Chemistry part 37, Julia Burdge,2e (2009) pdf

916 CHAPTER 24 Nonmetallic Elements and Their Compounds Figure 24.11 Structures of some common phosphorus-containing oxoacids. Pre gna nt women are advised not to consume lar ge qu anti t ies of soda because of the ph os ph ate content. • • • '0' . . II H-O-P-O-H . . I . . H '0' . . II H-O-P - O- H . . I '0' • • I H Phosp ho ri c ac id (H 3 P0 4 l '0' . . II H-O-P-H •• I H Hypophospho ro us acid (H 3 P0 2 l '0' '0' .t)o II II - II H-O-P-O-P-O-P-O-H " [ " 1 " 1 " :0: :0: I I H H '0' • • I H In the pure form, pho sphoric acid is a colorless solid (m.p. 42.2°C). The phosphoric acid we use in the laboratory is usually an 82% H 3 P0 4 solution (by mass). Pho sphoric acid and phosphates have many commerc ial applications in detergents, fertilizer s, flame retardant s, and toothpastes, and as buffers in carbonated be verages. Like nitrogen, pho sphorus is an element that is essential to life. It constitutes only about 1 percent by ma ss of the human bod y, but it is a v er y important 1 percent. About 23 percent of the human skeleton is mineral matter. The phosphorus content of this mineral matter, calcium phos- phate [Ca3( P0 4)2 ], is 20 percent. Our teeth are ba sically Ca 3( P04 )2 and Ca S (P0 4)3 0H. Phosphates are also important components of the genetic materials deoxyribonucleic acid (DNA) and ribo- nucleic acid ( RN A). Oxygen and Sulfur Oxygen Oxygen is by far the mo st abundant element in Earth's crust, constituting about 46 percent of its ma ss. In addition, the atmosphere contains about 21 percent molecular oxygen by volume (23 percent by ma ss). Like nitrogen, oxygen in the free state is a diatomic molecule (0 2 ), In the labo- rator y, oxygen gas can be obtained by heating potassium chlorate: 2KCI0 3 (s) +. 2KCl (s) + 30 2 (g) The reaction is usually catal yz ed by manganese(IV) dioxide ( Mn0 2)' Pure oxygen gas can be prepared by electrol yz ing water (page 781). Industrially, oxygen ga s is prepared by the fractional distillation of liquefied air. Oxygen ga s is colorless and odorles s. Ox ygen 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 i t, a human being cannot survive for more than a few minutes. Although oxygen ha s two allotropes, O 2 and 0 3 , when we speak of molecular oxygen, we normally mean O? Ozone (0 3 ) is less stable than O 2 , The O 2 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 chemi- cals. Its main uses are in the steel industry and in sewage treatment. Oxygen is also used as a bleaching agent for pulp and paper, in medicine to ease breathing difficulties, in oxyacetylene torche s, and as an oxidizing agent in man y inorganic and organic reactions. Oxygen forms three types of oxides: the normal oxide (or simply the oxide), which contains the 0 2 - ion; the peroxide, which contains the O ~ - ion; and the superoxide, which contains the O 2 ion: '0" .2- '0" '0" . 2- •• ••• • • • '0'0' - • • • • • • • • • • • • • Oxide Peroxide Superoxide The ions are all strong Br 0nsted ba ses and react with water as follows: Oxide: 0 2- (aq) + H? O(l) +. 20H - (aq) Pe ro xide: 20 ~ - ( aq) + 2H ?0(1) • 0 2(g) + 40H - (aq) Superoxide: 40 2 (aq) + 2H 2 0(l) • 30 ?(g) + 40H - (aq) The reaction of 0 2 - with water is a hy drol ys is reaction, but those involving O ~- and O 2 are redox processes. SECTION 24.5 Oxygen and Sulfur The nature of bonding in oxides changes across any period in the periodic table. 