Saturday, February 15, 2014

History of the Periodic Table Part 4: Lanthanides and Actinides - Elemental Misfits?

Two series of elements - lanthanides and actinides - make up a group called the inner transition metals, falling between groups 2 and 4 on the periodic table (see the single and double asterisks below).

The border indicates the natural occurrence of the element and the colour of the atomic number indicates its state of matter.

These are screenshots from Wikipedia's "Periodic Table." For clearer larger graphics go to the page.

The lanthanides and actinides are typical metals, even the radioactive elements in these series. They all have a silvery colour.

All the actinides are unstable, making them radioactive. Some actinides are found naturally and in very low concentrations, such as uranium (92), thorium (90) as well as tiny amounts of plutonium (94). The decay and transmutation of uranium also produces transient amounts of actinium (89) and protactinium (91) as well as occasional atoms of neptunium (93), americium (95), curium (96), berkelium (97) and californium (98). The rest of the actinides (Ensteinium (99), Fermium (100), Mendelevium (101), Nobelium (102) and Lawrencium (103), are purely synthetic and can only be made in colliders.

All the lanthanides are stable, except for one element, promethium (61). This synthetically produced element has no less than 38 different radioisotopes, most of which have half-lives of less of than 30 seconds. The other lanthanides are all found in nature and are often called rare earth elements, a group which also typically includes two non-lanthanides - scandium (21) and yttrium (39). These two elements are grouped in with them because they are often found in the same mineral deposits as the lanthanides and they share similar chemical properties with them.

What makes both actinides and lanthanides unique, and what boots them out into their own section, is their electron configurations.

The Electron Configurations of the Actinides/Lanthanides

In order to understand electron configurations, we must first take a look at electron energy shells, subshells and orbitals. The electron shells are labeled K, L, M, N, O, P and Q (or n=1, n=2, n=3, n=4 etc.) as you go from the innermost shell closest to the nucleus outward. Electrons in outer shells have more energy than those in inner shells. The very outermost electrons are the ones that participate in chemical reactions, so their configuration is an especially important indicator of how the element will act chemically.

Each shell is composed of one or more subshells. Subshells are labeled s, p, d, f and g. The s subshell can hold a maximum of 2 electrons (in one orbital), p can hold 6 electrons (in three orbitals), d can hold 10 electrons (in five orbitals), f can hold 14 electrons (in seven orbitals) and g can hold 18 electrons (in nine orbitals). Each orbital can hold a maximum of two electrons, with opposite electron spins. This allows them to obey the Pauli exclusion principle, which means no two electrons in an atom can have identical quantum numbers.

The table below is another Wikipedia screenshot. I've added it to show you how the increasing number of lobes in orbitals allows them to hold more electrons. Click the "table below" link to see a much larger version.

The K energy shell (n=1) is just a tiny sphere, a 1s subshell (see left), so it can hold only two electrons. The L shell (n=2) consists of a higher energy 2s spherical subshell plus three 2p subshells, each one looking like a barbell that is oriented in a different direction. This means that this energy shell can hold 8 electrons in total. The M shell (n=3) can hold 18 electrons in total and the N shell (n=4) can hold up to 32 electrons in total. Each increasing energy shell adds a new orbital:  p, d, f and so on.

An atom GENERALLY fills up electrons in order, so the K shell fills first, then the L shell starts to fill up and then the M shell fills up and so on.

If you look at the periodic table below, you will see group 3B (same thing as group 3): scandium, yttrium, and then going down this group, you see lanthanum and actinium. These lowest two squares are the first elements that begin filling up the d subshell with electrons, just like scandium and yttrium do, as you expect them to. Lanthanum and actinium are also the first two elements of the f-block. In fact, the whole f-block fits into group 3B.

Adapted from Roshan220195;wikipedia)
Scandium's (Z=21) electron configuration is written as [Ar]3d14s2 or 2,8,9,2. Noble gases such as argon (2,8,8) are simply used as writing shortcuts. These notations take a little practice with reading them. In scandium, this notation means that there are 2 electrons in the 1s K shell, 8 electrons in the L shell (2-2s electrons + 6-2p electrons), 9 electrons in the M shell (2-3s electrons, 6-3p electrons and 1-3d electron) as well as two electrons in the N shell (2-4s electrons). The important thing to note here is that some of the N energy shell begins to fill before the M energy shell is filled. Why would this be?

