P.A. Cox - Inorganic chemistry (793955), страница 22
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The Lewis acid-base definition should not be confused with the Brønsted one (see Topic E2):Brønsted bases are also Lewis bases, and H+ is a Lewis acid, but Brønsted acids such as HCl are not Lewis acids.Lewis acids and bases may interact to give a donor-acceptor complex; for example, The bond formed is sometimesdenoted by an arrow (as in 1) and called a dative bond but it is not really different from any other polar covalentbond.
Thus the complex [SiF6]2− has a regular octahedral structure where the two ‘new’ Si—F bonds areindistinguishable from the other four. (It is isoelectronic with SF6; see Topic C2.)92SECTION C—STRUCTURE AND BONDING IN MOLECULESThe scope of the donor-acceptor concept is extremely broad and encompasses many types of chemical interaction,including the solvation and complexation of metal ions and the formation of coordination compounds by transitionmetals (see Topics E1, E3, H2 and H9). Many chemical reactions also depend on donor-acceptor interactions.
Forexample, the hydrolysis of SiCl4 to give Si(OH)4 in water begins with a step such aswhere H2O is acting as a donor to the SiCl4 acceptor.Models of interactionInteraction between a Lewis acid and a base may have an electrostatic contribution as donor atoms are oftenelectronegative and possess some partial negative change, whereas acceptor atoms may be positively charged.
There isalso an orbital interaction, which can be represented by the simple molecular orbital (MO) diagram of Figure 1(see Topic C5). On the left and right are represented respectively the lowest unoccupied MO (LUMO) of the acceptorA and the highest occupied MO (HOMO) of the donor D. The levels in the center show the formation of a more stablebonding MO and a destabilized antibonding MO in the complex. The electron pair from the donor occupies the bondingMO, which is partially shared between the two species.Interaction between the orbitals in Fig.
1 will be strongest when the energy difference between the acceptor LUMOand the donor HOMO is least. In this model the best acceptors will have empty orbitals at low energies, the best donorsfilled orbitals at high energies. By contrast, the strongest electrostatic interactions will take place between the smallestand most highly charged (positive) acceptor and (negative) donor atoms.Hard-soft classificationIt is found that the relative strength of donors depends on the nature of the acceptor and vice versa. The hard and softacid-base (HSAB) classification is often used to rationalize some of the differences.
When two acids (A1 and A2) arein competition for two bases (B1 and B2) the equilibriumFig. 1. Molecular orbital interaction between a donor (:D) and an acceptor (A).C9—LEWIS ACIDS AND BASES93will lie in the direction where the harder of the two acids is in combination with the harder base, and the softer acidwith the softer base. As a standard for comparison the prototype hard acid H+ and soft acid [(CH3)Hg]+ are often used.Thus the equilibriumwill lie to the left or right according to the degree of hardness of the base B.Examples of hard acids are H+, cations of very electropositive metals such as Mg2+, and nonmetal fluorides such asBF3.
Soft acids include cations of late transition and post-transition metals such as Cu+, Pd2+ and Hg2+ (see Topics G4,G6, H3 and H5). The hardness of bases increases with the group number of the donor atom (e.g. NH3<H2O<F−) anddecreases down any group (e.g. NH3>PH3, and F−>Cl−>Br−>I−).Although the hard-soft classification provides a useful systematization of many trends it does not by itself provide anexplanation of the different behavior. Generally it is considered that hard-hard interactions have a greater electrostaticcomponent and soft-soft ones depend more on orbital interactions, but many other factors may be involved. Softacceptor and donor atoms are often large and van der Waals’ forces may contribute to the bonding (see Topic C9);some soft bases such as CO also show π-acceptor behavior (see Topic H9).
It is also important to remember that hardand soft behavior is defined in a competitive situation. When reactions are studied in solution some competition withsolvation is always present (see Topics E1 and E3).PolymerizationThe tendency of many molecules to aggregate and form dimers (e.g. Al2Cl6 2), larger oligomers, or extendedpolymeric structures can be regarded as a donor-acceptor interaction. Thus in the reactiona chlorine atom bound to one AlCl3 uses nonbonding electrons to complex with the other aluminum atom; as in mostother examples of this type the bridging atoms are symmetrically disposed with identical bonds to each aluminum.Polymerization of AXn molecules is more likely to occur when n is small, and when the atom A has vacant orbitals and islarge enough to increase its coordination number.
