P.A. Cox - Inorganic chemistry (793955), страница 40
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Autoprotolysis gives the ions H3O+ and OH−, which are also known in solid salts,H3O+ with anions of strong acids (e.g. [H3O]+[NO3]−; hydrated species such as [H5O2]+ are also known), and OH− inhydroxides, which are formed by many metals.Oxides of most metallic elements have structures that may be broadly classed as ionic (see Topics D3 and D4).
Theclosed-shell O2− ion is unknown in the gas phase, the reactionbeing very endothermic. It is therefore only the large lattice energy obtained with the O2− ion that stabilizes it in solids(see Topic D6). The variety of coordination numbers (CN) of oxide is large, examples being:Oxide has a notable tendency for symmetrical coordination in ionic solids (linear, planar or tetrahedral with CN=2, 3or 4, respectively) and unlike sulfide rarely forms layer structures.The distinction between ionic and polymeric solids is not absolute, and oxides of metals with low electropositivecharacter (e.g. HgO) or in high oxidation states (e.g.
CrO3) are better described as having polar covalent bonds. A fewmetals in very high oxidation states form molecular oxides (e.g. Mn2O7, OsO4).Many ternary and more complex oxides are known. It is normal to distinguish complex oxides such as CaCO3, whichcontain discrete oxoanions, and mixed oxides such as CaTiO3, which do not (see Topic D5).In water, the very basic O2− ion reacts to form hydroxide:174SECTION F—CHEMISTRY OF NONMETALSTable 1.
Some oxoacids, showing their anhydrides and the anions formed by themaAnionwith a strong tendency to polymerize and form complex structures.acid with intermediate states of ionization possible.cParent anhydride unknown.bPolyproticand so ionic oxides are basic and either form alkaline solutions if soluble in water, or otherwise dissolve in acidsolution. Covalent oxides (including those such as CrO3 formed by metals in high oxidation states) are acidic and reactwith water to form oxoacids:(See Topic E2 for Pauling’s rules on acid strength.) Such oxides may therefore be regarded as acid anhydrides.Table 1 shows a selection of oxoacids with their anhydrides and illustrates the conventional nomenclature.
For example,sulfurous and sulfuric acids display the lower (+4) and higher (+6) oxidation state, respectively, and their anions arecalled sulfite and sulfate.Some oxides are amphoteric and have both acidic and basic properties; this often happens with a metal ion with ahigh charge/size ratio such as Be2+ or Al3+ (see examples in Topics E2 and G3–G5). A few nonmetallic oxides (e.g. CO)are neutral and have no appreciable acid or basic properties.Peroxides and superoxidesAdding one or two electrons to dioxygen gives the superoxideand peroxideions.
As the added electronsoccupy the π antibonding orbital (see Topic C4) the bond becomes progressively weaker and longer. Superoxides MO2,rather than simple oxides M2O are the normal products of reacting the heavier alkali metals with oxygen; peroxidesM2O2 are also formed. This may be explained by lattice energy arguments (see Topic D6). With most metal ions, theandions.
With large, lowhigher lattice energy obtained with O2− forces the disproportionation of the largercharged cations, however, the lattice energy gain is insufficient to cause disproportionation. The peroxide ion can alsobe stabilized in peroxo complexes, where it acts as a ligand to transition metals, as in [CrV(O2)4]3−.The simplest covalent peroxide is hydrogen peroxide H2O2, which is normally encountered in aqueous solution.Although kinetically fairly stable, it can act as either an oxidizing agent (giving H2O) or a reducing agent (giving O2),and many transition metal ions catalyze its decomposition. Organic peroxides (R2O2) and peroxoacids (e.g.
theF7—OXYGEN175percarbonate ion, 2) contain the fairly weak peroxo O—O linkage. Some covalent peroxides can be unpredictably anddangerously explosive.Positive oxidation statesReaction with strong oxidizing agents gives theion, which has a stronger and shorter bond than O2 (see Topic C4):Fluorides include F2O and F2O2. The latter has a considerably shorter O-O bond than in peroxides, a fact that mayindicate some contribution of ionic valence structures such as (3), which allow a degree of multiple bonding.
Allcompounds in positive oxidation states are very strongly oxidizing. Compounds with heavier halogens are normallyregarded as halogen oxides and are discussed in Topic F9.Section F—Chemistry of nonmetalsF8SULFUR, SELENIUM AND TELLURIUMKey NotesThe elementsChalcogenidesHalidesOxides and oxoacidsOther compoundsRelated topicsThe elements known as chalcogens show pronounced differences fromoxygen in the same group, being much less electronegative. Sulfides areimportant minerals for some elements.
