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Unformatted text preview: Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website CHAPTER 20 Metallurgy and the Chemistry of Metals INTRODUCTION UP 20.1 OCCURRENCE OF METALS TO THIS POINT WE HAVE CONCENTRATED MAINLY ON FUNDAMENTAL PRINCIPLES: THEORIES OF CHEMICAL BONDING, INTERMOLECULAR FORCES, 20.2 METALLURGICAL PROCESSES RATES AND MECHANISMS OF CHEMICAL REACTIONS, EQUILIBRIUM, THE LAWS OF THERMODYNAMICS, AND ELECTROCHEMISTRY. AN 20.3 BAND THEORY OF CONDUCTIVITY UNDERSTAND- 20.4 PERIODIC TRENDS IN METALLIC PROPERTIES ING OF THESE TOPICS IS NECESSARY FOR THE STUDY OF THE PROPERTIES OF REPRESENTATIVE METALLIC ELEMENTS AND THEIR COMPOUNDS. 20.5 THE ALKALI METALS THE USE AND REFINEMENT OF METALS DATE BACK TO EARLY HU- MAN HISTORY. FOR 20.6 THE ALKALINE-EARTH METALS EXAMPLE, ARCHEOLOGISTS HAVE FOUND EVIDENCE THAT IN THE FIRST MILLENNIUM A.D. INHABITANTS OF SRI LANKA 20.7 ALUMINUM USED MONSOON WINDS TO RUN IRON-SMELTING FURNACES TO PRODUCE HIGHCARBON STEEL. THROUGH THE YEARS, THESE FURNACES COULD HAVE BEEN SOURCES OF STEEL FOR THE LEGENDARY DAMASCUS SWORDS, KNOWN FOR THEIR SHARPNESS AND DURABILITY. IN THIS CHAPTER WE WILL STUDY THE METHODS FOR EXTRACTING, REFINING, AND PURIFYING METALS AND EXAMINE THE PROPERTIES OF METALS THAT BELONG TO THE REPRESENTATIVE ELEMENTS. WE WILL EM- PHASIZE (1) THE OCCURRENCE AND PREPARATION OF METALS, (2) THE PHYSICAL AND CHEMICAL PROPERTIES OF SOME OF THEIR COMPOUNDS, AND (3) THEIR USES IN MODERN SOCIETY AND THEIR ROLES IN BIOLOGICAL SYSTEMS. 803 Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 804 METALLURGY AND THE CHEMISTRY OF METALS 20.1 OCCURRENCE OF METALS Most metals come from minerals. A mineral is a naturally occurring substance with a range of chemical composition. A mineral deposit concentrated enough to allow economical recovery of a desired metal is known as ore. Table 20.1 lists the principal types of minerals, and Figure 20.1 shows a classification of metals according to their minerals. The most abundant metals, which exist as minerals in Earth’s crust, are aluminum, iron, calcium, magnesium, sodium, potassium, titanium, and manganese (see p. 62). Seawater is a rich source of some metal ions, including Na , Mg2 and Ca2 . Furthermore, vast areas of the ocean floor are covered with manganese nodules, which are made up mostly of manganese, along with iron, nickel, copper, and cobalt in a chemically combined state (Figure 20.2). 20.2 METALLURGICAL PROCESSES Metallurgy is the science and technology of separating metals from their ores and of compounding alloys. An alloy is a solid solution either of two or more metals, or of a metal or metals with one or more nonmetals. The three principal steps in the recovery of a metal from its ore are (1) preparation of the ore, (2) production of the metal, and (3) purification of the metal. PREPARATION OF THE ORE In the preliminary treatment of an ore, the desired mineral is separated from waste materials—usually clay and silicate minerals — which are collectively called the gangue. One very useful method for carrying out such a separation is called flotation. In this process the ore is finely ground and added to water containing oil and detergent. The liquid mixture is then beaten or blown to form a froth. The oil preferentially wets the TABLE 20.1 Principal Types of Minerals TYPE MINERALS Uncombined metals Carbonates Ag, Au, Bi, Cu, Pd, Pt BaCO3 (witherite), CaCO3 (calcite, limestone), MgCO3 (magnesite), CaCO3 MgCO3 (dolomite), PbCO3 (cerussite), ZnCO3 (smithsonite) CaF2 (fluorite), NaCl (halite), KCl (sylvite), Na3AlF6 (cryolite) Al2O3 2H2O (bauxite), Al2O3 (corundum), Fe2O3 (hematite), Fe3O4 (magnetite), Cu2O (cuprite), MnO2 (pyrolusite), SnO2 (cassiterite), TiO2 (rutile), ZnO (zincite) Ca3(PO4)2 (phosphate rock), Ca5(PO4)3OH (hydroxyapatite) Be3Al2Si6O18 (beryl), ZrSiO4 (zircon), NaAlSi3O8 (albite), Mg3(Si4O10)(OH)2 (talc) Ag2S (argentite), CdS (greenockite), Cu2S (chalcocite), FeS2 (pyrite), HgS (cinnabar), PbS (galena), ZnS (sphalerite) BaSO4 (barite), CaSO4 (anhydrite), PbSO4 (anglesite), SrSO4 (celestite), MgSO4 7H2O (epsomite) Halides Oxides Phosphates Silicates Sulfides Sulfates Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 20.2 1 1A METALLURGICAL PROCESSES 805 18 8A Sulfides Uncombined 2 2A Chlorides Other compounds; see caption Li Be Oxides Na Mg 3 3B 4 4B 5 5B 6 6B 7 7B 8 9 8B 10 11 1B 12 2B Al K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb 13 3A 14 4A 15 5A 16 6A 17 7A Bi FIGURE 20.1 Metals and their best-known minerals. Lithium is found in spodumene (LiAlSi2O6), and beryllium in beryl (see Table 20.1). The rest of the alkaline earth metals are found in minerals that are carbonates and sulfates. The minerals for Sc, Y, and La are the phosphates. Some metals have more than one type of important mineral. For example, in addition to the sulfide, iron is found as the oxides hematite (Fe2O3) and magnetite (Fe3O4); and aluminum, in addition to the oxide, is found in beryl (Be3Al2Si6O18). Technetium (Tc) is a synthetic element. mineral particles, which are then carried to the top in the froth, while the gangue settles to the bottom. The froth is skimmed off, allowed to collapse, and dried to recover the mineral particles. Another physical separation process makes use of the magnetic properties of certain minerals. Ferromagnetic metals are strongly attracted to magnets. The mineral FIGURE 20.2 Manganese nodules on the ocean floor. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 806 METALLURGY AND THE CHEMISTRY OF METALS magnetite (Fe3O4), in particular, can be separated from the gangue by using a strong electromagnet. Cobalt is another ferromagnetic metal. Mercury forms amalgams with a number of metals. An amalgam is an alloy of mercury with another metal or metals. Mercury can therefore be used to extract metal from ore. Mercury dissolves the silver and gold in an ore to form a liquid amalgam, which is easily separated from the remaining ore. The gold or silver is recovered by distilling off mercury. PRODUCTION OF METALS Because metals in their combined forms always have positive oxidation numbers, the production of a free metal is always a reduction process. Preliminary operations may be necessary to convert the ore to a chemical state more suitable for reduction. For example, an ore may be roasted to drive off volatile impurities and at the same time to convert the carbonates and sulfides to the corresponding oxides, which can be reduced more conveniently to yield the pure metals: CaCO3(s) 88n CaO(s) 2PbS(s) 3O2(g) 88n 2PbO(s) CO2(g) 2SO2(g) This last equation points up the fact that the conversion of sulfides to oxides is a major source of sulfur dioxide, a notorious air pollutant (p. 711). How a pure metal is obtained by reduction from its combined form depends on the standard reduction potential of the metal (see Table 19.1). Table 20.2 outlines the reduction processes for several metals. Most major metallurgical processes now in use involve pyrometallurgy, procedures carried out at high temperatures. The reduction in these procedures may be accomplished either chemically or electrolytically. Chemical Reduction A more electropositive metal has a more negative standard reduction potential (see Table 19.1). We can use a more electropositive metal as a reducing agent to separate a less electropositive metal from its compound at high temperatures: V2O5(s) TiCl4(g) Cr2O3(s) 3Mn3O4(s) 5Ca(l ) 88n 2V(l ) 2Mg(l ) 88n Ti(s) 2Al(s) 88n 2Cr(l ) 8Al(s) 88n 9Mn(l ) 5CaO(s) 2MgCl2(l ) Al2O3(s) 4Al2O3(s) In some cases even molecular hydrogen can be used as a reducing agent, as in the preparation of tungsten (used as filaments in light bulbs) from tungsten(VI) oxide: WO3(s) TABLE 20.2 3H2(g) 88n W(s) 3H2O(g) Reduction Processes for Some Common Metals m7777777777777777777777777777 Decreasing activity of metals METAL Back REDUCTION PROCESS Lithium, sodium, magnesium, calcium Aluminum Chromium, manganese, titanium, vanadium, iron, zinc Mercury, silver, platinum, copper, gold Electrolytic reduction of the molten chloride Electrolytic reduction of anhydrous oxide (in molten cryolite) Reduction of the metal oxide with a more electropositive metal, or reduction with coke and carbon monoxide These metals occur in the free (uncombined) state or can be obtained by roasting their sulfides Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 20.2 METALLURGICAL PROCESSES 807 Electrolytic Reduction Electrolytic reduction is suitable for very electropositive metals, such as sodium, magnesium, and aluminum. The process is usually carried out on the anhydrous molten oxide or halide of the metal: 2MO(l ) 88n 2M (at cathode) O2 (at anode) 2MCl(l ) 88n 2M (at cathode) Cl2 (at anode) We will describe