CHEMISTRY THE CENTRAL SCIENCE
22 CHEMISTRY OF THE NONMETALS
CHAPTER SUMMARY AND KEY TERMS
INTRODUCTION AND SECTION 22.1 The periodic table is useful for organizing and remembering the descriptive chemistry of the elements. Among elements of a given group, size increases with increasing atomic number, and electronegativity and ionization energy decrease. Nonmetallic character parallels electronegativity, so the most nonmetallic elements are found in the upper right portion of the periodic table.
Among the nonmetallic elements, the first member of each group differs dramatically from the other members; it forms a maximum of four bonds to other atoms and exhibits a much greater tendency to form π bonds than the heavier elements in its group.
Because O2 and H2O are abundant in our world, we focus on two important and general reaction types as we discuss the nonmetals: oxidation by O2 and proton-transfer reactions involving H2O or aqueous solutions.
SECTION 22.2 Hydrogen has three isotopes: protium (), deuterium (), and tritium () Hydrogen is not a member of any particular periodic group, although it is usually placed above lithium. The hydrogen atom can either lose an electron, forming H+, or gain one, forming H– (the hydride ion). Because the H — H bond is relatively strong, H2 is fairly unreactive unless activated by heat or a catalyst. Hydrogen forms a very strong bond to oxygen, so the reactions of H2 with oxygen-containing compounds usually lead to the formation of H2O. Because the bonds in CO and CO2 are even stronger than the O — H bond, the reaction of H2O with carbon or certain organic compounds leads to the formation of H2. The H+ (aq) ion is able to oxidize many metals, forming H2(g). The electrolysis of water also forms H2(g).
The binary compounds of hydrogen are of three general types: ionic hydrides (formed by active metals), metallic hydrides (formed by transition metals), and molecular hydrides (formed by non-metals). The ionic hydrides contain the H– ion; because this ion is extremely basic, ionic hydrides react with H2O to form H2 and OH–.
SECTIONS 22.3 AND 22.4 The noble gases (group 8A) exhibit a very limited chemical reactivity because of the exceptional stability of their electron configurations. The xenon fluorides and oxides and KrF2 are the best-established compounds of the noble gases.
The halogens (group 7A) occur as diatomic molecules. All except fluorine exhibit oxidation states varying from –1 to +7. Fluorine is the most electronegative element, so it is restricted to the oxidation states 0 and – 1. The oxidizing power of the element (the tendency to form the –1 oxidation state) decreases as we proceed down the group. The hydrogen halides are among the most useful compounds of these elements; these gases dissolve in water to form the hydrohalic acids, such as HCl(aq). Hydrofluoric acid reacts with silica. The interhalogens are compounds formed between two different halogen elements. Chlorine, bromine, and iodine form a series of oxyacids, in which the halogen atom is in a positive oxidation state. These compounds and their associated oxyanions are strong oxidizing agents.
SECTIONS 22.5 AND 22.6 Oxygen has two allotropes, O2 and O3 (ozone). Ozone is unstable compared to O2, and it is a stronger oxidizing agent than O2. Most reactions of O2 lead to oxides, compounds in which oxygen is in the –2 oxidation state. The soluble oxides of non-metals generally produce acidic aqueous solutions; they are called acidic anhydrides or acidic oxides. In contrast, soluble metal oxides produce basic solutions and are called basic anhydrides or basic oxides. Many metal oxides that are insoluble in water dissolve in acid, accompanied by the formation of H2O. Peroxides contain O — O bonds and oxygen in the –1 oxidation state. Peroxides are unstable, decomposing to O2 and oxides. In such reactions peroxides are simultaneously oxidized and reduced, a process called disproportionation. Superoxides contain the O2– ion in which oxygen is in the oxidation state.
Sulfur is the most important of the other group 6A elements. It has several allotropic forms; the most stable one at room temperature consists of S8 rings. Sulfur forms two oxides, SO2 and SO3, and both are important atmospheric pollutants. Sulfur trioxide is the anhydride of sulfuric acid, the most important sulfur compound and the most-produced industrial chemical. Sulfuric acid is a strong acid and a good dehydrating agent. Sulfur forms several oxyanions as well, including the SO32– (sulfite), SO42– (sulfate), and S2O32– (thiosulfate) ions. Sulfur is found combined with many metals as a sulfide, in which sulfur is in the –2 oxidation state. These compounds often react with acids to form hydrogen sulfide (H2S), which smells like rotten eggs.
