Volume 2
2.8. ELECTROLYSIS. REDOX REACTIONS
How Solutions of Electrolytes Conduct Electricity. What Happens When Ions Reach the Electrodes. Electrolysis of HCl and CuCl2 • Oxidation and Reduction. Redox Reactions. Oxidants and Reductants • Exercises
How Solutions of Electrolytes Conduct Electricity. What Happens When Ions Reach the Electrodes. Electrolysis of HCl and CuCl2 • Oxidation and Reduction. Redox Reactions. Oxidants and Reductants • Exercises
2.8.1. How Solutions of Electrolytes Conduct Electricity. What Happens When Ions Reach the Electrodes. Electrolysis of HCl and CuCl2. Both metals and solutions of electrolytes conduct electricity. An electric current through a metal object is a flow of free electrons (electron gas, see section 2.4.5). However, there are no free electrons in aqueous solutions of HCl, NaOH, H2SO4, NaCl, etc. Yet such solutions conduct electricity well. How do they do that?
Solutions of electrolytes are electrical conductors due to the movement of ions. Figure 2-94 schematically represents a simple apparatus for electrical conductivity tests on liquids. The setup comprises a cell (a glass beaker or bowl) filled with a solution for the test, two electrodes immersed in the solution and wired to a battery, and a light bulb included in the circuit. Electrodes for such a setup are conventionally made of graphite that is an excellent conductor of electricity while being chemically inert, nontoxic, and inexpensive.
Solutions of electrolytes are electrical conductors due to the movement of ions. Figure 2-94 schematically represents a simple apparatus for electrical conductivity tests on liquids. The setup comprises a cell (a glass beaker or bowl) filled with a solution for the test, two electrodes immersed in the solution and wired to a battery, and a light bulb included in the circuit. Electrodes for such a setup are conventionally made of graphite that is an excellent conductor of electricity while being chemically inert, nontoxic, and inexpensive.
Figure 2-94. Solutions of strong electrolytes such as aqueous HCl conduct electricity due to the movement of ions.
If we place hydrochloric acid in the cell, the bulb will light up to show that there is an electric current present in the system (Figure 2-94). The observed conductivity of HCl in water is due to the movement of the H+ and Cl- ions in the solution. Since oppositely charged objects attract each other, the positively charged H+ ions in our solution are attracted by, and move toward, the negatively charged electrode. Likewise and simultaneously, the negatively charged Cl- ions move toward the positively charged electrode. The positively charged electrode that attracts anions is called the anode. The negative one attracting cations is called the cathode.
So, it is the electric forces that cause the H+ and Cl- ions move, which makes the solution a conductor. Two questions arise, however, as follows.
1. What happens to the ions when they finally arrive at the surface of the electrodes?
2. The electric current through the wiring of our system (Figure 2-94) is a flow of electrons from the anode to the cathode. Where do these electrons come from to arrive at the anode and where do they go from the cathode? There are no free electrons in the HCl solution.
Once a Cl- anion hits the positively charged anode, the ion gets discharged by giving away one electron to become a chlorine atom. At exactly the same time, an H+ cation that reaches the surface of the cathode picks up an electron to become a hydrogen atom. These processes are shown in Figure 2-95.
If we place hydrochloric acid in the cell, the bulb will light up to show that there is an electric current present in the system (Figure 2-94). The observed conductivity of HCl in water is due to the movement of the H+ and Cl- ions in the solution. Since oppositely charged objects attract each other, the positively charged H+ ions in our solution are attracted by, and move toward, the negatively charged electrode. Likewise and simultaneously, the negatively charged Cl- ions move toward the positively charged electrode. The positively charged electrode that attracts anions is called the anode. The negative one attracting cations is called the cathode.
So, it is the electric forces that cause the H+ and Cl- ions move, which makes the solution a conductor. Two questions arise, however, as follows.
1. What happens to the ions when they finally arrive at the surface of the electrodes?
2. The electric current through the wiring of our system (Figure 2-94) is a flow of electrons from the anode to the cathode. Where do these electrons come from to arrive at the anode and where do they go from the cathode? There are no free electrons in the HCl solution.