the specific procedures later in this chapter. THE METALLURGY OF IRON Iron exists in Earth’s crust in many different minerals, such as iron pyrite (FeS2), siderite (FeCO3), hematite (Fe2O3), and magnetite (Fe3O4, often represented as FeO Fe2O3). Of these, hematite and magnetite are particularly suitable for the extraction of iron. The metallurgical processing of iron involves the chemical reduction of the minerals by carbon (in the form of coke) in a blast furnace (Figure 20.3). The concentrated iron ore, limestone (CaCO3), and coke are introduced into the furnace from the top. A blast of hot air is forced up the furnace from the bottom—hence the name blast furnace. The oxygen gas reacts with the carbon in the coke to form mostly carbon monoxide and some carbon dioxide. These reactions are highly exothermic, and as the hot CO and CO2 gases rise, they react with the iron oxides in different temperature zones as shown in Figure 20.3. The key steps in the extraction of iron are FIGURE 20.3 A blast furnace. Iron ore, limestone, and coke are introduced at the top of the furnace. Iron is obtained from the ore by reduction with carbon. CO, CO2 Charge (ore, limestone, coke) 200°C Hot gases rise Solid charge descends 3Fe2O3 + CO CaCO3 Fe3O4 + CO 2Fe3O4 + CO2 CaO + CO2 3FeO + CO2 700°C C + CO2 FeO + CO 1200°C 1500°C 2CO Fe + CO2 Iron melts Molten slag forms 2C + O2 2CO 2000°C Hot air blast Slag Molten iron Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 808 METALLURGY AND THE CHEMISTRY OF METALS 3Fe2O3(s) Fe3O4(s) FeO(s) CaCO3 and other compounds used to form a molten mixture with the impurities in the ore for easy removal are called flux. CO(g) 88n 2Fe3O4(s) CO(g) 88n 3FeO(s) CO(g) 88n Fe(l ) CO2(g) CO2(g) CO2(g) The limestone decomposes in the furnace as follows: CaCO3(s) 88n CaO(s) CO2(g) The calcium oxide then reacts with the impurities in the iron, which are mostly sand (SiO2) and aluminum oxide (Al2O3): CaO(s) CaO(s) SiO2(s) 88n CaSiO3(l ) Al2O3(s) 88n Ca(AlO2)2(l ) The mixture of calcium silicate and calcium aluminate that remains molten at the furnace temperature is known as slag. By the time the ore works its way down to the bottom of the furnace, most of it has already been reduced to iron. The temperature of the lower part of the furnace is above the melting point of impure iron, and so the molten iron at the lower level can be run off to a receiver. The slag, because it is less dense, forms the top layer above the molten iron and can be run off at that level, as shown in Figure 20.3. Iron extracted in this way contains many impurities and is called pig iron; it may contain up to 5 percent carbon and some silicon, phosphorus, manganese, and sulfur. Some of the impurities stem from the silicate and phosphate minerals, while carbon and sulfur come from coke. Pig iron is granular and brittle. It has a relatively low melting point (about 1180°C), so it can be cast in various forms; for this reason it is also called cast iron. STEELMAKING Steel manufacturing is one of the most important metal industries. In the United States, the annual consumption of steel is well above 100 million tons. Steel is an iron alloy that contains from 0.03 to 1.4 percent carbon plus various amounts of other elements. The wide range of useful mechanical properties associated with steel is primarily a function of chemical composition and heat treatment of a particular type of steel. Whereas the production of iron is basically a reduction process (converting iron oxides to metallic iron), the conversion of iron to steel is essentially an oxidation process in which the unwanted impurities are removed from the iron by reaction with oxygen gas. One of several methods used in steelmaking is the basic oxygen process. Because of its ease of operation and the relatively short time (about 20 minutes) required for each large-scale (hundreds of tons) conversion, the basic oxygen process is by far the most common means of producing steel today. Figure 20.4 shows the basic oxygen process. Molten iron from the blast furnace is poured into an upright cylindrical vessel. Pressurized oxygen gas is introduced via a water-cooled tube above the molten metal. Under these conditions, manganese, phosphorus, and silicon, as well as excess carbon, react with oxygen to form oxides. These oxides are then reacted with the appropriate fluxes (for example, CaO or SiO2) to form slag. The type of flux chosen depends on the composition of the iron. If the main impurities are silicon and phosphorus, a basic flux such as CaO is added to the iron: SiO2(s) P4O10(l ) Back Forward Main Menu TOC CaO(s) 88n CaSiO3(l ) 6CaO(s) 88n 2Ca3(PO4)2(l ) Study Guide TOC Textbook Website MHHE Website 20.2 FIGURE 20.4 The basic oxygen process of steelmaking. The capacity of a typical vessel is 100 tons of cast iron. 809 METALLURGICAL PROCESSES O2 CO2, SO2 CaO or SiO2 Molten steel + slag Slag Vertical position Horizontal position Molten steel On the other hand, if manganese is the main impurity, then an acidic flux such as SiO2 is needed to form the slag: MnO(s) SiO2(s) 88n MnSiO3(l ) The molten steel is sampled at intervals. When the desired blend of carbon and other impurities has been reached, the vessel is rotated to a horizontal position so that the molten steel can be tapped off (Figure 20.5). FIGURE 20.5 Back Forward Steelmaking. Main Menu TOC Study Guide TOC Textbook Website MHHE Website 810 METALLURGY AND THE CHEMISTRY OF METALS The properties of steel depend not only on its chemical composition but also on the heat treatment. At high temperatures, iron and carbon in steel combine to form iron carbide, Fe3C, called cementite: 3Fe(s) C(s) 34 Fe3C(s) The forward reaction is endothermic, so that the formation of cementite is favored at high temperatures. When steel containing cementite is cooled slowly, the above equilibrium shifts to the left, and the carbon separates as small particles of graphite, which give the steel a gray color. (Very slow decomposition of cementite also takes place at room temperature.) If the steel is cooled rapidly, equilibrium is not attained and the carbon remains largely in the form of cementite, Fe3C. Steel containing cementite is light in color, and it is harder and more brittle than that containing graphite. Heating the steel to some appropriate temperature for a short time and then cooling it rapidly in order to give it the desired mechanical properties is known as “tempering.” In this way, the ratio of carbon present as graphite and as cementite can be varied within rather wide limits. Table 20.3 shows the composition, properties, and uses of various types of steel. PURIFICATION OF METALS Metals prepared by reduction usually need further treatment to remove impurities. The extent of purification, of course, depends on how the metal will be used. Three common purification procedures are distillation, electrolysis, and zone refining. Distillation Metals that have low boiling points, such as mercury, magnesium, and zinc, can be separated from other metals by fractional distillation. One well-known method of fractional distillation is the Mond† process for the purification of nickel. Carbon monoxide gas is passed over the impure nickel metal at about 70°C to form the volatile tetracarbonylnickel (b.p. 43°C), a highly toxic substance, which is separated from the less volatile impurities by distillation: Ni(s) 4CO(g) 88n Ni(CO)4(g) † Ludwig Mond (1839–1909). British chemist of German origin. Mond made many important contributions to industrial chemistry. His method for purifying nickel by converting it to the volatile Ni(CO)4 compound has been described as having given “wings” to the metal. TABLE 20.3 Types of Steel COMPOSITION (PERCENT BY MASS)* TYPE C Mn P Plain High-strength 1.35 0.25 1.65 1.65 0.04 0.04 Stainless S Si Ni Cr OTHERS USES 0.05 0.06 — — Cu (0.2 – 0.6) Sheet products, tools 0.05 0.15 – 0.9 0.4 – 1.0 0.3 – 1.3 Cu (0.01 – 0.08) Construction, steam turbines 0.03 – 1.2 1.0 – 10 0.04 – 0.06 0.03 1 – 3 1 – 22 4.0 – 27 — Kitchen utensils, razor blades *A single number indicates the maximum amount of the substance present. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 20.2 811 Pure metallic nickel is recovered from Ni(CO)4 by heating the gas at 200°C: Battery Impure copper anode METALLURGICAL PROCESSES Pure copper cathode Ni(CO)4(g) 88n Ni(s) 4CO(g) The carbon monoxide that is released is recycled back into the process. Electrolysis Electrolysis is another important purification technique. The copper metal obtained by roasting copper sulfide usually contains impurities such as zinc, iron, silver, and gold. The more electropositive metals are removed by an electrolysis process in which the impure copper acts as the anode and pure copper acts as the cathode in a sulfuric acid solution containing Cu2 ions (Figure 20.6). The reactions are Cu2+ 2 SO4 – FIGURE 20.6 Electrolytic purification of copper. Cu(s) 88n Cu2 (aq) Anode (oxidation): 2 Cathode (reduction): Cu (aq) The metal impurities separated from the copper anode are valuable by-products, the sale of which often pays for the electricity needed to drive the electrolysis. 