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Unformatted text preview: Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website CHAPTER 21 Nonmetallic Elements and Their Compounds INTRODUCTION OF THE 112 KNOWN ELEMENTS, ONLY TWENTY-FIVE ARE NONMETALLIC. UNLIKE THE METALS, THE CHEMISTRY OF THESE ELEMENTS IS DIVERSE. DESPITE 21.1 GENERAL PROPERTIES OF NONMETALS THEIR RELATIVELY SMALL NUMBER, MOST OF THE ESSENTIAL EL- EMENTS IN BIOLOGICAL SYSTEMS ARE NONMETALS THIS (O, C, H, N, P, 21.2 HYDROGEN 21.3 CARBON AND S). 21.4 NITROGEN AND PHOSPHORUS GROUP ALSO INCLUDES THE MOST UNREACTIVE ELEMENTS — THE NOBLE GASES. THE 21.5 OXYGEN AND SULFUR 21.6 THE HALOGENS UNIQUE PROPERTIES OF HYDROGEN SET IT APART FROM THE REST OF THE ELEMENTS IN THE PERIODIC TABLE. A WHOLE BRANCH OF CHEMISTRY — ORGANIC CHEMISTRY — IS BASED ON CARBON COMPOUNDS. IN THIS CHAPTER WE CONTINUE OUR SURVEY OF THE ELEMENTS WITH A DISCUSSION OF NONMETALS. THE EMPHASIS WILL BE ON IM- PORTANT CHEMICAL PROPERTIES AND ON THE ROLES OF NONMETALS AND THEIR COMPOUNDS IN INDUSTRIAL, CHEMICAL, AND BIOLOGICAL PROCESSES. 831 Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 832 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS 21.1 Recall that there is no totally suitable position for hydrogen in the periodic table. 21.2 1 Properties of nonmetals are more varied than those of metals. A number of nonmetals are gases in the elemental state: hydrogen, oxygen, nitrogen, fluorine, chlorine, and the noble gases. Only one, bromine, is a liquid. All the remaining nonmetals are solids at room temperature. Unlike metals, nonmetallic elements are poor conductors of heat and electricity; they exhibit both positive and negative oxidation numbers. A small group of elements, called metalloids, have properties characteristic of both metals and nonmetals. The metalloids boron, silicon, germanium, and arsenic are semiconducting elements (see Section 20.3). Nonmetals are more electronegative than metals. The electronegativity of elements increases from left to right across any period and from bottom to top in any group in the periodic table (see Figure 9.5). With the exception of hydrogen, the nonmetals are concentrated in the upper right-hand corner of the periodic table (Figure 21.1). Compounds formed by a combination of metals and 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: hydrogen; carbon (Group 4A); nitrogen and phosphorus (Group 5A); oxygen and sulfur (Group 6A); and fluorine, chlorine, bromine, and iodine (Group 7A). HYDROGEN Hydrogen is the simplest element known — its most common atomic form contains only one proton and one electron. The atomic form of hydrogen exists only at very high temperatures, however. Normally, elemental hydrogen is a diatomic molecule, the product of an exothermic reaction between H atoms: FIGURE 21.1 Representative nonmetallic elements (blue) and metalloids (gray). 1 1A H GENERAL PROPERTIES OF NONMETALS H(g) H(g) 88n H2(g) H° 436.4 kJ 18 8A 2 2A 13 3A 14 4A 15 5A 16 6A 17 7A 2 He 3 4 5 6 7 8 9 10 Li Be B C N O F Ne 13 14 15 16 17 18 Al Si P S Cl Ar 11 12 Na Mg 3 3B 4 4B 5 5B 6 6B 7 7B 8 9 8B 10 11 1B 12 2B 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 110 111 112 87 Back 88 89 104 105 106 107 108 109 Fr Ra Ac Rf Ha Sg Ns Hs Mt Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.2 HYDROGEN 833 Molecular hydrogen is a colorless, odorless, and nonpoisonous gas. At 1 atm, liquid hydrogen has a boiling point of 252.9°C (20.3 K). Hydrogen is the most abundant element in the universe, accounting for about 70 percent of the universe’s total mass. It is the tenth most abundant element in Earth’s crust, where it is found in combination with other elements. Unlike Jupiter and Saturn, Earth does not have a strong enough gravitational pull to retain the lightweight H2 molecules, so hydrogen is not found in our atmosphere. The ground-state electron configuration of H is 1s1. It resembles the alkali metals in that it can be oxidized to the H ion, which exists in aqueous solutions in the hydrated form. On the other hand, hydrogen resembles the halogens in that it forms the uninegative hydride ion (H ), which is isoelectronic with helium (1s2). Hydrogen is found in a large number of covalent compounds. It also has the unique capacity for hydrogen-bond formation (see Section 11.2). Hydrogen gas plays an important role in industrial processes. About 95 percent of the hydrogen produced is used captively; that is, it is produced at or near the plant where it is used for industrial processes, such as the synthesis of ammonia. The largescale industrial preparation is the reaction between propane (from natural gas and also as a product of oil refineries) and steam in the presence of a catalyst at 900°C: C3H8(g) 3H2O(g) 88n 3CO(g) 7H2(g) In another process, steam is passed over a bed of red-hot coke: C(s) H2O(g) 88n CO(g) H2(g) The mixture of carbon monoxide and hydrogen gas produced in this reaction is commonly known as water gas. Because both CO and H2 burn in air, water gas was used as a fuel for many years. But because CO is poisonous, water gas has been replaced by natural gases, such as methane and propane. Small quantities of hydrogen gas can be prepared conveniently in the laboratory by reacting zinc with dilute hydrochloric acid (Figure 21.2): FIGURE 21.2 Apparatus for the laboratory preparation of hydrogen gas. The gas is collected over water, as is also the case of oxygen gas (see Figure 5.13). HCl H2 gas Water Zn Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 834 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS Zn(s) 2HCl(aq) 88n ZnCl2(aq) H2(g) Hydrogen gas can also be produced by the reaction between an alkali metal or an alkaline earth metal (Ca or Ba) and water (see Section 4.4), but these reactions are too violent to be suitable for the laboratory preparation of hydrogen gas. Very pure hydrogen gas can be obtained by the electrolysis of water, but this method consumes too much energy to be practical on a large scale. BINARY HYDRIDES Binary hydrides are compounds containing hydrogen and another element, either a metal or a nonmetal. Depending on structure and properties, these hydrides are broadly divided into three types: (1) ionic hydrides, (2) covalent hydrides, and (3) interstitial hydrides. Ionic Hydrides Ionic hydrides are formed when molecular hydrogen combines directly with any alkali metal or with the alkaline earth metals Ca, Sr, or Ba: 2Li(s) H2(g) 88n 2LiH(s) Ca(s) H2(g) 88n CaH2(s) All ionic hydrides are solids that have the high melting points characteristic of ionic compounds. The anion in these compounds is the hydride ion, H , which is a very strong Brønsted base. It readily accepts a proton from a proton donor such as water: H (aq) H2O(l ) 88n OH (aq) H2(g) Due to their high reactivity with water, ionic hydrides are frequently used to remove traces of water from organic solvents. Covalent Hydrides This is an example of the diagonal relationship between Be and Al (see p. 308). Back Forward In covalent hydrides the hydrogen atom is covalently bonded to the atom of another element. There are two types of covalent hydrides — those containing discrete molecular units, such as CH4 and NH3, and those having complex polymeric structures, such as (BeH2)x and (AlH3)x, where x is a very large number. Figure 21.3 shows the binary ionic and covalent hydrides of the representative elements. The physical and chemical properties of these compounds change from ionic to covalent across a given period. Consider, for example, the hydrides of the secondperiod elements: LiH, BeH2, B2H6, CH4, NH3, H2O, and HF. LiH is an ionic compound with a high melting point (680°C). The structure of BeH2 (in the solid state) is polymeric; it is a covalent compound. The molecules B2H6 and CH4 are nonpolar. In contrast, NH3, H2O, and HF are all polar molecules in which the hydrogen atom is the positive end of the polar bond. Of this group of hydrides (NH3, H2O, and HF), only HF is acidic in water. As we move down any group in Figure 21.3, the compounds change from covalent to ionic. In Group 2A, for example, BeH2 and MgH2 are covalent, but CaH2, SrH2, and BaH2 are ionic. Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.2 1 1A 835 HYDROGEN 18 8A Discrete molecular units 2 2A 13 3A Ionic compound 14 4A 15 5A 16 6A 17 7A B2H6 Polymeric structure; covalent compound CH4 NH3 H2O HF AlH3 SiH4 PH3 H2S HCl LiH BeH2 NaH MgH2 KH CaH2 GaH3 GeH4 AsH3 H2Se HBr RbH SrH2 InH3 SnH4 SbH3 H2Te HI CsH BaH2 TlH3 PbH4 BiH3 3 3B 4 4B 5 5B 6 6B 7 7B 8 9 8B 10 11 1B 12 2B FIGURE 21.3 Binary hydrides of the representative elements. Where hydrogen forms more than one compound with the same element, only the formula of the simplest hydride is shown. The properties of many of the transition metal hydrides are not well characterized. Interstitial Hydrides These compounds are sometimes called nonstoichiometric compounds. Note that they do not obey the law of definite proportions (see Section 2.1). Molecular hydrogen forms a number of hydrides with transition metals. In some of these compounds, the ratio of hydrogen atoms to metal atoms is not a constant. Such compounds are called interstitial hydrides. For example, depending on conditions, the formula for titanium hydride can vary between TiH1.8 and TiH2. Many of the interstitial hydrides have metallic properties such as electrical conductivity. Yet it is known that hydrogen is definitely bonded to the metal in these compounds, although the exact nature of the bonding is often not clear. Molecular hydrogen interacts in a unique way with palladium (Pd). Hydrogen gas is readily adsorbed onto the surface of the palladium metal, where it dissociates into atomic hydrogen. The H atoms then “dissolve” into the metal. On heating and under the pressure of H2 gas on one side of the metal, these atoms diffuse through the metal and recombine to form molecular hydrogen, which emerges as the gas from the other side. Because no other gas behaves in this way with palladium, this process has been used to separate hydrogen gas from other gases on a small scale. ISOTOPES OF HYDROGEN 1 1H The isotope is also called protium. Hydrogen is the only element whose isotopes are given different names. Back Forward Main Menu Hydrogen has three isotopes: 1H (hydrogen), 2H (deuterium, symbol D), and 3H (tri1 1 1 tium, symbol T). The natural abundances of the stable hydrogen isotopes are hydrogen, 99.985 percent; and deuterium, 0.015 percent. Tritium is a radioactive isotope with a half-life of about 12.5 years. Table 21.1 compares some of the common properties of H2O with those of D2O. Deuterium oxide, or heavy water as it is commonly called, is used in some nuclear reactors as a coolant and a moderator of nuclear reactions (see Chapter 23). D2O can be TOC Study Guide TOC Textbook Website MHHE Website 836 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS TABLE 21.1 PROPERTY Properties of H2O and D2O H2O D2O Molar mass (g/mol) Melting point (°C) Boiling point (°C) Density at 4°C (g/cm3) 18.0 0 100 1.000 20.0 3.8 101.4 1.108 separated from H2O by fractional distillation because H2O boils at a lower temperature, as Table 21.1 shows. Another technique for its separation is electrolysis of water. Since H2 gas is formed about eight times as fast as D2 during electrolysis, the water remaining in the electrolytic cell becomes progressively enriched with D2O. Interestingly, the Dead Sea, which for thousands of years has entrapped water that has no outlet other than through evaporation, has a higher [D2O]/[H2O] ratio than water found elsewhere. Although D2O chemically resembles H2O in most respects, it is a toxic substance. The reason is that deuterium is heavier than hydrogen; thus, its compounds often react more slowly than those of the lighter isotope. Regular drinking of D2O instead of H2O could prove fatal because of the slower rate of transfer of D compared with that of H in the acid-base reactions involved in enzyme catalysis. This kinetic isotope effect is also manifest in acid ionization constants. For example, the ionization constant of acetic acid CH3COOH(aq) 34 CH3COO (aq) H (aq) Ka 1.8 Ka 6 10 5 is about three times as large as that of deuterated acetic acid: CH3COOD(aq) 34 CH3COO (aq) D (aq) 10 6 HYDROGENATION Hydrogenation is the addition of hydrogen to compounds containing multiple bonds, especially CPC and CqC bonds. A simple hydrogenation reaction is the conversion of ethylene to ethane: Platinum catalyst on alumina (Al2O3), used in hydrogenation. Back Forward HH AA HOCO COH AA HH ethylene H2 H H G D C PC G D H H ethane This reaction is quite slow under normal conditions, but the rate can be greatly increased by the presence of a catalyst such as nickel or platinum. As in the Haber synthesis of ammonia (see Section 13.6), the main function of the catalyst is to weaken the HOH bond and facilitate the reaction. Hydrogenation is an important process in the food industry. Vegetable oils have considerable nutritional value, but some oils must be hydrogenated before we can use them because of their unsavory flavor and their inappropriate molecular structures (that is, there are too many CPC bonds present). Upon exposure to air, these polyunsaturated molecules (that is, molecules with many CPC bonds) undergo oxidation to yield unpleasant-tasting products (oil that has oxidized is said to be rancid). In the hydrogenation process, a small amount of nickel (about 0.1 percent by mass) is added to the oil and the mixture is exposed to hydrogen gas at high temperature and pressure. Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.3 CARBON 837 Afterward, the nickel is removed by filtration. Hydrogenation reduces the number of double bonds in the molecule but does not completely eliminate them. If all the double bonds are eliminated, the oil becomes hard and brittle. Under controlled conditions, suitable cooking oils and margarine may be prepared by hydrogenation from vegetable oils extracted from cottonseed, corn, and soybeans. THE HYDROGEN ECONOMY The world’s fossil fuel reserves are being depleted at an alarmingly fast rate. Faced with this dilemma, scientists have made intensive efforts in recent years to develop a method of obtaining hydrogen gas as an alternative energy source. Hydrogen gas could replace gasoline to power automobiles (after considerable modification of the engine, of course) or be used with oxygen gas in fuel cells to generate electricity (see p. 778). One major advantage of using hydrogen gas in these ways is that the reactions are essentially free of pollutants; the end product formed in a hydrogen-powered engine or in a fuel cell would be water, just as in the burning of hydrogen gas in air: 2H2(g) The total volume of ocean water is about 1 1021 L. Thus, the ocean contains an almost inexhaustible supply of hydrogen. 21.3 The carbon cycle is discussed on p. 705. Back Forward Main Menu O2(g) 88n 2H2O(l ) Of course, success of a hydrogen economy would depend on how cheaply we could produce hydrogen gas and how easily we could store it. Although electrolysis of water consumes too much energy for large-scale application, if scientists can devise a more practical method of “splitting” water molecules, we could obtain vast amounts of hydrogen from seawater. One approach that is currently in the early stages of development would use solar energy. In this scheme a catalyst (a complex molecule containing one or more transition metal atoms, such as ruthenium) absorbs a photon from solar radiation and becomes energetically excited. In its excited state the catalyst is capable of reducing water to molecular hydrogen. Some of the interstitial hydrides we have discussed would make suitable storage compounds for hydrogen. The reactions that form these hydrides are usually reversible, so hydrogen gas can be obtained simply by reducing the pressure of the hydrogen gas above the metal. The advantages of using interstitial hydrides are as follows: (1) many metals have a high capacity to take up hydrogen gas — sometimes up to three times as many hydrogen atoms as there are metal atoms; and (2) because these hydrides are solids, they can be stored and transported more easily than gases or liquids. The Chemistry in Action essay on p. 838 describes what happens to hydrogen under pressure. CARBON Although it constitutes only about 0.09 percent by mass of Earth’s crust, carbon is an essential element of living matter. It is found free in the form of diamond and graphite (see Figure 8.17), and it is also a component of natural gas, petroleum, and coal. (Coal is a natural dark-brown to black solid used as a fuel; it is formed from fossilized plants and consists of amorphous carbon with various organic and some inorganic compounds.) Carbon combines with oxygen to form carbon dioxide in the atmosphere and occurs as carbonate in limestone and chalk. Diamond and graphite are allotropes of carbon. Figure 21.4 shows the phase diagram of carbon. Although graphite is the stable form of carbon at 1 atm and 25°C, owners of diamond jewelry need not be alarmed, for the rate of the spontaneous process TOC Study Guide TOC Textbook Website MHHE Website 838 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chem Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry Metallic Hydrogen Scientists have long been interested in how nonmetal- trically conductive fluid in its interior. (For example, lic substances, including hydrogen, behave under ex- Earth’s magnetic field is due to the heat-driven motion ceedingly high pressure. It was predicted that when of liquid iron within its core.) Jupiter is composed of atoms or molecules are compressed, their bonding elec- an outer layer of nonmetallic molecular hydrogen that trons might be delocalized, producing a metallic state. continuously transforms hydrogen within the core to In 1996, physicists at the Lawrence Livermore metallic fluid hydrogen. It is now believed that this Laboratory used a 60-foot-long gun to generate a shock metallic layer is much closer to the surface (because compression onto a thin (0.5 mm) layer of liquid hy- the pressure needed to convert molecular hydrogen to drogen. For an instant, at pressures between 0.9 and metallic hydrogen is not as high as previously thought), 1.4 million atmospheres, they were able to measure which would account for Jupiter’s unusually strong the electrical conductivity of the hydrogen sample and magnetic field. found that it was comparable to that of cesium metal Insulating molecular hydrogen at 2000 K. (The temperature of the hydrogen sample rose as a result of compression, although it remained Metallic molecular hydrogen in the molecular form.) As the pressure fell rapidly, the metallic state of hydrogen disappeared. Metallic atomic hydrogen The Livermore experiment suggested that metallic Rock core hydrogen, if it can be kept in a stable state, may act as a room-temperature superconductor. The fact that hydrogen becomes metallic at pressures lower than previously thought possible also has provided new insight into planetary science. For many years scientists were puzzled by Jupiter’s strong magnetic field, which is 20 times greater than that of Earth. A planet’s magnetic field results from the convection motion of elec- Interior composition of Jupiter. C(diamond) 88n C(graphite) G° 2.87 kJ is extremely slow. Millions of years may pass before a diamond turns to graphite. Synthetic diamond can be prepared from graphite by applying very high pressures and temperatures. Figure 21.5 shows a synthetic diamond and its starting material, graphite. Synthetic diamonds generally lack the optical properties of natural diamonds. They are useful, however, as abrasives and in cutting concrete and many other hard substances, including metals and alloys. The uses of graphite are described on p. 439. Diamond P (atm) FIGURE 21.4 Phase diagram of carbon. Note that under atmospheric conditions, graphite is the stable form of carbon. Liquid 2× 104 Graphite Vapor 3300 t (°C) Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.3 CARBON 839 FIGURE 21.5 A synthetic diamond and the starting material, graphite. Carbon has the unique ability to form long chains (consisting of more than 50 C atoms) and stable rings with five or six members. This phenomenon is called catenation, the linking of like atoms. Carbon’s versatility is responsible for the millions of organic compounds (made up of carbon and hydrogen and other elements such as oxygen, nitrogen, and halogens) found on Earth. The chemistry of organic compounds is discussed in Chapter 24. Carbides and Cyanides Carbon combines with metals to form ionic compounds called carbides, such as CaC2 and Be2C, in which carbon is in the form of C2 or C4 ions. These ions are strong 2 Brønsted bases and react with water as follows: C2 (aq) 2 4 C (aq) 2H2O(l ) 88n 2OH (aq) C2H2(g) 4H2O(l ) 88n 4OH (aq) CH4(g) Carbon also forms a covalent compound with silicon. Silicon carbide, SiC, is called carborundum and is prepared as follows: SiO2(s) HCN is the gas used in gas execution chambers. Forward Main Menu 2CO(g) Carborundum is also formed by heating silicon with carbon at 1500°C. Carborundum is almost as hard as diamond and it has the diamond structure; each carbon atom is tetrahedrally bonded to four Si atoms, and vice versa. It is used mainly for cutting, grinding, and polishing metals and glasses. Another important class of carbon compounds, the cyanides, contain the anion group :C q N: . Cyanide ions are extremely toxic because they bind almost irreversibly to the Fe(III) ion in cytochrome oxidase, a key enzyme in metabolic processes. Hydrogen cyanide, which has the aroma of bitter almonds, is even more dangerous because of its volatility (b.p. 26°C). A few tenths of 1 percent by volume of HCN in air can cause death within minutes. Hydrogen cyanide can be prepared by treating sodium cyanide or potassium cyanide with acid: NaCN(s) Back 3C(s) 88n SiC(s) TOC HCl(aq) 88n NaCl(aq) Study Guide TOC HCN(aq) Textbook Website MHHE Website 840 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS Because HCN (in solution, called hydrocyanic acid) is a very weak acid (Ka 4.9 10 10), most of the HCN produced in this reaction is in the nonionized form and leaves the solution as hydrogen cyanide gas. For this reason acids should never be mixed with metal cyanides in the laboratory without proper ventilation. Cyanide ions are used to extract gold and silver. Although these metals are usually found in the uncombined state in nature, in other metal ores they may be present in relatively small concentrations and are more difficult to extract. In a typical process, the crushed ore is treated with an aqueous cyanide solution in the presence of air to dissolve the gold by forming the soluble complex ion [Au(CN)2] : 4Au(s) 8CN (aq) O2(g) 2H2O(l ) 88n 4[Au(CN)2] (aq) 4OH (aq) The complex ion [Au(CN)2] (along with some cation, such as Na ) is separated from other insoluble materials by filtration and treated with an electropositive metal such as zinc to recover the gold: Zn(s) 2[Au(CN)2] (aq) 88n [Zn(CN)4]2 (aq) 2Au(s) Figure 21.6 shows an aerial view of a “cyanide pond” used for the extraction of gold. Oxides of Carbon Of the several oxides of carbon, the most important are carbon monoxide, CO, and carbon dioxide, CO2. Carbon monoxide is a colorless, odorless gas formed by the incomplete combustion of carbon or carbon-containing compounds: 2C(s) The role of CO as an indoor air pollutant is discussed on p. 718. O2(g) 88n 2CO(g) Carbon monoxide is used in metallurgical process for extracting nickel (see p. 810), in organic synthesis, and in the production of hydrocarbon fuels with hydrogen. Industrially, it is prepared by passing steam over heated coke. Carbon monoxide burns readily in oxygen to form carbon dioxide: FIGURE 21.6 A cyanide pond for extracting gold from metal ore. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 2CO(g) Carbon dioxide is a major greenhouse gas (see p. 705). CARBON 841 H° 21.3 566 kJ O2(g) 88n 2CO2(g) Carbon monoxide is not an acidic oxide (it differs from carbon dioxide in that regard), and it is only slightly soluble in water. Carbon dioxide is a colorless and odorless gas. Unlike carbon monoxide, CO2 is nontoxic. It is an acidic oxide (see p. 629). Carbon dioxide is used in beverages, in fire extinguishers, and in the manufacture of baking soda, NaHCO3, and soda ash, Na2CO3. Solid carbon dioxide, called dry ice, is used as a refrigerant (see Figure 11.42). Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry in Action Chemistry Back Synthetic Gas from Coal The very existence of our technological society depends on an abundant supply of energy. Although the United States has only 5 percent of the world’s population, we consume about 20 percent of the world’s energy! At present, the two major sources of energy are fossil fuels and nuclear fission (discussed in Chapters 23 and 24, respectively). Coal, oil (which is also known as petroleum), and natural gas (mostly methane) are collectively called fossil fuels because they are the end result of the decomposition of plants and animals over tens or hundreds of millions of years. Forward Oil and natural gas are cleaner-burning and more efficient fuels than coal, so they are preferred for most purposes. However, supplies of oil and natural gas are being depleted at an alarming rate, and research is under way to make coal a more versatile source of energy. Coal consists of many high-molar-mass carbon compounds that also contain oxygen, hydrogen, and small amounts of nitrogen and sulfur. Coal constitutes about 90 percent of the world’s fossil fuel reserves. For centuries coal has been used as a fuel both in Underground coal mining. Main Menu TOC Study Guide TOC Textbook Website MHHE Website 842 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS Chemistry in Action Chemistry in Action Chemistry homes and in industry. However, underground coal mining is expensive and dangerous, and strip mining (that is, mining in an open pit after removal of the overlaying earth and rock) is tremendously harmful to the environment. Another problem, this one associated with the burning of coal, is the formation of sulfur dioxide (SO2) from the sulfur-containing compounds. This process leads to the formation of “acid rain,” discussed on p. 711. One of the most promising methods for making coal a more efficient and cleaner fuel involves the conversion of coal to a gaseous form, called syngas for “synthetic gas.” This process is called coal gasification. In the presence of very hot steam and air, coal decomposes and reacts according to the following simplified scheme: 21.4 C(s) H2O(g) 88n CO(g) C(s) 2H2(g) 88n CH4(g) H2(g) The main component of syngas is methane. In addition, the first reaction yields hydrogen and carbon monoxide gases and other useful by-products. Under suitable conditions, CO and H2 combine to form methanol: CO(g) 2H2(g) 88n CH3OH(l ) Methanol has many uses, for example, as a solvent and a starting material for plastics. Syngas is easier than coal to store and transport. What’s more, it is not a major source of air pollution because sulfur is removed in the gasification process. NITROGEN AND PHOSPHORUS NITROGEN The nitrogen cycle is discussed on p. 694. Molecular nitrogen will boil off before molecular oxygen does during the fractional distillation of liquid air. About 78 percent of air by volum is nitrogen. The most important mineral sources of nitrogen are saltpeter (KNO3) and Chile saltpeter (NaNO3). Nitrogen is an essential element of life; it is a component of proteins and nucleic acids. Molecular nitrogen is obtained by fractional distillation of air (the boiling points of liquid nitrogen and liquid oxygen are 196°C and 183°C, respectively). In the laboratory, very pure nitrogen gas can be prepared by the thermal decomposition of ammonium nitrite: NH4NO2(s) 88n 2H2O(g) N2(g) The N2 molecule contains a triple bond and is very stable with respect to dissociation into atomic species. However, nitrogen forms a large number of compounds with hydrogen and oxygen in which the oxidation number of nitrogen varies from 3 to 5 (Table 21.2). Most nitrogen compounds are covalent; however, when heated with certain metals, nitrogen forms ionic nitrides containing N3 ions: 6Li(s) N2(g) 88n 2Li3N(s) The nitride ion is a strong Brønsted base and reacts with water to produce ammonia and hydroxide ions: N3 (aq) 3H2O(l) 88n NH3(g) 3OH (aq) Ammonia Ammonia is one of the best-known nitrogen compounds. It is prepared industrially from nitrogen and hydrogen by the Haber process (see Section 13.6 and p. 585). It can be prepared in the laboratory by treating ammonium chloride with sodium hydroxide: NH4Cl(aq) Back Forward Main Menu TOC NaOH(aq) 88n NaCl(aq) Study Guide TOC H2O(l ) NH3(g) Textbook Website MHHE Website 21.4 TABLE 21.2 NITROGEN AND PHOSPHORUS 843 Common Compounds of Nitrogen OXIDATION NUMBER COMPOUND FORMULA STRUCTURE 3 Ammonia NH3 HOOOH N A H 2 Hydrazine N2H4 HOOOOOH NN AA HH 1 Hydroxylamine NH2OH HOOOOOH NQ O A H Nitrogen* (dinitrogen) N2 SNqNS 1 Nitrous oxide (dinitrogen monoxide) N2O O SNqNOOS Q 2 Nitric oxide (nitrogen monoxide) NO SPPO NQ O 3 Nitrous acid HNO2 4 Nitrogen dioxide NO2 OPOOOOH O QNO Q OPO SOONPO Q Q 5 Nitric acid HNO3 0 OPNOOOH O O Q Q A SOS Q *We list the element here as a reference. Ammonia is a colorless gas (b.p. 33.4°C) with an irritating odor. About threequarters of the ammonia produced annually in the United States (18 million tons in 1996) is used in fertilizers. Liquid ammonia, like water, undergoes autoionization: 2NH3(l ) 34 NH4 NH2 or simply NH3(l ) 34 H The amide ion is a strong Brønsted base and does not exist in water. NH2 where NH2 is called the amide ion. Note that both H and NH2 are solvated with the NH3 molecules. (Here is an example of ion-dipole interaction.) At 50°C, the ion product [H ][NH2 ] is about 1 10 33, considerably smaller than 1 10 14 for water at 25°C. Nevertheless, liquid ammonia is a suitable solvent for many electrolytes, especially when a more basic medium is required or if the solutes react with water. The ability of liquid ammonia to dissolve alkali metals was discussed in Section 20.5. Hydrazine Another important hydride of nitrogen is hydrazine: H H G D ON NOO D G H H Each N atom is sp3-hybridized. Hydrazine is a colorless liquid that smells like ammonia. It melts at 2°C and boils at 114°C. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 844 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS Hydrazine is a base that can be protonated to give the N2H5 and N2H2 ions. A 6 reducing agent, it can reduce Fe3 to Fe2 , MnO4 to Mn2 , and I2 to I . Its reaction with oxygen is highly exothermic: N2H4(l ) O2(g) 88n N2(g) 2H2O(l) H° 666.6 kJ Hydrazine and its derivative methylhydrazine, N2H3(CH3), together with the oxidizer dinitrogen tetroxide (N2O4), are used as rocket fuels. Hydrazine also plays a role in polymer synthesis and in the manufacture of pesticides. Oxides and Oxoacids of Nitrogen There are many nitrogen oxides, but the three particularly important ones are: nitrous oxide, nitric oxide, and nitrogen dioxide. Nitrous oxide, N2O, is a colorless gas with a pleasing odor and sweet taste. It is prepared by heating ammonium nitrate to about 270°C: NH4NO3(s) 88n N2O(g) 2H2O(g) Nitrous oxide resembles molecular oxygen in that it supports combustion. It does so because it decomposes when heated to form molecular nitrogen and molecular oxygen: 2N2O(g) 88n 2N2(g) O2(g) It is chiefly used as an anesthetic in dental procedures and other minor surgery. Nitrous oxide is also called “laughing gas” because a person inhaling the gas becomes somewhat giddy. No satisfactory explanation has yet been proposed for this unusual physiological response. Nitric oxide, NO, is a colorless gas. The reaction of N2 and O2 in the atmosphere N2(g) According to Le Chatelier ’s principle, the forward endothermic reaction is favored by heating. O2(g) 34 2NO(g) G° 173.4 kJ is a form of nitrogen fixation (see p. 694). The equilibrium constant for the above reaction is very small at room temperature: KP is only 4.0 10 31 at 25°C, so very little NO will form at that temperature. However, the equilibrium constant increases rapidly with temperature, for example, in a running auto engine. An appreciable amount of nitric oxide is formed in the atmosphere by the action of lightning. In the laboratory the gas can be prepared by the reduction of dilute nitric acid with copper: 3Cu(s) 8HNO3(aq) 88n 3Cu(NO3)2(aq) 4H2O(l ) 2NO(g) The nitric oxide molecule is paramagnetic, containing one unpaired electron. It can be represented by the following resonance structures: PPO NQ QO OPP NQ QO As we noted in Chapter 9, this molecule does not obey the octet rule. The properties of nitric oxide are discussed on p. 352. Unlike nitrous oxide and nitric oxide, nitrogen dioxide is a highly toxic yellowbrown gas with a choking odor. In the laboratory nitrogen dioxide is prepared by the action of concentrated nitric acid on copper (Figure 21.7): Cu(s) FIGURE 21.7 The production of NO2 gas when copper reacts with concentrated nitric acid. Back Forward 4HNO3(aq) 88n Cu(NO3)2(aq) 2H2O(l ) 2NO2(g) Nitrogen dioxide is paramagnetic. It has a strong tendency to dimerize to dinitrogen tetroxide, which is a diamagnetic molecule: Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.4 NITROGEN AND PHOSPHORUS 845 2NO2 34 N2O4 This reaction occurs in both the gas phase and the liquid phase. Nitrogen dioxide is an acidic oxide; it reacts rapidly with cold water to form both nitrous acid, HNO2, and nitric acid: Neither N2O nor NO reacts with water. On standing, a concentrated nitric acid solution turns slightly yellow as a result of NO2 formation. 2NO2(g) H2O(l ) 88n HNO2(aq) HNO3(aq) This is a disproportionation reaction (see p. 130) in which the oxidation number of nitrogen changes from 4 (in NO2) to 3 (in HNO2) and 5 (in HNO3). Note that this reaction is quite different from that between CO2 and H2O, in which only one acid (carbonic acid) is formed. Nitric acid is one of the most important inorganic acids. It is a liquid (b.p. 82.6°C), but it does not exist as a pure liquid because it decomposes spontaneously to some extent as follows: 4HNO3(l ) 88n 4NO2(g) 2H2O(l ) O2(g) The major industrial method of producing nitric acid is the Ostwald process, discussed in Section 13.6. The concentrated nitric acid used in the laboratory is 68 percent HNO3 by mass (density 1.42 g/cm3), which corresponds to 15.7 M. Nitric acid is a powerful oxidizing agent. The oxidation number of N in HNO3 is 5. The most common reduction products of nitric acid are NO2 (oxidation number of N 4), NO (oxidation number of N 2), and NH4 (oxidation number of N 3). Nitric acid can oxidize metals both below and above hydrogen in the activity series (see Figure 4.15). For example, copper is oxidized by concentrated nitric acid, as discussed earlier. In the presence of a strong reducing agent, such as zinc metal, nitric acid can be reduced all the way to the ammonium ion: 4Zn(s) 10H (aq) NO3 (aq) 88n 4Zn2 (aq) NH4 (aq) 3H2O(l ) 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 HNO3 to three parts HCl), the resulting solution, called aqua regia, can oxidize gold, as follows: Au(s) 3HNO3(aq) 4HCl(aq) 88n HAuCl4(aq) 3H2O(l ) 3NO2(g) The oxidation of Au is promoted by the complexing ability of the Cl ion (to form the AuCl4 ion). Concentrated nitric acid also oxidizes a number of nonmetals to their corresponding oxoacids: P4(s) S(s) 20HNO3(aq) 88n 4H3PO4(aq) 6HNO3(aq) 88n H2SO4(aq) 20NO2(g) 6NO2(g) 4H2O(l ) 2H2O(l ) Nitric acid is used in the manufacture of fertilizers, dyes, drugs, and explosives. The Chemistry in Action essay on p. 849 describes a nitrogen-containing fertilizer that can be highly explosive. PHOSPHORUS Like nitrogen, phosphorus