Once a Cl- anion hits the positively charged anode, the ion gets discharged by giving away one electron to become a chlorine atom. At exactly the same time, an H+ cation that reaches the surface of the cathode picks up an electron to become a hydrogen atom. These processes are shown in Figure 2-95.
Figure 2-95. On passing an electric current through a solution of HCl, the Cl- anions give up electrons at the anode and, simultaneously, the same number of the H+ cations acquire the same number of electrons at the cathode.
As the anode accepts electrons from the discharging chloride anions, exactly the same number of electrons is donated to the hydrogen ions at the cathode. In this way, the anode is continuously pumped up with electrons from the chloride ions, while the cathode synchronously donates the same number of electrons to the protons. Together, these processes enable the flow of electrons through the circuit.
What happens to the chlorine and hydrogen atoms after their formation at the anode and cathode? Both are much less stable species than the parent anion and cation (remember the octet rule!), so they immediately seek opportunities to get stabilized. If a chlorine atom and a hydrogen atom were in close proximity, they would react to give a molecule of HCl. But they are too far away from each another to meet and react. As a result, the H and Cl atoms dimerize to give H2 at the cathode and Cl2 at the anode (Figure 2-96). Indeed, as an electric current passes through a solution of HCl, colorless H2 bubbles off the cathode and greenish-yellow Cl2 off the anode.
As the anode accepts electrons from the discharging chloride anions, exactly the same number of electrons is donated to the hydrogen ions at the cathode. In this way, the anode is continuously pumped up with electrons from the chloride ions, while the cathode synchronously donates the same number of electrons to the protons. Together, these processes enable the flow of electrons through the circuit.
What happens to the chlorine and hydrogen atoms after their formation at the anode and cathode? Both are much less stable species than the parent anion and cation (remember the octet rule!), so they immediately seek opportunities to get stabilized. If a chlorine atom and a hydrogen atom were in close proximity, they would react to give a molecule of HCl. But they are too far away from each another to meet and react. As a result, the H and Cl atoms dimerize to give H2 at the cathode and Cl2 at the anode (Figure 2-96). Indeed, as an electric current passes through a solution of HCl, colorless H2 bubbles off the cathode and greenish-yellow Cl2 off the anode.
Figure 2-96. The formation of Cl2 (anode) and H2 (cathode) on passing an electric current through aqueous HCl.
The overall reaction that takes place is:
2 HCl = H2 + Cl2
This is the reverse of the well-known spontaneous and energetic reaction of Cl2 with H2 to give HCl. A spontaneous reaction is like rolling down a hill on a bike (Figure 2-97, left). There is no need to use your muscle energy as you do not have to pedal. On the contrary, going back to H2 and Cl2 from HCl is like pedaling uphill (Figure 2-97, right), in that the decomposition of HCl requires energy. This energy is provided by the electric battery used in the experiment. A process of chemical change brought about by passing an electric current through a substance in solution (or in the molten state) is known as electrolysis.
The overall reaction that takes place is:
2 HCl = H2 + Cl2
This is the reverse of the well-known spontaneous and energetic reaction of Cl2 with H2 to give HCl. A spontaneous reaction is like rolling down a hill on a bike (Figure 2-97, left). There is no need to use your muscle energy as you do not have to pedal. On the contrary, going back to H2 and Cl2 from HCl is like pedaling uphill (Figure 2-97, right), in that the decomposition of HCl requires energy. This energy is provided by the electric battery used in the experiment. A process of chemical change brought about by passing an electric current through a substance in solution (or in the molten state) is known as electrolysis.
Figure 2-97. The reaction of H2 with Cl2 to give HCl is spontaneous, like rolling down a hill on a bike. The reverse process, decomposition of HCl, requires energy, like pedaling uphill. Paintings by Jiro Osuga.