Many oxides and halides of stoichiometry AB2 and AB3 formstructures that may be regarded as polymeric, although the distinction between this (polar covalent) description and anionic one is not clear-cut (see Topics B1, D4 and F7).Hydrogen bonding (see Topic F2) can also be regarded as a donor-acceptor interaction in which the acceptor LUMOis the (unoccupied) antibonding orbital of hydrogen bonded to an electronegative element.Section C—Structure and bonding in moleculesC10MOLECULES IN CONDENSED PHASESKey NotesMolecular solids andliquidsIntermolecular forcesMolecular polarityRelated topicsIntermolecular forces cause molecular substances to condense toform solids and liquids.
Trouton’s rule provides an approximaterelationship between the normal boiling point of a liquid and thestrength of intermolecular forces.Polar molecules have forces between permanent dipoles. Withnonpolar molecules London dispersion (or van der Waals’) forcesarise between fluctuating dipoles; their magnitude is related tomolecular polarizability, which generally increases with size.Molecules may also have more specific donor-acceptor interactionsincluding hydrogen bonding.The polarity of a molecule arises from charge separation caused byelectronegativity differences in bonds, although contributions fromlone-pairs and the consequences of molecular symmetry are alsoimportant.
High polarity gives strong intermolecular forces, and alsoprovides a major contribution to the dielectric constant.Electronegativity and bondSolvent types and propertiestype (B1)(E1)Molecular solids and liquidsThe condensation of molecular substances into liquid and solid forms is a manifestation of intermolecular forces.The enthalpies of fusion (i.e. melting) and vaporization provide a direct measure of the energy required toovercome such forces.We speak of molecular solids when molecules retain their identity, with geometries similar to those in the gasphase. The structures of molecular solids sometimes resemble those formed by close-packing of spheres (see Topic D2),although with highly unsymmetrical and polar molecules the directional nature of intermolecular forces may play a role.Molecular liquids are more disorganized, but the structural changes between solid and liquid can be subtle and the meltingpoint of a molecular solid is not in general a good guide to the strength of intermolecular forces.
A better correlation isfound with the normal boiling point, as molecules become isolated in the vapor and the influence of intermolecularinteractions is lost.The enthalpy of vaporization ΔHvap divided by the normal boiling point in kelvin (Tb) gives the standard entropyof vaporization (see Topic B3)SECTION C—STRUCTURE AND BONDING IN MOLECULES95and according to Trouton’s rule its magnitude is normally around 90 J K−1 mol−1.
Trouton’s rule is not quantitativelyreliable, and breaks down when molecules have an unusual degree of organization in either the liquid or vapor phase(e.g. because of hydrogen bonding); it does, however, express a useful qualitative relationship between the boilingpoint and the strength of intermolecular forces.
Figure 1 shows the normal boiling points for noble gas elements and somemolecular hydrides.Intermolecular forcesBetween charged ions (whether simple or complex) the Coulomb attraction is the dominant force, as discussed inTopic D6.
Even with neutral molecules, intermolecular forces are essentially electrostatic in origin. With polarmolecules the force between permanent electric dipoles is the dominant one (see below). When polarity isabsent the force arises from the interaction between instantaneous (fluctuating) dipoles, and is known as the Londondispersion or van der Waals’ force. Its strength is related to the polarizability of the molecules concerned.Polarizability generally increases with the size of atoms, and the sequence of boiling points He<Ne<Ar<Kr shown inFig.
1 reflects this. The boiling point increases down the group in most series of nonpolar molecules, for example,CH4<SiH4<GeH4 (also in Fig. 1), the diatomic halogens F2<Cl2<…, and the order CF4<CCl4<CBr4<CI4 found withother molecular halides. (Ionic halides tend to show the reverse order, reflecting the decrease in lattice energy expectedas the size of ions increases; see Topic D6.)Fig.
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