Chemistry part 36, Julia Burdge,2e (2009) ppsx

890 CHAPTER 23 Metallurgy and the Chemistry of Metals Figure 23.11 (a) Silicon crystal doped with phosphorus. (b) Silicon crystal doped with boron. Note the formation of a negative center in (a) and a positive center in (b) . • • Figure 23.12 Main group metals (green) and Group 2B metals (blue) according to their po sitions in the periodic table. (a) (b) for example, 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; that is, each Si atom is covalently bonded to four other Si atoms. Phos- phorus ([Ne]3/3 p 3) has one more valence electron than silicon ([Ne ]3 s 2 3p 2), so there is a valence electron left over after four of th em are used to form covalent bonds with silicon (Figure 23.11). 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, because 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 Oi2/2pl), one less than silicon. Thus, for every boron atom in the silicon crystal, th ere is a single vacancy in a bonding orbital. It is possible, though, 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 elec- tron 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 impuritie ~ 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 only a small amount of energy is needed to excite the electron s. Typicall y, 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 vo lume a point of particular importance in space travel, as well as in handheld calculators and microprocessors (computers-on-a-chip). 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 fiat), and ductile (can be drawn into wires). Figure 23.12 shows the positions of the representative metals and the Group lA 1 H Li Na K Rb Cs Fr 2A 2 Be Mg Ca Sr Ba Ra 3B 4B 3 4 Sc Ti Y Zr La Hi Ac Rf 5B 6B 7B1s8 B~ IB 5 6 7 8 9 10 11 V Cr Mn Fe Co Ni Cu Nb Mo Tc Ru Rh Pd Ag Ta W Re Os Ir Pt Au Db Sg Bh Hs Mt Ds Rg 3A 4A 5A 6A 7A 13 14 15 16 17 B C N 0 F 2B 12 AI Si P S Cl Zn Ga Ge As Se Br Cd In Sn Sb Te I Hg Tl Pb Bi Po At 8A 18 He Ne Ar Kr Xe Rn SECTION 23.5 The Alkali Metals 891 2B metals in the periodic table. (The transition metals are discussed in Chapter 22 .) As we saw in Chapter 8, the electronegativity of elements increases from left to right across a period and from bottom to top in a group [ ~ ~ Section 8.4, Figure 8.6 ]. 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 mag- nesium in Group 2A and the metals in Group 3A and beyond also form covalent compounds. In Sections 23.5 through 23.7 we will study the chemistry of selected metals from Group lA (the alkali metals), Group 2A (the alkaline earth metals), and Group 3A (aluminum). 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 23.4 Based on their electron configurations, we expect the oxidation number of these ele- ments in their compounds to be + 1 because the cations would be isoelectronic with the preceding noble gases. This is indeed the case. The alkali metals have low melting points and are soft enough to be sliced with a knife. These metals all possess a body-centered crystal structure 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; instead, they are found combined with halide, sulfate, carbonate, and silicate ions. In this section we will describe the chemistry of two members of Group lA 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 (NaAISi30g) and orthoclase (KAISi30g). 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. Thus, 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 23.13, Chile saltpeter (NaN0 3 ), and sylvite (KCI). Sodium chloride is also obtained from rock salt. Metallic sodium is most conveniently obtained from molten sodium chloride by electroly-' sis in the Downs cell (review Figure 19.10). 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 CaCl 2 , lowers the melting point to about 600°C a more convenient temperature for the electrolysis process. Li Na K Rb Cs Valence electron configuration 2S1 3s 1 4S1 5s 1 6s 1 Density (g/cm 3 ) 0.534 0.97 0.86 1.53 1.87 Melting point (OC) 179 97.6 63 39 28 Boiling point (0C) 1317 892 770 688 678 Atomic radius (pm) 155 190 235 248 267 Ionic radius (pm)* 60 95 133 148 169 Ionization energy (kJ/mol) 520 496 419 403 375 Electronegativity 1.0 0.9 0.8 0.8 0.7 Standard reduction potential (V)t -3.05 -2.71 -2.93 -2.93 -2.92 * Refers to the cation M +, where M denotes an alkali metal atom. t The hal f-reaction is M+(aq) + e- • M(s). Figure 23.13 Halite (NaCt). , 892 CHAPTER 23 Metallurg