Electrolysis of an aqueous solution of CuCl2 occurs similarly. Being a salt and, consequently, a strong electrolyte, CuCl2 is dissociated into Cu2+ and Cl- in water. On passing an electric current through a solution of CuCl2, the positively charged Cu2+ ions move toward the negatively charged cathode and the negatively charged Cl- ions toward the positively charged anode. The reactions occurring at the cathode and at the anode are presented in Figure 2-98. Video 2-8 demonstrates the electrolysis of CuCl2 in a rather simple setup. [Note that in this video the litmus paper turns white because the acid-base indicator dye is decomposed by the chlorine produced at the anode.]
Electrolysis of an aqueous solution of CuCl2 occurs similarly. Being a salt and, consequently, a strong electrolyte, CuCl2 is dissociated into Cu2+ and Cl- in water. On passing an electric current through a solution of CuCl2, the positively charged Cu2+ ions move toward the negatively charged cathode and the negatively charged Cl- ions toward the positively charged anode. The reactions occurring at the cathode and at the anode are presented in Figure 2-98. Video 2-8 demonstrates the electrolysis of CuCl2 in a rather simple setup. [Note that in this video the litmus paper turns white because the acid-base indicator dye is decomposed by the chlorine produced at the anode.]
Figure 2-98. Electrolysis of CuCl2 produces chlorine gas and copper metal.
Video 2-8. Laboratory demonstration of CuCl2 electrolysis (source).
The electrolytic decomposition reactions of CuCl2 (Figure 2-98) and HCl (Figures 2-94, 2-95, and 2-96) are very much alike, representing two simple examples of electricity-driven chemical transformations. In both cases, it is the ions produced directly upon dissociation of the electrolyte that are discharged at the electrodes. This is not always the case. Quite often the ions of an electrolyzed compound are reluctant to discharge and, as a result, it is the H2O present as the solvent that is forced by the electricity to accept or donate electrons. A typical example is the chloralkali process for the production of NaOH, Cl2, and H2 from brine, an aqueous solution of NaCl. Studying this process and other more complex examples of electrolysis is beyond the scope of our course.
The electrolytic decomposition reactions of CuCl2 (Figure 2-98) and HCl (Figures 2-94, 2-95, and 2-96) are very much alike, representing two simple examples of electricity-driven chemical transformations. In both cases, it is the ions produced directly upon dissociation of the electrolyte that are discharged at the electrodes. This is not always the case. Quite often the ions of an electrolyzed compound are reluctant to discharge and, as a result, it is the H2O present as the solvent that is forced by the electricity to accept or donate electrons. A typical example is the chloralkali process for the production of NaOH, Cl2, and H2 from brine, an aqueous solution of NaCl. Studying this process and other more complex examples of electrolysis is beyond the scope of our course.
Digression. Gold is currently about 23,000 times more expensive than aluminum. This is not surprising, especially given the fact that gold is very rare, whereas aluminum is the third most abundant element in the Earth's Crust (after oxygen and silicon). It might be hard to believe that there was a time when aluminum was much more expensive than gold. For decades since the first preparation of aluminum metal by the Danish physicist Hans Christian Ørsted in 1825, aluminum was much more expensive than gold. Although naturally occurring Al compounds were abundant, the challenge was to convert them to aluminum metal. In the mid-1850s, the French chemist Henri Étienne Sainte-Claire Deville solved the problem by developing the first industrial method to make aluminum. His method was based on electrolysis. In 1886, the advanced Hall-Héroult aluminum smelting process was developed, which was also based on electrolysis. This process is still used nowadays to produce aluminum on a tremendous scale, over 60 million metric tons a year worldwide.
Electrolysis has played a very important role in the development of chemistry, particularly in the discovery of many most highly reactive simple substances. For example, potassium and sodium metals were prepared by Humphry Davy using electrolysis in 1807. Likewise, electrolysis was used to make and isolate fluorine gas (F2), the most aggressive and reactive simple substance. This discovery was made by Henri Moissan in 1886, who was awarded the 1906 Nobel Prize for his work. For exactly 100 years since Moissan's electrochemical synthesis of fluorine, scientists were convinced that F2 could be prepared only by electrolysis. In 1986, however, the discovery of the first electricity-free, purely chemical method to make fluorine was reported.
