Colour is light that didn't make it through
A red apple is not adding red to the world. White light, which contains every colour, falls on the apple; the skin *absorbs* the greens and blues and lets the red bounce back to your eye. The colour you see is whatever is *left over* after some colours have been taken away. So colour is a story about absorption: a substance looks coloured precisely because it swallows certain colours of visible light and not others.
An instrument can do this more carefully than your eye. UV–visible spectroscopy shines a beam through a sample, sweeping smoothly across the ultraviolet and visible parts of the electromagnetic spectrum, and records how much light is absorbed at each wavelength. The result is an absorption curve — usually a few broad humps showing exactly which colours the sample drinks in.
Electrons jumping between levels
Ultraviolet and visible photons are fairly energetic, and what they move inside a molecule are its electrons. Recall the ladder of molecular energy levels: for electrons, the rungs are widely spaced, so it takes a strong, high-energy photon to bump an electron from a low rung to a high one. UV and visible light carry just about the right energy for these big electron jumps — which is why this technique reports on a molecule's electrons specifically.
Not every part of a molecule absorbs visible light, though. The piece that does — a group of atoms whose electrons have a gap small enough to catch a visible photon — is called a chromophore, from Greek words for *colour-bearer*. The orange of a carrot, the deep red of blood, the green of a leaf: each comes from a specific chromophore. Change the chromophore and you change the colour; that is exactly how dyes are designed and how indicators flip colour when acidity changes.
Turning colour into a number: the Beer–Lambert law
Now for the practical magic. Pour ink into a glass of water and it looks faintly grey; pour in twice as much and it looks twice as dark. Make the glass twice as wide so the light travels through twice the liquid, and again it looks darker. These two intuitions — *more stuff* and *longer path* — are captured exactly by the Beer–Lambert law. It says the amount of light absorbed grows in step with both the concentration of the absorbing substance and the distance the light travels through it.
Chemists measure absorption as a quantity called absorbance. The law states it plainly: absorbance equals a constant for that substance, times the concentration, times the path length. Written compactly, A = ε·c·ℓ, where ε (epsilon) is how greedily that chromophore drinks light, c is concentration, and ℓ is the width of the container. The beauty is that absorbance rises in a straight line with concentration — double the concentration, double the absorbance.
- Make a few solutions of known concentration and measure the absorbance of each.
- Plot absorbance against concentration — the points fall on a straight line through the origin. This is your calibration line.
- Measure the absorbance of an unknown sample, find where it lands on the line, and read off its concentration.
Where this shows up in real life
This little law does an enormous amount of work. A hospital blood-oxygen clip on your finger uses two colours of light because oxygen-rich and oxygen-poor blood are different shades of red — pure Beer–Lambert at work. Water-quality labs measure pollutants by adding a reagent that turns the sample a colour, then reading how deep the colour is. Biologists count cells and measure DNA the same way. Whenever you hear that something was measured 'by absorbance', this straight-line relationship is doing the counting.