2e 2e 88n Cu(s) Reactive metals in the copper anode, such as iron and zinc, are also oxidized at the anode and enter the solution as Fe2 and Zn2 ions. They are not reduced at the cathode, however. The less electropositive metals, such as gold and silver, are not oxidized at the anode. Eventually, as the copper anode dissolves, these metals fall to the bottom of the cell. Thus, the net result of this electrolysis process is the transfer of copper from the anode to the cathode. Copper prepared this way has a purity greater than 99.5 percent (Figure 20.7). Zone Refining Another often-used method of obtaining extremely pure metals is zone refining. In this process a metal rod containing a few impurities is drawn through an electrical heating FIGURE 20.7 Copper cathodes used in the electrorefining process. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 812 METALLURGY AND THE CHEMISTRY OF METALS FIGURE 20.8 Zone-refining technique for purifying metals. Top to bottom: An impure metal rod is moved slowly through a heating coil. As the metal rod moves forward, the impurities dissolve in the molten portion of the metal while pure metal crystallizes out in front of the molten zone. Eventually the end portion of the rod, which contains most of the impurities, is allowed to cool and is cut off from the rest of the rod. Heating coil Metal rod coil that melts the metal (Figure 20.8). Most impurities dissolve in the molten metal. As the metal rod emerges from the heating coil, it cools and the pure metal crystallizes, leaving the impurities in the molten metal portion that is still in the heating coil. (This is analogous to the freezing of seawater, in which the solid that separates is mostly pure solvent—water. In zone refining, the liquid metal acts as the solvent and the impurities as the solutes.) When the molten zone carrying the impurities, now at increased concentration, reaches the end of the rod, it is allowed to cool and is then cut off. Repeating this procedure a number of times results in metal with a purity greater than 99.99 percent. 20.3 BAND THEORY OF CONDUCTIVITY In Section 11.6 we saw that the ability of metals to conduct heat and electricity can be explained with molecular orbital theory. To gain a better understanding of the conductivity properties of metals we must also apply our knowledge of quantum mechanics. The model we will use to study metallic bonding is band theory, so called because it states that delocalized electrons move freely through “bands” formed by overlapping molecular orbitals. We will also apply band theory to certain elements that are semiconductors. CONDUCTORS Metals are characterized by high electrical conductivity. Consider magnesium, for example. The electron configuration of Mg is [Ne]3s2, so each atom has two valence electrons in the 3s orbital. In a metallic crystal, the atoms are packed closely together, so the energy levels of each magnesium atom are affected by the immediate neighbors of the atom as a result of orbital overlaps. In Chapter 10 we saw that, in terms of molecular orbital theory, the interaction between two atomic orbitals leads to the formation of a bonding and an antibonding molecular orbital. Since the number of atoms in even a small piece of magnesium is enormously large (on the order of 1020 atoms), the Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 20.3 3p 3s Energy FIGURE 20.9 Formation of conduction bands in magnesium. The electrons in the 1s, 2s, and 2p orbitals are localized on each Mg atom. However, the 3s and 3p orbitals overlap to form delocalized molecular orbitals. Electrons in these orbitals can travel throughout the metal, and this accounts for the electrical conductivity of the metal. BAND THEORY OF CONDUCTIVITY 813 Conduction band Valence band 2p 2s 1s 12 + 12 + 12 + 12 + 12 + Mg Mg Mg Mg Mg number of molecular orbitals they form is also very large. These molecular orbitals are so closely spaced on the energy scale that they are more appropriately described as a “band” (Figure 20.9). The closely spaced filled energy levels make up the valence band. The upper half of the energy levels corresponds to the empty, delocalized molecular orbitals formed by the overlap of the 3p orbitals. This set of closely spaced empty levels is called the conduction band. We can imagine a metallic crystal as an array of positive ions immersed in a sea of delocalized valence electrons (see Figure 11.30). The great cohesive force resulting from the delocalization is partly responsible for the strength noted in most metals. Because the valence band and the conduction band are adjacent to each other, the amount of energy needed to promote a valence electron to the conduction band is negligible. There the electron can travel freely through the metal, since the conduction band is void of electrons. This freedom of movement accounts for the fact that metals are good conductors, that is, they are capable of conducting electric current. Why don’t substances like wood and glass conduct electricity as metals do? Figure 20.10 provides an answer to this question. Basically, the electrical conductivity of a solid depends on the spacing and the state of occupancy of the energy bands. In magnesium and other metals the valence bands are adjacent to the conduction bands, and, therefore, these metals readily act as conductors. In wood and glass, on the other hand, the gap between the valence band and the conduction band is considerably greater than that in a metal. Consequently, much more energy is needed to excite an electron into the conduction band. Lacking this energy, electrons cannot move freely. Therefore, glass and wood are insulators, ineffective conductors of electricity. Back Forward Main Menu Conduction band Energy Valence band Energy gap Energy Conduction band Conduction band Energy FIGURE 20.10 Comparison of the energy gaps between valence band and conduction band in a metal, a semiconductor, and an insulator. In a metal the energy gap is virtually nonexistent; in a semiconductor the energy gap is small; and in an insulator the energy gap is very large, thus making the promotion of an electron from the valence band to the conduction band difficult. Energy gap Valence band Valence band Metal TOC Study Guide TOC Semiconductor Textbook Website Insulator MHHE Website 814 METALLURGY AND THE CHEMISTRY OF METALS SEMICONDUCTORS A number of elements are semiconductors, that is, they normally are not conductors, but will conduct electricity at elevated temperatures or when combined with a small amount of certain other elements. The Group 4A elements silicon and germanium are especially suited for this purpose. The use of semiconductors in transistors and solar cells, to name two applications, has revolutionized the electronic industry in recent decades, leading to increased miniaturization of electronic equipment. The energy gap between the filled and empty bands of these solids is much smaller than that for insulators (see Figure 20.10). If the energy needed to excite electrons from the valence band into the conduction band is provided, the solid becomes a conductor. Note that this behavior is opposite that of the metals. A metal’s ability to conduct electricity decreases with increasing temperature, because the enhanced vibration of atoms at higher temperatures tends to disrupt the flow of electrons. The ability of a semiconductor to conduct electricity can also be enhanced by adding small amounts of certain impurities to the element, a process called doping. Let us consider what happens when a trace amount of boron or phosphorus is added to solid silicon. (Only about five out of every million Si atoms are replaced by B or P atoms.) The structure of solid silicon is similar to that of diamond; each Si atom is covalently bonded to four other Si atoms. Phosphorus ([Ne]3s23p3) has one more valence electron than silicon ([Ne]3s23p2), so there is a valence electron left over after four of them are used to form covalent bonds with silicon (Figure 20.11). This extra electron can be removed from the phosphorus atom by applying a voltage across the solid. The free electron can move through the structure and function as a conduction electron. Impurities of this type are known as donor impurities, since they provide conduction electrons. Solids containing donor impurities are called n-type semiconductors, where n stands for negative (the charge of the “extra” electron). The opposite effect occurs if boron is added to silicon. A boron atom has three valence electrons (1s22s22p1). Thus, for every boron atom in the silicon crystal there is a single vacancy in a bonding orbital. It is possible to excite a valence electron from a nearby Si into this vacant orbital. A vacancy created at that Si atom can then be filled by an electron from a neighboring Si atom, and so on. In this manner, electrons can move through the crystal in one direction while the vacancies, or “positive holes,” move in the opposite direction, and the solid becomes an electrical conductor. Impurities that are electron deficient are called acceptor impurities. Semiconductors that contain acceptor impurities are called p-type semiconductors, where p stands for positive. In both the p-type and n-type semiconductors the energy gap between the valence band and the conduction band is effectively reduced, so that only a small amount of FIGURE 20.11 (a) Silicon crystal doped with phosphorus. (b) Silicon crystal doped with boron. Note the formation of a negative center in (a) and of a positive center in (b). e– (a) Back Forward Main Menu TOC Study Guide TOC + (b) Textbook Website MHHE Website 20.4 815 PERIODIC TRENDS IN METALLIC PROPERTIES energy is needed to excite the electrons. Typically, the conductivity of a semiconductor is increased by a factor of 100,000 or so by the presence of impurity atoms. The growth of the semiconductor industry since the early 1960s has been truly remarkable. Today semiconductors are essential components of nearly all electronic equipment, ranging from radios and television sets to pocket calculators and computers. One of the main advantages of solid-state devices over vacuum-tube electronics is that the former can be made on a single “chip” of silicon no larger than the cross section of a pencil eraser. Consequently, much more equipment can be packed into a small volume—a point of particular importance in space travel, as well as in hand-held calculators and microprocessors (computers-on-a-chip). 20.4 PERIODIC TRENDS IN METALLIC PROPERTIES Metals are lustrous in appearance, solid at room temperature (with the exception of mercury), good conductors of heat and electricity, malleable (can be hammered flat), and ductile (can be drawn into wire). Figure 20.12 shows the positions of the representative metals and the group 2B metals in the periodic table. (The transition metals are discussed in Chapter 22.) As we saw in Figure 9.5, the electronegativity of elements increases from left to right across a period and from bottom to top in a group. The metallic character of metals increases in just the opposite directions, that is, from right to left across a period and from top to bottom in a group. Because metals generally have low electronegativities, they tend to form cations and almost always have positive oxidation numbers in their compounds. However, beryllium and magnesium in Group 2A and metals in Group 3A and beyond also form covalent compounds. In the next three sections we will study the chemistry of selected metals from Group 1A (the alkali metals), Group 2A (the alkaline earth metals), and Group 3A (aluminum). FIGURE 20.12 Representative metals (green) and Group 2B metals (blue) according to their positions in the periodic table. 18 8A 1 1A H 13 3A 14 4A 15 5A 16 6A 17 7A He Li Be B C N O F Ne Na Mg 3 3B 4 4B 5 5B 6 6B 7 7B 8 9 8B 10 11 1B 12 2B Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Back 2 2A Ra Ac Rf Ha Sg Ns Hs Mt Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 816 METALLURGY AND THE CHEMISTRY OF METALS 20.5 le lab ai Av t No “ in FIGURE 20.13 n” io rs Ve t ex T e- Halite (NaCl). Remember that Ca2 is harder to reduce than Na . THE ALKALI METALS As a group, the alkali metals (the Group 1A elements) are the most electropositive (or the least electronegative) elements known. They exhibit many similar properties, some of which are listed in Table 20.4. From their electron configurations we expect the oxidation number of these elements in their compounds to be 1 since the cations would be isoelectronic with the noble gases. This is indeed the case. The alkali metals have low melting points and are soft enough to be sliced with a knife (see Figure 8.14). These metals all possess a body-centered crystal structure (see Figure 11.29) with low packing efficiency. This accounts for their low densities among metals. In fact, lithium is the lightest metal known. Because of their great chemical reactivity, the alkali metals never occur naturally in elemental form; they are found combined with halide, sulfate, carbonate, and silicate ions. In this section we will describe the chemistry of two members of Group 1A—sodium and potassium. The chemistry of lithium, rubidium, and cesium is less important; all isotopes of francium, the last member of the group, are radioactive. Sodium and potassium are about equally abundant in nature. They occur in silicate minerals such as albite (NaAlSi3O8) and orthoclase (KAlSi3O8). Over long periods of time (on a geologic scale), silicate minerals are slowly decomposed by wind and rain, and their sodium and potassium ions are converted to more soluble compounds. Eventually rain leaches these compounds out of the soil and carries them to the sea. Yet when we look at the composition of seawater, we find that the concentration ratio of sodium to potassium is about 28 to 1. The reason for this uneven distribution is that potassium is essential to plant growth, while sodium is not. Plants take up many of the potassium ions along the way, while sodium ions are free to move on to the sea. Other minerals that contain sodium or potassium are halite (NaCl), shown in Figure 20.13, Chile saltpeter (NaNO3), and sylvite (KCl). Sodium chloride is also obtained from rock salt (see p. 337). Metallic sodium is most conveniently obtained from molten sodium chloride by electrolysis in the Downs cell (see Section 19.8). The melting point of sodium chloride is rather high (801°C), and much energy is needed to keep large amounts of the substance molten. Adding a suitable substance, such as CaCl2, lowers the melting point to about 600°C—a more convenient temperature for the electrolysis process. TABLE 20.4 Properties of Alkali Metals Li Valence electron configuration Density (g/cm3) Melting point (°C) Boiling point (°C) Atomic radius (pm) Ionic radius (pm)* Ionization energy (kJ/mol) Electronegativity Standard reduction potential (V)† Na 1 2s 0.534 179 1317 155 60 520 1.0 3.05 1 3s 0.97 97.6 892 190 95 496 0.9 2.93 K Rb 1 4s 0.86 63 770 235 133 419 0.8 2.71 1 5s 1.53 39 688 248 148 403 0.8 2.93 Cs 6s1 1.87 28 678 267 169 375 0.7 2.92 *Refers to the cation M , where M denotes an alkali metal atom. † The half-reaction is M (aq) e 88n M(s). Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 20.5 THE ALKALI METALS 817 Metallic potassium cannot be easily prepared by the electrolysis of molten KCl because it is too soluble in the molten KCl to float to the top of the cell for collection. Moreover, it vaporizes readily at the operating temperatures, creating hazardous conditions. Instead, it is usually obtained by the distillation of molten KCl in the presence of sodium vapor at 892°C. The reaction that takes place at this temperature is Note that this is a chemical rather than electrolytic reduction. Na(g) KCl(l ) 34 NaCl(l ) K(g) This reaction may seem strange given that potassium is a stronger reducing agent than sodium (see Table 20.4). However, potassium has a lower boiling point (770°C) than sodium (892°C), so it is more volatile at 892°C and distills off more easily. According to Le Chatelier ’s principle, constant removal of potassium vapor shifts the above equilibrium from left to right, assuring recovery of metallic potassium. Sodium and potassium are both extremely reactive, but potassium is the more reactive of the two. Both react with water to form the corresponding hydroxides. In a limited supply of oxygen, sodium burns to form sodium oxide (Na2O). However, in the presence of excess oxygen, sodium forms the pale-yellow peroxide: 2Na(s) O2(g) 88n Na2O2(s) Sodium peroxide reacts with water to give an alkaline solution and hydrogen peroxide: Na2O2(s) 2H2O(l ) 88n 2NaOH(aq) H2O2(aq) Like sodium, potassium forms the peroxide. In addition, potassium also forms the superoxide when it burns in air: K(s) O2(g) 88n KO2(s) When potassium superoxide reacts with water, oxygen gas is evolved: 2KO2(s) 2H2O(l ) 88n 2KOH(aq) O2(g) H2O2(aq) This reaction is utilized in breathing equipment (Figure 20.14). Exhaled air contains both moisture and carbon dioxide. The moisture reacts with KO2 in the apparatus to FIGURE 20.14 Self-contained breathing apparatus. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 818 METALLURGY AND THE CHEMISTRY OF METALS FIGURE 20.15 Crystals of a salt composed of sodium anions and a complex of sodium cations. generate oxygen gas as shown above. Furthermore, KO2 also reacts with exhaled CO2, which produces more oxygen gas: 4KO2(s) 2CO2(g) 88n 2K2CO3(s) 3O2(g) Thus, a person using the apparatus can continue to breathe oxygen without being exposed to toxic fumes outside. Sodium and potassium metals dissolve in liquid ammonia to produce a beautiful blue solution: NH3 Na 88n Na NH3 K 88n K e e Both the cation and the electron exist in the solvated form; the solvated electrons are responsible for the characteristic blue color of such solutions. Metal-ammonia solutions are powerful reducing agents (because they contain free electrons); they are useful in synthesizing both organic and inorganic compounds. It was discovered recently that the hitherto unknown alkali metal anions, M , are also formed in such solutions. This means that an ammonia solution of an alkali metal contains ion pairs such as Na Na and K K ! (Keep in mind that in each case the metal cation exists as a complex ion with crown ether, an organic compound with a high affinity for cations.) In fact, these “salts” are so stable that they can be isolated in crystalline form (Figure 20.15). This finding is of considerable theoretical interest, for it shows clearly that the alkali metals can have an oxidation number of 1, although 1 is not found in ordinary compounds. Sodium and potassium are essential elements of living matter. Sodium ions and potassium ions are present in intracellular and extracellular fluids, and they are essential for osmotic balance and enzyme functions. We now describe the preparations and uses of several of the important compounds of sodium and potassium. Sodium Chloride The source, properties, and uses of sodium chloride were discussed in Chapter 9 (see p. 337). Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 20.5 THE ALKALI METALS 819 Sodium Carbonate Sodium carbonate (called soda ash) is used in all kinds of industrial processes, including water treatment and the manufacture of soaps, detergents, medicines, and food additives. Today about half of all Na2CO3 produced is used in the glass industry (in soda-lime glass; see Section 11.7). Sodium carbonate ranks eleventh among the chemicals produced in the United States (11 million tons in 1995). For many years Na2CO3 was produced by the Solvay† process, in which ammonia is first dissolved in a saturated solution of sodium chloride. Bubbling carbon dioxide into the solution results in the precipitation of sodium bicarbonate as follows: NH3(aq) NaCl(aq) H2CO3(aq) 88n NaHCO3(s) NH4Cl(aq) Sodium bicarbonate is then separated from the solution and heated to give sodium carbonate: 2NaHCO3(s) 88n Na2CO3(s) The last plant using the Solvay process in the United States closed in 1986. CO2(g) H2O(g) However, the rising cost of ammonia and the pollution problem resulting from by-products have prompted chemists to look for other sources of sodium carbonate. One is the mineral trona [Na5(CO3)2(HCO3) 2H2O], large deposits of which have been found in Wyoming. When trona is crushed and heated, it decomposes as follows: 2Na5(CO3)2(HCO3) 2H2O(s) 88n 5Na2CO3(s) CO2(g) 3H2O(g) The sodium carbonate obtained this way is dissolved in water, the solution is filtered to remove the insoluble impurities, and the sodium carbonate is crystallized as Na2CO3 10H2O. Finally, the hydrate is heated to give pure, anhydrous sodium carbonate. Sodium Hydroxide and Potassium Hydroxide The properties of sodium hydroxide and potassium hydroxide are very similar. These hydroxides are prepared by the electrolysis of aqueous NaCl and KCl solutions (see Section 19.8); both hydroxides are strong bases and very soluble in water. Sodium hydroxide is used in the manufacture of soap and many organic and inorganic compounds. Potassium hydroxide is used as an electrolyte in some storage batteries, and aqueous potassium hydroxide is used to remove carbon dioxide and sulfur dioxide from air. Sodium Nitrate and Potassium Nitrate Large deposits of sodium nitrate (chile saltpeter) are found in Chile. It decomposes with the evolution of oxygen at about 500°C: 2NaNO3(s) 88n 2NaNO2(s) O2(g) Potassium nitrate (saltpeter) is prepared beginning with the “reaction” KCl(aq) NaNO3(aq) 88n KNO3(aq) NaCl(aq) This process is carried out just below 100°C. Because KNO3 is the least soluble salt at room temperature, it is separated from the solution by fractional crystallization. Like NaNO3, KNO3 decomposes when heated. † Ernest Solvay (1838–1922). Belgian chemist. Solvay’s main contribution to industrial chemistry was the development of the process for the production of sodium carbonate that now bears his name. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 820 METALLURGY AND THE CHEMISTRY OF METALS TABLE 20.5 Properties of Alkaline Earth Metals Be Mg Ca Sr Ba 2s2 1.86 1280 2770 112 31 899 1757 1.5 1.85 Valence electron configuration Density (g/cm3) Melting point (°C) Boiling point (°C) Atomic radius (pm) Ionic radius (pm)* First and second ionization energies (kJ/mol) Electronegativity Standard reduction potential (V)† 3s2 1.74 650 1107 160 65 738 1450 1.2 2.37 4s2 1.55 838 1484 197 99 590 1145 1.0 2.87 5s2 2.6 770 1380 215 113 548 1058 1.0 2.89 6s2 3.5 714 1640 222 135 502 958 0.9 2.90 *Refers to the cation M2 , where M denotes an alkaline earth metal atom. † The half-reaction is M2 (aq) 2e 88n M(s). Gunpowder consists of potassium nitrate, wood charcoal, and sulfur in the approximate proportions of 6:1:1 by mass. When gunpowder is heated, the reaction is 2KNO3(s) 3C(s) 88n K2S(s) S(s) N2(g) 3CO2(g) The sudden formation of hot expanding gases causes an explosion. 20.6 THE ALKALINE EARTH METALS The alkaline earth metals are somewhat less electropositive and less reactive than the alkali metals. Except for the first member of the family, beryllium, which resembles aluminum (a Group 3A metal) in some respects, the alkaline earth metals have similar chemical properties. Because their M2 ions attain the stable electron configuration of the preceding noble gas, the oxidation number of alkaline earth metals in the combined form is almost always 2. Table 20.5 lists some common properties of these metals. Radium is not included in the table because all radium isotopes are radioactive and it is difficult and expensive to study the chemistry of this Group 2A element. MAGNESIUM n ei bl a ail Av ot “N ” on rsi e tV ex T e- FIGURE 20.16 Dolomite (CaCO3 MgCO3). Back Forward Magnesium (see Figure 8.15) is the sixth most plentiful element in Earth’s crust (about 2.5 percent by mass). Among the principal magnesium ores are brucite, Mg(OH)2; dolomite, CaCO3 MgCO3 (Figure 20.16); and epsomite, MgSO4 7H2O. Seawater is a good source of magnesium; there are about 1.3 g of magnesium in each kilogram of seawater. As is the case with most alkali and alkaline earth metals, metallic magnesium is obtained by electrolysis, in this case from its molten chloride, MgCl2 (obtained from seawater, see p. 143). The chemistry of magnesium is intermediate between that of beryllium and the heavier Group 2A elements. Magnesium does not react with cold water but does react slowly with steam: Mg(s) H2O(g) 88n MgO(s) H2(g) It burns brilliantly in air to produce magnesium oxide and magnesium nitride (see Figure 4.11): Main Menu TOC Study Guide TOC Textbook Website MHHE Website 20.6 THE ALKALINE EARTH METALS 821 FIGURE 20.17 A commercially available milk of magnesia. 