is a member of the Group 5A family; in some respects the chemistry of phosphorus resembles that of nitrogen. Phosphorus occurs most com- Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 846 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS FIGURE 21.8 ing. Phosphate min- monly in nature as phosphate rocks, which are mostly calcium phosphate, Ca3(PO4)2, and fluoroapatite, Ca5(PO4)3F (Figure 21.8). Elemental phosphorus can be obtained by heating calcium phosphate with coke and silica sand: 2Ca3(PO4)2(s) 10C(s) 6SiO2(s) 88n 6CaSiO3(s) 10CO(g) P4(s) There are several allotropic forms of phosphorus, but only white phosphorus and red phosphorus (see Figure 8.18) are of importance. White phosphorus consists of discrete tetrahedral P4 molecules (Figure 21.9). A solid (m.p. 44.2°C), white phosphorus is insoluble in water but quite soluble in carbon disulfide (CS2) and in organic solvents such as chloroform (CHCl3). White phosphorus is a highly toxic substance. It bursts into flames spontaneously when exposed to air; hence it is used in incendiary bombs and grenades: P4(s) 5O2(g) 88n P4O10(s) The high reactivity of white phosphorus is attributed to structural strain: The POP bonds are compressed in the tetrahedral P4 molecule. White phosphorus was once used in matches, but because of its toxicity it has been replaced by tetraphosphorus trisulfide, P4S3. When heated in the absence of air, white phosphorus is slowly converted to red phosphorus at about 300°C: FIGURE 21.9 The structures of white and red phosphorus. Red phosphorus is believed to have a chain structure, as shown. White phosphorus Red phosphorus Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.4 NITROGEN AND PHOSPHORUS 847 nP4 (white phosphorus) 88n (P4)n (red phosphorus) Red phosphorus has a polymeric structure (see Figure 21.9) and is more stable and less volatile than white phosphorus. Hydride of Phosphorus The most important hydride of phosphorus is phosphine, PH3, a colorless, very poisonous gas formed by heating white phosphorus in concentrated sodium hydroxide: P4(s) 3NaOH(aq) 3H2O(l ) 88n 3NaH2PO2(aq) PH3(g) Phosphine is moderately soluble in water and more soluble in carbon disulfide and organic solvents. Its aqueous solution is neutral, unlike that of ammonia. In liquid ammonia, phosphine dissolves to give NH4 PH2 . Phosphine is a strong reducing agent; it reduces many metal salts to the corresponding metal. The gas burns in air: PH3(g) 2O2(g) 88n H3PO4(s) Halides of Phosphorus Phosphorus forms binary compounds with halogens: the trihalides, PX3, and the pentahalides, PX5, where X denotes a halogen atom. In contrast, nitrogen can form only trihalides (NX3). Unlike nitrogen, phosphorus has a 3d subshell, which can be used for valence-shell expansion. We can explain the bonding in PCl5 by assuming that phosphorus undergoes sp3d hybridization of its 3s, 3p, and 3d orbitals (see Example 10.4). The five sp3d hybrid orbitals also account for the trigonal bipyramidal geometry of the PCl5 molecule (see Table 10.4). Phosphorus trichloride is prepared by heating white phosphorus in chlorine: P4(l ) 6Cl2(g) 88n 4PCl3(g) A colorless liquid (b.p. 76°C), PCl3 is hydrolyzed according to the equation: PCl3(l ) 3H2O(l ) 88n H3PO3(aq) 3HCl(g) In the presence of an excess of chlorine gas, PCl3 is converted to phosphorus pentachloride, which is a light-yellow solid: PCl3(l ) Cl2(g) 88n PCl5(s) X-ray studies have shown that solid phosphorus pentachloride exists as [PCl4 ][PCl6 ], in which the PCl4 ion has a tetrahedral geometry and the PCl6 ion has an octahedral geometry. In the gas phase, PCl5 (which has trigonal bipyramidal geometry) is in equilibrium with PCl3 and Cl2: PCl5(g) 34 PCl3(g) Cl2(g) Phosphorus pentachloride reacts with water as follows: PCl5(s) 4H2O(l ) 88n H3PO4(aq) 5HCl(aq) Oxides and Oxoacids of Phosphorus The two important oxides of phosphorus are tetraphosphorus hexaoxide, P4O6, and tetraphosphorus decaoxide, P4O10 (Figure 21.10). The oxides are obtained by burning white phosphorus in limited and excess amounts of oxygen gas, respectively: Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 848 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS FIGURE 21.10 The structures of P4O6 and P4O10. Note the tetrahedral arrangement of the P atoms in P4O10. Phosphorus Oxygen P4O10 P4O6 P4(s) 3O2(g) 88n P4O6 (s) P4(s) 5O2(g) 88n P4O10(s) Both oxides are acidic; that is, they are converted to acids in water. The compound P4O10 is a white flocculent powder (m.p. 420°C) that has a great affinity for water: P4O10(s) 6H2O(l ) 88n 4H3PO4(aq) For this reason it is often used for drying gases and for removing water from solvents. There are many oxoacids containing phosphorus. Some examples are phosphorous acid, H3PO3; phosphoric acid, H3PO4; hypophosphorous acid, H3PO2; and triphosphoric acid, H5P3O10 (Figure 21.11). Phosphoric acid, also called orthophosphoric acid, is a weak triprotic acid (see p. 619). It is prepared industrially by the reaction of calcium phosphate with sulfuric acid: Ca3(PO4)2(s) 3H2SO4(aq) 88n 2H3PO4(aq) 3CaSO4(s) In the pure form phosphoric acid is a colorless solid (m.p. 42.2°C). The phosphoric acid we use in the laboratory is usually an 82 percent H3PO4 solution (by mass). Phosphoric acid and phosphates have many commercial applications in detergents, fertilizers, flame retardants, and toothpastes, and as buffers in carbonated beverages. Like nitrogen, phosphorus is an element that is essential to life. It constitutes only about 1 percent by mass of the human body, but it is a very important 1 percent. About 23 percent of the human skeleton is mineral matter. The phosphorus content of this mineral matter, calcium phosphate, Ca3(PO4)2, is 20 percent. Our teeth are basically Ca3(PO4)2 and Ca5(PO4)3OH. Phosphates are also important components of the genetic materials deoxyribonucleic acid (DNA) and ribonucleic acid (RNA). FIGURE 21.11 Structures of some common phosphoruscontaining oxoacids. S O¼ B þ O P O þOH HO³ O O ³ A H SO¼ B þO P OH HO O ³ A H Phosphorous acid (H 3PO 3 ) Hypophosphorous acid (H 3 PO 2 ) SO¼ B þO P OþOH HO O O ³ ³ A SO¼ A H Phosphoric acid (H 3 PO 4 ) Back Forward Main Menu TOC S O¼ S O¼ S O¼ B B B þ O P O þO P O O O P O þ OH þ O O HO³ O ³ ³ ³ A A A S O¼ S O¼ S O¼ A A A H H H Triphosphoric acid (H 5 P 3 O 10 ) Study Guide TOC Textbook Website MHHE Website 21.4 NITROGEN AND PHOSPHORUS 849 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 Ammonium Nitrate, The Explosive Fertilizer About 1.46 kilojoules of heat are generated per gram of the compound decomposed. When it is combined with a combustible material, such as fuel oil, the energy released increases almost threefold. Ammonium nitrate can also be mixed with charcoal, flour, sugar, sulfur, rosin, and paraffin to form an explosive. Intense heat from the explosion causes the gases to expand rapidly, generating shock waves that destroy most objects in their path. Federal law regulates the sale of explosive-grade ammonium nitrate, which is used for 95 percent of all commercial blasting in road construction and mining. However, the wide availability of large quantities of ammonium nitrate and other substances that enhance its explosive power make it possible for anyone who is so-inclined to construct a bomb. The bomb that destroyed the federal building in Oklahoma City is estimated to have contained 4000 pounds of ammonium nitrate and fuel oil, which was set off by another small explosive device. How can the use of ammonium nitrate by terrorists be prevented? The most logical approach is to desensitize or neutralize the compound’s ability to act as an explosive, but to date no satisfactory way has been found to do so without diminishing its value as a fertilizer. A more passive method is to add to the fertilizer an agent known as a taggant, which would allow law enforcement to trace the source of an ammonium nitrate explosive. A number of European countries now forbid the sale of ammonium nitrate without taggants, although the U.S. Congress has yet to pass such a law. The Alfred P. Murrah building after a deadly explosion caused by an ammonium nitrate bomb. A bag of ammonium nitrate fertilizer, which is labeled as an explosive. Ammonium nitrate is the most important fertilizer in the world (see p. 96). It ranked fifteenth among the industrial chemicals produced in the United States in 1995 (8 million tons). Unfortunately, it is also a powerful explosive. In 1947 an explosion occurred aboard a ship being loaded with the fertilizer in Texas. The fertilizer was in paper bags and apparently blew up after sailors tried to stop a fire in the ship’s hold by closing a hatch, thereby creating the compression and heat necessary for an explosion. More than six hundred people died as a result of the accident. More recent disasters involving ammonium nitrate took place at the World Trade Center in New York City in 1993 and at the Alfred P. Murrah Federal Building in Oklahoma City in 1995. A strong oxidizer, ammonium nitrate is stable at room temperature. At 250°C, it begins to decompose as follows: NH4NO3(g) 88n N2O(g) 2H2O(g) At 300°C, different gaseous products and more heat are produced: 2NH4NO3(g) 88n 2N2(g) Forward Main Menu 4H2O(g) TOC O2(g) Study Guide TOC Textbook Website MHHE Website 850 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS 21.5 OXYGEN AND SULFUR OXYGEN The oxygen cycle is discussed on p. 694. Oxygen is by far the most abundant element in Earth’s crust, constituting about 46 percent of its mass. In addition, the atmosphere contains about 21 percent molecular oxygen by volume (23 percent by mass). Like nitrogen, oxygen in the free state is a diatomic molecule (O2). In the laboratory, oxygen gas can be obtained by heating potassium chlorate (see Figure 5.12): 2KClO3(s) 88n 2KCl(s) 3O2(g) The reaction is usually catalyzed by manganese(IV) dioxide, MnO2. Pure oxygen gas can be prepared by electrolyzing water (p. 784). Industrially, oxygen gas is prepared by the fractional distillation of liquefied air (p. 842). Oxygen gas is colorless and odorless. Oxygen is a building block of practically all biomolecules, accounting for about a fourth of the atoms in living matter. Molecular oxygen is the essential oxidant in the metabolic breakdown of food molecules. Without it, a human being cannot survive for more than a few minutes. Properties of Diatomic Oxygen Although oxygen has two allotropes, O2 and O3, when we speak of molecular oxygen, we normally mean O2. Ozone, O3, is less stable than O2 (see p. 852). The O2 molecule is paramagnetic because it contains two unpaired electrons (see Section 10.7). A strong oxidizing agent, molecular oxygen is one of the most widely used industrial chemicals. Its main uses are in the steel industry (see Section 20.2) and in sewage treatment. Oxygen is also used as a bleaching agent for pulp and paper, in medicine to ease breathing difficulties, in oxyacetylene torches, and as an oxidizing agent in many inorganic and organic reactions. Oxides, Peroxides, and Superoxides Oxygen forms three types of oxides: the normal oxide (or simply the oxide), which contains the O2 ion; the peroxide, which contains the O2 ion; and the superoxide, 2 which contains the O2 ion: O SOS2 Q oxide OO SOSOS2 QQ peroxide OO SOSPS QQ superoxide The ions are all strong Brønsted bases and react with water as follows: Oxide: O2 (aq) Peroxide: O2 (aq) 2 Superoxide: 4O2 (aq) H2O(l ) 88n 2OH (aq) 2H2O(l ) 88n O2(g) 2H2O(l ) 88n 3O2(g) 4OH (aq) 4OH (aq) 2 Note that the reaction of O with water is a hydrolysis reaction, but those involving O2 and O2 are redox processes. 