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 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: MgO J \, 1 • , " Acidic Bas ic Amphoteric The basicity of the oxides increases as we move down a paIticular group. MgO does not react with water, for example, but reacts with acid as follows: On the other hand, BaO , which is more basic, undergoes hydrolysis to yield the corresponding hydroxid~: BaO(s) + H 2 0(l) +. Ba(OH Maq) The be st-known peroxide is hydrogen peroxide (H 2 0 ?). It is a colorless, syrupy liquid (m.p. -O.9 °C), prepared in the laboratory by the action of cold dilute sulfuric acid on barium peroxide octahydrate: The structure of hydrogen peroxide is shown in Figure 24.12. Using the VSEPR method, we see that the H -0 and 0-0 bonds are bent about each oxygen atom in a configuration similar to the struc- ture of water. The lone-pair-bonding-pair repUlsion is greater in H 2 0 2 than in H 2 0, so the H-O-O angle is only 97° (compared with 104.5° for H-O-H in H 2 0) . Hydrogen peroxide is a polar mol- ecule (f.L = 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: t).H O = -196.4 kJ/mol This is a disproportionation reaction. The oxidation number of oxygen changes from - 1 to - 2 and O. 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 H 2 0 2 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 Fe 2+ ions to Fe 3 + ions in an acidic solution: H 2 0 2 (aq) + 2Fe 2+ (aq) + 2H +(aq) -_I 2Fe 3 +(aq) + 2H 2 0(l) It also oxidizes SO~- ions to SO~- ions: In addition, hydrogen peroxide can act as a reducing agent toward substances that are stronger oxi- dizing agents than itself. For example, hydrogen peroxide reduces silver oxide to metallic silver, • Figure 24 .12 The structure of H 2 0 2 · c- - j 918 CHAPTER 24 Nonmetallic Elements and Their Compounds Fig u re 24.13 The preparation of 0 3 from O 2 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 high- voltage source. (The metal foil on the inside of the inner tube is not shown.) During the electrical discharge, O 2 gas is passed through the tube. The 0 3 ga l' formed exits from the upper right-hand tube, along with some umeacted O 2 gas. lute~ tube Metal foil on outer tube 0, Inner tube , :- _ 0 3 plus some unreacted 02 High-voltage source and permanganate (M n0 4) to manganese(II) in an acidic solution, 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 (i.e., compounds containing the O 2 ion). In general, only the most reactive alkali metals (K, Rb, and Cs) form superoxides. 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 is a rather toxic, light-blue gas (b.p. -111.3 °C). Its pungent odor is noticeable around sources of significant electrical discharges (s uch as a subway train). Ozone can be prepared frommolec- ular oxygen, either photochemically or by subjecting O 2 to an electrical discharge (Figure 24.13): D G O = 326.8 kJ/mol Because the standard free energy of formation of ozone is a large positive quantity (D G'f 163.4 kJ /mol), ozone is less stable than molecular oxygen. The ozone molecule has a bent struc- ture in which the bond angle is 116.5°: •• o . / ~ :0. 0: • • •• • • • • /P :01" :0: • • • • 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: 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: 0 3(g) + 3Hg(l) +. 3HgO(s) The beneficial effect of ozone in the stratosphere and its undesirable action in smog formation were discussed in Chapter 21. Sulfur 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. SECTION 24.5 Oxygen and Sulfur 919 The largest known reserves of sulfur are found in sedimentary deposits. In addition, sulfur occurs widely in gypsum (CaS04 . 2H 2 0) and various sulfide minerals such as pyrite (FeS2) (Figure 24.14). Sulfur is also present in natural gas as H 2 S, S02, and other sulfur-containing compounds. Sulfur is extracted from underground deposits by the Frasch l process, shown in Figure 24.15. In this process, superheated water (liquid water heated to about 160°C under high pressure to pre- vent it from boiling) is pumped down the outermost pipe to melt the sulfur. Next, compressed air is forced down the innermost pipe. 