4.5.4 Resonance structures
Gaseous oxygen occurs as O2 molecules. But ultraviolet light or an electric discharge converts some of the oxygen to ozone (Box 6). This has the molecular formula O3.
Box 6: Ozone is blue
Many people know that gaseous ozone in the stratosphere protects us from harmful solar radiation, and that at low altitude it is a source of photochemical smog. But few know that the gas is blue. In the laboratory, ozone is made by exposing O2 gas to an electric discharge. This yields oxygen containing only 10-15% ozone, and the colour of ozone is then almost imperceptible. But if the gas is passed through a vessel immersed in liquid O2, it condenses to a liquid mixture of O2 and ozone; this is cornflower blue. If the liquid is kept cold, a vacuum pump will suck the more volatile O2 out of it, and the liquid soon separates into two layers. The upper layer is deep blue, and is a 30% solution of ozone in liquid O2. The lower layer is 30% O2 in liquid ozone, and has a dark violet colour.
Continued pumping on the lower layer eventually leaves pure liquid ozone, with a deep indigo colour and a boiling temperature of −112 °C (Figure 34). Evaporation normally leads to a violent explosion caused by the decomposition reaction:
2O3(g) = 3O2(g) (5.4)
However, clean procedures that exclude dust and organic matter allow slow uninterrupted evaporation. The product is a deep blue gas, which is almost 100% ozone.
A Lewis structure for the O2 molecule is shown in Figure 35a. For ozone too, a Lewis structure can be written which gives each atom a noble gas shell structure (Figure 35b); Figures 35c and 35d give the corresponding structural formulae with alternative representations of the dative bond.
Do Figures 35c and d suggest that the lengths of the two bonds in the ozone molecule should be equal or unequal?
Unequal; one is a double bond and the other a single dative bond. We would expect this difference to influence their lengths.
But experimental measurement shows that both bonds have the same length (127.8 pm). To account for this, we note that the structures shown in Figures 35c and d have companions in which the double and single bonds have simply been exchanged. Figure 36, for example, shows Figure 35d and its partner. The real structure of the molecule with its equal bond lengths is a sort of average of the two. In situations like this, where a molecule is not adequately represented by a single Lewis structure and seems like a composite of two or more, the competing structures are written down and linked by a double-headed arrow, as shown in Figure 36.
The two structures are called resonance structures, and the real structure of ozone is said to be a resonance hybrid of the two. The significance of the representations in Figure 36 is that in ozone, each bond is a mixture of one-half of a double bond and one-half of a single dative bond. Note that Figure 36 is not meant to imply that the molecule is constantly changing from one resonance structure to the other. It is a hybrid in the same sense that a mule is a hybrid: it does not oscillate between a horse and a donkey.
To clarify this, we turn to benzene, C6H6. Like ozone, it can be represented as a resonance hybrid of two resonance structures in which all atoms have noble gas configurations (Figure 37).
A typical C—C single bond length in an alkane hydrocarbon such as ethane, C2H6 (Structure 5.21), is 154 pm; in contrast, a typical C—C bond length in an alkene hydrocarbon such as ethene, C2H4 (Structure 5.22), is 134 pm. The individual resonance structures in Figure 37 therefore suggest that the carbon-carbon bond lengths in benzene should alternate between about 134 pm and 154 pm around the ring.
But what does the whole of Figure 37 suggest?
The real structure of benzene is a hybrid of the individual structures, and each carbon-carbon bond will be a mixture of one-half single and one-half double bonds; all carbon-carbon bond lengths should be equal and lie between 134 and 154 pm.
This is precisely the case: all carbon—carbon bond lengths in benzene are 140 pm!
Number the carbon—carbon bonds in a benzene ring of Figure 37 clockwise from 1-6. All bonds contain at least one pair of electrons. However, in one of the resonance structures, bonds 1, 3 and 5 are double bonds, each containing a second electron pair; in the other resonance structure, the double bonds and extra pair of electrons are found at bonds 2, 4 and 6. The implication of Figure 37 is that in the resonance hybrid these three extra pairs of electrons are not confined to, or localised within, just half of the bonds in the ring. Instead, they are delocalised around the ring and equally shared within all six bonds. Although, in this course, we shall draw benzene and its derivatives as a single resonance hybrid (Structure 5.23), remember that this delocalization makes the bond lengths in the ring equal, contrary to the implications of Structure 5.23.
We conclude with a resonance hybrid which is an ion. Structure 5.20 suggests unequal bond lengths in the carbonate ion. In fact, X-ray crystallography of carbonates suggests that all three bond lengths are equal, at about 129 pm; standard values for C—O and C=O bond lengths are around 143 pm and 120 pm, respectively. Three resonance structures, all equivalent to Structure 5.20 contribute to a resonance hybrid that accounts for the bond length (Figure 38).