Simple way to understand why chromophores often have double bonds?

Chromophores — atoms or groups of atoms within a molecule that absorb some visible wavelengths better than others — result in compounds that have color. For example, see this answer about azo dyes, and this mention of bleaching stains by oxidation of double bonds.

Without going too deep into quantum mechanics and/or molecular orbital theory (MO), is there a way to understand in a qualitative way why chromophores are often related to the presence of double carbon or double nitrogen bonds?

• – Klaus-Dieter Warzecha Feb 5 '17 at 7:24
• @KlausWarzecha I've been admiring those answers. Here I'm asking for something a little simpler to start with. Rather than a long discussion of the HOMO-LUMO gap or wavelength shift, I'm asking what it is about a double bond that gives it a wavelength dependent photon cross-section (i.e. color) to begin with, as compared to a single bond. – uhoh Feb 5 '17 at 7:43

Now in π orbitals, the atomic p orbitals are oriented perpendicular to the bond axis, so any overlap is doomed to be much smaller than the corresponding single bond’s σ overlap. Therefore, the energy difference from the bonding to the antibonding π orbital is not far into the ultraviolet region; typical values would correspond to a wavelength of $150$ to $200~\mathrm{nm}$. Extending the π network will move those orbitals closer together energetically (again, the complete answer requires much more MO theory) which is why benzene has an energy difference corresponding to a wavelength of $250~\mathrm{nm}$. Thenceforth, it is not a long way until you reach visual wavelengths ($400{-} 700~\mathrm{nm}$).