P.A. Cox - Inorganic chemistry (793955), страница 36
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They are very strongreducing agents and react with water to give dihydrogen:The hydride ion can act as a ligand and form hydride complexes similar in some ways to those of halides, althoughtheir stability is often limited by the reducing properties of the H− ion. The most important complexes are thetetrahedral ionsandnormally found as the salts NaBH4 and LiAlH4. They may be made by the action ofNaH or LiH on a halide or similar compound of B or Al, and are used as reducing agents and for the preparation ofhydrides of other elements.F2—HYDROGEN155Many transition metal complexes containing hydrogen are known, including the unusual nine-coordinate ion [ReH9]2(see Topic H5).
Hydride is a very strong σ-donor ligand and is often found in conjunction with π-acid ligands and inorganometallic compounds (see Topics H9 and H10).−The hydrogen bondA hydrogen atom bound to an electronegative atom such as N, O or F may interact in a noncovalent way with anotherelectronegative atom. The resulting hydrogen bond has an energy in the range 10–60 kJ mol−1, weak by standards ofcovalent bonds but strong compared with other intermolecular forces (see Topic C10). The strongest hydrogen bondsare formed when a fluoride ion is involved, for example in the symmetrical [F-H-F]− ion. Symmetrical bonds areoccasionally formed with oxygen but in most cases the hydrogen is not symmetrically disposed, a typical example beingin liquid water where the normal O-H bond has a length of 96 pm and the hydrogen bond a length around 250 pm.
Hydrogenbonding arises from a combination of electrostatic (ion-dipole and dipole-dipole) forces and orbital overlap; the lattereffect may be treated by a three-center molecular orbital approach (see Topic C6).Hydrogen bonding is crucial for the secondary structure of biological molecules such as proteins and nucleic acids,and for the operation of the genetic code. Its influence can be seen in the boiling points of simple hydrides (see Table 1and Topic C10, Fig. 1). The exceptional values for NH3, H2O and HF result from strong hydrogen bonding in the liquid.Deuterium and tritiumDeuterium (2D) and tritium (3T) are heavier isotopes of hydrogen (see Topic A1). The former is stable and makes upabout 0.015% of all normal hydrogen.
Its physical and chemical properties are slightly different from those of the lightisotope 1H. For example, in the electrolysis of water H is evolved faster and this allows fairly pure D2 to be prepared.Tritium is a radioactive β-emitter with a half-life of 12.35 years, and is made when some elements are bombarded withneutrons. Both isotopes are used for research purposes. They also undergo very exothermic nuclear fusionreactions, which form the basis for thermonuclear weapons (‘hydrogen bombs’) and could possibly be used as a futureenergy source.Section F—Chemistry of nonmetalsF3BORONKey NotesThe elementHydridesHalidesOxygen compoundsOther compoundsRelated topicsBoron has an unusual chemistry characterized by electron deficiency. Itoccurs in nature as borates.
Elemental structures are very complex.There is a vast range of neutral compounds and anions. Except in theion, the compounds show complex structures, which cannot beinterpreted using simple electron pair bonding models.BX3 compounds are Lewis acids, with acceptor strength in the orderBI3>BBr3> BCl3>BF3.B2O3 and the very weak acid B(OH)3 give rise to a wide range of metalborates with complex structures containing both three- and fourcoordinate boron.Some boron-nitrogen compounds have similar structures to those ofcarbon.
Structurally complex borides are formed with many metals.Rings and clusters (C7)Lewis acids and bases (C9)The elementThe only nonmetallic element in group 13 (see Topic B2), boron has a strong tendency to covalent bonding. Itsuniquely complex structural chemistry arises from the (2s)2(2p)1 configuration, which gives it one less valence electronthan the number of orbitals in the valence shell. Simple compounds such as BCl3 have an incomplete octet and arestrong Lewis acids (see Topics C1 and C9), but boron often accommodates its electron deficiency by formingclusters with multicenter bonding.Boron is an uncommon element on the Earth overall (about 9 p.p.m.
in the crust) but occurs in concentrateddeposits of borate minerals such as borax Na2[B4O5(OH)4].8H2O, often associated with former volcanic activity or hotsprings. It is used widely, mostly as borates in glasses, enamels, detergents and cosmetics, and in lesser amounts inmetallurgy.Boron is not often required in its elemental form, but it can be obtained by electrolysis of fused salts, or by reductioneither of B2O3 with electropositive metals or of a halide with dihydrogen, the last method giving the purest boron. Theelement has many allotropic structures of great complexity; their dominant theme is the presence of icosahedral B12units connected in different ways.
Multicenter bonding models are required to interpret these structures.F3—BORON157HydridesThe simplest hydrogen compounds are salts of the tetrahydroborate ionwhich is tetrahedral and isoelectronicwith methane (see Topic C1). LiBH4 is prepared by reducing BF3 with LiH. It is more widely used as the sodiumsalt, which is a powerful reducing agent with sufficient kinetic stability to be used in aqueous solution. Reaction ofNaBH4 with either I2 or BF3 in diglyme (CH3OCH2)2O gives diborane B2H6, the simplest molecular hydride. Itsstructure with bridging hydrogen atoms requires three-center two-electron bonds (see Topics C1 and C6):Heating B2H6 above 100°C leads to pyrolysis and generates a variety of more complex boranes of which tetraborane(10) B4H10 and decaborane(14) B10H14 are the most stable. Other reactions can lead to anionic species, such as theicosahedral dodecahydrododecaborate(2−) [B12H12]2−, prepared at 180°C:The structural classification and bonding in boranes is described in Topic C7; especially striking are the anions [BnHn]2−with closed polyhedral structures.