If electrical energy can be used to force a reaction that otherwise does not occur, can a reverse process be created to make electricity from the energy released by a spontaneous chemical reaction? Such a process could be compared to cycling downhill in contrast with electrolysis that is like cycling up the hill (Figure 2-97). The answer is "yes", and I bet you have many times seen and used such chemical reactors producing electricity. These reactors are batteries. There are a number of different chemical reactions that are used in batteries. However, not every chemical reaction can be used as a source of electricity. For instance, the reverse of the electrolytic decomposition of HCl (H2 + Cl2 = 2 HCl) and CuCl2 (Cu + Cl2 = CuCl2) cannot be used to make electricity for a number of reasons. Perhaps the oldest chemical system to generate an electric current is the galvanic cell employing Zn and Cu metals and their sulfates. As the Cu2+ discharges at the copper cathode by accepting two electrons (Cu2+ + 2 e- = Cu), the Zn metal that the anode is made of dissolves while releasing two electrons (Zn - 2 e- = Zn2+). The electric current in the Cu-Zn galvanic cell is a flow of electrons from the Zn metal to the Cu2+ cations. [Note that in a galvanic cell the anode is the negatively charged electrode and the cathode is the positively charged electrode.]
Electrolysis has played a very important role in the development of chemistry, particularly in the discovery of many most highly reactive simple substances. For example, potassium and sodium metals were prepared by Humphry Davy using electrolysis in 1807. Likewise, electrolysis was used to make and isolate fluorine gas (F2), the most aggressive and reactive simple substance. This discovery was made by Henri Moissan in 1886, who was awarded the 1906 Nobel Prize for his work. For exactly 100 years since Moissan's electrochemical synthesis of fluorine, scientists were convinced that F2 could be prepared only by electrolysis. In 1986, however, the discovery of the first electricity-free, purely chemical method to make fluorine was reported.
If electrical energy can be used to force a reaction that otherwise does not occur, can a reverse process be created to make electricity from the energy released by a spontaneous chemical reaction? Such a process could be compared to cycling downhill in contrast with electrolysis that is like cycling up the hill (Figure 2-97). The answer is "yes", and I bet you have many times seen and used such chemical reactors producing electricity. These reactors are batteries. There are a number of different chemical reactions that are used in batteries. However, not every chemical reaction can be used as a source of electricity. For instance, the reverse of the electrolytic decomposition of HCl (H2 + Cl2 = 2 HCl) and CuCl2 (Cu + Cl2 = CuCl2) cannot be used to make electricity for a number of reasons. Perhaps the oldest chemical system to generate an electric current is the galvanic cell employing Zn and Cu metals and their sulfates. As the Cu2+ discharges at the copper cathode by accepting two electrons (Cu2+ + 2 e- = Cu), the Zn metal that the anode is made of dissolves while releasing two electrons (Zn - 2 e- = Zn2+). The electric current in the Cu-Zn galvanic cell is a flow of electrons from the Zn metal to the Cu2+ cations. [Note that in a galvanic cell the anode is the negatively charged electrode and the cathode is the positively charged electrode.]
2.8.2. Oxidation and Reduction. Redox Reactions. Oxidants and Reductants. We have just learned that on contact with the anode, a chloride anion is discharged by losing one electron to become a chlorine atom (Figures 2-95, 2-96, 2-98). A process in which one or more electrons are lost by an ion, atom or molecule is called oxidation.
As the Cl- anion loses electrons during the electrolysis of HCl and CuCl2, the H+ or Cu2+ cations acquire electrons at the cathode to give H2 and Cu metal, respectively. A process of acquiring one or more electrons by an ion, atom or molecule is called reduction. Figure 2-99 provides a few simple examples of oxidation and reduction.
As the Cl- anion loses electrons during the electrolysis of HCl and CuCl2, the H+ or Cu2+ cations acquire electrons at the cathode to give H2 and Cu metal, respectively. A process of acquiring one or more electrons by an ion, atom or molecule is called reduction. Figure 2-99 provides a few simple examples of oxidation and reduction.
Figure 2-99. Examples of oxidation and reduction.
