Volume 3
3.7. ALUMINUM
Natural Occurrence and Physical Properties • Production and Uses of Aluminum • Chemical Properties and Amphoterism of Aluminum. Is Aluminum a Metal or Nonmetal? • Exercises
Natural Occurrence and Physical Properties • Production and Uses of Aluminum • Chemical Properties and Amphoterism of Aluminum. Is Aluminum a Metal or Nonmetal? • Exercises
3.7.1. Natural Occurrence and Physical Properties. Aluminum is a silvery-white, quite soft malleable metal that melts at 660 oC. Aluminum is the most abundant metal and the third most abundant element in the Earth's crust (after oxygen and silicon), occurring in nature in the form of various oxide, hydroxide, and silicate derivatives. Rubies and sapphires are Al2O3 colored by impurities of other metals. Beryl and some types of garnet are aluminum compounds. Feldspar rocks, the most common group of minerals in the Earth's crust, are composed of Al, Si, O, and an alkali (Na, K) or alkaline earth (Ca) metal.
3.7.2. Production and Uses of Aluminum. While there are many different aluminum minerals on Earth, only one group of them, bauxites, is used to make aluminum metal. Making aluminum from other minerals would be less economical. To produce aluminum, pure aluminum oxide (Al2O3), also known as alumina, is needed. The main components of bauxite, Al(OH)3 and AlO(OH), could be quite easily converted to alumina by dehydration at heating. First, however, some iron and titanium impurities present in bauxites must be removed. Without going into details, the Bayer alumina process furnishes pure Al2O3 suitable for making aluminum.
In Volume 2, it was mentioned that aluminum is made by the Hall-Héroult electrolytic process (section 2.8.1). In this highly energy demanding process, Al2O3 is electrolyzed to reduce Al3+ to Al0 at the cathode. The electrolysis is conducted in an Al2O3 solution in molten cryolite (Na3AlF6) at 950-1,000 oC. Note that pure Al2O3 melts at approximately 2,070 oC, too high a temperature for a large-scale operation. On passing electricity through a solution of alumina in molten cryolite, the aluminum ions move toward the cathode and the oxide ions toward the anode. Both electrodes are made of carbon, the anode of fused coke and the cathode of graphite. Figure 3-88 displays equations for the processes occurring at the anode and cathode in the electrolysis of Al2O3.
3.7.2. Production and Uses of Aluminum. While there are many different aluminum minerals on Earth, only one group of them, bauxites, is used to make aluminum metal. Making aluminum from other minerals would be less economical. To produce aluminum, pure aluminum oxide (Al2O3), also known as alumina, is needed. The main components of bauxite, Al(OH)3 and AlO(OH), could be quite easily converted to alumina by dehydration at heating. First, however, some iron and titanium impurities present in bauxites must be removed. Without going into details, the Bayer alumina process furnishes pure Al2O3 suitable for making aluminum.
In Volume 2, it was mentioned that aluminum is made by the Hall-Héroult electrolytic process (section 2.8.1). In this highly energy demanding process, Al2O3 is electrolyzed to reduce Al3+ to Al0 at the cathode. The electrolysis is conducted in an Al2O3 solution in molten cryolite (Na3AlF6) at 950-1,000 oC. Note that pure Al2O3 melts at approximately 2,070 oC, too high a temperature for a large-scale operation. On passing electricity through a solution of alumina in molten cryolite, the aluminum ions move toward the cathode and the oxide ions toward the anode. Both electrodes are made of carbon, the anode of fused coke and the cathode of graphite. Figure 3-88 displays equations for the processes occurring at the anode and cathode in the electrolysis of Al2O3.
Figure 3-88. Transformations occurring in the Hall-Héroult process.
At the cathode, aluminum metal is produced due to the reduction of the Al3+ cations. At the anode, the oxide anions, O2-, are discharged to form oxygen. The O2 formed oxidizes the carbon, the material that the anode is made of, to CO and CO2. As the electrochemical process occurs, the carbon anode is consumed and has to be periodically replaced.