Boranes with heteroatoms can also be prepared, such as B10C2H12, which isisoelectronic with [B12H12]2−.Boranes are strong reducing agents and the neutral molecules inflame spontaneously in air, although the anions[BnHn]2− have remarkable kinetic stability. Diborane itself reacts with Lewis bases (see Topic C9). The simplestproducts can be regarded as donor-acceptor complexes with BH3, which is a ‘soft’ Lewis acid and forms adducts with softbases such as CO (1). More complex products often result from unsymmetrical cleavage of B2H6, for example,HalidesMolecular BX3 compounds are formed with all halogens.
They have the trigonal planar structure (D3h) predicted byVSEPR (see Topics C2, C3), although there appears to be a certain degree of π bonding (strongest in BF3) involvinghalogen lone-pairs and the empty boron 2p orbital (see 2 for one of the possible resonance forms). The halides arestrong Lewis acids, BF3 and BCl3 being used as catalysts (e.g. in organic Friedel-Crafts acylations). Interaction with adonor gives a tetrahedral geometry around boron as with the analogous BH3 complex 1.
The π bonding in the parentmolecule is lost and for this reason BF3, where such bonding is strongest, is more resistant to adopting the tetrahedralgeometry than are the heavier halides. Thus the acceptor strengths follow the orderwhich is the reverse of that found with halides of most other elements (see Topic158SECTION F—CHEMISTRY OF NONMETALSF9).
Strongest interaction occurs with hard donors such as F− (forming the stable tetrafluoroborate ion [BF4]−) and withoxygen donors such as water. Except with BF3 (where the B—F bonds are very strong) complex formation often leadsto solvolysis, forming B(OH)3 in water. BF3 itself forms a 1:2 aduct with water, which in the solid state can beformulated as [BF3(H2O)].H2O, one water molecule being coordinated to boron by an oxygen lone pair and the otherheld separately by hydrogen bonding. On melting at 6°C an ionic liquid containing [H3O]+ and [BF3(OH)]− is obtained.Pyrolysis of BX3 compounds leads to halides with B—B bonds, for example, B2X4 (3 with X=F or Cl) and polyhedralBnCln molecules (n=4, 8, 9).Oxygen compoundsBoric oxide B2O3 is very hard to crystallize; the glass has a linked covalent network in which both bridging B—O—Band terminal B=O bonds may be present.
The hydroxide boric acid B(OH)3 is formed by the hydrolysis of manyboron compounds. It has a layer structure made up of planar molecules linked by hydrogen bonding. It is a Lewis acidthat acts as a Brønsted acid in protic solvents. In water the equilibriumgives a pKa=9.25 but complexing can increase the acidity; for example, in anhydrous H2SO4 it forms [B(HSO4)4]− and isone of the few species that can act as a strong acid in that solvent (see Topic F8).Borates can be formed with all metals, although those of groups 1 and 2 are best known.
The structural features arecomplex and rival those of silicates (see Topic D5). Boron can occur as planar BO3 or tetrahedral BO4 groups, oftenlinked by B—O—B bonds as in silicates. For example, 4 shows the ion found in borax Na2[B4O5(OH)4].8H2O, whereboth three- and four-coordinate boron is present. Borosilicate glasses (such as ‘Pyrex’) have lower coefficients ofthermal expansion than pure silicate glasses and so are more resistant to thermal shock.F3—BORON159Other compoundsBoron forms many compounds with nitrogen. Some of these are structurally analogous to carbon compounds, the pair ofatoms BN being isoelectronic with CC.
(For example, the ion [NH3BH2NH3]+ is analogous to propane,CH3CH2CH3.) Boron nitride BN can form two solid structures, one containing hexagonal BN layers similar tographite, and the other with tetrahedral sp3 bonding like diamond (see Topic D2). Borazine B3N3H6 has a 6-π-electronring like benzene (5 shows one resonance form; see Topic C7). Although BN is very hard and resistant to chemicalattack, borazine is much more reactive than benzene and does not undergo comparable electrophilic substitutionreactions.
The difference is a result of the polar B-N bond, and the more reactive B-H bonds.Boron forms a binary carbide, often written B4C but actually nonstoichiometric, and compounds with most metals. Thestoichiometries and structures of these solids mostly defy simple interpretation. Many types of chains, layers andpolyhedra of boron atoms are found. Simple examples are CaB6 and UB12, containing linked octahedra and icosahedra,respectively.Section F—Chemistry of nonmetalsF4CARBON, SILICON AND GERMANIUMKey NotesThe elementsHydrides and organiccompoundsHalidesOxygen compoundsOther compoundsRelated topicsCarbonates and reduced forms of carbon are common on Earth, andsilicates make up the major part of the crust; germanium is much lesscommon. All elements can form the diamond structure; graphite andother allotropes are unique to carbon.Silanes and germanes are less stable than hydrocarbons.
Double bondsinvolving Si and Ge are very much weaker than with C.Halides of all the elements have similar formulae and structures.Those of Si and Ge (but not of C) are Lewis acids and are rapidlyhydrolyzed by water.Carbon oxides are molecular with multiple bonds, those of Si and Gepolymeric in structure. Carbonates contain simpleions, butsilicates and germanates have very varied and often polymericstructures.Compounds with S and N also show pronounced differences betweencarbon and the other elements. Many compounds with metals areknown but these are not highly ionic.
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