P.A. Cox - Inorganic chemistry (793955), страница 41
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Theliquid is a good solvent for reactions with strong oxidizing agents. SO2 dissolves in water giving acid solutions containingthe pyramidal hydrogensulfite (HSO3−) and sulfite (SO32−) ions. The expected sulfurous acid H2SO3, however, ispresent only in very low concentrations. SeO2 and TeO2 have polymeric structures and give oxoacid salts similar tothose from sulfur.Sulfur trioxide SO3 is made industrially as a route to sulfuric acid, by oxidizing SO2 with oxygen using a vanadium oxidecatalyst.
It can exist as a monomeric planar molecule but readily gives cyclic S3O9 trimers and linear polymers withcorner-sharing SO4 units (see 2 and Topic D3). The highly exothermic reaction with water gives sulfuric acid H2SO4,which is the world’s major industrial chemical, being used in many large-scale processes for making fertilizers,dyestuffs, soaps and detergents, and synthetic fibers (see Topic J4). Anhydrous sulfuric acid undergoes a series of acidbase equilibria such as(see Topic E1). It is a very strongly acid medium, in which HNO3 (a strong acid in water) acts as a base:F8—SULFUR, SELENIUM AND TELLURIUM179The resulting ‘nitrating mixture’ is used for preparing aromatic nitro compounds by electrophilic reactions of.Reaction of HF with SO3 gives fluorosulfonic acid HSO3F, which is even more strongly acidic than sulfuric acid.In mixtures with SO3 and powerful fluoride acceptors such as SbF5 it gives superacid media, which are capable ofprotonating even most organic compounds (see Fig.
2 for examples).Fig. 2. Reactions in ‘superacid’ solutions.SeO3 and selenic acid H2SeO4 are similar to the sulfur analogs except that they are more strongly oxidizing. Telluriumbehaves differently, as telluric acid has the octahedral Te(OH)6 structure, which, as expected from Pauling’s rules, is avery weak acid (see Topic E2).which has a peroxoThere are many other oxoacids of sulfur, of which the most important are peroxodisulfate(O—O) bond, and compounds with S-S bonds including thiosulfatedithioniteand tetrathionate.The reactionis used for the quantitative estimation of I2 in aqueous solution.Other compoundsOxidation of the elements (e.g. by AsF5) in a suitable solvent such as SO2 or H2SO4 gives a series of polyatomiccations such as [S8]2+ and [S4]2+.
The latter (and its Se and Te analogs) has a square-planar structure and can beregarded as a 6π-electron ring (see Topic C7).Also of note are sulfur-nitrogen compounds. The cage-like S4N4 (see Topic C7) is formed by the reaction of S2Cl2with ammonia or NH4Cl. Passing the heated vapor over silver wool gives the planar S2N2 with the same valenceelectron count as [S4]2+. Polymerization forms polythiazyl (SN)x, a linear polymer with metallic conductivity arisingfrom delocalization of the one odd electron per SN unit.Section F—Chemistry of nonmetalsF9HALOGENSKey NotesThe elementsHalides and halidecomplexesOxides and oxoacidsInterhalogen andpolyhalogencompoundsRelated topicsThe halogens are electronegative and oxidizing elements, fluorineexceptionally so. They occur in nature as halides, and form highlyreactive diatomic molecules.Molecular halides are formed with most nonmetals, ionic halides withmetals.
Some halides are good Lewis acids, and many halidecomplexes are known.Most halogen oxides are of low stability, but several oxoacids areknown except for fluorine. Redox stability depends on pH, Cl2 and Br2disproportionating in alkaline solution.Halogens form an extensive range of neutral and ionic compoundswith each other, including some cationic species.Introduction to nonmetals(F1)Binary compounds: simplestructures (D3)Binary compounds: factorsinfluencing structure (D4)The elementsThe halogen group (17) is the most electronegative in the periodic table, and all elements readily form halide ions X−.Trends in chemistry resemble those found in other groups (see Topic F1).