Importantly, atoms, molecules, and ions that are capable of losing electrons do so only if there are species around to accept them. Electrons are matter. They neither disappear nor come from nowhere, but can only be transferred from one species to another. Consequently, an oxidation process is always coupled with a reduction process and vice versa in such a manner that the number of electrons given away by one species always equals the number of electrons accepted by another species. A solely oxidation or solely reduction process (such as those shown in Figure 2-99) that cannot occur on its own, but only in the presence of a partner process for electron transfer, is often referred to as a half-reaction.
The combination of an oxidation half-reaction and a reduction half-reaction is a redox reaction, an abbreviation for "reduction-oxidation". A typical redox reaction is that between sodium and chlorine, a combination of the two half-reactions presented in the top two lines of Figure 2-99. Sodium metal and chlorine gas do react very energetically to give NaCl, as shown in Video 2-9.
Importantly, atoms, molecules, and ions that are capable of losing electrons do so only if there are species around to accept them. Electrons are matter. They neither disappear nor come from nowhere, but can only be transferred from one species to another. Consequently, an oxidation process is always coupled with a reduction process and vice versa in such a manner that the number of electrons given away by one species always equals the number of electrons accepted by another species. A solely oxidation or solely reduction process (such as those shown in Figure 2-99) that cannot occur on its own, but only in the presence of a partner process for electron transfer, is often referred to as a half-reaction.
The combination of an oxidation half-reaction and a reduction half-reaction is a redox reaction, an abbreviation for "reduction-oxidation". A typical redox reaction is that between sodium and chlorine, a combination of the two half-reactions presented in the top two lines of Figure 2-99. Sodium metal and chlorine gas do react very energetically to give NaCl, as shown in Video 2-9.
Video 2-9. Redox reaction of chlorine gas (Cl2) with sodium metal (Na) (source).
To write a balanced equation for the reaction of sodium with chlorine, we add together the two half-reaction equations (Figure 2-100).
To write a balanced equation for the reaction of sodium with chlorine, we add together the two half-reaction equations (Figure 2-100).
Figure 2-100. Deriving a chemical equation from oxidation and reduction half-reaction equations.
Since one Na atom releases one electron and one Cl atom accepts one electron, the balancing is easy (Figure 2-100). First, the electrons can be cancelled out. Second, the product should be written in its molecular (NaCl) rather than ionic form (Na+ + Cl-) because the reaction gives rise to solid sodium chloride. Finally, we have to take into account the fact that chlorine exists in the form of diatomic molecules (Cl2) not individual chlorine atoms (Cl).
The two half-reactions at the bottom of Figure 2-99 can also be combined into one full equation in exactly the same way (Figure 2-101).
Since one Na atom releases one electron and one Cl atom accepts one electron, the balancing is easy (Figure 2-100). First, the electrons can be cancelled out. Second, the product should be written in its molecular (NaCl) rather than ionic form (Na+ + Cl-) because the reaction gives rise to solid sodium chloride. Finally, we have to take into account the fact that chlorine exists in the form of diatomic molecules (Cl2) not individual chlorine atoms (Cl).
The two half-reactions at the bottom of Figure 2-99 can also be combined into one full equation in exactly the same way (Figure 2-101).
Video 2-10. Calcium metal (Ca) reacts with oxygen (O2) of the air to produce calcium oxide (CaO) (source).
To avoid confusion, you need to master the terminology of redox reactions.
Reduction is a process of the gaining of one or more electrons by an atom, molecule or ion in a chemical reaction.
Oxidation is the loss of one or more electrons by an atom, molecule or ion in a chemical reaction.
A reductant or reducing agent or reducer is an atom, molecule or ion that donates electrons in a redox reaction.
An oxidant or oxidizing agent or oxidizer is an atom, molecule or ion that gains electrons in a redox reaction.
An atom, molecule or ion is oxidized if it loses electrons or, in other words, serves as a reductant in a chemical reaction.
When an atom, molecule or ion gains electrons or, in other words, serves as an oxidant in a chemical reaction, it is said that this atom, molecule, or ion is reduced.
A reductant reduces an oxidant and an oxidant oxidizes a reductant.
Review Figure 2-102 to understand and memorize the terminology of redox reactions.
To avoid confusion, you need to master the terminology of redox reactions.