Aluminum is produced on a huge scale and is the second most widely used metal (after iron). Being non-toxic, non-carcinogenic, very light (density 2.7 g/cm3), resistant to corrosion, greatly conductive to electricity, readily available, and inexpensive aluminum is a very attractive material for a broad variety of applications.
Pure aluminum is rarely used as a material because of its softness. However, mechanical properties of aluminum can be vastly improved by alloying it with just a fraction to a few percent of Cu, Mg, Zn, Mn, and Si. Aluminum alloys are used to make cans and foil, parts and bodies of cars, trucks, aircraft skins, spacecraft, and railway cars, electric motors and generators, windows and doors, cooking utensils, and electric wire. The list of applications of aluminum compounds is as impressive. Anhydrous aluminum chloride (AlCl3) is broadly employed as a catalyst and promoter in petrochemical and chemical industries as well as in the manufacturing of pharmaceuticals and agrochemicals. Alumina (Al2O3) is also used as a catalyst and support for a broad variety of catalysts. Being just slightly less hard than diamond, alumina is used to make abrasives. Powdered alumina is an excellent adsorbent and desiccant. Synthetic ruby and sapphire are made of Al2O3. Aluminum compounds are also used in the production of deodorants, in water treatment, paper manufacturing, and leather tanning.
3.7.3. Chemical Properties and Amphoterism of Aluminum. Is Aluminum a Metal or Nonmetal? An aluminum atom has three valence electrons in the 3rd electron shell: 1s2 2s2 2p6 3s2 3p1. These three electrons are easily discarded by an aluminum atom in order to attain the stable electron configuration of neon. Aluminum is very easily oxidized by O2 and even by water. Would you question this statement, given the fact that aluminum wire, foil, and other aluminum items are virtually indefinitely stable in moist air and even in water, exhibiting no sign of oxidation?
As a matter of fact, aluminum is immediately oxidized to alumina on exposure to air (Figure 3-89). However, the very thin film of the Al2O3 produced on the metal surface is strikingly good at blocking access of the oxygen to the aluminum atoms underneath. In this way, the bulk of the aluminum metal is remarkably well protected from oxidation. Watch this video to see the melting of an aluminum rod with a torch in the air. Note the thin Al2O3 film on the surface of the molten metal. Also note that due to this film of alumina the aluminum rod does not catch a fire in spite of the very high temperature of the flame and plenty of oxygen around.
At the cathode, aluminum metal is produced due to the reduction of the Al3+ cations. At the anode, the oxide anions, O2-, are discharged to form oxygen. The O2 formed oxidizes the carbon, the material that the anode is made of, to CO and CO2. As the electrochemical process occurs, the carbon anode is consumed and has to be periodically replaced.
Aluminum is produced on a huge scale and is the second most widely used metal (after iron). Being non-toxic, non-carcinogenic, very light (density 2.7 g/cm3), resistant to corrosion, greatly conductive to electricity, readily available, and inexpensive aluminum is a very attractive material for a broad variety of applications.
Pure aluminum is rarely used as a material because of its softness. However, mechanical properties of aluminum can be vastly improved by alloying it with just a fraction to a few percent of Cu, Mg, Zn, Mn, and Si. Aluminum alloys are used to make cans and foil, parts and bodies of cars, trucks, aircraft skins, spacecraft, and railway cars, electric motors and generators, windows and doors, cooking utensils, and electric wire. The list of applications of aluminum compounds is as impressive. Anhydrous aluminum chloride (AlCl3) is broadly employed as a catalyst and promoter in petrochemical and chemical industries as well as in the manufacturing of pharmaceuticals and agrochemicals. Alumina (Al2O3) is also used as a catalyst and support for a broad variety of catalysts. Being just slightly less hard than diamond, alumina is used to make abrasives. Powdered alumina is an excellent adsorbent and desiccant. Synthetic ruby and sapphire are made of Al2O3. Aluminum compounds are also used in the production of deodorants, in water treatment, paper manufacturing, and leather tanning.
