P.A. Cox - Inorganic chemistry (793955), страница 14
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As suchdimensions are characteristic of particular crystalline substances, XRPD is a valuable fingerprinting technique for solids,and may be used to follow the course of solid-state reactions. More detailed structural information depends on findingthe positions of atoms within the unit cell, which can be done from the intensities of the different diffraction lines.Although simple structures can be determined from powder diffraction, most structural information comes fromsingle-crystal X-ray diffraction.
A good quality crystal of around 0.3 mm dimension is required, and additionalinformation over that from XRPD comes from its specific orientation in the X-ray beam. Detailed calculations arerequired to match the structure against the observed diffraction intensities but modern computational methodscombined with automated data collection allow this to be done routinely in many laboratories.B7—METHODS OF CHARACTERIZATION53For the compound X discussed above, the structure 1 can be established with reasonable certainly on the basis of itsMS, IR and NMR spectra.
None of these techniques however gives any information about the bond lengths and angles.Single crystal X-ray diffraction on this compound confirms its structure, and gives this further information. All C-Cbond lengths in the benzene ring are equal to 140 pm, and the C=O lengths equal to 114 pm; both lengths are veryslightly longer than ones in the free ligands (139 and 113 pm, respectively).Although X-ray diffraction is usually reliable, some difficulties can arise. Crystals with disorder—sometimes notsuspected—can give misleading results. Problems can be caused by the fact the X-rays are scattered by electrons so thatthe scattering power of an atom is proportional to its atomic number.
It can be hard to locate light atoms such ashydrogen in the neighborhood of heavy elements. It may be also be impossible from X-ray measurements alone todistinguish between elements of nearly the same atomic number. In principle some of these problems can be overcomeby using neutron diffraction but that is a much more expensive technique not routinely available. In general it isimportant that X-ray structures should be backed up by other information, especially a good elemental analysis.Section C—Structure and bonding in moleculesC1ELECTRON PAIR BONDSKey NotesLewis and valencestructuresOctets and‘hypervalence’ResonanceFormal chargesLimitationsRelated topicsA Lewis structure shows the valence electrons in a molecule.
Twoshared electrons form a single bond, with correspondingly more formultiple bonds. Some atoms may also have nonbonding electrons(lone-pairs). Valence structures show the bonds simply as lines.In most stable molecules and ions of the elements C-F, each of theseatoms has eight electrons (an octet) in its valence shell. Expansion ofthe octet and increased valency is possible with elements in periods 3and below.When several alternative valence structures are possible, the bondingmay be described in terms of resonance between them.Formal charges are assigned by apportioning bonding electronsequally between the two atoms involved.
They can be useful torationalize apparent anomalies in bonding, and to assess the likelystability of a proposed valence structure.Many covalent molecules and ions cannot be understood in terms ofelectron pair bonds between two atoms. They include electrondeficient boron hydrides and transition metal compounds.Electronegativity and bondMolecular shapes: VSEPRtype (B1)(C2)Introduction to nonmetals(F1)Lewis and valence structuresA single covalent bond is formed when two atoms share a pair of electrons.
Double and triple bonds can be formedwhen two or three such pairs are shared. A Lewis structure is a representation of a molecule or complex ion thatshows the disposition of valence electrons (inner shells are not drawn) around each atom. 1–4 show Lewis structuresof CH4, H2O, O2 and N2, the last two molecules having a double and triple bond, respectively. These representationsare entirely equivalent to the valence structures (1′–4′) in which each bonding pair of electrons is represented by aline.A molecule such as H2O has nonbonding or lone-pair electrons localized on one atom rather than shared.
Thepresence of these has important consequences for both the shape of a molecule and its chemical properties (see TopicsC2 and C9).56SECTION C—STRUCTURE AND BONDING IN MOLECULESSimple complex ions such as ammoniumand tetrahydroboratecan be drawn in a similar way; thevalence structures shown are essentially identical to those for CH4 as the total number of valence electrons is the samein all examples. The isoelectronic principle suggests that molecules or ions having the same number of valenceelectrons should have similar valence structures, although this idea has limitations.Octets and ‘hypervalence’A great majority of simple molecules containing the elements C-F of the second period can be represented by Lewisstructures with eight electrons around each of these atoms, including all shared electrons and lone-pairs.