y and the Chemistry of Metals , . ' •• >': • • Multimedia Periodic Table properties of the alkal i and alkaline earth metals. Figure 23.14 Self-contained breathing apparatus. Metallic potassium cannot be ea sily 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 Na (g) + KCl(l) +. NaCl(l) + K(g) This reaction may seem strange given that potassium is a stronger reducing agent than sodium (s ee Table 23.4). Potass ium has a lower boiling point (770°C) than sodium (892°C), however, so it is more volatile at 892°C and distills off more easily. According to Le Chfltelier's principle, con- stantly removing the potassium vapor drives the reaction to the right, ensuring metallic potassium is recovered. Sodium and potassium are both extremely reactive, but potassium is the more reactive ofthe tw o. Both react with water to form the corresponding hydroxides. In a limited supply of oxygen, sodium bum s to form sodium oxide ( Na ?O). In the presence of excess oxygen, however, sodium forms the pale-yellow peroxide: Sodium peroxide reacts with water to give an alkaline solution and hydrogen peroxide: Like sodium, potassium forms the peroxide. In addition, potassium also forms the superoxide when it bum s in air: K(s) + 0 2(g) +. K0 2 (s) When potassium superoxide reacts with water, oxygen gas is evolved: This reaction is utilized in breathing equipment (Figure 23.14). Exhaled air contains both mois- ture and carbon dioxide. The moisture reacts with K0 2 in the apparatus to generate oxygen gas as shown in the preceding reaction. Furthermore, K0 2 also reacts with exhaled CO 2 , which produces more oxygen gas: 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 beautiful blue solutions: Na NH3 .Na + + e- . Both the cation and the electron exist in the solvated form, and the solvated electrons are respon- sible 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 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- ! (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. This finding is of considerable theoretical inter- est, because 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 com- pounds of sodium and potassium. Sodium Chloride The source, properties, and uses of sodium chloride were discussed in Chapter 7. SECTION 23.6 The Alkaline Earth Metals 893 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 Na2C03 produced is used in the glass industry. Sodium carbonate ranks eleventh among the chemicals produced in the United States. For many years, Na2C03 was produced by the Solvay2 process, in which ammonia is first dissolved in a saturated solution of sodium chloride. Bubbling carbon dioxide into the solution precipitates sodium bicarbonate as follows: Sodium bicarbonate is then separated from the solution and heated to give sodium carbonate: However, the rising cost of ammonia and the pollution problem resulting from the by-products have prompted chemists to look for other sources of sodium carbonate. One is the mineral trona [N as(C0 3 MHC0 3 ) • 2H 2 0], large deposits of which have been found in Wyoming. When trona is crushed and h.,ated, it decomposes as follows: 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 Na2C03 . lOH 2 0. 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 NaCI and KCI solutions; 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 evo- lution of oxygen at about 500° C: Potassium nitrate (saltpeter) is prepared beginning with the "reaction" KCI(aq) + NaN0 3 (aq) +. KN0 3 (aq) + NaCI(aq) This process is carried out just below 100° C. Because KN0 3 is the least soluble salt at room re mperature, it is separated from the solution by fractional crystallization. Like NaN0 3 , KN0 3 decomposes when heated. Gunpowder consists of potassium nitrate, wood charcoal, and sulfur in the approximate pro- portions of 6: 1: 1 by mass. When gunpowder is heated, the reaction is The sudden formation of hot nitrogen and carbon dioxide gases causes an explosion. The Alkaline Earth Metals The alkaline earth metals are somewhat less electropositive and less reactive than the alkali met- als. 