Bonding and Electronegativity 2
Looking back at types of bonding: we left it with consideration of BF 3 which did not have an 'octet' of electrons around boron, only six electrons; such a bonding mode is sometimes called "electron deficient". Compare that with ammonia (NH3) which has a 'lone pair', and an octet around nitrogen.
If BF3 and NH3 are put together (reacted), then boron's deficiency can be satisfied with nitrogen's lone pair by the formation of a donor or co-ordinate bond. with both electrons originating from one source (the nitrogen). Once formed, the resulting bond is like any other covalent bond, the idea of a 'source' for the electrons is again only our accounting procedure. Such a complex is called a 'donor-acceptor' or Lewis base, Lewis acid complex.
The existence of charge separation across a bond means that most bonds have dipoles (except homonuclear diatomic molecules like Cl2). Whether a whole molecule has a dipole depends on its geometry. We shall shortly see that BF3 is flat (trigonal planar) so that the bond dipoles cancel out. However, NH3 is not flat (because of the lone pair on nitrogen), so the bond dipoles have a net resultant, and the molecule has a dipole.
A knowledge of electronegativities, and the polarities of bonds, permits a definition of the Oxidation State of an atom in a molecule. The definition is "the formal charge remaining on an atom when all shared electron pairs are assigned to the more electronegative of the atoms sharing them". For example in ammonia NH3 the polarities are d+ on the hydrogens and a larger d- on the nitrogen. All the N-H shared pairs are assigned to the nitrogen, leaving a formal 3- charge on N and a formal 1+ on each H. These are not actual charges (we are not dealing with ions here); this is sometimes made clear by using Roman numerals of +III and -I. In H3 N-->BF3 we have 3H+I / N-III / B+III / 3F-I. Other examples of oxidation states will be given as they are met in subsequent chemistry.
So far, all covalent bonds have (by implication) been sigma bonds, that is ones in which the shared electrons are concentrated along the inter-nuclear line joining the atoms which share them. We should look how multiple bonds are displayed.
Molecular oxygen (di-oxygen - O2) can be written with four electrons shared - two from each oxygen: in this way each oxygen atom achieves an octet, and seemingly forms one sigma bond and one pi () bond - in total, a double bond (bond order = 2).
In molecular nitrogen (di-nitrogen - N2) each nitrogen achieves an octet by sharing three of its electrons - six shared in total leading to a triple bond [one sigma () and two pi ()] (bond order = 3).
These bond orders of O2 and N2 are the same as in alkenes and alkynes - indeed such molecules are iso-electronic (O2 with ethene, and N2 with ethyne).
However, in the case of di-oxygen, this is only part of the story, since O2 is paramagnetic - which means that the electrons are not all paired! We need a better theory to explain that! Before we address that problem, we must develop a theory which allows us to decide on the shapes and geometries adopted by molecules, and to prove that BF3 really is flat, but the NH3is not.
The Lewis approach 'surrounds' a 'central' atom with pairs of electrons. What geometry do these electron pairs adopt? VSEPR theory (also known as Nyholm-Gillespie theory) proposes that electron pairs mutually repel, and that the basic molecular geometry seen is the result of that repulsion. So let us investigate some of the simple options (the theory will be further developed in other modules). The diagrammes below (thanks to Dr Mark Winter) serve to illustrate five of the many possible geometries that one can have as the number of electron pairs increase. The electron pairs are represented by large dots centralised around the centre of a sphere.
As an illustrative example of what the above diagrams mean I have included two molecules. Firstly methane (CH4) which shows a tetrahedral geometry and secondly boron trifluoride which exhibits trigonal planar geometry.
The table below contains further information on some of these structures as well as providing examples of molecules that adopt these geometries in a non theoretical enviroment.
| 2 electron pairs | e.g. BeH2180° angular separation. | Linear geometry |
|---|---|---|
| 3 electron pairs | e.g. BF3 120°angular separation. | Trigonal geometry
|
| 4 electron pairs | e.g. CH4 109.47°angular separatioN. | Tetrahedral geometry |
Note how a tetrahedron can be drawn on a cubic framework: it allows easy calculation of the 'tetrahedral angle' [2xsin-1(half face-diagonal)/(half body-diagonal) = 2xsin-1(
2/3)].
These examples all imply that the electron pairs are equivalent - indeed they were all 'bonding electron pairs'. However ammonia has three bond pairs and one lone pair, and many of the Lewis structures contain one or more lone pairs. We need to ask if they are equivalent in their 'space requirements", that is, whether they repel to the same extent at the same angular separation. If, in a molecule ABn, A also carries some lone pairs, we can display their angular distributions as follows:-
The angular space required by a lone pair is greater than the angular space required by a bond pair. Thus, at a given angle, a lone pair-lone pair interaction is greater than a lone pair-bond pair interaction which is greater than a bond pair-bond pair interaction (LP-LP > LP-BP > BP-BP). Thus the positioning of lone pairs is the more critical, and they tend to push bond pairs closer together than in the ideal co-ordination polyhedon.
For example, in ammonia with one lone pair and three bond pairs, the observed H-N-H bond angle is less than the ideal tetrahedral angle (at 107°), due to the greater repulsive influence of the lone pair: the unobservable LP-N-H angles would be greater than this ideal tetrahedral angle. In water (H2O), which has two each of lone pairs and bond pairs, the H-O-H angle is further reduced from tetrahedral (to 104.5°) by the greater repulsive interactions of both LP-LP and LP-BP.
Interesting comparisons can be made if a series of related molecules with varying 
values are considered, for example, the trihydrides of the Group 15 elements MH3 (M is N, P, As, Sb, Bi).
As M becomes less electronegative (on descending the Group), the polarity of the bond reduces, and then reverses, so the maximum in the shared electron density of the bond pair moves progressively further out from the Group 15 atom (in fraction of an increasingly long bond); so the angle subtended at the Group 15 atom diminishes.
The progressively diminishing H-M-H angle confirms this, as the repulsive power of the lone pair exerts more influence.