Electronegativities of elements and valence-bond theory

With modem data the electronegativity values can be used to estimate the charge transfer (8 or AN/2 AN, differential charge of dipole) in diatomic molecules, as illustrated in Table 3-2 for HCl, metal-chlorine molecules, and metal-oxygen molecules. The charge excesses (5) for the chlorine and oxygen of several diatomic molecules [estimated on the basis of their respective 

The valence-bond concept of chemical bonding begins with the proposition that two neutral atoms combine to form covalent electron-pair 

This confusion in nomenclature may be reduced by the use of covalence in place of valence and defined as the number of covalent bonds of an atom in a molecule. This term was first suggested by Langmuir to give a quantitative measure of the number of covalent bonds for an atom in a molecule. Thus, for the molecules FeCh, FeO, Fe(Cp)2, (porphyrin)Fe, and Fe(Me)2 the covalence of iron is two for FeCls, Fe2O3, Fe(OH)3, and (porphyrin)Fe(OH) the covalence of 

FeO4 (ferrate) and Fe(catecholate)3 the covalence of iron is six and for Fe(CO)5 the covalence of iron is eight. In each of these compounds the iron atom is uncharged and has eight valence electrons (3d 4s2 — d sp - d sp. For these examples, the traditionally used formal oxidation states of iron [II, III, IV, VI, and VIII (or 0), respectively] are the same as their covalences (number of covalent bonds). However, the iron in (porphyrin)Fe( (OH2) (d5sp2) has a covalence of three, a formal oxidation state of three, and a charge of 1-t- via the covalently bound H2O. In the present discussion Roman numeral superscripts associated with the metals in the formulas for their compounds and complexes indicate their covalence (number of covalent bonds), not their oxidation state or number. 

With simple organic compounds covalent bonds result from the combination of the four unpaired electrons in the sp -hybridized orbitals of carbon with the impaired electrons of hydrogen or oxygen. 

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