Transition Metals (Part 2)

Why is copper red? Why is it so soft compared to, say, nickel—the element right next to it in the periodic table? Why is it such a good conductor of electricity?

All of this stems from a violation of Hund’s rules. Let me explain.

In Part 1, I explained the basic math of transition metals. Now I just want to talk about how the first row of transition metals fill up the 10 orbitals in the 3d subshell, and what’s special about copper:

These elements have all the electrons that argon does: that’s the [Ar] in this chart. Most also have two electrons in the 4s subshell: one spin up, and one spin down. So the action mainly happens in the 3d subshell. This has 10 slots for electrons: 5 spin up, and 5 spin down. If you don’t know why 5, read Part 1.

Hund’s rules predict how these 10 slots are filled. It predicts that we first get 5 metals with 1, 2, 3, 4, and 5 spin-up electrons, and then 5 more metals that add in 1, 2, 3, 4, and 5 spin-down electrons. And that’s almost what we see.

But notice: when we hit chromium we get an exception! Chromium steals an electron from the 4s subshell to get 5 spin-up electrons in the 4d subshell. And we get an another exception when we hit copper. Can you guess why?

The short answer is that in every atom, its electrons are arranged so as to minimize energy. Hund’s rules are a good guess about how this works. But they’re not the whole story. It turns out that electrons in the d subshell can lower their energy if we have the maximum number possible, namely 5, with spins pointing in the same direction. (We arbitrarily call this direction ‘up’, but there’s nothing special about the ‘up’ direction so don’t lose sleep over that.

So: Hund’s rules predict that we get 5 spin-up electrons when we hit manganese, with 5 electrons in the d subshell. And that’s true. But the energy-lowering effect is strong enough that chromium ‘jumps the gun’ and steals one electron from the somewhat lower-energy s subshell to put 5 spin-up electrons in the d subshell! So chromium also has 5 electrons in the d subshell.

Similarly, Hund’s rules predict that we get 5 spin-up and 5 spin-down electrons when we reach zinc, with 10 electrons in the d subshell. And that’s true. But the element before zinc, namely copper, jumps the gun and steals an electron from the s subshell, so it also has 10 electrons in the d subshell.

The lone electron in its 4s shell makes copper a great conductor of electricity: these electrons can easily hop from atom to atom. And that in turn means that blue and green light are energetic enough to push those electrons around, so copper absorbs blue and green light… while still reflecting the lower-energy red light!

Similar things happen with the elements directly below copper in the periodic table: silver and gold. Wikipedia explains it a bit more technically:

Copper, silver, and gold are in group 11 of the periodic table; these three metals have one s-orbital electron on top of a filled d-electron shell and are characterized by high ductility, and electrical and thermal conductivity. The filled d-shells in these elements contribute little to interatomic interactions, which are dominated by the s-electrons through metallic bonds. Unlike metals with incomplete d-shells, metallic bonds in copper are lacking a covalent character and are relatively weak. This observation explains the low hardness and high ductility of single crystals of copper”

Copper is one of a few metallic elements with a natural color other than gray or silver. Pure copper is orange-red and acquires a reddish tarnish when exposed to air. This is due to the low plasma frequency of the metal, which lies in the red part of the visible spectrum, causing it to absorb the higher-frequency green and blue colors.”

It would take more work to understand why copper, silver and gold have such different colors! People often blame the color of gold on relativistic effects, but this of course is not a full explanation:

• Physics FAQ, Relativity in chemistry: the color of gold.