Metals: An Introduction

When I first registered for Inorganic Chemistry, I was a little psyched. Since I’d been in high school, the majority of what I studied (not just in chemistry, but in science in general) had been biology related, and rightfully so, since I’m a biochem concentration. However, Inorganic was a chemistry major thing, not a biochem concentration thing, so I knew that taking this class would fill in some holes in my knowledge of chemistry quite nicely. Yes, I’ve seen quite a lot of carbon, thank you. Yes, water exists. Oh wait, that’s platinum in that cataly—oh, we’re just going to gloss over it. Okay.

Long story short, I wanted to learn stuff about metals, dangit, and organic didn’t deliver.

Our first section of Inorganic was, as I previously stated, an effective “Gen Chem III” course. As such, we reviewed lots of interesting things, including Localized and MO Bonding Theory, which were cool and pretty useful. Now, though, we’re doing Actual Inorganic Things, and one of the things we’ve started studying are—you guessed it—metals.

Well, actually, we’re studying semiconductors in depth, but that requires a brief understanding of how metals work, so…

(And yes, if you’ve had inorganic chem before, I’m skipping the crystal lattices right now. Why, you ask? Uh???)

Metallic solids have several characteristics intimately associated with them. They’re distinctively lustrous, conduct both electricity and heat, and are malleable and ductile. So, why is that? Well, turns out, we can explain the properties of metals with something called “band theory,” which is an offshoot of our friend, MO theory.

If you make molecular orbitals for a diatomic association of two metals, such as lithium, you’ll end up with your 2s-orbitals being turned into two orbitals: a bonding and an antibonding orbital. Since lithium has one valence electron, two lithiums will put their two valence electrons into the bonding orbital, leaving the antibonding one empty.

Now extend this to Li3. You get three orbitals, with two electrons in the bottom one and one in the middle. That leaves the highest orbital plus half of the middle one empty. Li4? Four orbitals, two filled with two electrons each and two empty. I think you can see where this is going.

Yup, turns out, for a metallic solid Lin, half of the total number of molecular orbitals will be filled, while the other half will be empty. That’s all well and good, but how does that relate to conductivity?

Well, that’s where the “bands” part comes in. You see, when you put so many of these atoms together, the energy levels of the orbitals are so similar that we think of them as blending together into bands. Thus, rather than saying it has half its orbitals filled, we would say that lithium has a half-filled band. What’s the significance of that? You got it! Because the band is half-filled, there’s room for electrons to, as my professor says, “be sloshed in from some end of the band.”

The same can be said for all metals: they all have partially-filled bands that can accept electrons. Therefore, all metals are conductive.

Ah, I bet you’re wondering about what appears to be the exception, the Group 2 Alkaline Earth Metals. If you understand this stuff pretty well, you’re probably sitting there thinking, “But won’t the molecular orbitals for magnesium, for example, be completely full? Couldn’t I stick a magnesium wire in an electrical socket and be fine?”

(Okay, so maybe I completely stole that example from my professor. It continues like this: “You try that, and then you get electrocuted, and the person who comes and finds your dead body goes, ‘Hm, I guess magnesium does conduct electricity after all.'”)

So the answer to that would be no. Why? Well, it turns out that, even though the molecular orbitals from the 2s orbitals are completely full in the alkaline earth metals, the orbitals from the p-orbitals overlap with them quite nicely, giving extra electrons plenty of room to move around.

Shiny, isn’t it? ([is shot]) I thought so too!

So that’s it for your intro to metals! Not too bad, right? Now we can apply it to a more interesting study: a study in semiconductors.

Questions? Comments? Put ’em below. 😉


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