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 Ss ring structure: • • • • • 'S' 'S' ./ ' ./ .S. .S. o. '5" ./ ' .S. .S . • • • • It is a yellow, tasteless, and odorless solid (m.p. 112°C) 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 Sg 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, so the viscosity decreases. Like nitrogen, sulfur shows a wide variety of oxidation numbers in its compounds (Table 24.3). The best-known hydrogen compound of sulfur is hydrogen sulfide, which is prepared by the action of an acid on a sulfide; for example, Compressed air f I ' = + Sulfur Superheated water ,, 110- ~ Molten su lf ur 1. Hennan Frasch (1 8 51 - 1914). German chemical engineer. Bes id es in venting the process for obtaining pure sulfur, Frasch developed methods for refining petroleum. Figure 24.14 Pyrite (FeS2)' . " ' ~ " ~ . Multimedia Nonmetallic elements -t he extraction of sulfur . Figure 24.15 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. , 920 CHAPTER 24 Nonmetallic Elements and Their Compounds , Oxidation Number -2 o + 1 +2 +4 +6 Compound Hydrogen sulfide Sulfur (for reference) Disulfur dichloride Sulfur dichloride Sulfur dioxide Sulfur tlioxide Formula Structure • • 'S' · . /' s 'S' • • Cl: /'" s-s ,./ :CI • • • • S .' / " ' . . Cl. .Cl. o. " • • -0 S 0- :0'/ "0: .' '. '0' II S . / " . :0. .0: • • • • Nowadays, hydrogen sulfide used in qualitative analysis is prepared by the hydrolysis of thioacetamide: S II CH3-C + 2H 20 + H+ \ NH 2 Thioacetamide o II -_. CH3-C + H 2 S + NHt \ O-H 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 weak diprotic acid (see Table 15.5). In basic solu- tion, H 2 S is a reducing agent. For example, it is oxidized by perrnanganate to elemental sulfur: 3H 2 S(aq) + 2Mn0 4 (aq) +. 3S(s) + 2Mn02(S) + 2H 2 0(l) + 20H - (aq) • Sulfur has two important oxides: sulfur dioxide (S02) and sulfur trioxide (S03)' Sulfur dioxide is formed when sulfur burns in air: In the laboratory, it can be prepared by the action of an acid on a sulfite; for example, or by the action of concentrated sulfuric acid on copper: Sulfur dioxide (b.p. -10 °C) is a pungent, colorless gas that is quite toxic. An acidic oxide, it reacts with water as follows: Sulfur dioxide is slowly oxidized to sulfur trioxide, but the reaction rate can be greatly enhanced by a platinum or vanadium oxide catalyst: '. SECTION 24.6 The Halogens Sulfur trioxide dissolves in water to form sulfuric acid: The contributing role of sulfur dioxide to acid rain is discussed on page 844. Sulfuric acid is the world's most important industrial chemical. It is prepared industrially by first burning sulfur in air: Next is the key step of converting sulfur dioxide to sulfur trioxide: Vanadium(V) oxide (V 205) is the catalyst used for the second step. Because the sulfur dioxide and oxygen molecules react in contact with the surface of solid V? OS, the process is referred to as the contact process. Sulfuric acid is a diprotic acid. It is a colorless, viscous liquid (m.p. 10.4 QC) . The concen- trated sulfuric acid we use in the laboratory is 98 percent H 2 S0 4 by mass (density = 1.84 g/cm\ 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 serie s, thereby liberating molecular hydrogen in a displacement reaction: 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). Thu s, copper reacts with concentrated sulfuric acid as follows: Depending on the nature of the reducing agents, the sulfate ion may be further reduced to elemen- tal sulfur or the sulfide ion. For example, the reduction of H 2 S0 4 by HI yields H 2 S and 1 2 : Concentrated sulfuric acid oxidizes nonmetals. For example, it oxidizes carbon to. carbon dioxide and sulfur to sulfur dioxide: C(s) + 2H 2 S0 4 (aq) +. CO 2 (g) + 2S0 2 (g) + 2H 2 0(l) S(s) + 2H 2 S0 4 (aq) • 3S0 2 (g) + 2H 2 0(t) Carbon disulfide, a colorless, flammable liquid (b.p. 46 Q C), is formed by heating carbon and sulfur to a high temperature: C(s) + 2S(l) +. CS 2 (l) It is only slightly soluble in water. Carbon disulfide is a good solvent for sulfur, phosphoru