2 The nature of bonding in oxides changes across any period in the periodic table (see Figure 15.8). Oxides of elements on the left side of the periodic table, such as those of the alkali metals and alkaline earth metals, are generally ionic solids with high Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.5 OXYGEN AND SULFUR 851 melting points. Oxides of the metalloids and of the metallic elements toward the middle of the periodic table are also solids, but they have much less ionic character. Oxides of nonmetals are covalent compounds that generally exist as liquids or gases at room temperature. The acidic character of the oxides increases from left to right. Consider the oxides of the third-period elements: Na2O MgO Al2O3 SiO2 P4O10 SO3 Cl2O7 basic amphoteric acidic The basicity of the oxides increases as we move down a particular group. MgO does not react with water but reacts with acid as follows: MgO(s) 2H (aq) 88n Mg2 (aq) H2O(l ) On the other hand, BaO, which is more basic, undergoes hydrolysis to yield the corresponding hydroxide: BaO(s) H2O(l ) 88n Ba(OH)2(aq) The best-known peroxide is hydrogen peroxide (H2O2). It is a colorless, syrupy liquid (m.p. 0.9°C), prepared in the laboratory by the action of cold dilute sulfuric acid on barium peroxide octahydrate: BaO2 8H2O(s) 97° 86° H2SO4(aq) 88n BaSO4(s) The structure of 8H2O(l ) The structure of hydrogen peroxide is shown in Figure 21.12. Using the VSEPR method we see that the HOO and OOO bonds are bent around each oxygen atom in a configuration similar to the structure of water. The lone-pair – bonding-pair repulsion is greater in H2O2 than in H2O, so that the HOO angle is only 97° (compared with 104.5° for HOH in H2O). Hydrogen peroxide is a polar molecule ( 2.16 D). Hydrogen peroxide readily decomposes when heated or exposed to sunlight or even in the presence of dust particles or certain metals, including iron and copper: 2H2O2(l ) 88n 2H2O(l ) FIGURE 21.12 H2O2. H2O2(aq) O2(g) H° 196.4 kJ Note that this is a disproportionation reaction. The oxidation number of oxygen changes from 1 to 2 and 0. Hydrogen peroxide is miscible with water in all proportions due to its ability to hydrogen-bond with water. Dilute hydrogen peroxide solutions (3 percent by mass), available in drugstores, are used as mild antiseptics; more concentrated H2O2 solutions are employed as bleaching agents for textiles, fur, and hair. The high heat of decomposition of hydrogen peroxide also makes it a suitable component in rocket fuel. Hydrogen peroxide is a strong oxidizing agent; it can oxidize Fe2 ions to Fe3 ions in an acidic solution: H2O2(aq) 2Fe2 (aq) 2H (aq) 88n 2Fe3 (aq) 2H2O(l ) It also oxidizes SO2 ions to SO2 ions: 3 4 H2O2(aq) SO2 (aq) 88n SO2 (aq) 3 4 H2O(l ) In addition, hydrogen peroxide can act as a reducing agent toward substances that are stronger oxidizing agents than itself. For example, hydrogen peroxide reduces silver oxide to metallic silver: H2O2(aq) Back Forward Main Menu TOC Ag2O(s) 88n 2Ag(s) Study Guide TOC H2O(l ) O2(g) Textbook Website MHHE Website 852 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS and permanganate, MnO4 , to manganese(II) in an acidic solution: 5H2O2(aq) 2MnO4 (aq) 6H (aq) 88n 2Mn2 (aq) 5O2(g) 8H2O(l ) If we want to determine hydrogen peroxide concentration, this reaction can be carried out as a redox titration, using a standard permanganate solution. There are relatively few known superoxides, that is, compounds containing the O2 ion. In general, only the most reactive alkali metals (K, Rb, and Cs) form superoxides. We should take note of the fact that both the peroxide ion and the superoxide ion are by-products of metabolism. Because these ions are highly reactive, they can inflict great damage on living cells. Fortunately, our bodies are equipped with the enzymes catalase, peroxidase, and superoxide dismutase which convert these toxic substances to water and molecular oxygen. Ozone Ozone is a rather toxic, light-blue gas (b.p. 111.3°C). Its pungent odor is noticeable around sources of significant electrical discharges (such as a subway train). Ozone can be prepared from molecular oxygen, either photochemically or by subjecting O2 to an electrical discharge (Figure 21.13): 3O2(g) 88n 2O3(g) Liquid ozone. G° 326.8 kJ Since the standard free energy of formation of ozone is a large positive quantity ( Go (326.8/2) kJ/mol or 163.4 kJ/mol), ozone is less stable than molecular oxygen. f The ozone molecule has a bent structure in which the bond angle is 116.5°: S S S S S S S S S O O JG O O S O O DM O O Ozone is mainly used to purify drinking water, to deodorize air and sewage gases, and to bleach waxes, oils, and textiles. Ozone is a very powerful oxidizing agent — its oxidizing power is exceeded only by that of molecular fluorine (see Table 19.1). For example, ozone can oxidize sulfides of many metals to the corresponding sulfates: FIGURE 21.13 The preparation of O3 from O2 by electrical discharge. The outside of the outer tube and the inside of the inner tube are coated with metal foils that are connected to a highvoltage source. (The metal foil on the inside of the inner tube is not shown.) During the electrical discharge, O2 gas is passed through the tube. The O3 gas formed exits from the upper righthand tube, along with some unreacted O2 gas. Outer tube Metal foil on outer tube O3 plus some unreacted O2 O2 Inner tube High voltage source Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.5 4O3(g) PbS(s) 88n PbSO4(s) OXYGEN AND SULFUR 853 4O2(g) Ozone oxidizes all the common metals except gold and platinum. In fact, a convenient test for ozone is based on its action on mercury. When exposed to ozone, mercury loses its metallic luster and sticks to glass tubing (instead of flowing freely through it). This behavior is attributed to the change in surface tension caused by the formation of mercury(II) oxide: O3(g) 3Hg(l ) 88n 3HgO(s) The beneficial effect of ozone in the stratosphere and its undesirable action in smog formation were discussed in Chapter 17. SULFUR FIGURE 21.14 Pyrite (FeS2), commonly called “fool’s gold” because of its gold luster. Although sulfur is not a very abundant element (it constitutes only about 0.06 percent of Earth’s crust by mass), it is readily available because it occurs commonly in nature in the elemental form. The largest known reserves of sulfur are found in sedimentary deposits. In addition, sulfur occurs widely in gypsum (CaSO4 2H2O) and various sulfide minerals such as pyrite (FeS2) (Figure 21.14). Sulfur is also present in natural gas as H2S, SO2, and other sulfur-containing compounds. Sulfur is extracted from underground deposits by the Frasch† process, shown in Figure 21.15. In this process, superheated water (liquid water heated to about 160°C under high pressure to prevent it from boiling) is pumped down the outermost pipe to melt the sulfur (Figure 21.16). Next, compressed air is forced down the innermost pipe. † Herman Frasch (1851–1914). German chemical engineer. Besides inventing the process for obtaining pure sulfur, Frasch developed methods for refining petroleum. FIGURE 21.15 The Frasch process. Three concentric pipes are inserted into a hole drilled down to the sulfur deposit. Superheated water is forced down the outer pipe into the sulfur, causing it to melt. Molten sulfur is then forced up the middle pipe by compressed air. Compressed air Sulfur Superheated water Molten sulfur Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 854 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS FIGURE 21.16 Steam injection in the Frasch process. S S S S S S S S EH S S S S S S S S SH ESH ES S S S S S8 S Liquid sulfur mixed with air forms an emulsion that is less dense than water and therefore rises to the surface as it is forced up the middle pipe. Sulfur produced in this manner, which amounts to about 10 million tons per year, has a purity of about 99.5 percent. There are several allotropic forms of sulfur, the most important being the rhombic and monoclinic forms. Rhombic sulfur is thermodynamically the most stable form; it has a puckered S8 ring structure: It is a yellow, tasteless, and odorless solid (m.p. 112°C) (see Figure 8.19) that is insoluble in water but soluble in carbon disulfide. When heated, it is slowly converted to monoclinic sulfur (m.p. 119°C), which also consists of the S8 units. When liquid sulfur is heated above 150°C, the rings begin to break up, and the entangling of the sulfur chains results in a sharp increase in the liquid’s viscosity. Further heating tends to rupture the chains, and the viscosity decreases. Like nitrogen, sulfur shows a wide variety of oxidation numbers in its compounds (Table 21.3). The best-known hydrogen compound of sulfur is hydrogen sulfide, which is prepared by the action of an acid on a sulfide; for example, FeS(s) H2SO4(aq) 88n FeSO4(aq) H2S(g) Nowadays, hydrogen sulfide used in qualitative analysis (see Section 16.10) is prepared by the hydrolysis of thioacetamide: J H3CO C G S 2H2O H NH2 thioacetamide O J H3COC G O OH H2S NH4 acetic acid Hydrogen sulfide is a colorless gas (b.p. 60.2°C) that smells like rotten eggs. (The odor of rotten eggs actually does come from hydrogen sulfide, which is formed by the bacterial decomposition of sulfur-containing proteins.) Hydrogen sulfide is a highly toxic substance that, like hydrogen cyanide, attacks respiratory enzymes. It is a very Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.5 Common Compounds of Sulfur S S8 S H2S Sulfur* S DG H H ESH E S S S SS S H S S S SS S 0 STRUCTURE S Hydrogen sulfide 2 FORMULA S COMPOUND S OXIDATION NUMBER S TABLE 21.3 855 OXYGEN AND SULFUR S Sulfur dichloride S 2 SCl2 S SO3 S S S S Sulfur trioxide Cl O S JG O O SO2 6 S DG S Sulfur dioxide Cl S S S 4 S O Q E Cl S OO S OS QQ E S O S Cl Q S S S2Cl2 S Disulfur dichloride S S 1 S EH S S S S S S S S OS B S DG O O *We list the element here as a reference. weak diprotic acid (see Table 15.5). In basic solution, H2S is a reducing agent. For example, it is oxidized by permanganate to elemental sulfur: 3H2S(aq) 2MnO4 (aq) 88n 3S(s) 2MnO2(s) 2H2O(l ) 2OH (aq) Oxides of Sulfur Sulfur has two important oxides: sulfur dioxide, SO2; and sulfur trioxide, SO3. Sulfur dioxide is formed when sulfur burns in air: S(s) O2(g) 88n SO2(g) In the laboratory, it can be prepared by the action of an acid on a sulfite; for example, 2HCl(aq) Na2SO3(aq) 88n 2NaCl(aq) H2O(l ) SO2(g) or by the action of concentrated sulfuric acid on copper: Cu(s) There is no evidence for the formation of sulfurous acid, H2SO3, in water. Back Forward Main Menu 2H2SO4(aq) 88n CuSO4(aq) 2H2O(l ) SO2(g) Sulfur dioxide (b.p. 10°C) is a pungent, colorless gas that is quite toxic. An acidic oxide, it reacts with water as follows: SO2(g) H2O(l ) 34 H (aq) HSO3 (aq) Sulfur dioxide is slowly oxidized to sulfur trioxide, but the reaction rate can be greatly enhanced by a platinum or vanadium oxide catalyst (see Section 13.6): TOC Study Guide TOC Textbook Website MHHE Website 856 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS 2SO2(g) O2(g) 88n 2SO3(g) Sulfur trioxide dissolves in water to form sulfuric acid: SO3(g) H2O(l ) 88n H2SO4(aq) The contributing role of sulfur dioxide to acid rain is discussed on p. 711. Sulfuric Acid Sulfuric acid is the world’s most important industrial chemical. It is prepared industrially by first burning sulfur in air: S(s) O2(g) 88n SO2(g) Next is the key step of converting sulfur dioxide to sulfur trioxide: 2SO2(g) O2(g) 88n 2SO3(g) Vanadium(V) oxide (V2O5) is the catalyst used for the second step. Because the sulfur dioxide and oxygen molecules react in contact with the surface of solid V2O5, the process is referred to as the contact process. Sulfuric acid is a diprotic acid (see Table 15.5). It is a colorless, viscous liquid (m.p. 10.4°C). The concentrated sulfuric acid we use in the laboratory is 98 percent H2SO4 by mass (density: 1.84 g/cm3), which corresponds to a concentration of 18 M. The oxidizing strength of sulfuric acid depends on its temperature and concentration. A cold, dilute sulfuric acid solution reacts with metals above hydrogen in the activity series (see Figure 4.15), thereby liberating molecular hydrogen in a displacement reaction: Mg(s) Vanadium oxide on alumina (Al2O3). H2SO4(aq) 88n MgSO4(aq) H2(g) This is a typical reaction of an active metal with an acid. The strength of sulfuric acid as an oxidizing agent is greatly enhanced when it is both hot and concentrated. In such a solution, the oxidizing agent is actually the sulfate ion rather than the hydrated proton, H (aq). Thus, copper reacts with concentrated sulfuric acid as follows: Cu(s) 2H2SO4(aq) 88n CuSO4(aq) SO2(g) 2H2O(l ) Depending on the nature of the reducing agents, the sulfate ion may be further reduced to elemental sulfur or the sulfide ion. For example, reduction of H2SO4 by HI yields H2S and I2: 8HI(aq) H2SO4(aq) 88n H2S(aq) 4I2(s) 4H2O(l ) Concentrated sulfuric acid oxidizes nonmetals. For example, it oxidizes carbon to carbon dioxide and sulfur to sulfur dioxide: C(s) 2H2SO4(aq) 88n CO2(g) S(s) 2H2SO4(aq) 88n 3SO2(g) 2SO2(g) 2H2O(l ) 2H2O(l ) Other Compounds of Sulfur Carbon disulfide, a colorless, flammable liquid (b.p. 46°C), is formed by heating carbon and sulfur to a high temperature: Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.6 C(s) THE HALOGENS 857 2S(l ) 88n CS2(l ) It is only slightly soluble in water. Carbon disulfide is a good solvent for sulfur, phosphorus, iodine, and nonpolar substances such as waxes and rubber. Another interesting compound of sulfur is sulfur hexafluoride, SF6, which is prepared by heating sulfur in an atmosphere of fluorine: S(l ) 3F2(g) 88n SF6(g) Sulfur hexafluoride is a nontoxic, colorless gas (b.p. 63.8°C). It is the most inert of all sulfur compounds; it resists attack even by molten KOH. The structure and bonding of SF6 were discussed in Chapters 9 and 10 and its critical phenomenon illustrated in Chapter 11 (see Figure 11.37). 