Postponed, but not forgotten

Happy chemistry Wednesday, everyone! Wow, can you believe we’re already at the midpoint of the week? Speaking of midpoints.... The astute of you may have noticed that we apparently skipped a bunch or elements back when we took a look at the transition metals. The trouble crops up in Group III, the first group of the transition metals. Going down the group, we see scandium, yttrium...and then a weird sort of cut-out, as lanthanum and actinium, the next two elements of the group, have been removed and placed below the table. Additionally, there are fourteen additional elements after lanthanum and actinium which have been removed from the table proper. What’s with the excision of these thirty elements? Are they not transition metals?

Look down

Not exactly. Today, we’re going to take a look at the world of the lanthanides and the actinides. Lanthanides are sometimes known by the supremely cool moniker of “rare earth elements” - this isn’t quite accurate, but we’ll get there in a moment. In many ways, this group of elements are similar to the transition metals, given that much of their behavior is the result of a partially filled electron orbital. However, the lanthanides and actinides are marked by a partly filled f orbital, in contrast to the partly filled d orbital of the transition metals. F orbitals are the final orbital in the “s, p, d, f” progression. They only come into play in elements with an atomic weight above 58, and never act as a valence shell. Similar to the transition metals, where s orbitals of the next level fill before the d orbital (it’s really strange, but I think the transition metals post a few weeks ago did a passable job of explaining the whole thing), s orbitals of the next level in the lanthanides and actinides fill before the f orbitals. Actually, several orbitals further out from the f orbitals fill first, leaving the f orbitals buried, and unlikely to be involved in any chemical reactions. 

A theoretical construct, true, but f orbitals are pretty weird-looking

Regardless of the order in which they fill, f orbitals contain a maximum of 14 electrons. As we move across the lanthanide and actinide series, the f orbitals fill in, while the number of valence electrons does not change. Again, as in the transition metals, the lanthanides and actinides are both relatively likely to shed the two electrons in the outermost s orbital. They are also likely to shed an electron from the d-orbital - thus, lanthanides and actinides tend to exist in a +3 ionic state.

Lanthanides and actinides have some interesting and useful chemical properties. We’ll start with the lanthanides and, as you’ll soon see why, magnetism.

This will make sense in a moment, I promise.

While a full examination of magnetism is somewhat beyond the scope of today’s post, I’ll see what I can cover in a paragraph. Remember the EM spectrum? Light, UV radiation, gamma rays - it’s come up a few times already, and no doubt will continue to do so. Now, the “EM” stands for electromagnetic (with “radiation” implied). All of those different kinds of radiation are fundamentally expressions of a coupled electric and magnetic force. EM radiation can be broken down into two opposite but equal waves, one magnetic, and one electric. Interestingly, for much of human history we understood electric and magnetic phenomena as two unrelated forces - it wasn’t until the 1850’s that Maxwell understood that the two were expressions of the same force, and produced the unifying mathematical proof. Now, electric force refers to the behavior of particles carrying an electric charge - positive versus negative, based on the difference in the number of cations and electrons within the atoms making up said particles. Magnetic force operates under the same principle of opposites attracting and likes repelling, but rather than the number of electrons, has to do with the direction of spin on electrons. Yes, electrons spin. Yes, the more you think about the universe, the more surreal it becomes. The behavior of magnetic fields makes for its own fascinating entry, so we’ll stick with permanent magnets. Any element with an incomplete electron shell can exhibit magnetic behavior, and elements that form certain kinds of solid structures can display permanent magnetism - all of the associated electrons spin in the same way, giving the entire structure a stable magnetic orientation. Magnetism can be stable on Earth because the entire planet is covered by a magnetic field (North and South poles, anyone?).  Metals tend to form these kinds of structures, and thus, metals can become stable magnets.

As a consequence of their incomplete f-orbitals and stable valence configurations, certain elements in the lanthanide series can form strong permanent magnets. These rare earth elements can complex with hydrogen and other metal atoms in order to form structures capable of generating an extremely strong magnetic field. If you want a more detailed explanation of how rare earth magnetism works, I invite you to peruse this publication by the Niels Bohr Institute. However, for practical purposes, just keep in mind that combining a rare earth element like neodymium or samarium with a transition element like iron or cobalt makes an unbelievably strong magnet with a small field size - the field is intense, but limited. One could use a rare earth magnet to apply a strong, targeted magnetic field in a situation where you would not want to accidentally expose other areas to magnetism (say, your computer hard drive).

The obligatory photo of a ridiculously strong magnet

For this reason, the rare earth metals are very important, commercially. The good news is, they actually aren’t that rare. Quite a few of the rare earth elements are actually more abundant than that perennial environmental bugbear, lead. The bad news is that the rare earth elements tend to be found in mineral complexes, and need to be painstakingly isolated before being used. It’s worth it, though. In addition to its use in rare earth magnets, samarium (Sa) turns up in the specialized lighting used on film sets, as a radiation-absorber in glass, and even in the flints in lighters. Neodymium (Nd) also turns up in lighters, as well as glass (where it imparts a violet or red tinge), welders’ goggles, artificial rubies and in magnets. Gadolinium (Gd) turns up in microwaves, metal alloys, and cathode ray tubes. Erbium (Er) is also used in metal alloys, as well as fiber optic cables and nuclear reactors. Other rare earth metals are less commercially important, but no less interesting.

The we have the actinides. You don’t hear as much about the actinides because, while all of them are theoretically possible and have existed at some point or another, most of them have only existed for short periods of time inside of supercolliders. Nothing past uranium (U), the fourth element out of the fifteen actinides, occurs naturally in usable quantities. All of the actinides are radioactive. Fans of astronomy will likely be amused by the progression of uranium, neptunium and plutonium. Plutonium, interestingly, while artificially created, is actually fairly stable - while it decays into U, it does through over a period of about 82 million years. Plutonium has been used to power several long-distance spacecraft. Plutonium’s precursor and destination, U, is also used as a power source, and less commonly, as a nuclear weapon. Uranium was also used at one point in the coloring of Fiestaware dishes - similar to the glowing radium watch fiasco of the early 20th century, this form of pigment was retired, as radioactive dinnerware is seldom considered a good idea. 

 If it can do this...

 It probably shouldn't be used for this

So, this neglected bit of the periodic table contains both the ingredients for unimaginably strong magnets, and nuclear weapons. They may be shunted off to the bottom of the table, but it really doesn’t pay to underestimate the lanthanides and actinides.