and paper.

Production of aluminum

The production of aluminum by the Hall process was one of the earliest applications of electrolysis on a large scale, and is still the major method for obtaining that very useful metal. Charles M. Hall, a 21-year-old student at Oberlin College in Ohio, who had been searching for a way to reduce aluminum oxide to the metal, discovered the process in 1886. Aluminum was a rare and expensive luxury at that time, because the metal is very reactive and therefore difficult to reduce from its compounds by chemical means. On the other hand, electrolysis of a molten aluminum salt or oxide is difficult because the salts are hard to obtain in anhydrous (dry) form and the oxide, Al₂ O₃, does not melt until 3, 762°F (2, 072°C).

Hall discovered that Al₂ O₃, in the form of the mineral bauxite, dissolves in another aluminum mineral called cryolite, Na₃ AlF₆, and that the resulting mixture could be melted easily. When an electric current is passed through this molten mixture, the aluminum ions migrate to the cathode, where they are reduced to metal:

At the anode, oxide ions are oxidized to oxygen gas:

The molten aluminum metal sinks to the bottom of the cell and can be drawn off.

Notice that three moles of electrons (three faradays of electricity) are needed to produce each mole of aluminum, because there are three positive charges on each aluminum ion that must be neutralized by electrons. The production of aluminum by the Hall process, therefore, consumes huge amounts of electrical energy. The recycling of beverage cans and other aluminum objects has become an important energy conservation measure.

Refining of copper

Unlike aluminum, copper metal is fairly easy to obtain chemically from its ores. But by electrolysis, it can be refined and made very pure—up to 99.999%. Pure copper is important in making electrical wire, because copper’s electrical conductivity is reduced by impurities. These impurities include such valuable metals as silver, gold, and platinum; when they are removed by electrolysis and recovered, they go a long way toward paying electricity bill.

In the electrolytic refining of copper, the impure copper is made from the anode in an electrolyte bath of copper sulfate, CuSO₄, and sulfuric acid H₂ SO₄ . The cathode is a sheet of pure copper. As current is passed through the solution, positive copper ions, Cu²⁺, in the solution are attracted to the negative cathode, where they take on electrons and deposit themselves as neutral copper atoms, thereby building up more and more pure copper on the cathode. Meanwhile, copper atoms in the positive anode give up electrons and dissolve into the electrolyte solution as copper ions. However, the impurities in the anode do not go into solution because silver, gold, and platinum atoms are not as easily oxidized (converted into positive ions) as copper is oxidized. So the silver, gold, and platinum simply fall from the anode to the bottom of the tank, where they can be scraped up.

Electroplating

Another important use of electrolytic cells is in the electroplating of silver, gold, chromium, and nickel. Electroplating produces a thin coating of these expensive metals on the surfaces of cheaper metals in order to give them the appearance and the chemical resistance of the expensive ones.

In silver plating, the object to be plated (i.e., a spoon) is made from the cathode of an electrolytic cell. The anode is a bar of silver metal, and the electrolyte (the liquid in between the electrodes) is a solution of silver cyanide, AgCN, in water. When a direct current is passed through the cell, positive silver ions (Ag⁺) from the silver cyanide migrate to the negative anode (the spoon), where they are neutralized by electrons and stick to the spoon as silver metal:

Meanwhile, the silver anode bar gives up electrons to become silver ions:

Thus, the anode bar gradually dissolves to replenish the silver ions in the solution. The net result is that silver metal has been transferred from the anode to the cathode, in this case the spoon. This process continues until the desired coating thickness is built up on the spoon—usually only a few thousandths of an inch—or until the silver bar has completely dissolved.

In electroplating with silver, silver cyanide is used in the electrolyte rather than other compounds of silver such as silver nitrate, AgNO₃, because the cyanide ion, CN^-, reacts with silver ion, Ag⁺, to form the complex ion Ag(CN)₂^-. This limits the supply of free Ag⁺ ions in the solution, so they can deposit themselves only gradually onto the cathode. This produces shinier and more adherent silver plating. Gold plating

KEY TERMS

Complex ion— A large ion that is made up of smaller ions, combined with each other or with other atoms or molecules

Faraday— A unit of electrical charge equal to the amount of charge carried by one mole of electrons. One faraday is equivalent to 96, 485 coulombs.

Oxidation— The process in which an atom’s oxidation state is increased, by its losing one or more electrons.

Reduction— The process by which an atom’s oxidation state is decreased, by its gaining one or more electrons.

is done in much the same way, using a gold anode and an electrolyte containing gold cyanide, AuCN.

Resources

BOOKS

Chang, Raymond. Chemistry. Boston, MA: McGraw-Hill, 2002.

Moog, Richard Samuel. Chemistry: A Guide Inquiry. New York: Wiley, 2005.

Tro, Nivaldo J. Introductory Chemistry. Upper Saddle River, NJ: Pearson Education, 2006.

Robert L. Wolke