Inorganic Chemistry
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Inorganic Chemistry

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with the 
elements C, N and O than with others in the same group. The stability of O2 (double bonds) versus 
S8 (single bonds) can be rationalized from the fact that B(O=O) is more than twice as large as B(O-
O) but the same is not true of sulfur. In a similar way we have CO2 (molecular with C=O) and SiO2 
Table 2. Variation in bond enthalpy (kJ mol\u22121) with bond order 
A-B Single Double Triple 
C-C 347 612 838 
C-O 358 805 1077 
Si-O 466 638 \u2013 
N-N 167 247 942 
N-O 214 587 \u2013 
O-O 144 498 \u2013 
S-S 266 429 \u2013 
(vi) Other exceptions to rule (i) above occur with A-O and A-X bonds (X being a halogen), which 
generally increase in strength between periods 2 and 3 (e.g. C-O<Si-O). This may be partly 
due to the increased electronegativity difference when A is period 3 (see (iv) above), but 
repulsion between lone-pairs electrons on nonbonded atoms may also play a role (e.g. F-F 
repulsion in CF4, where the atoms are closer together than in SiF4). 
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(polymeric with single Si-O). The formation of multiple bonds is one of the main factors leading 
to differences in chemistry between 2p series elements and those in lower periods (see Topic F1). 
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Changes with valence state are important in understanding the stability of \u2018hypervalent\u2019 
compounds. Thus SH4 and SH6 are unknown, whereas they would be thermodynamically stable 
compounds if their S-H bonds were as strong as in H2S. The common formation of fluorides in high 
valency states (e.g. SF6, IF7) can be understood from a combination of factors. The F-F bond is itself 
rather weak, E-F bonds are generally strong, and they decline less rapidly with increasing n in EFn 
molecules than in other compounds. 
Other measures 
Thermochemical bond energies may be hard to determine, either for experimental reasons or because 
of the limitations in the assumption of transferability that is often required. Alternative measures of 
comparative bond strength that are often useful include the following: 
\u2022 the bond length, which for a given pair of elements decreases with increasing strength (e.g. 
with increasing bond order, as in the sequence N-N 145 pm, N=N 125 pm, N\u2261N 110 pm); bond 
length measurements are often useful for showing the existence of metal-metal bonds in 
transition metal compounds (see Topic H5); 
\u2022 the bond stretching frequency measured by vibrational spectroscopy (e.g. IR, see Topic B7) is 
related to the stretching force constant and increases with bond strength; IR measurements 
have been particularly useful in the study of CO as a ligand in transition metal carbonyl 
compounds (Topic H9). 
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Section C\u2014Structure and bonding in molecules 
Definition and scope 
A Lewis acid is any species capable of accepting a pair of electrons, and a Lewis base is a species 
with a pair of electrons available for donation. The terms acceptor and donor are also commonly 
used. Lewis acids include H+ and metal cations, molecules such as BF3 with incomplete octets, and 
ones such as SiF4 where octet expansion is possible (see Topic C1). Any species with nonbonding 
electrons is potentially a Lewis base, including molecules such as NH3 and anions such as F
\u2212. 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 
Lewis acids and bases may interact to give a donor-acceptor complex; for example, 
The bond formed is sometimes denoted by an arrow (as in 1) and called a dative bond but it is not 
really different from any other polar covalent bond. Thus the complex [SiF6]
2\u2212 has a regular 
octahedral structure where the two \u2018new\u2019 Si\u2014F bonds are indistinguishable from the other four. (It is 
isoelectronic with SF6; see Topic C2.) 
Key Notes 
Definition and 
A Lewis acid (or acceptor) can accept an electron pair from a Lewis base (or donor) to 
form a donor-acceptor complex. The formation of solvated ions, complexes in solution and 
coordination compounds are examples of this type of interaction. 
Models of 
Contributions to the donor-acceptor interaction may come from electrostatic forces, and 
from the overlap between the highest occupied MO (HOMO) of the donor and the lowest 
unoccupied MO (LUMO) of the acceptor. 
Hard donors interact more strongly with hard acceptors, soft donors with soft acceptors. 
Harder acids tend to be more electropositive, and harder bases more electronegative. Softer 
donor and acceptor atoms tend to be larger and more polarizable. 
Polymerization Formation of dimers and polymeric structures is a manifestation of donor-acceptor 
interaction between molecules of the same kind. 
Related topics Solvent types and properties (E1) Brønsted acids and bases (E2) 
Complex formation (E3) 
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The 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 transition metals (see Topics E1, E3, H2 and H9). Many chemical 
reactions also depend on donor-acceptor interactions. For example, the hydrolysis of SiCl4 to give Si
(OH)4 in water begins with a step such as 
where H2O is acting as a donor to the SiCl4 acceptor. 
Models of interaction 
Interaction between a Lewis acid and a base may have an electrostatic contribution as donor atoms 
are often electronegative and possess some partial negative change, whereas acceptor atoms may be 
positively charged. There is also 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 acceptor A and the highest occupied MO 
(HOMO) of the donor D. The levels in the center show the formation of a more stable bonding MO 
and a destabilized antibonding MO in the complex. The electron pair from the donor occupies the 
bonding MO, 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 LUMO and the donor HOMO is least. In this model the best acceptors will have empty 
orbitals at low energies, the best donors filled orbitals at high energies. By contrast, the strongest 
electrostatic interactions will take place between the smallest and most highly charged (positive) 
acceptor and (negative) donor atoms. 
Hard-soft classification 
It is found that the relative strength of donors depends on the nature of the acceptor and vice versa. 
The hard and soft acid-base (HSAB) classification is often used to rationalize some of the 
differences. When two acids (A1 and A2) are in competition for two bases (B1 and B2) the 
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Fig. 1. Molecular orbital interaction between a donor (:D) and an acceptor (A). 
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will lie in the direction where the harder of the two acids is in combination with the harder base, and 
the softer