2Mg(s) O2(g) 88n 2MgO(s) 3Mg(s) N2(g) 88n Mg3N2(s) This property makes magnesium (in the form of thin ribbons or fibers) useful in flash photography and flares. Magnesium oxide reacts very slowly with water to form magnesium hydroxide, a white solid suspension called milk of magnesia (Figure 20.17), which is used to treat acid indigestion: MgO(s) H2O(l ) 88n Mg(OH)2(s) “N ot Av ail ab le in e-T ex t Magnesium is a typical alkaline earth metal in that its hydroxide is a strong base. [The only alkaline earth hydroxide that is not a strong base is Be(OH)2, which is amphoteric.] The major uses of magnesium are in lightweight structural alloys, for cathodic protection (see Section 19.7), in organic synthesis, and in batteries. Magnesium is essential to plant and animal life, and Mg2 ions are not toxic. It is estimated that the average adult ingests about 0.3 g of magnesium ions daily. Magnesium plays several important biological roles. It is present in intracellular and extracellular fluids. Magnesium ions are essential for the proper functioning of a number of enzymes. Magnesium is also present in the green plant pigment chlorophyll, which plays an important part in photosynthesis. CALCIUM Earth’s crust contains about 3.4 percent calcium (see Figure 8.15) by mass. Calcium occurs in limestone, calcite, chalk, and marble as CaCO3; in dolomite as CaCO3 MgCO3 (see Figure 20.16); in gypsum as CaSO4 2H2O; and in fluorite as CaF2 (Figure 20.18). Metallic calcium is best prepared by the electrolysis of molten calcium chloride (CaCl2). As we read down Group 2A from beryllium to barium, we note an increase in metallic properties. Unlike beryllium and magnesium, calcium (like strontium and barium) reacts with cold water to yield the corresponding hydroxide, although the rate of reaction is much slower than those involving the alkali metals (see Figure 4.13): Ca(s) FIGURE 20.18 Back Forward Fluorite (CaF2). Main Menu 2H2O(l ) 88n Ca(OH)2(aq) H2(g) Calcium hydroxide [Ca(OH)2] is commonly known as slaked lime or hydrated lime. TOC Study Guide TOC Textbook Website MHHE Website 822 METALLURGY AND THE CHEMISTRY OF METALS Lime (CaO), which is also referred to as quicklime, is one of the oldest materials known to mankind. Quicklime is produced by the thermal decomposition of calcium carbonate (see Section 18.3): CaCO3(s) 88n CaO(s) CO2(g) while slaked lime is produced by the reaction between quicklime and water: CaO(s) H2O(l ) 88n Ca(OH)2(s) Quicklime is used in metallurgy (see Section 20.2) and the removal of SO2 when fossil fuel is burned (see p. 712). Slaked lime is used in water treatment. For many years, farmers have used lime to lower the acidity of soil for their crops (a process called liming). Nowadays lime is also applied to lakes affected by acid rain (see Section 17.6). Metallic calcium has rather limited uses. It serves mainly as an alloying agent for metals like aluminum and copper and in the preparation of beryllium metal from its compounds. It is also used as a dehydrating agent for organic solvents. Calcium is an essential element in living matter. It is the major component of bones and teeth; the calcium ion is present in a complex phosphate salt, hydroxyapatite, Ca5(PO4)3OH. A characteristic function of Ca2 ions in living systems is the activation of a variety of metabolic processes. Calcium plays a vital role in heart action, blood clotting, muscle contraction, and nerve impulse transmission. 20.7 ALUMINUM Aluminum (see Figure 8.16) is the most abundant metal and the third most plentiful element in Earth’s crust (7.5 percent by mass). The elemental form does not occur in nature; its principal ore is bauxite (Al2O3 2H2O). Other minerals containing aluminum are orthoclase (KAlSi3O8), beryl (Be3Al2Si6O18), cryolite (Na3AlF6), and corundum (Al2O3) (Figure 20.19). Aluminum is usually prepared from bauxite, which is frequently contaminated with silica (SiO2), iron oxides, and titanium(IV) oxide. The ore is first heated in sodium hydroxide solution to convert the silica into soluble silicates: SiO2(s) 2OH (aq) 88n SiO2 (aq) 3 H2O(l ) At the same time, aluminum oxide is converted to the aluminate ion (AlO2 ): Al2O3(s) FIGURE 20.19 (Al2O3). Corundum e in abl l vai tA “No Back Forward Main Menu TOC 2OH (aq) 88n 2AlO2 (aq) H2O(l ) n” rsio Ve ext e-T Study Guide TOC Textbook Website MHHE Website 20.7 ALUMINUM 823 Iron oxide and titanium oxide are unaffected by this treatment and are filtered off. Next, the solution is treated with acid to precipitate the insoluble aluminum hydroxide: AlO2 (aq) H3O (aq) 88n Al(OH)3(s) After filtration, the aluminum hydroxide is heated to obtain aluminum oxide: 2Al(OH)3(s) 88n Al2O3(s) Molten cryolite provides a good conducting medium for electrolysis. Anhydrous aluminum oxide, or corundum, is reduced to aluminum by the Hall† process. Figure 20.20 shows a Hall electrolytic cell, which contains a series of carbon anodes. The cathode is also made of carbon and constitutes the lining inside the cell. The key to the Hall process is the use of cryolite, or Na3AlF6 (m.p. 1000°C), as the solvent for aluminum oxide (m.p. 2045°C). The mixture is electrolyzed to produce aluminum and oxygen gas: 3[2O2 88n O2(g) Anode (oxidation): Cathode (reduction): 4[Al3 Carbon cathode Molten aluminum Al2O3 in molten cryolite FIGURE 20.20 Electrolytic production of aluminum based on the Hall process. 4e ] 3e 88n Al(l)] 2Al2O3 88n 4Al(l ) Overall: Carbon anodes 3H2O(g) 3O2(g) Oxygen gas reacts with the carbon anodes (at elevated temperatures) to form carbon monoxide, which escapes as a gas. The liquid aluminum metal (m.p. 660.2°C) sinks to the bottom of the vessel, from which it can be drained from time to time during the procedure. Aluminum is one of the most versatile metals known. It has a low density (2.7 g/cm3) and high tensile strength (that is, it can be stretched or drawn out). Aluminum is malleable, it can be rolled into thin foils, and it is an excellent electrical conductor. Its conductivity is about 65 percent that of copper. However, because aluminum is cheaper and lighter than copper, it is widely used in high-voltage transmission lines. Although aluminum’s chief use is an aircraft construction, the pure metal itself is too soft and weak to withstand much strain. Its mechanical properties are greatly improved by alloying it with small amounts of metals such as copper, magnesium, and manganese, as well as silicon. Aluminum is not used by living systems and is generally considered to be nontoxic. As we read across the periodic table from left to right in a given period, we note a gradual decrease in metallic properties. Thus, although aluminum is considered an active metal, it does not react with water as do sodium and calcium. Aluminum reacts with hydrochloric acid and with strong bases as follows: 2Al(s) 2Al(s) 2NaOH(aq) 6HCl(aq) 88n 2AlCl3(aq) 3H2(g) 2H2O(l ) 88n 2NaAlO2(aq) 3H2(g) Aluminum readily forms the oxide Al2O3 when exposed to air: 4Al(s) 3O2(g) 88n 2Al2O3(s) A tenacious film of this oxide protects metallic aluminum from further corrosion and accounts for some of the unexpected inertness of aluminum. † Charles Martin Hall (1863–1914). American inventor. While Hall was an undergraduate at Oberlin College, he became interested in finding an inexpensive way to extract aluminum. Shortly after graduation, when he was only 22 years old, Hall succeeded in obtaining aluminum from aluminum oxide in a backyard woodshed. Amazingly, the same discovery was made at almost the same moment in France, by Paul Héroult, another 22-year-old inventor working in a similar makeshift laboratory. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 824 METALLURGY AND THE CHEMISTRY OF METALS Aluminum oxide has a very large exothermic enthalpy of formation (∆Ho f 1670 kJ/mol). This property makes aluminum suitable for use in solid propellants for rockets such as those used for the space shuttle Columbia. When a mixture of aluminum and ammonium perchlorate (NH4ClO4) is ignited, aluminum is oxidized to Al2O3, and the heat liberated in the reaction causes the gases that are formed to expand with great force. This action lifts the rocket. The great affinity of aluminum for oxygen is illustrated nicely by the reaction of aluminum powder with a variety of metal oxides, particularly the transition metal oxides, to produce the corresponding metals. A typical reaction is 2Al(s) Fe2O3(s) 88n Al2O3(l ) ∆H° 2Fe(l ) 852 kJ which can result in temperatures approaching 3000°C. This reaction, which is used in the welding of steel and iron, is called the thermite reaction (Figure 20.21). Aluminum chloride exists as a dimer: Cl FIGURE 20.21 reaction. A thermite Cl GD Al D Cl Cl D Al D G Cl Cl Each of the bridging chlorine atoms forms a normal covalent bond and a coordinate covalent bond (indicated by 88n ) with two aluminum atoms. Each aluminum atom is assumed to be sp3-hybridized, so the vacant sp3 hybrid orbital can accept a lone pair from the chlorine atom (Figure 20.22). Aluminum chloride undergoes hydrolysis as follows: AlCl3(s) 3H2O(l ) 88n Al(OH)3(s) 3HCl(aq) Aluminum hydroxide, like Be(OH)2, is amphoteric: Al(OH)3(s) 3s 3p Promotion of electron 3p 3s sp3Hybridized state sp 3 orbitals FIGURE 20.22 The sp3 hybridization of an Al atom in Al2Cl6. Each Al atom has one vacant sp3 hybrid orbital that can accept a lone pair from the bridging Cl atom. 3H (aq) 88n Al3 (aq) Al(OH)3(s) Ground state OH (aq) 88n Al(OH)4 (aq) 3H2O(l ) In contrast to the boron hydrides, which are a well-defined series of compounds, aluminum hydride is a polymer in which each aluminum atom is surrounded octahedrally by bridging hydrogen atoms (Figure 20.23). When an aqueous mixture of aluminum sulfate and potassium sulfate is evaporated slowly, crystals of KAl(SO4)2 12H2O are formed. Similar crystals can be formed by substituting Na or NH4 for K , and Cr3 or Fe3 for Al3 . These compounds are called alums, and they have the general formula M M3 (SO4)2 12H2O M : K , Na , NH4 M3 : Al3 , Cr3 , Fe3 Alums are examples of double salts, that is, salts that contain two different cations. FIGURE 20.23 Structure of aluminum hydride. Each Al atom is surrounded octahedrally by six bridging H atoms. