2012: Advent Chemistry: Benzonitrile

by on December 7, 2012

A flat drawing of the structure of benzonitrile

This is benzonitrile, a benzene ring (remember those?) with a nitrile group attached. Hermann Fehling discovered it in 1844. It smells of almonds, and is the substance which gave its name to all “nitrile” compounds. Benzonitrile is an excellent example of the kinds of thing that can happen when elements besides hydrogen interact.

The benzene ring, as we’ve previously discussed, is drawn as a series of alternating single and double bonds, but isn’t’ really. It’s a ring of single (sigma) bonds, with a ring of pi electrons – the double bonds – above and below.

A covalent bond has to fulfil certain conditions before it can form. Both atoms must either have electrons available for it or, more usually, be able to break one of the bonds they’ve already got and use those electrons for the new bond. There must be the right kind of symmetry and overlap between the electron orbitals that will make the bond. And there can never, ever, be two electrons of the same kind in the same place at the same time.

“Of the same kind” is because electrons come in two varieties, spinning upwards and spinning downwards. An up electron and a down electron can share an orbital, and the structure of atoms is based on electrons pairing off like this. One electron alone is lonely, and wants to form a bond.

Bonds themselves are really just orbitals, ones that involve two atoms. And each bond can therefore only hold two electrons – one up electron, and one down electron.

This is called the Pauli Exclusion Principle, and it’s the reason there’s a difference between the sigma and pi bonds in benzene rings. You can’t have two sigma (single) bonds between the same pair of atoms, because then you’d have two up electrons and two down electrons in the sigma bond orbital and that’s impossible. So the second pair of electrons have to do something else. They become a pi bond, so now they’re in a different orbital from the sigma bond electrons, which is allowed, and we get a second bond – a double bond.

Double bonds have a more complicated shape, with part of the orbital going above the sigma bond and part below. In benzene rings, all the upper parts overlap around the circle, and so do all the lower parts, and the pi electrons whizz around the ring and hold everything together.

So we know what single bonds and double bonds are, but what’s that bit on the end, with three bonds between the carbon and the nitrogen?

That’s a triple bond, and it’s actually really simple now that you know how double bonds work. A triple bond is just another pi bond, but at right angles to the first. So between those two atoms there’s a sigma bond, a vertical pi bond, and a horizontal pi bond.

Of course, an atom has three p orbitals, all at right angles to each other, and each of them can form a pi bond. And after the p orbitals, there are five d orbitals, to form delta bonds with, and then orbitals can hybridise to create bonds with some sigma chatacter and some pi character… So when you get more complicated than hydrogen, you can get really complicated. Tomorrow we’ll look at something we haven’t seen before, but which is much easier than all this sigma-pi stuff: ionic bonds.

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