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 ~ ~.1 ions attain the stable electron configuration ofthe preceding noble gas, the oxidation number 0 alk aline earth metals in the combined form is almost always +2. Table 23.5 lists some common :: F::nest Solvay (1838-1922). Belgian chemist. Solvay's main contribution to industrial chemistry was the development of !he process for the production of sodium carbonate that now bears his name. • 894 CHAPTER 23 Metal l urgy and the Chemistry of Metals • Figure 23.15 Dolomite ( CaC0 3 • MgC0 3 ) · Be Mg Ca Sr Ba Valence electron configuration 2i 3i 4i 5s 2 6i Density (g/ cm 3) 1.86 1.74 1.55 2.6 3.5 Melting point (DC) 1280 650 838 770 714 Boiling point (DC) 2770 1107 1484 1380 1640 Atomic radius (pm) 112 160 197 215 222 Ionic radius ( pm )* 31 65 99 113 135 First ionization energy (k J/ mol ) 899 738 590 548 502 Second ionization energy (k J/mol ) 1757 1450 1145 1058 958 Electronegativity 1.5 1.2 1.0 1.0 0.9 Standard reduction potential (V)t -1.85 -2.37 -2.87 -2 .89 -2.90 * Refers to the cation M 2 +, w here M denotes an alkali earth metal atom . t The half-reaction is M 2 +(aq) + 2e- • M(s). properties of these metal s. 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 Magnesium 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)z], dolomite (CaC0 3 • MgC0 3 ) (Figure 23.l5), and epsomite (MgS04 . 7H z O). 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 met- als, metallic magnesium is obtained by electrolysis, in this case from its molten chloride, MgCl 2 (obtained from seawater). 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) + HzO(g) - -+. MgO(s) + HzCg) It burns brilliantly in air to produce magnesium oxide and magnesium nitride: 2Mg(s) + Oz(g) - -+. 2MgO(s) 3Mg (s) + Nz(g) • Mg 3 N 2 (s) This property makes magnesium (in the form of thin ribbons or fibers) usef ul in flash photography and flares. Magnesium oxide reacts very slowly with water to form magnesi um hydroxide, a white solid suspension called milk of magnesia, which is used to treat acid indigestion: MgO (s) + HzO(I) - -+. Mg(OHh(s) Magnesium is a typical alkaline earth metal in that its hydroxide is a strong base. (The only alka- line earth hydroxide that is not a strong base is Be(OHh which is amphoteric.) The major uses of magnesium are in lightweight structural alloys, for ca thodic protection; in organic synthesis; and in batteries. Magnesium is essential to plant and animal life, and Mg 2+ ions are not toxic. It is estimated that the average adult ingests about 0.3 g of magnesium ions daily. Magne s ium plays several important biological roles. It is present, for instance, in intracellular and extracellular fluids, and 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 by mass. Calcium occ ur s in limestone, calcite, chalk, and marble as CaC0 3 ; in dolomite as CaC0 3 • MgC0 3 (see Figure 23.l5); in gypsum as CaS04 . 2H 2 0; and in fluorite as CaF2 (Figure 23.16). Metallic calcium is best prepared by the electrolysis of molten calcium chloride (CaCI 2 ). As we read down Group 2A from beryllium to barium, metallic properties increase. 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: Ca(s) + 2H 2 0(l) - Ca(OHMaq) + H 2 (g) Calcium hydroxide [Ca(OH )z J is commonly known as slaked lime or hydrated lime. Lime (CaO ), which is also referred to as quicklime, is one of the oldest materials known to humankind. Quick- lime is produced by the thermal decomposition of calcium carbonate: whereas slaked lime is produced by the reaction between quicklime and water: CaO(s) + H 2 0 (I) - Ca ( OH Maq) Quicklime is used in metallurgy (see Section 23.2) and in the removal of S0 2 when fossil fuel is burned. Slaked lime is used in water treatment. For many years, farmers have used lime to lower the acidity of the soil for their crops (a process called liming). Nowadays lime is also applied to lakes affected by acid rain. 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 bone s and teeth; the calcium ion is present in a complex phosphate salt called hydroxyapatite [Ca5(P04)3 0HJ. A characteristic function of Ca 2+ ions in living systems is the activation of a variety of metabolic processes, including a vital role in heart action, blood clotting, muscle contraction, and nerve impulse transmission. Aluminum Aluminum is the most abundant metal and the third most plentiful element in Earth's crust (7.5 percent by ma ss). The elemental fOlln does not occur in nature; instead, its principal ore is bauxite (A1 2 0 3 . 