3.7.3. Chemical Properties and Amphoterism of Aluminum. Is Aluminum a Metal or Nonmetal? An aluminum atom has three valence electrons in the 3rd electron shell: 1s2 2s2 2p6 3s2 3p1. These three electrons are easily discarded by an aluminum atom in order to attain the stable electron configuration of neon. Aluminum is very easily oxidized by O2 and even by water. Would you question this statement, given the fact that aluminum wire, foil, and other aluminum items are virtually indefinitely stable in moist air and even in water, exhibiting no sign of oxidation?
As a matter of fact, aluminum is immediately oxidized to alumina on exposure to air (Figure 3-89). However, the very thin film of the Al2O3 produced on the metal surface is strikingly good at blocking access of the oxygen to the aluminum atoms underneath. In this way, the bulk of the aluminum metal is remarkably well protected from oxidation. Watch this video to see the melting of an aluminum rod with a torch in the air. Note the thin Al2O3 film on the surface of the molten metal. Also note that due to this film of alumina the aluminum rod does not catch a fire in spite of the very high temperature of the flame and plenty of oxygen around.
Figure 3-89. Reaction of aluminum with oxygen.
If we cover aluminum with mercury and scratch off the protecting layer of alumina underneath the liquid mercury layer, the two metals blend together to form an aluminum amalgam, an Al-Hg alloy. Alumina does not stick to the amalgam as tightly as it does to pure aluminum. As a result, the exposed aluminum is air-oxidized quite efficiently, as shown in this nice video demonstration. (Unfortunately, it is incorrectly implied in the open captions to this video that the mercury is "converted" during the process. While the aluminum is indeed air-oxidized to Al2O3 in the experiment, the mercury metal is used only for the amalgamation and is not converted to any mercury compound.)
The ability of aluminum metal to react with water is similar. While the Al2O3 film protects the underlying aluminum atoms from oxidation by H2O, amalgamated aluminum reacts with water to give H2 and Al(OH)3 (Figure 3-90), as demonstrated in this video. The experimenter does not tell us in the description that the reaction actually occurs due to the presence of mercury on the surface. However, he admits in the comments that the aluminum wire used for the experiment was pretreated with HgCl2. The latter reacts with the aluminum to give mercury metal, as shown in the equation below.
3 HgCl2 + 2 Al = 2 AlCl3 + 3 Hg
The metallic mercury produced in this reaction amalgamates the aluminum, thereby activating its surface.
If we cover aluminum with mercury and scratch off the protecting layer of alumina underneath the liquid mercury layer, the two metals blend together to form an aluminum amalgam, an Al-Hg alloy. Alumina does not stick to the amalgam as tightly as it does to pure aluminum. As a result, the exposed aluminum is air-oxidized quite efficiently, as shown in this nice video demonstration. (Unfortunately, it is incorrectly implied in the open captions to this video that the mercury is "converted" during the process. While the aluminum is indeed air-oxidized to Al2O3 in the experiment, the mercury metal is used only for the amalgamation and is not converted to any mercury compound.)
The ability of aluminum metal to react with water is similar. While the Al2O3 film protects the underlying aluminum atoms from oxidation by H2O, amalgamated aluminum reacts with water to give H2 and Al(OH)3 (Figure 3-90), as demonstrated in this video. The experimenter does not tell us in the description that the reaction actually occurs due to the presence of mercury on the surface. However, he admits in the comments that the aluminum wire used for the experiment was pretreated with HgCl2. The latter reacts with the aluminum to give mercury metal, as shown in the equation below.
3 HgCl2 + 2 Al = 2 AlCl3 + 3 Hg
The metallic mercury produced in this reaction amalgamates the aluminum, thereby activating its surface.
