Bonding and Electronegativity
There are two extreme, recognised forms of bonding. The first is ionic, and a common example of such a bonding type occurs in sodium chloride, which contains Na+ and Cl-. There is total electron transfer from sodium to chlorine - from Na {[Ne]3s1} and Cl {[Ne]3s23p5} to Na+ {[Ne]} and Cl- {[Ne]3s23p6 or [Ar]}. Each species achieves a 'rare gas configuration' - in these cases a complete octet of electrons: from this arises the octet rule. If the rare gas is helium (as in the case of Li+), it is a 'duplet rule'. The second extreme type of bonding is covalent, and is shown by any homonuclear diatomic molecule (and strictly speaking not by anything else), for example H2 or Cl2.
with the two electrons shared by each hydrogen, each of which therefore achieves the electron configuration of the rare gas helium. A similar treatment accounts for the covalent bond in Cl2
Each chlorine atom carries three 'lone pairs' which are not involved in the bond formation, although they may mutually repel, thereby weakening slightly the bond strength. Note that we cannot identify which electron in the covalent bond "comes from" which atom; this is our accounting procedure. This method of representing covalent bonds is known as the Lewis "dot and cross" representation. The same method can be used for heteronuclear diatomic molecules such as hydrogen chloride - HCl.
In this case, the two different atoms have 'different tendencies to attract' electrons, and this must be addressed. If you think back to NaCl. Na+ is smaller than Na (the outer 3s electron has been lost), whereas Cl- is larger than Cl (an extra electron has been added).
Electrons are attracted back from the anion towards the cation, concentrating them in the internuclear region. This is partial electron sharing, i.e. covalency. Even NaCl has approximately 5% covalency: nothing is 100% ionic, and only homonuclear diatomic molecules are wholly covalent (e.g. H2 , Cl2, etc.). Any heteronuclear diatomic has some degree of charge separation, depending on the tendency of each atom to 'grab' the available electrons. For example ClF has a charge separation as shown, because F is smaller, and Zeff is high for both halogens. Similarly boron trifluoride has charge separations as shown:-
The extent of charge separation depends on the relative electronegativities of the atoms concerned. Electronegativity depends on how it is defined. We shall look at three scales.
Electronegativity is a measure of the relative tendency of atoms, when bonded in a molecule, to attract electrons to themselves. There are many scales of electronegativity. One of the earliest was due to Mulliken who defined electronegativity (
) as the average of ionisation energy (
Hion the energy required for A ==> A+ + e-) and electron affinity or electron attachment enthalpy (H EA - the energy released when B + e- ===> B-): these processes are respectively endothermic and exothermic, and the energies are positive and negative, so the average of the magnitudes of these energies is written as = (Hion - HEA )/2. The difficulty here is that HEA is not easy to measure directly.
Pauling produced a scale based on thermodynamic data (i.e. on experimental measurements) which yielded differences in electronegativities (
for atom pairs. He noticed that the enthalpy of dissociation (Hdiss ) of a heteronuclear diatomic molecule A-B was frequently greater than the mean of the corresponding enthalpies of the related heteronuclear diatomic molecules A-A and B-B. He defined this excess energy using an arithmetic mean:-
E = Hdiss (A-B) - 0.5Hdiss (A2) +
Hdiss (B2)]
He reasoned that the excess energy derived from an ionic contribution to the bonding in A-B, resulting from the charge separation between its ends (its dipole), and that this [[Delta]]E should lead to a scale of electronegativities. Pauling found that internally consistent electronegativity differences were proportional to the square-root of E: the equation used these days is:-
The constant 0.102 accounts for the present tabulation of enthalpies in kJ mol-1, rather than the electron-volts of Pauling's time [1 ev = 96.5 kJ mol-1 and 1/96.5 = (0.102)2]. For the molecule ClF the enthalpies required are 255 (ClF), 242 (Cl2), 158 (F2), leading to E = 55 and = 0.76. Notice that this procedure does not say which of A or B is the more electronegative.
A third scale of electronegativities which is now much in favour is the Allred-Rochow scale, about which you will hear later. It is based on calculated values of Zeff using shielding constants: it therefore suffers from the disadvantage that it is not based on experimental measurements, but it has acquired great respectability. It also contains 'scaling constants' so that it covers the same range of values as Pauling's electronegativities.
The result of all these scales is that electronegativities are high for small atoms and/or atoms with high values of Zeff; is therefore larger at the tops of Groups and at the (right-hand) end of Periods. However, there are some irregular variations within the Periodic Table; again, more about those in Module CHM102.