s, 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 fluorin e: Set) + 3Fig) +. SF6(g) Sulfur hexafluoride is a nontoxic, colorless gas (b.p. 63.8 Q C). It is the most inert of all sulfur com- pound s; it resists attack even by molten KOH. The structure and bonding of SF 6 were discussed in Chapters 8 and 9. The Halogens The halogens fluorine, chlorine, bromine, and iodine-are reactive nonmetals. Table 24.4 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: 1. Fluorine is the most reactive of all the halogens. The difference in reactivity between flu o- rine and chlorine is greater than that between chlorine and bromine. Table 24.4 shows that • 922 CHAPTER 24 Nonmetallic Elements and Their Compounds • Property F (I Br 1 Valence electron 2i2l 3i3 p s 4i4 p 5 5s 2 5l configuration Melting point (0C) -223 -102 -7 114 Boiling point (0C) -18 7 -35 59 183 Appearance* Pale-yellow Yellow~green Red-brown Dark-violet vapor gas gas liquid Dark metallic- looking solid Atomic radius (pm) 72 99 114 133 Ionic radius ( pm ) -'- 136 181 195 216 Ionization energy 1680 1251 1139 1003 (kJ/mol) Electronegativity 4.0 3.0 2.8 2.5 Standard reduction 2.87 1.36 1.07 0.53 potential (V)* Bond enthalpy 150.6 242.7 192.5 151.0 (kJ/mol)* * The se va lues and descriptions apply to the diatomic species X" where X represents a halogen atom. The half- reaction is X 2 (g) + 2e- • 2X-(aq ). t Refers to the anion X- . the F-F bond is considerably weaker than the Cl-Cl bond. The weak bond in F2 can be explained in terms of the lone pairs on the F atoms: •• •• :F-F: •• •• The small size of the F atoms (see Table 24.4) allows a close approach of the three lone pairs on each of the F atoms, resulting in a greater repulsion than that found in C1 2 , which consists of larger atoms. 2. Hydrogen fluoride (HF) has a relatively high boiling point (19.5°C) as a result of strong intermolecular hydrogen bonding, whereas all other hydrogen halides have much lower boiling points. 3. Hydrofluoric acid is a weak acid, whereas all other hydrohalic acids (HCl, HBr, and HI) are strong acids. 4. Fluorine reacts with cold sodium hydroxide solution to produce oxygen difluoride as follows: 2F2(g) + 2NaOH(aq) +. 2NaF(aq) + H 2 0(l) + OF 2 (g) The same reaction with chlorine or bromine, on the other hand, produces a halide and a hypohalite: Xig) + 2NaOH(aq) +. NaX(aq) + NaXO(aq) + H 2 0(l) where X stands for Cl or Br. Iodine does not react under the same conditions. 5. Silver fluoride (AgF) is soluble. All other silver halides (AgCl, AgBr, and AgI) are insoluble. The element astatine also belongs to the Group 7 A family. However, all isotopes of astatine are radioactive; its longest-lived isotope is astatine-21O, which has a half-life of 8.3 h. As a result, 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 (X 2 ). 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 (CaF 2 ) and cryolite (Na3AlF6)' SECTION 24.6 The Halogens 923 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 oxy- gen. From Table 19.1. we find that F 2 (g) + 2e - +. 2F - (aq) 02(g) + 4H + (aq) + 4e - • 2H 2 0(I) EO = 2.87 V EO = 1.23 V IfF2 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 24.16): Anode (oxidation): 2F - +. Fig) + 2e- Cathode (reduction): 2H + + 2e - + ~ Hi g) Overall reaction: 2HF(l) +. H 2 (g) + F2(g) Chlorine gas (CI 2 ) is prepared industrially by the electrolysis of molten NaCI or by the chlor-alkali process, the electrolysis of a concentrated aqueous NaCI 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 2NaCI(aq) + 2H 2 0(l) e l ec trol ys i ~ 2NaOH(aq ) + H2(g) + Cl i g) As you can see, this reaction yields two useful by-products, NaOH and H 2 . The cells are designed to separate the molecular chlorine from the sodium hydroxide solution and the molecular hydro- gen to prevent side reactions such as 2NaOH(aq) + CI 2 (g) +. NaOCI(aq) + NaCI(aq) + H 2 0(l) and Hig) + C1 2 (g) +. 