Electrolysis
Electrolysis
Production of sodium and chlorine
Production of sodium hydroxide, chlorine and hydrogen
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
Electrolysis
Electrolysis
Electrolysis is a process by which electrical energy is used to produce a chemical change. Perhaps the most familiar example of electrolysis is the decomposition (breakdown) of water into hydrogen and oxygen by means of an electric current. The same process can be used to decompose compounds other than water. Sodium, chlorine, magnesium, and aluminum are four elements produced commercially by electrolysis.
Principles
The electrolysis of water illustrates the changes that take place when an electric current passes through a chemical compound. Water consists of water molecules, represented by the formula H2O. In any sample of water, some small fraction of molecules exist in the form of ions, or charged particles. Ions are formed in water when water molecules break apart to form positively charged hydrogen ions and negatively charged hydroxide ions. Chemists describe that process with the following chemical equation:
H2O → H+ + OH−
In order for electrolysis to occur, ions must exist. Seawater can be electrolyzed, for example, because it contains many positively charged sodium ions (Na+) and negatively charged chloride ions (Cl−). Any liquid, like seawater, that contains ions is called an electrolyte.
Water is not usually considered an electrolyte because it contains so few hydrogen and hydroxide ions. Normally, only one water molecule out of two billion ionizes. In contrast, sodium chloride (table salt) breaks apart completely when dissolved in water. A salt water solution consists entirely of sodium ions and chloride ions.
In order to electrolyze water, then, one prior step is necessary. Some substance, similar to sodium chloride, must be added to water to make it an electrolyte. The substance that is usually used is sulfuric acid.
The electrolysis process
The equipment used for electrolysis of a compound consists of three parts: a source of DC (direct) current; two electrodes; and an electrolyte. A common arrangement consists of a battery (the source of current) whose two poles are attached to two strips of platinum metal (the electrodes), which are immersed in water to which a few drops of sulfuric acid have been added (the electrolyte).
Electrolysis begins when electrical current (a flow of electrons) flows out of one pole of the battery into one electrode, the cathode. Positive hydrogen ions (H+) in the electrolyte pick up electrons from that electrode and become neutral hydrogen molecules (H2):
2 H+ + 2 e− → H2
(Hydrogen molecules are written as H2 because they always occur as pairs of hydrogen atoms. The same is true for molecules of oxygen, O2.)
As the electrolysis of water occurs, one can see tiny bubbles escaping from the electrolyte at the cathode. These are bubbles of hydrogen gas.
Words to Know
Anode: The electrode in an electrolytic cell through which electrons move from the electrolyte to the battery.
Cathode: The electrode in an electrolytic cell through which electrons move from the battery to the electrolyte.
Electrolyte: Any substance that, when dissolved in water, conducts an electric current.
Electrolytic cell: A system in which electrical energy is used to bring about chemical changes.
Electroplating: A process that uses an electrolytic cell to deposit a thin layer of metal on some kind of surface.
Ion: Any particle, such as an atom or molecule, that carries an electric charge.
Bubbles can also be seen escaping from the second electrode, the anode. The anode is connected to the second pole of the battery, the pole through which electrons enter the battery. At this electrode, electrons are being taken out of the electrolyte and fed back into the battery. The electrons come from negatively charged hydroxide ions (OH−), which have an excess of electrons. The anode reaction is slightly more complicated than the cathode reaction, as shown by this chemical equation:
4 OH− − 4 e− → O2 + 2 H2O
Essentially this equation says that electrons are taken away from hydroxide ions and oxygen gas is produced in the reaction. The oxygen gas bubbles off at the anode, while the extra water formed remains behind in the electrolyte.
The overall reaction that takes place in the electrolysis of water is now obvious. Electrons from the battery are given to hydrogen ions in the electrolyte, changing them into hydrogen gas. Electrons are taken from hydroxide ions in the electrolyte and transferred to the battery. Over time, water molecules are broken down to form hydrogen and oxygen molecules:
2 H2O → 2 H2 + O2
Commercial applications
Preparing elements. Electrolysis is used to break down compounds that are very stable. For example, aluminum is a very important metal in modern society. It is used in everything from pots and pans to space shuttles. But the main natural source of aluminum, aluminum oxide, is a very stable compound. A compound that is stable is difficult to break apart. You can't get aluminum out of aluminum oxide just by heating the compound—you need more energy than heat can provide.
Aluminum is prepared by an electrolytic process first discovered in 1886 by a 21-year-old student at Oberlin College in Ohio, Charles Martin Hall (1863–1914). Hall found a way of melting aluminum oxide and then electrolyzing it. Once melted, aluminum oxide forms ions of aluminum and oxygen, which behave in much the same way as hydrogen and hydroxide ions in the previous example. Pure aluminum metal is obtained at the cathode, while oxygen gas bubbles off at the anode. Sodium, chlorine, and magnesium are three other elements obtained commercially by an electrolytic process similar to the Hall process.