2H 2 0). Other minerals containing aluminum are orthoclase (KA1Si 3 0 8 ), beryl (Be3AI2Si6018), cryolite (Na3AlF6), and corundum ( A1 2 0 3 ) (Figure 23.17). Aluminum is usually prepared from bauxite, which is frequently contaminated with silica (Si0 2 ), iron oxide s, and titanium(IV) oxide. The ore is first heated in sodium hydroxide solution to convelt the silica into soluble silicates: Si0 2 (s) + 20H - (aq) - SiO j- (aq) + H 2 0 (l) At the same time, aluminum oxide is converted to the aluminate ion (AI0 2 ): Iron oxide and titanium oxide are unaffected by this treatment and are filtered off. Next, the solu- tion is treated with acid to precipitate the insoluble aluminum hydroxide: After filtration, the aluminum hydroxide is heated to obtain aluminum oxide: Anhydrous aluminum oxide, or corundum, is reduced to aluminum by the Ha1l 3 process. Figure 23.18 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 3. Charles Martin Hall (1863-1914). American invent or. 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. Amazingl y, the same discovery was made at almost the same moment in France by Paul Heroult, another 22-year-old inventor working in a similar makeshift laboratory. SECTION 23.7 Aluminum 895 Figure 23.16 Fluorite (CaF 2 ) . Figure 23.17 Corundum (A1 2 0 3 ). 896 CHAPTER 23 Metallurgy and the Chemistry of Metals Figure 23.18 Electrolytic production of aluminum ba s ed on the Hall process. • • Al 2 0 3 in molten cryolite Carbon anodes _ ' "- ,,- cathode Molten aluminum cryolite (Na3AIF6; m.p. lOOO °C) as the solv ent for aluminum oxide (m.p. 2045°C). The mixture is electrolyzed to produce aluminum and oxygen ga s: Anode (oxidation): 3[20 2 - 0 2(g) + 4e- ] Cat hod e (reduction): 4[AI 3+ + 3e- • AI(l)] Overall: 2Al 2 0 3 • 4AI(l) + 30 2 (g) Oxygen ga s reacts with the carbon anodes (at elevated temperatures) to form carbon monoxide, which es cape s as a ga s. The liquid aluminum metal (m.p. 660.2°C) sinks to the bottom of the ves- sel, from which it can be drained from time to time during the procedure. Aluminum is one of the most v er satile metals known. It has a low density (2.7 g/cm 3 ) and high tensile strength (i.e., 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. Howe ver, because aluminum is cheaper and lighter than copper, it is widely used in high- voltage transmission lines. Although aluminum's chief use is in 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, metallic properties grad- ually decreas e. Thu s, 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: 2AI(s) + 6HCI(aq) +. 2AICI 3 (aq) + 3Hig) • 2AI(s) + 2NaOH(aq) + 2H 2 0(/) • 2NaAI0 2 (aq) + 3H 2 (g) Aluminum readily forms the oxide Al 2 0 3 when exposed to air: A tenacious film of this oxide protects metallic aluminum from further corrosion and accounts for some of the unexpected inertness of aluminum. Aluminum oxide has a very large exothermic enthalpy of formation (!1H 'f = -1670 kJ /mol). This property makes aluminum suitable for use in solid propellants for rockets such as those used for some space shuttles. When a mixture of aluminum and ammonium perchlorate (NH 4 CI0 4 ) is ignited, aluminum is oxidized to A1 2 0 3 , 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 cor- responding metals. A typical reaction is !1H o = -852 kJ/mol which can result in temperatures approaching 3000° C. This transformation, which is used in the welding of steel and iron, is called the thermite reaction (Figure 23.19). Aluminum chloride exists as a dimer: Each of the bridging chlorine atoms forms a normal covalent bond and a coordinate covalent bond (each indicated by an arrow) with two aluminum atoms. Each aluminum atom is assumed to be Sp 3 -hybridized, so the vacant si hybrid orbital can accept a lone pair from the chlorine atom (Fig- ure 23.20). Aluminum chloride undergoes hydrolysis as follows: AICI 3 (s) + 3H 2 0(l) +. AI(OHMs) + 3HCI(aq) Aluminum hydroxide, like Be(OHh, is amphoteric: AI(OHMs) + 3H+(aq) • AI 3+ (aq) + 3H 2 0(I) AI(OHMs) + OH - (aq) • AI(OH)4 (aq) 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 23.21). When an aqueous mixture of aluminum sulfate and potassium sulfate is evaporated slowly, crystals of KAI(S04h . 12H 2 0 are formed. Similar crystals can be formed by substituting Na + or NHt for K+, and Cr 3+ or Fe 3+ for AI 3 +. These compounds are called alums, and they have the general formula M+: K+, Na +, NH 4 + M 3+ : AI 3+ , Cr 3+ , Fe 3 + Alums are examples of double salts that is, salts that contain two different cations. SECTION 23.7 Aluminum 897 Figure 23.19 The temperature of a thermite reaction can reach 3000 °C. Ground state Promotion of electron Sp3_ hybridiz ed state 11 1 1 3s 3p IT] 11 11 1 1 3s 3p 1111111 I I sp3 orbitals Figure 23.20 The Sp 3 hybridization of an Al atom in A1 2 Cl 6 . Each Al atom has one vacant sp3 hybrid orbital that can accept a lone pair from the bridging CI atom. Figure 23.21 Structure of aluminum hydride. Note that this compound is a polymer. Each Al atom is s urrounded in an octahedral arrangement by six bridging H atoms. 