2HCI(g) These reactions must be prevented because they cons ume the desired products and can be danger- ous because a mixture of H2 and Cl 2 is explosive. Figure 24.17 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 Fz gas 1, H2 gas ( . 11 ,Carbon anode Hz ga s • • Diaphragm to prevent mixing of Hz and F z ga ses Steel cathode Liquid HF Figure 24.16 Electrolytic cell for the preparation of fluorine gas. Note that because Hz and F z form an explosive mixture, these gases must be separated from each other. , 924 CHAPTER 24 Nonmetallic Elements and Their Compounds Figu re 24.17 Mercury cell used in the chlor-alkai process. The cathode contains mercury. The sodium-mercury amalgam is treated with water outside the cell to produce sodium hydroxide and hydrogen gas. Figu re 24 . 18 The industrial manufacture of chlorine gas. • Figure 24.19 Diaphragm ce ll used in the chlor-alkali process. Graphite anode I • Brine • • • Brine • ,tr Hg cathode Hg plus NalHg with platinum. Brine is continuously passed through the cell as shown in the diagram. The elec- trode reactions are Anode (oxidation): 2CI- (aq)' Clig) + 2e- Cathode (reduction): 2Na(aq) + 2e- H g(l). 2NaHg Overall reaction: 2NaCI(aq). 2NafHg + Clig) where NafHg 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: 2NafHg + 2H 2 0(l) 2NaOH(aq) + Hig) + 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 24.18 shows the industrial manufacture of chlorine gas. The half-cell reactions in a diaphragm cell are shown in Figure 24.19. The asbestos dia- phragm 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 com- partment 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. Anode Brine Battery Oxidation -_. Cl, (g) + 2e - Cathode Asbestos :;:::::.~J= ~r, diaphragm • 1 ' - - 1 • NaOH solution Reduction SECTION 24.6 The Halogens 925 In the laboratory, chlorine, bromine, and iodine can be prepared by heating the alkali halides (NaCI, KBr, or KI) in concentrated sulfuric acid in the presence of manganese(IV) oxide. A rep- resentative reaction is Mn02(s) + 2H 2 SOiaq) + 2NaCI(aq) -_. MnS0 4(aq) + Na2S0iaq) + 2H 2 0(l ) + Cl i g) Compounds of the Halogens Most of the halides can be categorized as either ionic or covalent. 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. 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 F 2 ) and - 1, in all its compounds. The hydrogen halides, an important class of halogen compounds, can be formed by the direct combination of the elements: H2(g) + Xi g) :;:::. =::!:' 2HX(g) where X denotes a halogen atom. These reactions (especially the ones involving F2 and C1 2 ) can occur with explosive violence. Industrially, hydrogen chloride is produced as a by-product in the manufacture of chlorinated hydrocarbons: In the laboratory, hydrogen fluoride and hydrogen chloride can be prepared by combining the metal halides with concentrated sulfuric acid: CaP 2 (s) + H 2 SOiaq) +. 2HF (g) + CaS04(s) 2NaCl(s) + H 2 SOiaq) • 2HCl(g) + Na2 S0 i 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 H 2 S0 4 is Instead, hydrogen bromide is prepared by first reacting bromine with phosphorus to form phos- phorus tribromide: Next, PBr3 is treated with water to yield HBr: Hydrogen iodide can be prepared in a similar manner. HF is so highly reactive that it attacks silica and silicates: 6HF(aq) + Si0 2 (g) +. H 2 SiF 6 (aq) + 2H 2 0(l ) This property makes HF suitable for etching glass and is the reason that hydrogen fluoride must be kept in plastic or inert metal (e.g., Pt) containers. Hydrogen fluoride is us ed in the manufacture of Freons (see Chapter 17); for example, CClil) + HF(g) +. CFCI 3 (g) + HCl (g) CFC1 3 (g) + HF(g) • CF 2 C1 2 (g) + HCl(g) It is also important in the production of aluminum. 