SECTIONS 22.7 AND 22.8 Nitrogen is found in the atmosphere as N2 molecules. Molecular nitrogen is chemically very stable because of the strong N≡N bond. Molecular nitrogen can be converted into ammonia via the Haber process. Once the ammonia is made, it can be converted into a variety of different compounds that exhibit nitrogen oxidation states ranging from –3 to +5. The most important industrial conversion of ammonia is the Ostwald process, in which ammonia is oxidized to nitric acid (HNO3). Nitrogen has three important oxides: nitrous oxide (N2O), nitric oxide (NO), and nitrogen dioxide (NO2). Nitrous acid (HNO2) is a weak acid; its conjugate base is the nitrite ion (NO2–). Another important nitrogen compound is hydrazine (N2H4).
Phosphorus is the most important of the remaining group 5A elements. It occurs in nature as phosphate minerals. Phosphorus has several allotropes, including white phosphorus, which consists of P4 tetrahedra. In reaction with the halogens, phosphorus forms trihalides PX3 and pentahalides PX5. These compounds undergo hydrolysis to produce an oxyacid of phosphorus and HX. Phosphorus forms two oxides, P4O6 and P4O10. Their corresponding acids, phosphorous acid and phosphoric acid, undergo condensation reactions when heated. Phosphorus compounds are important in biochemistry and as fertilizers.
SECTIONS 22.9 AND 22.10 The allotropes of carbon include diamond, graphite, fullerenes, carbon nanotubes, and graphene. Amorphous forms of graphite include charcoal and carbon black. Carbon forms two common oxides, CO and CO2. Aqueous solutions of CO2 produce the weak diprotic acid carbonic acid (H2CO3), which is the parent acid of hydrogen carbonate and carbonate salts. Binary compounds of carbon are called carbides. Carbides may be ionic, interstitial, or covalent. Calcium carbide (CaC2) contains the strongly basic acetylide ion (C22–), which reacts with water to form acetylene. Other important inorganic carbon compounds include hydrogen cyanide (HCN) and carbon disulfide (CS2).
The other group 4A elements show great diversity in physical and chemical properties. Silicon, the second most abundant element, is a semiconductor. It reacts with Cl2 to form SiCl4, a liquid at room temperature, a reaction that is used to help purify silicon from its native minerals. Silicon forms strong Si — O bonds and therefore occurs in a variety of silicate minerals. Silica is SiO2; silicates consist of SiO4 tetrahedra, linked together at their vertices to form chains, sheets, or three-dimensional structures. The most common three-dimensional silicate is quartz (SiO2).Glass is an amorphous (noncrystalline) form of SiO2. Silicones contain O — Si — O chains with organic groups bonded to the Si atoms. Like silicon, germanium is a metalloid; tin and lead are metallic.
SECTION 22.11 Boron is the only group 3A element that is a non-metal. It forms a variety of compounds with hydrogen called boron hydrides, or boranes. Diborane (B2H6) has an unusual structure with two hydrogen atoms that bridge between the two boron atoms. Bo-ranes react with oxygen to form boric oxide (B2O3), in which boron is in the +3 oxidation state. Boric oxide is the anhydride of boric acid (H3BO3). Boric acid readily undergoes condensation reactions.
• Be able to use periodic trends to explain the basic differences between the elements of a group or period (Section 22.1).
• Explain the ways in which the first element in a group differs from subsequent elements in the group (Section 22.1).
• Be able to determine electron configurations, oxidation numbers, and molecular shapes of elements and compounds (Sections 22.2–22.11).
• Know the sources of the common nonmetals, how they are obtained, and how they are used (Sections 22.2–22.11).
• Understand how phosphoric and phosphorous acids undergo condensation reactions (Section 22.8).
• Explain how the bonding and structures of silicates relate to their chemical formulas and properties (Section 22.10).