Fluorine is limited to an octet of valenceelectrons. It is the most electronegative and reactive of all elements and often (as with oxygen) brings out the highestoxidation state in other elements: examples where no corresponding oxide is known include PtF6 and AuF5 (seeTopic H5).F and Cl are moderately abundant elements, principal sources being fluorite CaF2 and halite NaCl, from which thevery electronegative elements are obtained by electrolysis. Bromine is mainly obtained by oxidation of Br− found in saltwater; iodine occurs as iodates such as Ca(IO3)2.
Astatine is radioactive and only minute amounts are found in nature.Chlorine is used (as ClO− and ClO2) in bleaches and is an important industrial chemical, other major uses (as with allthe halogens) being in the manufacture of halogenated organic compounds (see Topic J4).The elements form diatomic molecules, F2 and Cl2 being gases at normal temperature and pressure, Br2 liquid and I2solid. They react directly with most other elements and are good oxidizing agents, although reactivity declines down thegroup.
X-X bond strengths follow the sequence F<Cl>Br>I (see Topic C8).F9—HALOGENS181Halides and halide complexesNearly all elements form thermodynamically stable halides. The normal stability sequence is F>Cl>Br>I, which incovalent compounds follows the expected order of bond strengths, and in ionic compounds that of lattice energies (seeTopics C8 and D6). The thermodynamic stability of fluorides (and the kinetic reactivity of F2) is also aided by the weakF-F bond. Many halides can be made by direct combination, but fluorinating agents such as ClF3 are sometimes used inpreference to F2, which is very difficult to handle (see Topic B6).The structural and bonding trends in halides follow similar patterns to those in oxides (see Topics B2 and F7).
Mostnonmetallic elements form simple molecular compounds in which halogen atoms each have a single bond to theother element. This is true also for metals in high oxidation states (e.g. TiCl4 and UF6). The compounds may be solids,liquids or gases, with volatility in the order F>Cl>Br> I as expected from the strength of van der Waals’ forces.
In thehydrogen halides HF is exceptional because of strong hydrogen bonding (see Topic C10). HF is a weak acid inwater, the other HX compounds being strong acids (see Topic E2).Covalent halides are less often polymeric in structure than oxides, a difference partly caused by the differentstoichiometries (e.g.
SiF4 versus SiO2), which provide a higher coordination number in the monomeric molecularhalides. However, the halides of some metals (e.g. beryllium; Topic G3) may be better regarded as polymeric thanionic. Some molecular halides of both metallic and nonmetallic elements form halogen-bridged dimers and higheroligomers (e.g. Al2Cl6; Topic G4).Most metallic elements form solid halides with structures expected for ionic solids (see Topics D3 and D4).Structural differences often occur with MX2 and MX3, fluorides more often having rutile, fluorite or rhenium trioxidestructures, and the heavier halides layer structures.
These differences reflect the more ionic nature of fluorides, and thehigher polarizability of the larger halide ions. Many halides are very soluble in water, but low solubilities are often foundwith fluorides of M2+ and M3+ ions (e.g. CaF2, AlF3), and with heavier halides of less electropositive metals (e.g. AgCl,TlCl). These differences are related to lattice energy trends (see Topics D6 and E4).Many halides of metals and nonmetals are good Lewis acids (see Topic C9). Such compounds are often hydrolyzed bywater, and also form halide complexes (e.g. AlCl42−, PF6−), which can make useful counterions in solids with large orstrongly oxidizing cations.
Both cationic and anionic complexes may be formed by halide transfer, for example, in solidPCl5 (Topic F6) and in liquid BrF3 (see below). Many metal ions also form halide complexes in aqueous solution. For amajority of elements the fluoride complexes are more stable but softer or class b metals form stronger complexes withheavier halides (see Topic E3).Oxides and oxoacidsI2O5 is the only halogen oxide of moderate thermodynamic stability. Other compounds include X2O (not I), X2O2 (Fand Cl), the odd-electron XO2 (Cl and Br), and Cl2O7. Most of these compounds are strongly oxidizing, have lowthermal stability and can decompose explosively.
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