Elemental structures are based onrings and chains with single bonds.The hydrides are toxic gases. Metal chalcogenides are much less ionic thanoxides, and often have different (e.g. layer) structures.Many halides are known in oxidation states up to +6. Most are molecularcompounds but some have polymeric structures.EO2 and EO3 compounds have structures that are increasingly polymericfor heavier elements.
They form oxoacids, of which sulfuric acid is themost important.Cationic species such ascan be prepared. Sulfur and nitrogen forman interesting range of binary compounds.Introduction to nonmetalsOxygen (F7)(E1)The elementsThe elements known collectively as the chalcogens are in the same group (16) as oxygen (Topic F7). They form somecompounds similar to those of oxygen, but show many differences characteristic of other nonmetal groups (seeTopic F1).Sulfur is widespread in the Earth’s crust, occurring as metal sulfides, sulfates, and native or elemental sulfur formedby bacterial oxidation of sulfides.
Many less electropositive metals known as chalcophiles are found commonly assulfide minerals (see Topic J2); some important examples are pyrites (FeS2), sphalerite (zinc blende, ZnS), molybdenite(MoS2), cinnabar (HgS) and galena (PbS). Volatile sulfur compounds such as H2S and organic compounds are also foundin petroleum and natural gas. The element is used in large amounts for the manufacture of sulfuric acid (see below).Selenium and tellurium are much rarer, found as minor components of sulfide minerals.Sulfur has several allotropic forms, the most stable of which are molecular solids containing S8 rings.
The elementalforms of Se and Te have spiral chains and are semiconductors. In all of these solids each atom forms two single bonds toneighbors (see Topic D2). Sulfur combines directly with oxygen and halogens (except I), and with many lesselectronegative elements to form sulfides. The other elements show similar properties although reactivity declinesdown the group.F8—SULFUR, SELENIUM AND TELLURIUM177ChalcogenidesMolecular compounds include H2S and its analogs, and many organic compounds. The hydrides are made by the actionof Brønsted acids on metal chalcogenides. They are extremely toxic gases, weakly acidic in water (e.g. for H2S, pK1=6.8, pK2=14.2).
Many polysulfanes H2Sn containing S-S bonds are also known.Solid chalcogenides are formed by all metallic elements and by many nonmetals. Only with the mostelectropositive metals do they commonly have the same structures as oxides (see Topics D3 and D4). With transitionmetals, compounds MX (which are frequently of variable stoichiometry) have the nickel arsenide or similar structures inwhich metal-metal bonding is present.
MX2 compounds either have layer structures (e.g. TiS2, TiSe2, TiTe2, all CdI2types) or structures containing diatomic ions (e.g. FeS2 has S22− units and so is formally a compound of FeII not FeIV).Chalcogenides of electropositive metals are decomposed by water giving hydrides such as H2S, but those of lesselectropositive elements (often the ones forming sulfide ores, see above) are insoluble in water.HalidesA selection of the most important halides is show in Table 1 and routes to the preparation of sulfur compounds areshown in Fig.
1. With sulfur the fluorides are most stable and numerous, but Se and Te show an increasing range of heavierhalides. Compounds such as S2Cl2 and S2F10 have S-S bonds; S2F2 has another isomer S=SF2. Sulfur halides aremolecular and monomeric with structures expected from VSEPR (e.g. SF4 ‘see-saw’, SF6 octahedral; see Topic C2).With the heavier elements increasing polymerization is found, as in (TeCl4)4 (1) and related tetramers.Table 1. Principal halides of S, Se and TeFig.
1. Routes to the preparation of sulfur halides.178SECTION F—CHEMISTRY OF NON-METALSThe hexahalides are kinetically inert, but most other halides are highly reactive and are hydrolyzed in water givingoxides and oxoacids. Intermediate hydrolysis products are oxohalides of which thionyl chloride SOCl2 andsulfuryl chloride SO2Cl2 are industrially important compounds.Some of the halides show donor and/or acceptor properties (see Topic C8). For example, SF4 reacts with both Lewisacids (forming compounds such as [SF3]+[BF4]−) and bases (forming either simple adducts such as C5H5N:SF4 withpyridine, or compounds containing the square pyramidal ion [SF5]−).
The complex ions [SeX6]2 and [TeX6]2− (X=Cl,Br, I) are interesting as they appear to have regular octahedral structures in spite of the presence of a nonbondingelectron pair on the central atom (see Topic C2).Oxides and oxoacidsThe major oxides of all three elements (E) are EO2 and EO3. Sulfur in addition forms many oxides of low thermodynamicstability, for example S8O with a structure containing an S8 ring. Sulfur dioxide SO2 is the major product of burningsulfur and organic sulfur compounds in air, and is a serious air pollutant giving rise (after oxidation to H2SO4; seeTopic J6) to acid rain. With one lone-pair, SO2 is a bent molecule and has both Lewis acid and basic properties.
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