In general chemistry class, you probably learned that electrons behave like waves. And then, later, when you take organic chemistry, you see resonance for the first time and might think that it’s basically a fancy word for drawing a bunch of silly looking Lewis structures with the electrons in valid but slightly atypical places. The interesting thing is, resonance and the wave property of electrons are two sides of the same coin. In a molecule, when neighbouring atoms keep their p orbitals aligned, an electron’s wavefunction spreads across them, so they don’t simply “sit” in the individual bonds you see drawn on paper. So in reality, molecules are closer to the weighted average of all the Lewis structures that can represent them, and more weight is given to the contributors that keep octets and place charge on better hosts. In practice, it’s not easy to draw the weighted average of a molecule, so we often stick with the highest-contributing Lewis structure as the “representative,” while remembering the real structure is the blend. To be clear, this is not electrons flipping between drawings in time; it’s one delocalised state.

When charge is spread over more space, it lowers potential energy because no single atom has to bear the full charge. It’s a bit like in physics: when you evenly distribute weight, a surface can hold a much larger load than if that weight is concentrated at a point. For example, lying on a bed of nails compared to a single nail. The pressure falls when the area grows. For resonance to work, all the p orbitals in the segment need to sit in the same plane so they can overlap side by side. If the molecule twists and the p orbitals lose alignment, the resonance stabilisation drops. Each p orbital is a member of a team, with no freeloaders. If one turns away, it stops sharing with the next group and the support ends there.

Let’s see an example of this concept in biology. Ordinary amides have a great resonance team: the nitrogen, carbon, and oxygen p orbitals work together to keep electron density smeared among them. As a result, amides are generally not very hungry for more electrons, so most incoming nucleophiles won’t succeed without strong activation. But an interesting exception arises in β-lactam antibiotics. The four-membered ring forces the amide away from planarity with the carbonyl, so the nitrogen’s lone pair can’t participate in resonance very much. The p orbital decided it doesn’t like teamwork. The carbonyl carbon ends up more electron-hungry than in a flat amide, which is why an enzyme’s serine nucleophile can attack it so readily.

As we’ve seen, resonance can hugely impact the reactivity of molecules. As you go further, you’ll meet the ring version of this sharing, called aromaticity, and its opposite, antiaromaticity; in those cases, planarity and the pi electron count decide whether delocalisation is unusually stabilising or, if the pattern is wrong, destabilising. A plain hydrocarbon like cyclopentadiene gives up a proton to strong base more easily than you would expect, because the leftover electrons can spread around a closed loop in the conjugate base, which makes that negative charge unusually comfortable. Turns out, this conjugate base is aromatic. Another beauty is azulene, a hydrocarbon that looks deep blue while its isomer, naphthalene, is colourless. Both molecules are aromatic! Azulene’s fused five- and seven-membered rings keep a continuous path of p orbitals, and from its resonance structures, you can see that there is some extra electron density on the five-membered ring and a slight shortage on the seven-membered ring. Because the rings stay flat and connected, the π electrons are shared across both rings. Extra negative character on the five ring means electrons there are more crowded, which pushes up the energy of the highest filled π level. Slight positive character on the seven ring means it is a better host for added electron density, which pulls down the lowest empty π level. Orange light then has enough energy to promote an electron, so it is absorbed and the molecule appears blue.

In organic chemistry, these initially mystifying cases usually have explanations, and when they don’t, the puzzle is part of the fun.