21.6 Recall that the first member of a group usually differs in properties from the rest of the members of the group (see p. 308). THE HALOGENS The halogens — fluorine, chlorine, bromine, and iodine — are reactive nonmetals (see Figure 8.20). Table 21.4 lists some of the properties of these elements. Although all halogens are highly reactive and toxic, the magnitude of reactivity and toxicity generally decreases from fluorine to iodine. The chemistry of fluorine differs from that of the rest of the halogens in the following ways: • Fluorine is the most reactive of all the halogens. The difference in reactivity between fluorine and chlorine is greater than that between chlorine and bromine. Table 21.4 shows that the FOF bond is considerably weaker than the ClOCl bond. The weak bond in F2 can be explained in terms of the lone pairs on the F atoms: OF SQOOS FQ The small size of the F atoms (see Table 21.4) allows a close approach of the three lone pairs on each of the F atoms, resulting in a greater repulsion than that found in Cl2, which consists of larger atoms. TABLE 21.4 Properties of the Halogens PROPERTY F Valence electron configuration Melting point (°C)* Boiling point (°C)* Appearance* Cl 2 5 2s 2p 223 187 Paleyellow gas Atomic radius (pm) 72 Ionic radius (pm)† 136 Ionization energy (kJ/mol) 1680 Electronegativity 4.0 Standard reduction potential (V)* 2.87 Bond energy (kJ/mol)* 150.6 Br 2 5 3s 3p 102 35 Yellowgreen gas 99 181 1251 3.0 1.36 242.7 2 I 5 4s 4p 7 59 Redbrown liquid 114 195 1139 2.8 1.07 192.5 5s25p5 114 183 Dark-violet vapor Dark metalliclooking solid 133 216 1003 2.5 0.53 151.0 *These values and descriptions apply to the diatomic species X2, where X represents a halogen atom. The half-reaction is X2(g) 2e 88n 2X (aq). † Refers to the anion X . Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 858 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS Hydrogen fluoride, HF, has a high boiling point (19.5°C) as a result of strong intermolecular hydrogen bonding, whereas all other hydrogen halides have much lower boiling points (see Figure 11.6). • Hydrofluoric acid is a weak acid, whereas all other hydrohalic acids (HCl, HBr, and HI) are strong acids. • Fluorine reacts with cold sodium hydroxide solution to produce oxygen difluoride as follows: • 2F2(g) 2NaOH(aq) 88n 2NaF(aq) H2O(l ) OF2(g) The same reaction with chlorine or bromine, on the other hand, produces a halide and a hypohalite: X2(g) • 2NaOH(aq) 88n NaX(aq) NaXO(aq) H2O(l ) where X stands for Cl or Br. Iodine does not react under the same conditions. Silver fluoride, AgF, is soluble. All other silver halides (AgCl, AgBr, and AgI) are insoluble (see Table 4.2). The element astatine also belongs to the Group 7A family. However, all isotopes of astatine are radioactive; its longest-lived isotope is astatine-210, which has a halflife of 8.3 hours. Therefore it is both difficult and expensive to study astatine in the laboratory. The halogens form a very large number of compounds. In the elemental state they form diatomic molecules, X2. In nature, however, because of their high reactivity, halogens are always found combined with other elements. Chlorine, bromine, and iodine occur as halides in seawater, and fluorine occurs in the minerals fluorite (CaF2) (see Figure 20.18) and cryolite (Na3AlF6). PREPARATION AND GENERAL PROPERTIES OF THE HALOGENS Because fluorine and chlorine are strong oxidizing agents, they must be prepared by electrolysis rather than by chemical oxidation of the fluoride and chloride ions. Electrolysis does not work for aqueous solutions of fluorides, however, because fluorine is a stronger oxidizing agent than oxygen. From Table 19.1 we find that O2(g) 4H (aq) 4e 88n 2H2O(l ) E° 1.23 V F2(g) 2e 88n 2F (aq) E° 2.87 V If F2 were formed by the electrolysis of an aqueous fluoride solution, it would immediately oxidize water to oxygen. For this reason, fluorine is prepared by electrolyzing liquid hydrogen fluoride containing potassium fluoride to increase its conductivity, at about 70°C (Figure 21.17): Anode (oxidation): 2F 88n F2(g) Cathode (reduction): 2H Overall reaction: 2e 2e 88n H2(g) 2HF(l ) 88n H2(g) F2(g) Chlorine gas, Cl2, is prepared industrially by the electrolysis of molten NaCl (see Section 19.8) or by the chlor-alkali process, the electrolysis of a concentrated aqueous NaCl solution (called brine). (Chlor denotes chlorine and alkali denotes an alkali metal, such as sodium.) Two of the common cells used in the chlor-alkali process are the mercury cell and the diaphragm cell. In both cells the overall reaction is Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.6 FIGURE 21.17 Electrolytic cell for the preparation of fluorine gas. Note that because H2 and F2 form an explosive mixture, these gases must be separated from each other. 859 THE HALOGENS F2 gas Carbon anode H2 gas H2 gas Diaphragm to prevent mixing of H2 and F2 gases Steel cathode Liquid HF electrolysis 2NaCl(aq) 2H2O(l ) 8877777777n 2NaOH(aq) H2(g) Cl2(g) As you can see, this reaction yields two useful by-products, NaOH and H2. The cells are designed to separate the molecular chlorine from the sodium hydroxide solution and the molecular hydrogen to prevent side reactions such as 2NaOH(aq) H2(g) Cl2(g) 88n NaOCl(aq) NaCl(aq) H2O(l ) Cl2(g) 88n 2HCl(g) These reactions consume the desired products and can be dangerous because a mixture of H2 and Cl2 is explosive. Figure 21.18 shows the mercury cell used in the chlor-alkali process. The cathode is a liquid mercury pool at the bottom of the cell, and the anode is made of either graphite or titanium coated with platinum. Brine is continuously passed through the cell as shown in the diagram. The electrode reactions are Anode (oxidation): 2Cl (aq) 88n Cl2(g) Cathode (reduction): 2Na (aq) Overall reaction: 2e Hg(l ) 2e 88n 2Na/Hg 2NaCl(aq) 88n 2Na/Hg Cl2(g) where Na/Hg denotes the formation of sodium amalgam. The chlorine gas generated this way is very pure. The sodium amalgam does not react with the brine solution but decomposes as follows when treated with pure water outside the cell: FIGURE 21.18 Mercury cell used in the chlor-alkali process. The cathode contains mercury. The sodium-mercury amalgam is treated with water outside the cell to produce sodium hydroxide and hydrogen gas. Cl2 Graphite anode Brine Brine Hg cathode Hg plus Na/Hg Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 860 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS 2Na/Hg From Table 19.1 we see that the oxidizing strength decreases from Cl2 to Br2 to I2. 2H2O(l ) 88n 2NaOH(aq) H2(g) 2Hg(l ) the by-products are sodium hydroxide and hydrogen gas. Although the mercury is cycled back into the cell for reuse, some of it is always discharged with waste solutions into the environment, resulting in mercury pollution. This is a major drawback of the mercury cell. Figure 21.19 shows the industrial manufacture of chlorine gas. The half-cell reactions in a diaphragm cell are shown in Figure 21.20. The asbestos diaphragm is permeable to the ions but not to the hydrogen and chlorine gases and so prevents the gases from mixing. During electrolysis a positive pressure is applied on the anode side of the compartment to prevent the migration of the OH ions from the cathode compartment. Periodically fresh brine solution is added to the cell and the sodium hydroxide solution is run off as shown. The diaphragm cell presents no pollution problems. Its main disadvantage is that the sodium hydroxide solution is contaminated with unreacted sodium chloride. The preparation of molecular bromine and iodine from seawater by oxidation with chlorine was discussed in Section 4.4. In the laboratory, chlorine, bromine, and iodine can be prepared by heating the alkali halides (NaCl, KBr, or KI) in concentrated sulfuric acid in the presence of manganese(IV) oxide. A representative reaction is MnO2(s) 2H2SO4(aq) 2NaCl(aq) 88n MnSO4(aq) Na2SO4(aq) 2H2O(l ) Cl2(g) COMPOUNDS OF THE HALOGENS Most of the halides can be classified into two categories. The fluorides and chlorides of many metallic elements, especially those belonging to the alkali metal and alkaline earth metal (except beryllium) families, are ionic compounds. Most of the halides of nonmetals such as sulfur and phosphorus are covalent compounds. As Figure 4.10 shows, the oxidation numbers of the halogens can vary from 1 to 7. The only exception is fluorine. Because it is the most electronegative element, fluorine can have only two oxidation numbers, 0 (as in F2) and 1, in its compounds. The Hydrogen Halides The hydrogen halides, an important class of halogen compounds, can be formed by the direct combination of the elements: FIGURE 21.19 The industrial manufacture of chlorine gas. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.6 FIGURE 21.20 Diaphragm cell used in the chlor-alkali process. THE HALOGENS 861 Battery e– e– Asbestos diaphragm Anode Cathode Brine NaOH solution Oxidation 2Cl–(aq) Cl2( g) + 2e – Reduction 2H2O(l) + 2e – H2( g) + 2OH–(aq) H2(g) X2(g) 34 2HX(g) where X denotes a halogen atom. These reactions (especially the ones involving F2 and Cl2) can occur with explosive violence. Industrially, hydrogen chloride is produced as a by-product in the manufacture of chlorinated hydrocarbons: C2H6(g) Cl2(g) 88n C2H5Cl(g) HCl(g) In the laboratory, hydrogen fluoride and hydrogen chloride can be prepared by reacting the metal halides with concentrated sulfuric acid: CaF2(s) H2SO4(aq) 88n 2HF(g) 2NaCl(s) H2SO4(aq) 88n 2HCl(g) CaSO4(s) Na2SO4(aq) Hydrogen bromide and hydrogen iodide cannot be prepared this way because they are oxidized to elemental bromine and iodine. For example, the reaction between NaBr and H2SO4 is 2NaBr(s) 2H2SO4(aq) 88n Br2(l ) SO2(g) Na2SO4(aq) 2H2O(l ) Instead, hydrogen bromide is prepared by first reacting bromine with phosphorus to form phosphorus tribromide: P4(s) 6Br2(l ) 88n 4PBr3(l ) Next, PBr3 is treated with water to yield HBr: PBr3(l ) 3H2O(l ) 88n 3HBr(g) H3PO3(aq) Hydrogen iodide can be prepared in a similar manner. The high reactivity of HF is demonstrated by the fact that it attacks silica and silicates: 6HF(aq) SiO2(g) 88n H2SiF6(aq) 2H2O(l ) This property makes HF suitable for etching glass and is the reason that hydrogen fluoride must be kept in plastic or inert metal (for example, Pt) containers. Hydrogen fluoride is used in the manufacture of Freons (see Chapter 17); for example, CCl4(l ) CFCl3(g) Back Forward Main Menu TOC HF(g) 88n CFCl3(g) HF(g) 88n CF2Cl2(g) Study Guide TOC HCl(g) HCl(g) Textbook Website MHHE Website 862 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS It is also important in the production of aluminum (see Section 20.7). Hydrogen chloride is used in the preparation of hydrochloric acid, inorganic chlorides, and in various metallurgical processes. Hydrogen bromide and hydrogen iodide do not have any major industrial uses. Aqueous solutions of hydrogen halides are acidic. The strength of the acids increases as follows: HF HCl HBr HI Oxoacids of the Halogens The halogens also form a series of oxoacids with the following general formulas: HXO HXO2 HXO3 HXO4 hypohalous acid halous acid halic acid perhalic acid Chlorous acid, HClO2, is the only known halous acid. All the halogens except fluorine form halic and perhalic acids. The Lewis structures of the chlorine oxoacids are OO HSQS ClS OQ OOO HSOS ClSQS QQO OOO OQO HSQS ClSQS O SQS hypochlorous acid chlorous acid chloric acid O SOS OOO HSOS ClSQS QQO O SQS perchloric acid For a given halogen, the acid strength decreases from perhalic acid to hypohalous acid; the explanation of this trend is discussed in Section 15.9. Table 21.5 lists some of the halogen compounds. Periodic acid, HIO4, does not appear because this compound cannot be isolated in the pure form. Instead the formula H5IO6 is often used to represent periodic acid. USES OF THE HALOGENS Fluorine Applications of the halogens and their compounds are widespread in industry, health care, and other areas. One is fluoridation, the practice of adding small quantities of fluorides (about 1 ppm by mass) such as NaF to drinking water to reduce dental caries. One of the most important inorganic fluorides is uranium hexafluoride, UF6, which is essential to the gaseous diffusion process for separating isotopes of uranium (U-235 TABLE 21.5 Common Compounds of Halogens* COMPOUND F Cl Br I Hydrogen halide Oxides HF ( 1) OF2 ( 1) HBr ( 1) Br2O ( 1) BrO2 ( 4) HI ( 1) I2O5 ( 5) Oxoacids HFO ( 1) HCl ( 1) Cl2O ( 1) ClO2 ( 4) Cl2O7 ( 7) HClO ( 1) HClO2 ( 3) HClO3 ( 5) HClO4 ( 7) HBrO ( 1) HIO ( 1) HBrO3 ( 5) HBrO4 ( 7) HIO3 ( 5) H5IO6 ( 7) *The number in parentheses indicates the oxidation number of the halogen. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 21.6 THE HALOGENS 863 and U-238). Industrially, fluorine is used to produce polytetrafluoroethylene, a polymer better known as Teflon: O CF2 O CF2 O ( )n where n is a large number. Teflon is used in electrical insulators, high-temperature plastics, cooking utensils, and so on. Chlorine Chlorine plays an important biological role in the human body, where the chloride ion is the principal anion in intracellular and extracellular fluids. Chlorine is widely used as an industrial bleaching agent for paper and textiles. Ordinary household laundry bleach contains the active ingredient sodium hypochlorite (about 5 percent by mass), which is prepared by reacting chlorine gas with a cold solution of sodium hydroxide: Cl2(g) 2NaOH(aq) 88n NaCl(aq) NaClO(aq) H2O(l ) Chlorine is also used to purify water and disinfect swimming pools. When chlorine dissolves in water, it undergoes the following reaction: Cl2(g) H2O(l ) 88n HCl(aq) HClO(aq) It is thought that the ClO ions destroy bacteria by oxidizing life-sustaining compounds within them. Chlorinated methanes, such as carbon tetrachloride and chloroform, are useful organic solvents. Large quantities of chlorine are used to produce insecticides, such as DDT. However, in view of the damage they inflict on the environment, the use of many of these compounds is either totally banned or greatly restricted in the United States. Chlorine is also used to produce polymers such as poly(vinyl chloride). Bromine So far as we know, bromine compounds occur naturally only in some marine organisms. Seawater is about 1 10 3 M Br ; therefore, it is the main source of bromine. Bromine is used to prepare ethylene dibromide (BrCH2CH2Br), which is used as an insecticide and as a scavenger for lead (that is, to combine with lead) in gasoline to keep lead deposits from clogging engines. Recent studies have shown that ethylene dibromide is a very potent carcinogen. Bromine combines directly with silver to form silver bromide (AgBr), which is used in photographic films. Iodine Iodine is not used as widely as the other halogens. A 50 percent (by mass) alcohol solution of iodine, known as tincture of iodine, is used medicinally as an antiseptic. Iodine is an essential constituent of the thyroid hormone thyroxine: I HO O I Back Forward Main Menu TOC I I Study Guide TOC H O A J CH2 O COC G A NH2 OH Textbook Website MHHE Website 864 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS FIGURE 21.21 Cloud seeding using AgI particles. Iodine deficiency in the diet may result in enlargement of the thyroid gland (known as goiter). Iodized table salt sold in the United States usually contains 0.01 percent KI or NaI, which is more than sufficient to satisfy the 1 mg of iodine per week required for the formation of thyroxine in the human body. A compound of iodine that deserves mention is silver iodide, AgI. It is a paleyellow solid that darkens when exposed to light. In this respect it is similar to silver bromide. Silver iodide is sometimes used in cloud seeding, a process for inducing rainfall on a small scale (Figure 21.21). The advantage of using silver iodide is that enormous numbers of nuclei (that is, small particles of silver iodide on which ice crystals can form) become available. About 1015 nuclei are produced from 1 g of AgI by vaporizing an acetone solution of silver iodide in a hot flame. The nuclei are then dispersed into the clouds from an airplane. SUMMARY OF FACTS AND CONCEPTS Back Forward 1. Hydrogen atoms contain one proton and one electron. They are the simplest atoms. Hydrogen combines with many metals and nonmetals to form hydrides; some hydrides are ionic and some are covalent. 2. There are three isotopes of hydrogen: 1H, 2H (deuterium), and 3H (tritium). Heavy water 1 1 1 contains deuterium. 3. The important inorganic compounds of carbon are the carbides; the cyanides, most of which are extremely toxic; carbon monoxide, also toxic and a major air pollutant; the carbonates and bicarbonates; and carbon dioxide, an end product of metabolism and a component of the global carbon cycle. 4. Elemental nitrogen, N2, contains a triple bond and is very stable. Compounds in which nitrogen has oxidation numbers from 3 to 5 are formed between nitrogen and hydrogen and/or oxygen atoms. Ammonia, NH3, is widely used in fertilizers. 5. White phosphorus, P4, is highly toxic, very reactive, and flammable; the polymeric red phosphorus, (P4)n, is more stable. Phosphorus forms oxides and halides with oxidation numbers of 3 and 5 and several oxoacids. The phosphates are the most important phosphorus compounds. Main Menu TOC Study Guide TOC Textbook Website MHHE Website QUESTIONS AND PROBLEMS 865 6. Elemental oxygen, O2, is paramagnetic and contains two unpaired electrons. Oxygen forms ozone (O3), oxides (O2 ), peroxides (O2 ), and superoxides (O2 ). The most abun2 dant element in Earth’s crust, oxygen is essential for life on Earth. 7. Sulfur is taken from Earth’s crust by the Frasch process as a molten liquid. Sulfur exists in a number of allotropic forms and has a variety of oxidation numbers in its compounds. 8. Sulfuric acid is the cornerstone of the chemical industry. It is produced from sulfur via sulfur dioxide and sulfur trioxide by means of the contact process. 9. The halogens are toxic and reactive elements that are found only in compounds with other elements. Fluorine and chlorine are strong oxidizing agents and are prepared by electrolysis. 10. The reactivity, toxicity, and oxidizing ability of the halogens decrease from fluorine to iodine. The halogens all form binary acids (HX) and a series of oxoacids. KEY WORDS Carbide, p. 839 Catenation, p. 839 Chlor-alkali process, p. 858 Cyanide, p. 839 Hydrogenation, p. 836 QUESTIONS AND PROBLEMS GENERAL PROPERTIES OF NONMETALS Review Questions 21.1 Without referring to Figure 21.1, state whether each of the following elements are metals, metalloids, or nonmetals: (a) Cs, (b) Ge, (c) I, (d) Kr, (e) W, (f) Ga, (g) Te, (h) Bi. 21.2 List two chemical and two physical properties that distinguish a metal from a nonmetal. 21.3 Make a list of physical and chemical properties of chlorine (Cl2) and magnesium. Comment on their differences with reference to the fact that one is a metal and the other is a nonmetal. 21.4 Carbon is usually classified as a nonmetal. However, the graphite used in “lead” pencils conducts electricity. Look at a pencil and list two nonmetallic properties of graphite. HYDROGEN Review Questions 21.5 Explain why hydrogen has a unique position in the periodic table. 21.6 Describe two laboratory and two industrial preparations for hydrogen. 21.7 Hydrogen exhibits three types of bonding in its compounds. Describe each type of bonding with an example. 21.8 What are interstitial hydrides? 21.9 Give the name of (a) an ionic hydride and (b) a co- Back Forward Main Menu TOC valent hydride. In each case describe the preparation and give the structure of the compound. 21.10 Describe what is meant by the “hydrogen economy.” Problems 21.11 Elements number 17 and 20 form compounds with hydrogen. Write the formulas for these two compounds and compare their chemical behavior in water. 21.12 Give an example of hydrogen as (a) an oxidizing agent and (b) a reducing agent. 21.13 Compare the physical and chemical properties of the hydrides of each of the following elements: Na, Ca, C, N, O, Cl. 21.14 Suggest a physical method that would allow you to separate hydrogen gas from neon gas. 21.15 Write a balanced equation to show the reaction between CaH2 and H2O. How many grams of CaH2 are needed to produce 26.4 L of H2 gas at 20°C and 746 mmHg? 21.16 How many kilograms of water must be processed to obtain 2.0 L of D2 at 25°C and 0.90 atm pressure? Assume that deuterium abundance is 0.015 percent and that recovery is 80 percent. 21.17 Predict the outcome of the following reactions: (a) CuO(s) H2(g) 88n (b) Na2O(s) H2(g) 88n Study Guide TOC Textbook Website MHHE Website 866 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS 21.18 Starting with H2, describe how you would prepare (a) HCl, (b) NH3, (c) LiOH. to burn in an atmosphere of CO2 even though CO2 does not support combustion. Explain. 21.34 Is carbon monoxide isoelectronic with nitrogen (N2)? CARBON NITROGEN AND PHOSPHORUS Review Questions 21.19 Give an example of a carbide and a cyanide. 21.20 How are cyanide ions used in metallurgy? 21.21 Briefly discuss the preparation and properties of carbon monoxide and carbon dioxide. 21.22 What is coal? 21.23 Explain what is meant by coal gasification. 21.24 Describe two chemical differences between CO and CO2. Problems Forward 21.35 Describe a laboratory and an industrial preparation of nitrogen gas. 21.36 What is meant by nitrogen fixation? Describe a process for fixation of nitrogen on an industrial scale. 21.37 Describe an industrial preparation of phosphorus. 21.38 Why is the P4 molecule unstable? Problems 21.25 Describe the reaction between CO2 and OH in terms of a Lewis acid-base reaction such as that shown on p. 632. 21.26 Draw a Lewis structure for the C2 ion. 2 21.27 Balance the following equations: (a) Be2C(s) H2O(l ) 88n (b) CaC2(s) H2O(l ) 88n 21.28 Unlike CaCO3, Na2CO3 does not yield CO2 when heated. On the other hand, NaHCO3 undergoes thermal decomposition to produce CO2 and Na2CO3. (a) Write a balanced equation for the reaction. (b) How would you test for the CO2 evolved? [Hint: Treat the gas with limewater, an aqueous solution of Ca(OH)2.] 