Why Things Are Different Colors

A few years ago, my mother, fascinated by all of the colorful and weird effects of pottery glazes, posited that a chemistry course or two could answer her questions about why things are different colors. In reality, it wasn't until three years into an undergraduate degree in chemistry that this question was fully answered in inorganic and physical chemistry courses.

Oh, don't get me wrong, freshman "general" chemistry covered it, too, but then it was hidden in a thicket of electronic states too abstract for the freshman to fully appreciate. And then we moved on to the concept of equilibrium, which is more important than color but not as colorful.

As sophomores, students take organic chemistry, the chemistry of carbon, and there color leaches back in again, although it was difficult to discern while fighting through the headaches of interpreting magnetic resonance spectra (don't ask (unless you're really interested), it's just a way to figure out what atoms are connected to what other atoms, among other things, and hence what kind of chemical transformations a molecule can undergo). Most organic chemicals, in my experience, are white crystalline powders, white oily powders, or colorless oils. Not necessarily boring, especially for how they smell (mint, bananas, rotting corpses) and how they fly in mass spectrometers (also a way to figure out how atoms are connected with each other, among other things). Some organics are colored, but there are more fun ways to make molecules be colored, and hence to find color in everyday objects.

And even in physical chemistry--that branch that approaches questions about how molecules move, and hence react and interrelate, and describes why things are the way they are--color is buried under mountains of stationary states, Hermitian operators and their eigenspectra, and semiclassical field approximations.

Okay, so you do get "color" in some sense in physical chemistry, usually right at the beginning. It's a problem known as the ultraviolet catastrophe that baffled the physicists of the day, whose solution contains one of the important clues that led scientists to paint over Newton's old white picket fence with the more shocking colors of quantum mechanics. In short: things that are hot emit light. Camp fires, incandescent (not an accidental epithet!) lightbulbs, stars. In the greatest spirit of physics, these kinds of things are idealized as "black bodies," with no inherent color of their own, which is why they're black. According to classical physics, hot things should emit a little bit of infrared and a lot of color. The hotter the black body, the less infrared and the more color. The problem is that classical physics also predicts that black bodies should give up an infinite amount of ultraviolet light, higher in energy than visible light (color) which is higher than infrared. This contradicts actual observations, such as, I don't immediately get a bad sunburn sitting next to a lamp, and the more quantitative observation that black bodies give out a little bit of infrared (low energy), a lot of "color" in the form of visible light (medium energy), and a little bit of ultraviolet (high energy).

It turns out that quantum mechanics was the right way to solve the black body problem, and that classical mechanics was the wrong way. This is exactly the kind of thing that led to the formulation and acceptance of quantum mechanics: it predicted and explained things that classical mechanics could not, and what's more it predicted classical mechanics.

This should give the reader the impression that color is due to quantum mechanics, which is the correct impression. To be more specific, color happens because of the funny ways that atoms behave and consequently how they interact with light. But this is physical chemistry, and physical chemistry does not really explain why things are colored; it just says how they could be colored if they were.

No, the best course to find out about color (other than this article, of course) is inorganic chemistry. (Bad news for the dabbler, but I'm the scientist here: dabblers can just ask me. Or continue reading.) Inorganic chemistry can officially be thought of as the chemistry of everything that isn't mostly carbon (inorganic means not organic which means not made of carbon), but unoffially and more realistically it is the study of metals.

"Metals" are the elements of the periodic table that, in their pure forms, tend to be shiny, solid (at ordinary temperatures, with mercury being the main exception), and conductive, allowing heat and electricity to flow. Combined with other elements--oxygen, nitrogen, carbon, hydrogen, and their ilk--metal atoms often cause color.

The reason that metals in compounds cause color is due to the way they interact with light. Ordinary atoms are made up of a positively charged nucleus, which is as small as a point, and a collection of negatively charged electrons whizzing here and there around the nucleus, attracted to it because of the opposite charge, but never getting there, as far as anyone can tell, because no one, least of all the electrons, is sure how much of an electron's energy is divided between "I'm whizzing around" (AKA kinetic energy), and "I want to go towards the nucleus" (AKA potential energy). Funny little things, electrons.

In fact, it's this uncertainty that leads, indirectly, to color: atoms can only absorb certain wavelengths of light. (This is very important, and if you take nothing else from this article, take the fact that atoms can only absorb certain wavelengths of light.) The closer to the nuclei that electrons are, the higher energy the light has to be in order to be absorbed. This explains why organic liquids tend to be clear: their electrons are all very close to the nuclei, so they absorb high-energy light, ultraviolet, that you can't see anyway, and let all the light you can see pass through.

The difference with metal atoms is that they have electrons that are futher from the nucleus than in many nonmetals. These far-from-the-nucleus electrons interact with lower-energy light--visible light, the light that makes things colored--absorbing some portions of it. What light gets absorbed depends on the chemical details: dichromate (chromium), for instance, absorbs blue light, so it looks yellow; hemoglobin (iron) absorbs blue and yellow, so it looks red; cupric ions (copper) absorb red and yellow, so solutions of cupric ions look blue.

It all has to do with which electrons are where, in the particular compound the atoms are in, which depends on chemical details like what other atoms are connected to the metal atoms. Other atoms tend to push and pull electrons from the metal atom, and the metal atoms tend to push and pull electrons from the other ("ligand") atoms. The particular balance of how that pushing and pulling is done determines where the electrons tend to be--near the metal nucleus, where they interact with ultraviolet light, or far where they interact with visible light.

Organic compounds (that is, not containing metal atoms, but the distinction becomes fuzzy in lots of places) can have these kinds of effects, as well, but in order to get electrons that are "further from the nuclei" requires tricks like conjugation, which spreads electrons around, making them easier to move and hence to interact with visible light. This is why rhodopsin (in mammalian eyes) is colored--biology has designed a molecule to interact with visible light, which enables seeing, and it does this by making conjugated compounds. This is related to the concept of "unsaturation" that everybody is familiar with in conjunction with dietary fat. Some fats are very unsaturated; most colored organic compounds are much more unsaturated (=conjugated) than that.

In order to understand the details of all the pushing and pulling that makes things different colors, quantum mechanical calculations are required. These are not too hard, but not too easy, either. Four years of a chemistry degree should do the trick. The simpler answer, if you wonder why something is colored, is that electrons interact with light, and the further from atomic nuclei the electrons are--the squishier they are, the more they move around--the less ultraviolet they absorb, and the more visible light. No degree required, just imagination: which better serves the potter and the artist, anyway.