Strike Up the Band

Welcome back Science Squaddies! We're almost done with our journey through the wild world of color. Bear with us, as we're almost done. I know, I know, what will you fill your time with when it's done? I'm sure you'll find something. Oh wait, more glorious science posts from your's truly! Huzzah!

Today we're talking about band theory and how it affects the colors of metals. Band theory explains a ton of things about metals, including electrical and thermal conductivity, but we won't be discussing that. We're strictly about the color, yo.

Imagine, if you will, 2 copper atoms just hanging about, doing their thing. As long as the atoms stay far enough apart from each other, the electrons of each atom occupy specific energy levels within the atom's orbital levels. When our two copper atoms get closer together, the two separate atoms become one 2-atom system, which, more importantly, means that we now have twice as many electrons that have to be arranged. In order to occupy all of these new electrons, the orbitals of each isolated atom have to split. Two atoms, mean twice as many orbitals. 5 atoms means 5x as many orbitals. Once you have enough atoms together to make up a solid, the number of orbital levels gets to be amazingly huge, which means that the energy difference between the levels is amazingly small. To quote Halliday and Resnick, "in this way, each level of the isolated atom becomes a band of levels in the solid". It is the movement of electrons between and within these bands that help define things like electrical conductivity and color.

Each metal has what is called the Fermi Level. At a temperature of absolute zero, the Fermi Level, which occurs mid-band by the way, is the highest occupied level in a metal. Above the Fermi Level, there are no electrons, below it, are all sorts of electrons. Because this level is mid-band, electrons can move above the Fermi Level when agitated. When an electron jumps from below this level, to above the level, you now have a negatively charged electron above the level and a postively charged hole below the level. These two particles, for lack of a better word, move in opposite directions and lo, we have an electric current.

When light strikes a metal, the electrons on the surface of the metal absorb the light as energy, become agitated and jump levels. By absorbing this light, alternating electric currents are generated on the metal's surface, by the method described above. These currents immediately cause the light to be re-emitted (remember that if an electron absorbs energy to jump a level, it must emit the same amount of energy to go back to its original level) and this causes the shiny surface you see on metals. The color of metal is a result of how efficient the light absorption is at all the energy levels. If all of a metal's energy levels absorb light with equal efficiency, the light will all be re-emitted and you get a silvery polish. Gold, for example, loses efficiency as the energy levels increase, causing an absorption of light at the blue end of the spectrum, hence the yellow color. Ditto for copper and its resulting orange color.

See, it's simple! And here you were all worried for nothing. Next week is the last topic on this discussion and is pretty damn big. I don't think I'll have to split it into 2 parts, but I'll wait and see. Try and contain your excitement.

Sources:
WebExhibits - Causes of Color
Wikipedia - Band Structure
Wikipedia - Fermi Surface
Georgia Mineral Society, Inc - Color in Minerals - Doug Daniels
Fundamentals of Physics, 4th Edition - David Halliday, Robert Resnick, Jearl Walker, 1993, John Wiley & Sons, Inc.