21.29 Two solutions are labeled A and B. Solution A contains Na2CO3 and solution B contains NaHCO3. Describe how you would distinguish between the two solutions if you were provided with a MgCl2 solution. (Hint: You need to know the solubilities of MgCO3 and MgHCO3.) 21.30 Magnesium chloride is dissolved in a solution containing sodium bicarbonate. On heating, a white precipitate is formed. Explain what causes the precipitation. 21.31 A few drops of concentrated ammonia solution added to a calcium bicarbonate solution cause a white precipitate to form. Write a balanced equation for the reaction. 21.32 Sodium hydroxide is hygroscopic — that is, it absorbs moisture when exposed to the atmosphere. A student placed a pellet of NaOH on a watch glass. A few days later, she noticed that the pellet was covered with a white solid. What is the identity of this solid? (Hint: Air contains CO2.) 21.33 A piece of red-hot magnesium ribbon will continue Back Review Questions Main Menu TOC 21.39 Nitrogen can be obtained by (a) passing ammonia over red-hot copper(II) oxide and (b) heating ammonium dichromate [one of the products is Cr(III) oxide]. Write a balanced equation for each preparation. 21.40 Write balanced equations for the preparation of sodium nitrite by (a) heating sodium nitrate and (b) heating sodium nitrate with carbon. 21.41 Sodium amide (NaNH2) reacts with water to produce sodium hydroxide and ammonia. Describe this reaction as a Brønsted acid-base reaction. 21.42 Write a balanced equation for the formation of urea, (NH2)2CO, from carbon dioxide and ammonia. Should the reaction be run at a high or low pressure to maximize the yield? 21.43 Some farmers feel that lightning helps produce a better crop. What is the scientific basis for this belief? 21.44 At 620 K the vapor density of ammonium chloride relative to hydrogen (H2) under the same conditions of temperature and pressure is 14.5, although, according to its formula mass, it should have a vapor density of 26.8. How would you account for this discrepancy? 21.45 Explain, giving one example in each case, why nitrous acid can act both as a reducing agent and as an oxidizing agent. 21.46 Explain why nitric acid can be reduced but not oxidized. 21.47 Write a balanced equation for each of the following processes: (a) On heating, ammonium nitrate produces nitrous oxide. (b) On heating, potassium nitrate produces potassium nitrite and oxygen gas. (c) On heating, lead nitrate produces lead(II) oxide, nitrogen dioxide (NO2), and oxygen gas. Study Guide TOC Textbook Website MHHE Website QUESTIONS AND PROBLEMS 21.48 Explain why, under normal conditions, the reaction of zinc with nitric acid does not produce hydrogen. 21.49 Potassium nitrite can be produced by heating a mixture of potassium nitrate and carbon. Write a balanced equation for this reaction. Calculate the theoretical yield of KNO2 produced by heating 57.0 g of KNO3 with an excess of carbon. 21.50 Predict the geometry of nitrous oxide, N2O, by the VSEPR method and draw resonance structures for the molecule. (Hint: The atoms are arranged as NNO.) 21.51 Consider the reaction N2(g) O2(g) 34 2NO(g) Given that the G° for the reaction at 298 K is 173.4 kJ, calculate (a) the standard free energy of formation of NO, (b) KP for the reaction, and (c) Kc for the reaction. 21.52 From the data in Appendix 3, calculate H° for the synthesis of NO (which is the first step in the manufacture of nitric acid) at 25°C: 4NH3(g) 5O2(g) 88n 4NO(g) 6H2O(l ) 21.53 Explain why two N atoms can form a double bond or a triple bond, whereas two P atoms normally can form only a single bond. 21.54 When 1.645 g of white phosphorus are dissolved in 75.5 g of CS2, the solution boils at 46.709°C, whereas pure CS2 boils at 46.300°C. The molal boiling-point elevation constant for CS2 is 2.34°C/m. Calculate the molar mass of white phosphorus and give the molecular formula. 21.55 Starting with elemental phosphorus, P4, show how you would prepare phosphoric acid. 21.56 Dinitrogen pentoxide is a product of the reaction between P4O10 and HNO3. Write a balanced equation for this reaction. Calculate the theoretical yield of N2O5 if 79.4 g of P4O10 are reacted with an excess of HNO3. (Hint: One of the products is HPO3.) 21.57 Explain why (a) NH3 is more basic than PH3, (b) NH3 has a higher boiling point than PH3, (c) PCl5 exists but NCl5 does not, (d) N2 is more inert than P4. 21.58 What is the hybridization of phosphorus in the phosphonium ion, PH4 ? 21.60 Give an account of the various kinds of oxides that exist and illustrate each type by two examples. 21.61 Hydrogen peroxide can be prepared by treating barium peroxide with sulfuric acid. Write a balanced equation for this reaction. 21.62 Describe the Frasch process for obtaining sulfur. 21.63 Describe the contact process for the production of sulfuric acid. 21.64 How is hydrogen sulfide generated in the laboratory? Problems 21.65 Draw molecular orbital energy level diagrams for O2, O2 , and O2 . 2 21.66 One of the steps involved in the depletion of ozone in the stratosphere by nitric oxide may be represented as NO(g) 2H2O2(aq) 88n 2H2O(l ) 21.68 21.69 21.70 21.71 Review Questions 21.59 Describe one industrial and one laboratory preparation of O2. Back Forward Main Menu TOC O3(g) 88n NO2(g) O2(g) From the data in Appendix 3 calculate G°, KP, and Kc for the reaction at 25°C. 21.67 Hydrogen peroxide is unstable and decomposes readily: 21.72 OXYGEN AND SULFUR 867 21.73 21.74 O2(g) This reaction is accelerated by light, heat, or a catalyst. (a) Explain why hydrogen peroxide sold in drugstores comes in dark bottles. (b) The concentrations of aqueous hydrogen peroxide solutions are normally expressed as percent by mass. In the decomposition of hydrogen peroxide, how many liters of oxygen gas can be produced at STP from 15.0 g of a 7.50 percent hydrogen peroxide solution? What are the oxidation numbers of O and F in HFO? Oxygen forms double bonds in O2, but sulfur forms single bonds in S8. Explain. In 1996, 48 million tons of sulfuric acid were produced in the United States. Calculate the amount of sulfur (in grams and moles) used to produce that amount of sulfuric acid. Sulfuric acid is a dehydrating agent. Write balanced equations for the reactions between sulfuric acid and the following substances: (a) HCOOH, (b) H3PO4, (c) HNO3, (d) HClO3. (Hint: Sulfuric acid is not decomposed by the dehydrating action.) Calculate the amount of CaCO3 (in grams) that would be required to react with 50.6 g of SO2 emitted by a power plant. SF6 exists but OF6 does not. Explain. Explain why SCl6, SBr6, and SI6 cannot be prepared. Study Guide TOC Textbook Website MHHE Website 868 NONMETALLIC ELEMENTS AND THEIR COMPOUNDS 21.75 Compare the physical and chemical properties of H2O and H2S. 21.76 The bad smell of water containing hydrogen sulfide can be removed by the action of chlorine. The reaction is H2S(aq) Cl2(aq) 88n 2HCl(aq) S(s) If the hydrogen sulfide content of contaminated water is 22 ppm by mass, calculate the amount of Cl2 (in grams) required to remove all the H2S from 2.0 102 gallons of water. (1 gallon 3.785 L.) 21.77 Describe two reactions in which sulfuric acid acts as an oxidizing agent. 21.78 Concentrated sulfuric acid reacts with sodium iodide to produce molecular iodine, hydrogen sulfide, and sodium hydrogen sulfate. Write a balanced equation for the reaction. THE HALOGENS Review Questions 21.79 Describe an industrial method for preparing each of the halogens. 21.80 Name the major uses of the halogens. Problems 21.81 Metal chlorides can be prepared in a number of ways: (a) direct combination of metal and molecular chlorine, (b) reaction between metal and hydrochloric acid, (c) acid-base neutralization, (d) metal carbonate treated with hydrochloric acid, (e) precipitation reaction. Give an example for each type of preparation. 21.82 Sulfuric acid is a weaker acid than hydrochloric acid. Yet hydrogen chloride is evolved when concentrated sulfuric acid is added to sodium chloride. Explain. 21.83 Show that chlorine, bromine, and iodine are very much alike by giving an account of their behavior (a) with hydrogen, (b) in producing silver salts, (c) as oxidizing agents, and (d) with sodium hydroxide. (e) In what respects is fluorine not a typical halogen element? 21.84 A 375-gallon tank is filled with water containing 167 g of bromine in the form of Br ions. How many liters of Cl2 gas at 1.00 atm and 20°C will be required to oxidize all the bromide to molecular bromine? 21.85 Draw structures for (a) (HF)2 and (b) HF2 . 21.86 Hydrogen fluoride can be prepared by the action of sulfuric acid on sodium fluoride. Explain why hy- Back Forward Main Menu TOC 21.87 21.88 21.89 21.90 drogen bromide cannot be prepared by the action of the same acid on sodium bromide. Aqueous copper(II) sulfate solution is blue. When aqueous potassium fluoride is added to the CuSO4 solution, a green precipitate is formed. If aqueous potassium chloride is added instead, a bright-green solution is formed. Explain what happens in each case. What volume of bromine (Br2) measured at 100°C and 700 mmHg pressure would be obtained if 2.00 L of dry chlorine (Cl2), measured at 15°C and 760 mmHg, were absorbed by a potassium bromide solution? Use the VSEPR method to predict the geometries of the following species: (a) I3 , (b) SiCl4, (c) PF5, (d) SF4. Iodine pentoxide, I2O5, is sometimes used to remove carbon monoxide from the air by forming carbon dioxide and iodine. Write a balanced equation for this reaction and identify species that are oxidized and reduced. ADDITIONAL PROBLEMS 21.91 Write a balanced equation for each of the following reactions: (a) Heating phosphorous acid yields phosphoric acid and phosphine (PH3). (b) Lithium carbide reacts with hydrochloric acid to give lithium chloride and methane. (c) Bubbling HI gas through an aqueous solution of HNO2 yields molecular iodine and nitric oxide. (d) Hydrogen sulfide is oxidized by chlorine to give HCl and SCl2. 21.92 (a) Which of the following compounds has the greatest ionic character? PCl5, SiCl4, CCl4, BCl3 (b) Which of the following ions has the smallest ionic radius? F , C4 , N3 , O2 (c) Which of the following atoms has the highest ionization energy? F, Cl, Br, I (d) Which of the following oxides is most acidic? H2O, SiO2, CO2 21.93 Both N2O and O2 support combustion. Suggest one physical and one chemical test to distinguish between the two gases. 21.94 What is the change in oxidation number for the following reaction? 3O2 88n 2O3 21.95 Describe the bonding in the C2 ion in terms of the 2 molecular orbital theory. 21.96 Starting with deuterium oxide (D2O), describe how you would prepare (a) NaOD, (b) DCl, (c) ND3, (d) C2D2, (e) CD4, (f) D2SO4. 21.97 Solid PCl5 exists as [PCl4 ][PCl6 ]. Draw Lewis Study Guide TOC Textbook Website MHHE Website QUESTIONS AND PROBLEMS structures for these ions. Describe the hybridization state of the P atoms. 21.98 Consider the Frasch process. (a) How is it possible to heat water well above 100°C without turning it into steam? (b) Why is water sent down the outermost pipe? (c) Why would excavating a mine and digging for sulfur be a dangerous procedure for obtaining the element? 21.99 Predict the physical and chemical properties of astatine, a radioactive element and the last member of Group 7A. Back Forward Main Menu TOC 869 21.100 Lubricants used in watches usually consist of longchain hydrocarbons. Oxidation by air forms solid polymers that eventually destroy the effectiveness of the lubricants. It is believed that one of the initial steps in the oxidation is removal of a hydrogen atom (hydrogen abstraction). By replacing the hydrogen atoms at reactive sites with deuterium atoms, it is possible to substantially slow down the overall oxidation rate. Why? (Hint: Consider the kinetic isotope effect.) Study Guide TOC Textbook Website MHHE Website ...
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This document was uploaded on 07/27/2009.

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