Figure 3-90. Reaction of aluminum with water.
The layer of Al2O3 on the surface of aluminum protects the metal from oxygen and water, but cannot defend it against acids and bases. Being an amphoteric oxide, alumina reacts with both acids and strong bases (Figure 3-91). As already mentioned above, the formula NaAlO2 is simplified for this introductory course; the correct formula is Na[Al(OH)4].
The layer of Al2O3 on the surface of aluminum protects the metal from oxygen and water, but cannot defend it against acids and bases. Being an amphoteric oxide, alumina reacts with both acids and strong bases (Figure 3-91). As already mentioned above, the formula NaAlO2 is simplified for this introductory course; the correct formula is Na[Al(OH)4].
Figure 3-91. Being amphoteric, aluminum oxide reacts with both acids and bases.
After the protecting film of alumina has been etched off (Figure 3-91), the exposed aluminum readily reacts with an acid or alkali (Figure 3-92). Watch aluminum cans dissolve in HCl and in NaOH here.
After the protecting film of alumina has been etched off (Figure 3-91), the exposed aluminum readily reacts with an acid or alkali (Figure 3-92). Watch aluminum cans dissolve in HCl and in NaOH here.
Figure 3-92. Reactions of aluminum with HCl and aqueous NaOH.
Naturally, aluminum hydroxide, Al(OH)3, is amphoteric too. Like Al2O3, Al(OH)3 is insoluble in water but readily dissolves in aqueous acids or alkali (Figure 3-93).
Naturally, aluminum hydroxide, Al(OH)3, is amphoteric too. Like Al2O3, Al(OH)3 is insoluble in water but readily dissolves in aqueous acids or alkali (Figure 3-93).
Figure 3-93. Amphoteric aluminum hydroxide reacts with both acids and bases.
The question arises as to how can we account for the fact that some hydroxides are bases, some are acids, and some amphoteric? A simplified explanation is provided in Figure 3-94. Virtually any compound featuring an O-H group bonded to an element (E) could, in principle, react both as a base and as an acid. Unless E = F (fluorine), which is the extreme exotic case, E is less electronegative than oxygen, and so is hydrogen. Therefore, both electron pairs, one shared by E and O and the other by O and H, are polarized (shifted) toward oxygen. Either bond may be cleaved. Which one?
The question arises as to how can we account for the fact that some hydroxides are bases, some are acids, and some amphoteric? A simplified explanation is provided in Figure 3-94. Virtually any compound featuring an O-H group bonded to an element (E) could, in principle, react both as a base and as an acid. Unless E = F (fluorine), which is the extreme exotic case, E is less electronegative than oxygen, and so is hydrogen. Therefore, both electron pairs, one shared by E and O and the other by O and H, are polarized (shifted) toward oxygen. Either bond may be cleaved. Which one?
Figure 3-94. A molecule bearing an OH group on an atom could be, in principle, a base or an acid.
A less electronegative element E (such as an alkali metal) has a lower affinity for electrons and is therefore prone to E-O bond heterolytic cleavage (Figure 3-94, left). Consequently, alkali metal hydroxides such as NaOH and KOH (E = Na, K) are bases producing OH- upon dissociation.
A more electronegative element E more successfully competes with the oxygen atom for the shared electron pair. This competition prompts the O atom to stronger pull toward itself the electron pair it shares with the H atom. This further polarizes the E-H bond away from the hydrogen, thereby making it more positively charged and facilitating the dissociation of the O-H bond to give rise to H+. Consequently, OH derivatives of nonmetals are acids.
Naturally, some elements E located between typical metals and nonmetals in the periodic table can and do exhibit both types of behavior, precisely what we call amphoterism. In the presence of a strong base, an amphoteric hydroxide acts as an acid. Conversely, an amphoteric hydroxide acts as a base in the presence of an acid. Aluminum hydroxide, Al(OH)3, is one such amphotheric compound.