898 CHAPTER 23 Metallurgy and the Chemistry of Metals • Applying What You've Learned Most health problems related to copper are the result of errors in copper metabolism. However, although it is rare, copper deficiency can result from a diet that is poor in cop- per. Symptoms of dietary copper deficiency include anemia (a deficiency of red blood cells) and neutropenia (a deficiency of a particular type of white blood cell). The fact that copper is essential to human health was first demonstrated with a group of children in Peru. One patient's ordeal was detailed by Cordano and Graham in the journal Pediatrics in 1966. During her first few years of life, the patient was hospital- ized several times with anemia, neutropenia, osteoporosis, and multiple fractures. At age 6, over a period of 3 months, she received 20 blood transfusions for her severe anemia, which had not responded to treatment. When Dr. Cordano became aware of the patient's history, he initiated treatment with copper supplementation. The patient never required another transfusion and after 6 months on copper supplements, at age 7, she walked for the first time in her life. Writing Prompt: Research the subject on the Web, and write a SOO-word essay on the causes, diagnosis, and treatment of dietary copper deficiency. Include a specific case study. • l CHAPTER SUMMARY Section 23.1 • Depending on their reactivities, metals exist in nature in either the free or combined state. (More reactive metals are found combined with other elements.) Most metals are found in minerals. Minerals with high metal content are called ores. Section 23.2 • Metallurgy involves recovering metal from ores. The three stages of metal recovery are preparation, separation, and purification. An alloy is a solid mixture of one or more metals, sometimes also containing one or more nonmetals. An amalgam is a mixture of mercury and one or more other metals. • The methods' commonly used for purifying metals are distillation, electrolysis, and zone refining. Pyrometallurgy refers to metallurgical processes carried out at high temperatures. Section 23.3 • 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. In an insulator, the energy gap between the valence band and the conduction band is so large that electrons cannot be promoted into the conduction band. • Semiconductors are substances that normally are not conductors but will conduct electricity at elevated temperatures or when combined KEyWORDS Alloy, 883 Amalgam, 883 Band theory, 888 Conductor, 889 Insulator, 889 Hall process, 895 Metallurgy, 883 QUESTIONS AND PROBLEMS Section 23.1: Occurrence of Metals Review Questions 23 .1 23.2 23 .3 Define the terms mineral and ore. 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. 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. QUESTIONS AND PROBLEMS 899 with a small amount of certain other elements. Semiconductors in which an electron-rich impurity is added to enhance conduction are known as n-type semiconductors. Semiconductors in which an electron-poor impurity is added to enhance conduction are known as p-type semiconductors. Section 23.4 • Metals typically are good conductors and are malleable and ductile. Metallic character increases from top to bottom in a group and decreases from left to right across a period. Section 23.5 • The alkali metals are the most reactive of all the metallic elements. They have an oxidation state of + I in their compounds. Under special conditions, some of them can form anions with an oxidation state of -1. Section 23.6 • 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 group. Section 23.7 • Aluminum ordinarily does not react with water due to a protective coating of aluminum oxide; its hydroxide is amphoteric. The Hall process is used to reduce aluminum oxide to aluminum. Mineral, 882 p-type semiconductor, 890 Pyrometallurgy, 883 Semiconductor, 889 n-type semiconductor, 890 Ore, 882 23.4 Name the following minerals: (a) MgC0 3 , (b) Na3AIF6, (c) A1 2 0 3 , (d) Ag 2 S; (e) HgS, (f) ZnS, (g) SrS04, (h) PbC0 3 , (i) Mn0 2, (j) Ti0 2 · • Section 23.2: Metallurgical Processes Review Questions 23.5 23.6 23.7 Define the terms metallurgy, alloy, and amalgam. Describe the main steps involved in the preparation of an ore. What does roasting mean in metallurgy? Why is roasting a major source of air pollution and acid rain? [...]