- Related Topics:
- theory of resonance
- chemical association
- electronegativity
- bond energy
- bond length
There is a sharp distinction between ionic and covalent bonds when the geometric arrangements of atoms in compounds are considered. In essence, ionic bonding is nondirectional, whereas covalent bonding is directional. That is, in ionic compounds there is no intrinsically preferred direction in which a neighbour should lie for the strength of bonding to be maximized. In contrast, in a covalently bonded compound, the atoms adopt specific locations relative to one another, as in the tetrahedral arrangement of hydrogen atoms around the central carbon atom in methane, CH4, or the angular arrangement of atoms in H2O.
The lack of directionality of ionic bonds stems from the isotropy (spherical symmetry) of the electrostatic forces between ions. As has already been pointed out, the result of this isotropy is that ions stack together in the locations necessary to achieve the lowest energy and in this way give rise to the common packing patterns characteristic of many ionic solids. When deviations from stacking schemes are observed that seem to indicate that the ions are being held in certain orientations relative to their neighbours, it is a sign that covalent bonding is beginning to influence the structure of the solid and that the bonding is not purely ionic. This is the case, for example, in the compound nickel arsenide (NiAs), which has a structure that suggests that a degree of covalent bonding is present (). It is fully apparent in the structure of diamond (), in which each carbon atom is in a tetrahedral position relative to its neighbour and in which the bonding is essentially purely covalent.
The rationalization of the structures adopted by purely ionic solids is essentially a straightforward exercise in the analysis of electrostatic interactions between ions. The problem of the structures of covalent compounds, both individual molecules, such as methane, and covalently bonded solids, such as diamond, is much more subtle, for it involves delving into the characteristics of the electron arrangements in individual atoms. Thus, if the formation of a covalent bond is regarded as corresponding to the accumulation of electrons in a particular region of an atom, then, to form a second bond, electrons can be accumulated into only certain parts of the atom relative to that first region of enhanced electron density. As a result, the bonds will lie in a geometric array that is characteristic of the atom. The remainder of this section focuses on this problem, but a detailed quantum mechanical analysis is required for a full understanding of the matter.
The theory of molecular shape known as valence-shell electron-pair repulsion (VSEPR) theory grew out of Lewis’s theory, and, like that approach to bonding, VSEPR focuses on the role of electron pairs. It stems from the work of the British chemists H.M. Powell and Nevil V. Sidgwick in the 1940s and was extensively developed by R.J. Gillespie in Canada and Ronald S. Nyholm in London during the 1960s. As such, it postdates quantum mechanical theories of bonding and shape but should be seen (as is so common a motivation in chemistry) as an attempt to identify the essential features of a problem and to formulate them into a simple qualitative procedure for rationalization and prediction.
A Lewis structure, as shown above, is a topological portrayal of bonding in a molecule. It ascribes bonding influences to electron pairs that lie between atoms and acknowledges the existence of lone pairs of electrons that do not participate directly in the bonding. The VSEPR theory supposes that all electron pairs, both bonding pairs and lone pairs, repel each other—particularly if they are close—and that the molecular shape is such as to minimize these repulsions. The approach is commonly applied to species in which there is an identifiable central atom (the oxygen atom in H2O, for instance), but it is straightforward to extend it to discussions of the local shape at any given atom in a polyatomic species.
Applying VSEPR theory to simple molecules
The methane molecule, CH4, can be used to illustrate the procedure for predicting molecular shape. The Lewis structure of this molecule ascribes four bonding electron pairs to the carbon atom (). These pairs repel one another, and their separation is maximized if they adopt a tetrahedral disposition around the central carbon atom. A hydrogen atom is attached by each bonding pair, so it can be predicted that CH4 is likely to be a tetrahedral species, which is in fact the case.
When applying VSEPR theory, attention is first focused on the electron pairs of the central atom, disregarding the distinction between bonding pairs and lone pairs. These pairs are then allowed to move around the central atom (at a constant distance) and to take up positions that maximize their mutual separations. As in the methane molecule, four pairs adopt a tetrahedral disposition. The arrangements adopted by two through six pairs are summarized in the table. At this stage, the atoms that are attached by the bonding pairs are introduced, and the shape of the molecule is reported on the basis of the arrangement of these atoms.
The water molecule, H2O, provides a simple example. The oxygen atom has four electron pairs, so these pairs adopt a tetrahedral arrangement. Two of the pairs are bonding, and hydrogen atoms are attached to them. Hence, the molecule is angular. (Note that the shape of the molecule is determined by the disposition of the atoms, not the disposition of the electron pairs.) The ammonia molecule, NH3, has four electron pairs in a tetrahedral arrangement around the nitrogen atom; three of these pairs are used to bond hydrogen atoms, so the molecule is predicted to be trigonal pyramidal, with a lone pair in the apical position. Some of the names of the shapes of simple molecules are summarized in the .
The angle between electron pairs in a tetrahedral arrangement is 109.5°. However, although H2O is indeed angular and NH3 is trigonal pyramidal, the angles between the bonds are 104° and 107°, respectively. In a sense, such close agreement is quite satisfactory for so simple an approach, but clearly there is more to explain. To account for variations in bond angle, it is supposed that electron pair repulsions are greatest between lone pairs, less between lone pairs and bonding pairs, and least between bonding pairs. The justification of this ordering has proved somewhat elusive; qualitatively it is presumed that lone pairs, being attached only to a single centre, spread over a greater volume than bonding pairs, which are pinned between two attracting centres. Whatever the reason may be, the order correlates quite well with observation. Thus, in H2O the two lone pairs move apart a little, and the two bonding pairs move away from them by closing the angle between one another. Likewise, in NH3 the three bonding pairs move back from the single lone pair to minimize their interaction with it. As a result, the H―N―H bond angle decreases slightly. In each case, the predicted angle is less than the tetrahedral angle, as is observed experimentally.
VSEPR theory is quite successful at predicting (or at least rationalizing) the overall shapes of molecules. Thus, the hypervalent species SF6 (sulfur hexafluoride), with six bonding pairs, is predicted and found to be a regular octahedron, and PCl5 (phosphorus pentachloride), with five bonding pairs, is predicted and found to be a trigonal bipyramid. The XeF4 (xenon tetrafluoride) molecule is hypervalent with six electron pairs around the central xenon (Xe) atom. These pairs adopt an octahedral arrangement. Four of the pairs are bonding pairs, and two are lone pairs. According to VSEPR theory, the repulsion between the lone pairs is minimized if they lie on opposite sides of the xenon atom, leaving the four equatorial pairs as bonding pairs.
This analysis suggests that XeF4 should be a planar species, which is found to be the case.
Molecules with multiple bonds
There are further rules in VSEPR theory that simplify the discussion of species with multiple bonds and of species in which resonance must be considered. An analysis of the shapes adopted by species with multiple bonds suggests that each multiple bond can be treated as a single “superpair” of electrons. This rule can be justified by considering the geometric shapes that stem from two atoms sharing two or more pairs of electrons (). Thus, the sulfate ion, SO42−, for which a Lewis structure is
can be treated as having the equivalent of four pairs (two ordinary pairs and two superpairs) around the sulfur atom in a tetrahedral arrangement. All four pairs are bonding, so the ion is predicted to be a regular tetrahedron, which it indeed is. The same conclusion about the shape of the molecule would be drawn from another possible Lewis structure, in which each bond is single:
The actual molecule is a resonance hybrid of these and related structures; but, as each one corresponds to the same geometry, no particular Lewis structure need be selected before one can make a prediction based on VSEPR theory. In other words, resonance does not affect the shapes of molecules.