Refining of copper. Electrolysis can be used for purposes other than preparing elements. One example is the refining of copper. Very pure copper is often required in the manufacture of electrical equipment. (A purity of 99.999 percent is not unusual.) The easiest way to produce a product of this purity is with electrolysis.
An electrolytic cell for refining copper contains very pure copper at the cathode, impure copper at the anode, and copper sulfate as the electrolyte. When the anode and cathode are connected to a battery, electrons flow into the cathode, where they combine with copper ions (Cu2+) in the electrolyte:
Cu2+ + 2 e− → Cu0
Pure copper metal (Cu0 in the above equation) is formed on the cathode.
At the anode, copper atoms (Cu0) lose electrons and become copper ions (Cu2+) in the electrolyte:
Cu0 − 2 e− → Cu2+
Overall, the only change that occurs in the cell is that copper atoms from the impure anode become copper ions in the electrolyte. Those copper ions are then plated out on the cathode. Any impurities in the anode are just left behind, and nearly 100 percent pure copper builds up on the cathode.
Electroplating. Another important use of electrolytic cells is in the electroplating of silver, gold, chromium, and nickel. Electroplating produces a very thin coating of these expensive metals on the surfaces of cheaper metals, giving them the appearance and the chemical resistance of the expensive ones.
In silver plating, the object to be plated (a spoon, for example) is used as the cathode. A bar of silver metal is used as the anode. And the electrolyte is a solution of silver cyanide (AgCN). When this arrangement is connected to a battery, electrons flow into the cathode where they combine with silver ions (Ag+) from the electrolyte to form silver atoms (Ag0):
Ag+ + 1 e− → Ag0
These silver atoms plate out as a thin coating on the cathode—in this case, the spoon. At the anode, silver atoms give up electrons and become silver ions in the electrolyte:
Ag0 − 1 e− → Ag0
Silver is cycled, therefore, from the anode to the electrolyte to the cathode, where it is plated out.
Electrolysis
Electrolysis
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 a 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 sucking 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 inventing 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, we 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 H2O molecules, which then cancel four of the H2O 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 , H2SO4, 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, we will first melt some sodium chloride by heating it above its melting point of 1,474°F (801°C). Then we 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, we 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 sea water. Sea water 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. It is produced at the rate of 25 billion pounds a year in the United States alone. 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 Cl2molecules.) Meanwhile, sodium ions, Na+, are drawn to the negative cathode. But they do not pick up electrons to become sodium metal atoms as they do in molten salt, 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 a 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 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. The process was discovered in 1886 by 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. 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, Al2O3, does not melt until 3,762°F (2,072°C).
Hall discovered that Al2O3, in the form of the mineral bauxite, dissolves in another aluminum mineral called cryolite, Na3AlF6, and that the resulting mixture could be melted fairly 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. Therecycling 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 the 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 H2SO4. The cathode is a sheet of very 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. But 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. 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 very thin coating of these expensive metals on the surfaces of cheaper metals, to give them the appearance and the chemical resistance of the expensive ones.
In silver plating, the object to be plated (e.g., 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) -. This limits the supply of free Ag+ ions in the solution, so they can deposit themselves only very gradually onto the cathode. This produces a shinier and more adherent silver plating. Gold plating is done in much the same way, using a gold anode and an electrolyte containing gold cyanide, AuCN.
Resources
books
Chang, Raymond. Chemistry. New York: McGraw-Hill, 1991.
Sherwood, Martin, and Christine Sutton, eds. The PhysicalWorld. New York: Oxford, 1991.
Robert L. Wolke
KEY TERMS
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .- Complex ion
—A big 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.
electrolysis
e·lec·trol·y·sis / ilekˈträləsis; ˌēlek-/ • n. 1. Chem. chemical decomposition produced by passing an electric current through a liquid or solution containing ions.2. the removal of hair roots or small blemishes on the skin by the application of heat using an electric current.DERIVATIVES: e·lec·tro·lyt·ic / iˌlektrəˈlitik/ adj.e·lec·tro·lyt·i·cal / iˌlektrəˈlitikəl/ adj.e·lec·tro·lyt·i·cal·ly / iˌlektrəˈlitik(ə)lē/ adv.
electrolysis
electrolysis
1. the chemical decomposition of a substance (see electrolyte) into positively and negatively charged ions (see anion, cation) when an electric current is passed through it.
2. destruction of tissue, especially hair follicles (see epilation), by the passage of an electric current.