... CHAPTER 24 Nonmetallic Elements and Their Compounds Concentrated nitric acid does not oxidize gold However, when the acid is added to concentrated hydrochloric acid in a 1:3 ratio by volume (one part HN0 3 to three parts HC1), the resulting solution, called aqua regia, can oxidize gold as follows: Au(s) + 3HN0 3 (aq) + 4HCl(aq) - - HAuCI4 (aq) + 3H20(I) + 3NOig) The oxidation of Au is promoted by the... Starting with rutile (TiOl ), explain how you would obtain pure titanium metal (Hint: First convert TiO z to TiCI 4 Next, reduce TiCl4 with Mg Look up physical properties of TiCI 4, Mg, and MgCI l in a chemistry handbook.) 23.19 + H10(1) (b) NaH(s) + H 2 0(l) (c) Na(s) + Oz(g) (d) K(s) + 0 2(g) Although iron is only about two-thirds as abundant as aluminum in Earth's crust, mass for mass it costs only... two ways of preparing magnesium chloride 23.41 23.54 23.57 23.39 23.42 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.) Additional Problems 23.44 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... a similar property? 23.67 Explain each of the following statements: (a) An aqueous solution of AlCl 3 is acidic (b) Al(OH)3 is soluble in NaOH solution but not in NH3 solution 902 Metallurgy and the Chemistry of Metals CHAPTER 23 23.68 Write balanced equations for the following reactions: (a) the heating of aluminum carbonate, (b) the reaction between AICI 3 and K, (c) the reaction between solutions... solution 23.70 23.75 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 23.76 To prevent the formation of oxides, peroxides, and superoxides, alkali metals are sometimes stored... peroxide and superoxide ion The accumulation of ROS can lead to DNA damage and initiate carcinogenic processes Arsenic is a metalloid The metalloids and the nonmetallic elements and their compounds exhibit chemistry that varies considerably In This Chapter, You Will Learn some of the properties of metalloids and nonmetals and which compounds they form Before you begin, you should review • General trends... corner ofthe periodic table (Figure 24.1) Compounds formed by a combination of metals with nonmetals tend to be ionic, having a metallic cation and a nonmetallic anion In this chapter we will discuss the chemistry of a number of common and important nonmetallic elements-namely, hydrogen; carbon (Group 4A); nitrogen and phosphorus (Group SA); oxygen and sulfur (Group 6A); and the halogens: fluorine, chlorine,... the oxidizer dinitrogen tetroxide (N20 4 ), are used as rocket fuels Hydrazine also plays a role in polymer synthesis and in the manufacture of pesticides There are many nitrogen oxides, but the three particularly important ones are nitrous oxide, nitric oxide, and nitrogen dioxide Nitrous oxide (N 20) is a colorless gas with a pleasing odor and sweet taste It is prepared by heating ammonium nitrate...900 CHAPTER 23 Metallurgy and the Chemistry of Metals State whether silicon would form n-type or p-type semiconductors with the following elements: Ga, Sb, AI, As 23.8 Describe with examples the chemical and electrolytic reduction processes... 6N0 2(g) + 2H20(I) Nitric acid is used in the manufacture of fertilizers, dyes , drugs, and explosives Phosphorus Like nitrogen, phosphorus is a member of the Group SA family, and in some respects the chemistry of phosphorus resembles that of nitrogen Phosphorus occurs most commonly in nature as phosphate rocks, which are mostly calcium phosphate [Ca3(P04)2] and fiuoroapatite [Cas(P04)3F] (Figure 24.8) . carbonate, and silicate ions. In this section we will describe the chemistry of two members of Group lA sodium and potassium. The chemistry of lithium, rubidium, and cesium is less important;. to industrial chemistry was the development of !he process for the production of sodium carbonate that now bears his name. • 894 CHAPTER 23 Metal l urgy and the Chemistry of Metals. Group 3A and beyond also form covalent compounds. In Sections 23.5 through 23.7 we will study the chemistry of selected metals from Group lA (the alkali metals), Group 2A (the alkaline earth