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 20.7 825 ALUMINUM Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry Back Recycling Aluminum Aluminum beverage cans were virtually unknown in 1960; yet by the early 1970s over 1.3 billion pounds of aluminum had been used for these containers. The reasons for aluminum’s popularity in the beverage industry are that it is nontoxic, odorless, tasteless, and lightweight. Furthermore, it is thermally conducting, so the fluid inside the container can be chilled rapidly. The tremendous increase in the demand for aluminum does have a definite drawback, however. More than 3 billion pounds of the metal cans and foils are discarded in the United States annually. They litter the countryside and clog landfills. The best solution to this environmental problem, and the way to prevent the rapid depletion of a finite resource, is recycling. What are the economic benefits of recycling aluminum? Let us compare the energy consumed in the production of aluminum from bauxite with that consumed when aluminum is recycled. The overall reaction for the Hall process can be represented as Al2O3 (in molten cryolite) 3C(s) 88n 2Al(l ) 3CO(g) standard free-energy change for the reaction is given by G° T S° 1340 kJ (1273 K) 586 J K 1 kJ 1000 J 594 kJ Equation (19.3) states that G° nFE°; therefore, the amount of electrical energy needed to produce 1 mole of Al from bauxite is 594 kJ/2, or 297 kJ. Recycling aluminum requires only enough energy to heat the metal to its melting point (660°C) plus the heat of fusion (10.7 kJ/mol). The heat change where 1 mole of aluminum is heated from 25°C to 660°C is heat input Ms t (27.0 g)(0.900 J/g °C)(660 25)°C 15.4 kJ where M is the molar mass, s is the specific heat of Al, and t is the temperature change. Thus, the total energy needed to recycle 1 mole of Al is given by for which H° 1340 kJ and S° 586 J/K. At 1000°C, which is the temperature of the process, the Forward H° total energy 15.4 kJ 10.7 kJ 26.1 kJ Left: Collecting aluminum cans for recycling. Right: Melting and purifying recycled aluminum. Main Menu TOC Study Guide TOC Textbook Website MHHE Website 826 METALLURGY AND THE CHEMISTRY OF METALS Chemistry in Action Chem To compare the energy requirements of the two methods we write energy needed to recycle 1 mol Al energy needed to produce 1 mol Al by electrolysis 26.1 kJ 297 kJ 100% 8.8% SUMMARY OF FACTS AND CONCEPTS Thus, by recycling aluminum cans we can save about 91 percent of the energy required to extract the metal from bauxite. Recycling most of the aluminum cans thrown away each year saves 20 billion kilowatt-hours of electricity—about 1 percent of the electric power used in the United States annually. (Watt is the unit for power, 1 watt 1 joule per second.) 1. Depending on their reactivities, metals exist in nature in either the free or combined state. 2. Recovering a metal from its ore is a three-stage process. First, the ore must be prepared. The metal is then separated, usually by a reduction process, and finally, it is purified. 3. The methods commonly used for purifying metals are distillation, electrolysis, and zone refining. 4. Metallic bonds can be thought of as the force between positive ions immersed in a sea of electrons. In terms of band theory, the atomic orbitals merge to form energy bands. A substance is a conductor when electrons can be readily promoted to the conduction band, where they are free to move through the substance. 5. In insulators, the energy gap between the valence band and the conduction band is so large that electrons cannot be promoted into the conduction band. In semiconductors, electrons can cross the energy gap at higher temperatures, and therefore conductivity increases with increasing temperature as more electrons are able to reach the conduction band. 6. n-Type semiconductors contain donor impurities and extra electrons. p-Type semiconductors contain acceptor impurities and “positive holes.” 7. The alkali metals are the most reactive of all the metallic elements. They have an oxidation number of 1 in their compounds. Under special conditions, some of them also form uninegative ions. 8. The alkaline earth metals are somewhat less reactive than the alkali metals. They almost always have an oxidation number of 2 in their compounds. The properties of the alkaline earth elements become increasingly metallic from top to bottom in their periodic group. 9. Aluminum does not react with water due to the formation of a protective oxide; its hydroxide is amphoteric. KEY WORDS Mineral, p. 804 n-Type semiconductor, p. 814 Ore, p. 804 p-Type semiconductor, p. 814 Conductor, p. 813 Donor impurity, p. 814 Ferromagnetic, p. 805 Insulator, p. 813 Metallurgy, p. 804 Acceptor impurity, p. 814 Alloy, p. 804 Amalgam, p. 806 Band theory, p. 812 Pyrometallurgy, p. 806 Semiconductors, p. 814 QUESTIONS AND PROBLEMS OCCURRENCE OF METALS Review Questions 20.1 Define mineral, ore, and metallurgy. 20.2 List three metals that are usually found in an uncombined state in nature and three metals that are always found in a combined state in nature. Back Forward Main Menu TOC 20.3 Write chemical formulas for the following minerals: (a) calcite, (b) dolomite, (c) fluorite, (d) halite, (e) corundum, (f) magnetite, (g) beryl, (h) galena, (i) epsomite, ( j) anhydrite. Study Guide TOC Textbook Website MHHE Website QUESTIONS AND PROBLEMS 20.4 Name the following minerals: (a) MgCO3, (b) Na3AlF6, (c) Al2O3, (d) Ag2S, (e) HgS, (f) ZnS, (g) SrSO4, (h) PbCO3, (i) MnO2, ( j) TiO2. METALLURGICAL PROCESSES 20.5 Describe the main steps involved in the preparation of an ore. 20.6 What does roasting mean in metallurgy? Why is roasting a major source of air pollution and acid rain? 20.7 Describe with examples the chemical and electrolytic reduction processes used in the production of metals. 20.8 Describe the main steps used to purify metals. 20.9 Describe the extraction of iron in a blast furnace. 20.10 Briefly discuss the steelmaking process. Problems 20.11 In the Mond process for the purification of nickel, CO is passed over metallic nickel to give Ni(CO)4: 20.12 20.13 20.14 20.15 20.16 20.17 Back Main Menu TOC Review Questions 20.19 Define the following terms: conductor, insulator, semiconducting elements, donor impurities, acceptor impurities, n-type semiconductors, p-type semiconductors. 20.20 Briefly discuss the nature of bonding in metals, insulators, and semiconducting elements. 20.21 Describe the general characteristics of n-type and p-type semiconductors. 20.22 State whether silicon would form n-type or p-type semiconductors with the following elements: Ga, Sb, Al, As. ALKALI METALS 4CO(g) 34 Ni(CO)4(g) Given that the standard free energies of formation of 137.3 kJ/mol and CO(g) and Ni(CO)4(g) are 587.4 kJ/mol, respectively, calculate the equilibrium constant of the reaction at 80°C. (Assume Go f to be independent of temperature.) Copper is purified by electrolysis (see Figure 20.6). A 5.00-kg anode is used in a cell where the current is 37.8 A. How long (in hours) must the current run to dissolve this anode and electroplate it onto the cathode? Consider the electrolytic procedure for purifying copper described in Figure 20.6. Suppose that a sample of copper contains the following impurities: Fe, Ag, Zn, Au, Co, Pt, and Pb. Which of the metals will be oxidized and dissolved in solution and which will be unaffected and simply form the sludge that accumulates at the bottom of the cell? How would you obtain zinc from sphalerite (ZnS)? Starting with rutile (TiO2), explain how you would obtain pure titanium metal. (Hint: First convert TiO2 to TiCl4. Next, reduce TiCl4 with Mg. Look up physical properties of TiCl4, Mg, and MgCl2 in a chemistry handbook.) A certain mine produces 2.0 108 kg of copper from chalcopyrite (CuFeS2) each year. The ore contains only 0.80 percent Cu by mass. (a) If the density of the ore is 2.8 g/cm3, calculate the volume (in cm3) of ore removed each year. (b) Calculate the mass (in kg) of SO2 produced by roasting (assume chalcopyrite to be the only source of sulfur). Which of the following compounds would require electrolysis to yield the free metals? Ag2S, CaCl2, NaCl, Fe2O3, Al2O3, TiCl4. Forward 20.18 Although iron is only about two-thirds as abundant as aluminum in Earth’s crust, mass for mass it costs only about one-quarter as much to produce. Why? BAND THEORY OF CONDUCTIVITY Review Questions Ni(s) 827 Review Questions 20.23 How is sodium prepared commercially? 20.24 Why is potassium usually not prepared electrolytically from one of its salts? 