The octetrule provides a systematization of the normal valencies of these elements: for example, a nitrogen atom has fiveelectrons in its valence shell and so must share three more to achieve an octet, thus forming three bonds. Hydrogen islimited to two electrons in its valence shell, and these differences may be understood from the valence atomic orbitalsavailable for electrons, 1s only for H, 2s and 2p in the second period; the exclusion principle then limits the number ofelectrons that can be accommodated (see Topics A3 and A4).Some molecules containing boron (e.g.
BF3 7) have an incomplete octet and this has implications for theirchemical reactivity (see Topics C8 and F3). Generally, however, structures with complete octets are preferred. Thusthe triple-bonded representation for carbon monoxide (8) is better than the double-bonded one (8′) where carbon onlyhas six valence-shell electrons.C1—ELECTRON PAIR BONDS57Nonmetallic elements of the third and subsequent periods form some compounds entirely analogous to those of thesame group in period 2. Thus we have H2S, H2Se and H2Te similar to H2O, all with octets. These heavier elements,however, are capable of octet expansion or hypervalence, the latter term implying a valency higher than ‘normal’.Examples are SF4 and SF6 (9, 10) where sulfur has respectively 10 and 12 electrons in its valence shell. Hypervalence issometimes considered to be a consequence of the availability of further orbitals for bonding (e.g.
3d in addition to 3s and3p for sulfur). Although this may play a part, it is generally thought that other differences between the periods areequally important, especially size and electronegativity (see Topic F1).ResonanceSometimes more than one valence structure is possible and there appears to be no unique assignment. A familiar organicexample is in the disposition of double and single C—C bonds in benzene (see Topic C6, Structure 6). In the carbonateion11) the three structures shown are equivalent by symmetry, and experimentally all three C—O bonds haveequal length. We describe this situation as resonance between the different structures, and represent it by the doubleheaded arrows shown in 11.
The term is misleading as it suggests a rapid oscillation between different structures, whichcertainly does not happen. It is better to think of a wavefunction that is formed by combining the structures, none of whichon their own describe the bonding correctly.Resonance may also be appropriate with different valence structures that are not equivalent but look equally plausible, asin nitrous oxide (N2O 12).Formal chargesAtoms are often found in bonding situations that do not correspond to their ‘normal’ valency. Such cases can berationalized by the concept of formal charge. A formal charge on an atom is essentially the charge that would remainif all covalent bonds were broken, with the electrons being assigned equally to the atoms involved.
Moremathematically, it is defined asformal charge=(no. of valence electrons in neutral atom)58SECTION C—STRUCTURE AND BONDING IN MOLECULES−(no. of nonbonding electrons)−(1/2) (no. of electrons in bonds formed)The formal charges in CO and in the two valence structures for N2O are shown in 13 and 14. The isoelectronicprinciple allows us to understand these structures by analogy. Thus C− and O+ are both isoelectronic to neutral N andcan similarly form three bonds.
The N2O structures can be understood with the isoelectronic relations N− and O (twobonds expected), N+ and C (four bonds) and O− and F (one bond).Formal charges are frequently drawn in organic structures; for example, ‘trivalent’ carbon can occur as a carbocation (C+ isoelectronic to B, and with an incomplete octet) or a carbanion (C− isoelectronic to N, with a nonbonding pair).They are not always written on inorganic valence structures, but the idea is useful in judging the viability of a proposedstructure. Some general principles are:• structures without formal charges are preferred if possible;• structures with formal charges outside the range −1 to +1 are generally unfavorable;• negative formal charges should preferably be assigned to more electronegative atoms, positive charges to moreelectropositive atoms.Thus in N2O (14), the structure with O− is probably more significant than that with N−. The BF molecule (15) isisoelectronic with CO but the corresponding triple-bonded structure appears very unlikely because it requires formalcharges B2− and F2+.
The single-bonded form without charges may best describe the bonding.Formal charge is very different from oxidation state, which is assigned by apportioning electrons in a bond to themore electronegative atom rather than equally (Topic B4). Both are artificial assignments, useful in their respectiveways, but neither is intended as a realistic judgment of the charges on atoms.LimitationsThe model described in this section can be justified theoretically using the quantum mechanical valence bondtheory. Nevertheless, there are many molecules where bonding cannot be described simply in terms of electron pairslocalized between two atoms.
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