Justin

Using Computational Analysis to Understand Colors of Compounds We See

Introduction

Different compounds have different colors that we see or may not see. The colors of compounds we see are due to some characteristics of the compounds that help that particular compound reflect a particular color of light due to the light it absorbs.

The diagram above shows the energy levels that electrons are promoted to when they absorb light in a particular compound. The diagram below shows the possible electron jumps when light passes through a compound.

In each possible case, an electron is excited from a full orbital into an empty anti-bonding orbital. Each jump takes energy from the light, and a big jump obviously needs more energy than a small one. Each wavelength of light has a particular energy associated with it. If that particular amount of energy is just right for making one of these energy jumps, then that wavelength will be absorbed - its energy will have been used in promoting an electron.

Basically, if the compound in question needs a higher energy jump, you will need to absorb light that has a high frequency (E=hv) which in turn would show at a lower wavelength in the UV-visible absorption spectra (lambda=c/v).

In terms of colors the Uv-vis spectrum we would be interested in would be the wavelengths that range from 380 to 700 nanometers because those are the wavelengths that the human eye can see. In this case the jumps that we would concern ourselves with would be;

pi bonding orbitals to pi anti-bonding orbitals; non-bonding orbitals to pi anti-bonding orbitals; non-bonding orbitals to sigma anti-bonding orbitals.

That means that in order to absorb light in the region from 200 - 800 nm (which is where the spectra are measured), the molecule must contain either pi bonds or atoms with non-bonding orbitals.