20.25 Describe the uses of the following compounds: NaCl, Na2CO3, NaOH, KOH, KO2. 20.26 Under what conditions do sodium and potassium form Na and K ions? Problems 20.27 Complete and balance the following equations: (a) K(s) H2O(l) 88n (b) NaH(s) H2O(l) 88n (c) Na(s) O2(g) 88n (d) K(s) O2(g) 88n 20.28 Write a balanced equation for each of the following reactions: (a) sodium reacts with water; (b) an aqueous solution of NaOH reacts with CO2; (c) solid Na2CO3 reacts with a HCl solution; (d) solid NaHCO3 reacts with a HCl solution; (e) solid NaHCO3 is heated; (f) solid Na2CO3 is heated. 20.29 Sodium hydride (NaH) can be used as a drying agent for many organic solvents. Explain how it works. 20.30 Calculate the volume of CO2 at 10.0°C and 746 mmHg pressure obtained by treating 25.0 g of Na2CO3 with an excess of hydrochloric acid. ALKALINE EARTH METALS Review Questions 20.31 List the common ores of magnesium and calcium. 20.32 How are magnesium and calcium obtained commercially? Study Guide TOC Textbook Website MHHE Website 828 METALLURGY AND THE CHEMISTRY OF METALS Problems 20.33 From the thermodynamic data in Appendix 3, calculate the ∆H° values for the following decompositions: (a) MgCO3(s) 88n MgO(s) CO2(g) (b) CaCO3(s) 88n CaO(s) CO2(g) Which of the two compounds is more easily decomposed by heat? 20.34 Starting with magnesium and concentrated nitric acid, describe how you would prepare magnesium oxide. [Hint: First convert Mg to Mg(NO3)2. Next, MgO can be obtained by heating Mg(NO3)2.] 20.35 Describe two ways of preparing magnesium chloride. 20.36 The second ionization energy of magnesium is only about twice as great as the first, but the third ionization energy is 10 times as great. Why does it take so much more energy to remove the third electron? 20.37 List the sulfates of the Group 2A metals in order of increasing solubility in water. Explain the trend. (Hint: You need to consult a chemistry handbook.) 20.38 Helium contains the same number of electrons in its outer shell as do the alkaline earth metals. Explain why helium is inert whereas the Group 2A metals are not. 20.39 When exposed to air, calcium first forms calcium oxide, which is then converted to calcium hydroxide, and finally to calcium carbonate. Write a balanced equation for each step. 20.40 Write chemical formulas for (a) quicklime, (b) slaked lime, (c) limewater. ALUMINUM Review Questions 20.41 Describe the Hall process for preparing aluminum. 20.42 What action renders aluminum inert? Problems 20.43 Before Hall invented his electrolytic process, aluminum was produced by the reduction of its chloride with an active metal. Which metals would you use for the production of aluminum in that way? 20.44 With the Hall process, how many hours will it take to deposit 664 g of Al at a current of 32.6 A? 20.45 Aluminum forms the complex ions AlCl4 and AlF3 . 6 Describe the shapes of these ions. AlCl3 does not 6 form. Why? (Hint: Consider the relative sizes of Al3 , F , and Cl ions.) 20.46 The overall reaction for the electrolytic production of aluminum by means of the Hall process may be represented as Al2O3(s) 3C(s) 88n 2Al(l ) 3CO(g) At 1000°C, the standard free-energy change for this Back Forward Main Menu TOC 20.47 20.48 20.49 20.50 20.51 20.52 process is 594 kJ. (a) Calculate the minimum voltage required to produce 1 mole of aluminum at this temperature. (b) If the actual voltage applied is exactly three times the ideal value, calculate the energy required to produce 1.00 kg of the metal. In basic solution, aluminum metal is a strong reducing agent and is oxidized to AlO2 . Give balanced equations for the reaction of Al in basic solution with the following: (a) NaNO3, to give ammonia; (b) water, to give hydrogen; (c) Na2SnO3, to give metallic tin. Write a balanced equation for the thermal decomposition of aluminum nitrate to form aluminum oxide, nitrogen dioxide, and oxygen gas. Describe some of the properties of aluminum that make it one of the most versatile metals known. The pressure of gaseous Al2Cl6 increases more rapidly with temperature than predicted by the ideal gas equation even though Al2Cl6 behaves like an ideal gas. Explain. Starting with aluminum, describe with balanced equations how you would prepare (a) Al2Cl6, (b) Al2O3, (c) Al2(SO4)3, (d) NH4Al(SO4)2 12H2O. Explain the change in bonding when Al2Cl6 dissociates to form AlCl3 in the gas phase. ADDITIONAL PROBLEMS 20.53 In steelmaking, nonmetallic impurities such as P, S, and Si are removed as the corresponding oxides. The inside of the furnace is usually lined with CaCO3 and MgCO3, which decompose at high temperatures to yield CaO and MgO. How do CaO and MgO help in the removal of the nonmetallic oxides? 20.54 When 1.164 g of a certain metal sulfide was roasted in air, 0.972 g of the metal oxide was formed. If the oxidation number of the metal is 2, calculate the molar mass of the metal. 20.55 An early view of metallic bonding assumed that bonding in metals consisted of localized, shared electron-pair bonds between metal atoms. What evidence would help you to argue against this viewpoint. 20.56 Referring to Figure 20.6, would you expect H2O and H to be reduced at the cathode and H2O oxidized at the anode? 20.57 A 0.450-g sample of steel contains manganese as an impurity. The sample is dissolved in acidic solution and the manganese is oxidized to the permanganate ion MnO4 . The MnO4 ion is reduced to Mn2 by reacting with 50.0 mL of 0.0800 M FeSO4 solution. The excess Fe2 ions are then oxidized to Fe3 by 22.4 mL of 0.0100 M K2Cr2O7. Calculate the percent by mass of manganese in the sample. 741.0 kJ/mol and that 20.58 Given that ∆Go(Fe2O3) f Study Guide TOC Textbook Website MHHE Website QUESTIONS AND PROBLEMS 20.59 20.60 20.61 20.62 20.63 20.64 20.65 ∆Go(HgS) 48.8 kJ/mol, calculate ∆G° for the f following reactions at 25°C: (a) 2Fe2O3(s) 88n 4Fe(s) 3O2(g) (b) 2Al2O3(s) 88n 4Al(s) 3O2(g) Use compounds of aluminum as example to explain what is meant by amphoterism. When an inert atmosphere is needed for a metallurgical process, nitrogen is frequently used. However, in the reduction of TiCl4 by magnesium, helium is used. Explain why nitrogen is not suitable for this process. It has been shown that Na2 species form in the vapor phase. Describe the formation of the “disodium molecule” in terms of a molecular orbital energy-level diagram. Would you expect the alkaline earth metals to exhibit a similar property? Explain each of the following statements: (a) An aqueous solution of AlCl3 is acidic. (b) Al(OH)3 is soluble in NaOH solutions but not in NH3 solution. Write balanced equations for the following reactions: (a) the heating of aluminum carbonate; (b) the reaction between AlCl3 and K; (c) the reaction between solutions of Na2CO3 and Ca(OH)2. Write a balanced equation for the reaction between calcium oxide and dilute HCl solution. What is wrong with the following procedure for obtaining magnesium? MgCO3 88n MgO(s) MgO(s) CO(g) 88n Mg(s) CO2(g) CO2(g) 20.66 Explain why an aqueous solution of aluminum chloride has a pH lower than 7. Back Forward Main Menu TOC 829 20.67 Predict the chemical properties of francium, the last member of Group 1A. 20.68 Describe a medicinal or health-related application for each of the following compounds: NaF, Li2CO3, Mg(OH)2, CaCO3, BaSO4, Al(OH)2NaCO3 (see p. 633). 20.69 The following are two reaction schemes involving magnesium. Scheme I: When magnesium burns in oxygen, a white solid (A) is formed. A dissolves in 1 M HCl to give a colorless solution (B). Upon addition of Na2CO3 to B, a white precipitate is formed (C). On heating, C decomposes to D and a colorless gas is generated (E). When E is passed through limewater [an aqueous suspension of Ca(OH)2], a white precipitate appears (F). Scheme II: Magnesium reacts with 1 M H2SO4 to produce a colorless solution (G). Treating G with an excess of NaOH produces a white precipitate (H). H dissolves in 1 M HNO3 to form a colorless solution. When the solution is slowly evaporated, a white solid (I) appears. On heating I, a brown gas is given off. Identify A – I and write equations representing the reactions involved. 20.70 Lithium and magnesium exhibit a diagonal relationship in some chemical properties. How does lithium resemble magnesium in its reaction with oxygen and nitrogen? Consult a handbook of chemistry and compare the solubilities of carbonates, fluorides, and phosphates of these metals. 20.71 To prevent the formation of oxides, peroxides and superoxides, alkali metals are sometimes stored in an inert atmosphere. Which of the following gases should not be used for lithium? Why? Ne, Ar, N2, Kr. Study Guide TOC Textbook Website MHHE Website ...
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