Is the amphoterism of Al(OH)3 an indication that aluminum should be classified as a nonmetal rather than a metal? No. Aluminum (i) is an excellent donor of electrons (and, consequently, a good reducing agent), (ii) has a typical metallic silvery color and luster; (iii) is ductile; (iv) is a great conductor of electricity and heat. All these characteristics of aluminum indicate that it falls in the category of metals.
Amphoterism is a relative concept. Even a strong acid could be forced to act as a base by a sufficiently stronger acid. For example, fuming HNO3 reacts with concentrated H2SO4 to give nitronium hydrogen sulfate (Figure 3-95). We will encounter this reaction again in Volume 4.
A less electronegative element E (such as an alkali metal) has a lower affinity for electrons and is therefore prone to E-O bond heterolytic cleavage (Figure 3-94, left). Consequently, alkali metal hydroxides such as NaOH and KOH (E = Na, K) are bases producing OH- upon dissociation.
A more electronegative element E more successfully competes with the oxygen atom for the shared electron pair. This competition prompts the O atom to stronger pull toward itself the electron pair it shares with the H atom. This further polarizes the E-H bond away from the hydrogen, thereby making it more positively charged and facilitating the dissociation of the O-H bond to give rise to H+. Consequently, OH derivatives of nonmetals are acids.
Naturally, some elements E located between typical metals and nonmetals in the periodic table can and do exhibit both types of behavior, precisely what we call amphoterism. In the presence of a strong base, an amphoteric hydroxide acts as an acid. Conversely, an amphoteric hydroxide acts as a base in the presence of an acid. Aluminum hydroxide, Al(OH)3, is one such amphotheric compound.
Is the amphoterism of Al(OH)3 an indication that aluminum should be classified as a nonmetal rather than a metal? No. Aluminum (i) is an excellent donor of electrons (and, consequently, a good reducing agent), (ii) has a typical metallic silvery color and luster; (iii) is ductile; (iv) is a great conductor of electricity and heat. All these characteristics of aluminum indicate that it falls in the category of metals.
Amphoterism is a relative concept. Even a strong acid could be forced to act as a base by a sufficiently stronger acid. For example, fuming HNO3 reacts with concentrated H2SO4 to give nitronium hydrogen sulfate (Figure 3-95). We will encounter this reaction again in Volume 4.
Figure 3-95. In the reaction of HNO3 with H2SO4, the former acts as a base and the latter as an acid.
Both H2SO4 and HNO3 are strong acids. However, H2SO4 is considerably stronger than HNO3, forcing HNO3 to act as a base when the two are mixed. In fact, the reaction shown in Figure 3-95 is formally a neutralization reaction, in which H2SO4 serves as a source of H+ and HNO3 as a source of OH-. Molecules of HNO3 do not dissociate to NO2+ and OH- in this reaction. The NO2+ is formed in a different way (Figure 3-96) that involves proton transfer from H2SO4, the stronger acid, to HNO3, the weaker acid. The proton generated upon dissociation of H2SO4 seeks a lone electron pair to bond to within the system. It finds one on the O atom of the OH group of HNO3 and coordinates to it much like a proton coordinates to the lone electron pair on the N atom of ammonia to give NH4+. The resultant cation (protonated HNO3) then collapses into a molecule of water and the nitronium cation, NO2+. All of these processes are reversible.
Both H2SO4 and HNO3 are strong acids. However, H2SO4 is considerably stronger than HNO3, forcing HNO3 to act as a base when the two are mixed. In fact, the reaction shown in Figure 3-95 is formally a neutralization reaction, in which H2SO4 serves as a source of H+ and HNO3 as a source of OH-. Molecules of HNO3 do not dissociate to NO2+ and OH- in this reaction. The NO2+ is formed in a different way (Figure 3-96) that involves proton transfer from H2SO4, the stronger acid, to HNO3, the weaker acid. The proton generated upon dissociation of H2SO4 seeks a lone electron pair to bond to within the system. It finds one on the O atom of the OH group of HNO3 and coordinates to it much like a proton coordinates to the lone electron pair on the N atom of ammonia to give NH4+. The resultant cation (protonated HNO3) then collapses into a molecule of water and the nitronium cation, NO2+. All of these processes are reversible.
