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Elementary valence theory 261

shapes of molecules. The more sophisticated valence theories yield information about the
electrical, magnetic, and spectroscopic properties of molecules.
Elementary valence theories invoke two principal bond types. In ionic bonding,
electrostatic interactions generate bonds between ions formed by electron transfer from
one element to the other. In covalent bonding two elements are held together by shared
electrons in order that both may adopt an energetically favorable electron configuration.
In reality, both are extreme forms of the same bonding phenomenon. Pure covalent bonds
are formed by elements with identical electronegativities, with more ionic bonding
character being introduced to the bond as the electronegativity difference between the
elements increases (see Topic H4). Even in extreme cases of ionic bonding, the degree of
covalent character may still be quite high.
Two complementary theories were originally developed to explain the number and
nature of covalent bonds (Lewis theory) and the shapes of molecules (VSEPR theory).
More sophisticated theories have superseded these approaches for detailed investigations,
but they remain useful in semi-empirical and non-rigorous discussions of molecular
bonding.

Lewis theory

The Lewis theory of covalent bonding may be regarded as an elementary form of valence
bond theory. It is nonetheless useful for describing covalent molecules with simple
covalent bonds, and works successfully in describing the majority of, for example,
organic compounds. Lewis theory recognizes both the free energy gains made in the
formation of complete atomic electron shells, and the ability of atoms to achieve this state
by sharing electrons. The sharing process is used as a description of covalent bonds.
The atoms are firstly drawn so as to represent their relative arrangement, with electron
pairs (marked as pairs of dots) between neighboring atoms to indicate a shared bonding
electron pair. No attempt is made to describe the three-dimensional geometric shape of
the molecule. Multiple bonds are represented by two or three electron pairs as appropriate
(Fig. 1a). Further electrons are added to each atom, so as to represent the non-bonding
electrons and so complete the electron configuration of all the atoms (Fig. 1b). It is
customary to replace the bonding pairs of shared electrons with one line for each pair—
each line representing a bond—with multiple lines representing multiple bonds (Fig. 1c).
Although main group elements tend to adopt inert gas configurations, which may be
represented by eight valence electrons (an octet), or two in the case of helium, a number
of elements are energetically stable with incomplete octets. The most commonly cited
example is boron, which is stable with six valence electrons as in BF 3, or Be with four as
in BeCl 2. Larger elements are capable of hypervalency, where it is energetically
favorable for more than eight valence electrons to be held in an expanded octet.
Examples of this are PF 5 (ten valence electrons) and XeF 4 (twelve valence electrons)
(Fig. 1d).