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 The halogens also form a series of oxoacids with the following general formulas: HXO Hypohalous acid HX0 2 Halous acid HX0 3 Halic acid HX0 4 Perhalic acid [...]... Materials Dental Implants Soft Tissue Materials Artificial]oints Nanotechnology Graphite, Buckyballs, and Nanotubes Semiconductors Superconductors • aterla s Modern Materials Chemistry and the 2005 Nobel Prize • The 2005 Nobel Prize in Chemistry was awarded to Yves Chauvin, Robert H Grubbs, and Richard R Schrock for the development of the metathesis method in organic synthesis In the metathesis method,... Chapter, You Will Learn about some of the chemistry involved in the development of modem materials Before you begin, you should review • Functional groups [ I~~ Section 10.2] • Organic polymers [ ~~ Section 10.6] , Media Player/ MPEG Content J- Chapter in Review Yves Chauvin (France), Robert H Grubbs (USA), and Richard R Schrock (USA) shared the 2005 Nobel Prize in Chemistry for their work on the metathesis... longer without interstate highways b) the Morton Salt Company didn't exist before 1924 c) public health officials had to be convinced that iodization of salt was safe and effective d) Michigan was not yet part of the United States 4 According to the passage, goiter a) can be caused by a low-salt diet b) can be cured by the addition of iodide to the diet c) is potentially debilitating d) was a more serious... 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 I Figure 24.20 Agi particles I IP H I CH 2 -C-C I \ NH 2 0H 0 The Halogens I 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... bromide Silver iodide is sometimes used in cloud seeding, a process for inducing rainfall on a small scale (Figure 24.20) The advantage of using silver iodide is that enormous numbers of nuclei (i.e., 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... CHAPTER 25 Modern Materials Polymers Many biological molecules, such as DNA, starch, and proteins, have very large molecular masses These molecules are polymers because they are made up of many smaller parts linked together [ ~~ Section 10.6] The prefix poly comes from a Greek root meaning "many," and the small molecules that make up the individual building blocks of polymers are called monomers Many... implants, and prosthetics Some of these polymers are thermoplastic, which means that they can be melted and reshaped, or heated and bent Others are thermosetting, which means that their shape is determined as part of the chemical process that fOllIled the polymer Thermosetting polymers cannot be reshaped easily and are not easily recycled, whereas thermoplastic polymers can be melted down and cast into new... chain contains a carbon atom with an unpaired electron (i.e., -CHzCHz'), though, so it must be closed by bonding with another atom This might be a hydrogen atom (e.g., -CHzCHz-H), but it might also be part of the initiator molecule (e.g., -CHzCHz-OR) Although the residue of the initiator molecule has very different chemical properties than the main polymer chain, it will not significantly affect the... systematic name of C2 H4 is ethene, but the common name ethylene is still widely used Recall that species with an unpaired electron are called free radicals [~ Section 8.8] Multimedia Organic and Biochemistry- natural and synthetic polymers ~C=C~~C=C~~C=C~ - _ H Figure 25.1 / \ / HH \ / HH \ H ~~~~~~ -C-C-C-C-C-CI11111 HHHHHH Addition polymerization to form polyethylene from ethylene Each curved, . positive pressure is applied on the anode side of the com- partment to prevent the migration of the OH- ions from the cathode compartment. Periodically, fresh brine solution is added to the. 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:. peroxide readily decomposes when heated or exposed to sunlight or even in the presence of dust particles or certain metals, including iron and copper: t).H O = -196.4 kJ/mol This is a disproportionation