Figure 3-96. Mechanism of formation of NO2+ from HNO3 and H2SO4.
3.7.4. Exercises.
1. Aluminum metal is (a) nontoxic; (b) light (low density); (c) resistant to corrosion; (d) conductive to electricity and heat (e) ductile. Answer
2. What is the name of the mineral that is used to make aluminum? Answer
3. Describe the electrochemical process for making aluminum. Explain the needs for (A) using cryolite in the process and (B) frequently replacing the anode. [Answer: See subsection 3.7.2]
4. Aluminum is used to make motor vehicles, aircraft, electrical equipment, etc. because it is (a) not flammable; (b) flammable but is strongly protected from oxidation by a thin film of Al2O3 on the surface; (c) flammable, but the aluminum alloys, which are exclusively used in the manufacturing, are not. Answer
5. Aluminum reacts with both acids and bases to give H2. Would more H2 be produced if 1 g of pure aluminum was treated with excess of aqueous NaOH or excess aqueous HCl? Solve this problem mentally, without writing equations and doing calculations. Answer
6. Write balanced chemical equations for reactions of (a) Al with NaOH; (b) Al with H2SO4; (c) Al2O3 with KOH; (d) Al2O3 with HNO3; (e) Al(OH)3 with HCl; (f) Al(OH)3 with KOH. [Answer: See subsection 3.7.3]
7. Aluminum does not react with water. True or false? Answer
8. There are two roughly half-full test tubes, one containing an aqueous solution of AlCl3 and the other an aqueous solution of NaOH of about the same concentration. Both solutions are odorless, colorless, and clear. The test tubes are not marked. That is all you have and know about the samples. Find out which solution is in which test tube (Figure 3-97). Touching and tasting solutions is not allowed (and could be dangerous).
3.7.4. Exercises.
1. Aluminum metal is (a) nontoxic; (b) light (low density); (c) resistant to corrosion; (d) conductive to electricity and heat (e) ductile. Answer
2. What is the name of the mineral that is used to make aluminum? Answer
3. Describe the electrochemical process for making aluminum. Explain the needs for (A) using cryolite in the process and (B) frequently replacing the anode. [Answer: See subsection 3.7.2]
4. Aluminum is used to make motor vehicles, aircraft, electrical equipment, etc. because it is (a) not flammable; (b) flammable but is strongly protected from oxidation by a thin film of Al2O3 on the surface; (c) flammable, but the aluminum alloys, which are exclusively used in the manufacturing, are not. Answer
5. Aluminum reacts with both acids and bases to give H2. Would more H2 be produced if 1 g of pure aluminum was treated with excess of aqueous NaOH or excess aqueous HCl? Solve this problem mentally, without writing equations and doing calculations. Answer
6. Write balanced chemical equations for reactions of (a) Al with NaOH; (b) Al with H2SO4; (c) Al2O3 with KOH; (d) Al2O3 with HNO3; (e) Al(OH)3 with HCl; (f) Al(OH)3 with KOH. [Answer: See subsection 3.7.3]
7. Aluminum does not react with water. True or false? Answer
8. There are two roughly half-full test tubes, one containing an aqueous solution of AlCl3 and the other an aqueous solution of NaOH of about the same concentration. Both solutions are odorless, colorless, and clear. The test tubes are not marked. That is all you have and know about the samples. Find out which solution is in which test tube (Figure 3-97). Touching and tasting solutions is not allowed (and could be dangerous).