Reduction is a process of the gaining of one or more electrons by an atom, molecule or ion in a chemical reaction.
Oxidation is the loss of one or more electrons by an atom, molecule or ion in a chemical reaction.
A reductant or reducing agent or reducer is an atom, molecule or ion that donates electrons in a redox reaction.
An oxidant or oxidizing agent or oxidizer is an atom, molecule or ion that gains electrons in a redox reaction.
An atom, molecule or ion is oxidized if it loses electrons or, in other words, serves as a reductant in a chemical reaction.
When an atom, molecule or ion gains electrons or, in other words, serves as an oxidant in a chemical reaction, it is said that this atom, molecule, or ion is reduced.
A reductant reduces an oxidant and an oxidant oxidizes a reductant.
Review Figure 2-102 to understand and memorize the terminology of redox reactions.
Figure 2-102. Redox reaction terminology, as exemplified by the reaction between Na and Cl2.
2.8.3. Exercises.
1. Solutions of electrolytes (a) conduct electricity exclusively due to the presence of positively charged ions; (b) conduct electricity exclusively due to the presence of negatively charged ions; (c) conduct electricity due to the presence of both positively and negatively charged ions; (d) do not conduct electricity; (e) conduct electricity due to the formation of free electrons from negatively charged ions on dissociation. Answer
2. Dissociation of an electrolyte is a reversible process, whereas the discharge of the resultant ions at the electrodes during electrolysis is not. True or false? Answer
3. During electrolysis, positively charged ions in the electrolyte solution (a) move toward the anode; (b) move toward the cathode; (c) do not move toward either electrode; only lighter and faster negatively charged ions do. Answer
4. It is the ions produced upon electrolytic dissociation that always get discharged in any electrolysis reaction. True or false? Answer
5. An oxidant is a molecule or ion that (a) donates electrons; (b) accepts electrons; (c) either accepts or donates electrons. Answer
6. There are chemical reactions that involve only oxidation or only reduction. True or false? Answer
7. In the vigorous reaction between magnesium metal and bromine (Mg + Br2 = MgBr2), (a) Mg is the oxidizer that reduces bromine; (b) Mg is the reductant that oxidizes bromine; (c) Mg is oxidized and Br is reduced; (d) Br2 is the oxidant that oxidizes Mg; (e) Br2 is the oxidant that reduces Mg; (f) Br2 is the oxidant that undergoes reduction; (g) Br2 is the reductant that undergoes reduction; (h) Mg is the reductant that undergoes oxidation. Answer
2.8.3. Exercises.
1. Solutions of electrolytes (a) conduct electricity exclusively due to the presence of positively charged ions; (b) conduct electricity exclusively due to the presence of negatively charged ions; (c) conduct electricity due to the presence of both positively and negatively charged ions; (d) do not conduct electricity; (e) conduct electricity due to the formation of free electrons from negatively charged ions on dissociation. Answer
2. Dissociation of an electrolyte is a reversible process, whereas the discharge of the resultant ions at the electrodes during electrolysis is not. True or false? Answer
3. During electrolysis, positively charged ions in the electrolyte solution (a) move toward the anode; (b) move toward the cathode; (c) do not move toward either electrode; only lighter and faster negatively charged ions do. Answer
4. It is the ions produced upon electrolytic dissociation that always get discharged in any electrolysis reaction. True or false? Answer
5. An oxidant is a molecule or ion that (a) donates electrons; (b) accepts electrons; (c) either accepts or donates electrons. Answer
6. There are chemical reactions that involve only oxidation or only reduction. True or false? Answer
7. In the vigorous reaction between magnesium metal and bromine (Mg + Br2 = MgBr2), (a) Mg is the oxidizer that reduces bromine; (b) Mg is the reductant that oxidizes bromine; (c) Mg is oxidized and Br is reduced; (d) Br2 is the oxidant that oxidizes Mg; (e) Br2 is the oxidant that reduces Mg; (f) Br2 is the oxidant that undergoes reduction; (g) Br2 is the reductant that undergoes reduction; (h) Mg is the reductant that undergoes oxidation. Answer