All the f-block elements have just one electron in their d-subshell. Electrons, instead, are filling up the 4f-subshell in the N energy shell. That's what places them all in Group 3 (or 3B). The number of electrons in the outermost shell determines the group. Notice from the table above that the 4f-subshell can hold a maximum of 14 electrons. Each series in the f-block contains, just as you expect, 14 elements. When the 4f-subshell is completely filled, atoms add electrons to the 4d-subshell. This subshell can hold ten electrons, and that's why the d-block is ten elements across. In Group 4, next door, there are two electrons in the d-subshell. In Group 5, there are three electrons in the d-subshell, and so on.

This periodic table shows electron energy shell diagrams (Bohr diagrams) of the elements. It's really useful to count the electrons and test your predictive skills but it is almost impossible to see, so click here to see a greatly expanded version on Wikipedia. If you scroll down, I've captured a screenshot of the first three elements in each of the two series. If you count the energy shells outward you have electrons in the K, L, M, N, O, P (lanthanides) and Q (actinides) energy shells. The O energy shell can hold up to 50 electrons, the P energy shell can hold up to 72 electrons, and the Q energy shell can theoretically hold up to 98 electrons. The largest atom in the current table contains 118 electrons - a partly filled P energy shell. A theoretical (behemoth!) atom having all shells filled including the Q shell would have the atomic number 260, with 260 electrons.

For the larger f-block atoms we use Xenon (2,8,18,18,8) as a shortcut. Lanthanum, like scandium, has one electron in the d-subshell. Its electron configuration is [Xe]5d16s2 or 2,8,18,18,9,2. Cerium, the next element over IN THE LANTHANIDES, not in group 4, also has just one electron in the d subshell, and it is written as [Xe]4f15d16s2 or 2,8,18,19,9,2. Notice that an electron is added instead to the f subshell. If you look at the diagram below, you can see where the new electron is added to cerium - to the N energy shell - and to the 4d-subshell in this energy shell in particular.

Lanthanum (2,8,18,18,9,2) has 1-5d, 2-5s and 6-5p (a total of 9) electrons in the O shell and 2-6s electrons in the outermost P shell. Cerium has exactly the same numbers of electrons in its outer O and P shells as lanthanum does. Next over, praseodymium has the same outermost shell as the other two do, but one less electron in its O shell and two more electrons in its N shell. You can begin to get a hint about why these elements look and behave so similarly to each other. They all share identical outermost electron shells, the shell that most determines their chemical nature. You can see the same trend in the actinides - they all fill an f-subshell first. The f-subshell contains seven orbitals, with each one holding a maximum of two electrons. The lanthanides fill up the 4f-subshell and the actinides below them fill up the 5f-subshell before the 6d energy shell is filled.

All the lanthanide elements have just one d1 electron until you get to hafnium (72) (which is not a lanthanide). Its electron configuration is [Xe}4f15d26s2 or 2,8,18,32,10,2 where the d2 electron is finally filled.

Americium and Curium Offer First Clues to the Existence of the F-block

The first clue that two series of elements should form rows separate from the rest of the periodic table came from Glenn Seaborg (below left) thanks to his work on the Manhattan Project in 1943.

He had found it unexpectedly difficult to isolate two radioactive actinide elements - americium (Z=95) and curium (Z=96). Most americium, a silvery white metal (a small disc under a microscope is shown below right), is produced by the fission of uranium or plutonium in nuclear reactors. Each tonne of spent nuclear fuel contains 100 grams of highly radioactive americium. It is most commonly used in smoke detectors.

Most curium is produced the same way as americium. 1 tonne of spent nuclear fuel contains 20 grams of it. It is one of the most radioactive elements known, emitting radiation so strong that a sample of curium glows purple in the dark, shown below left.

Its most common use is in X-ray spectrometers. Curium, harder and denser than americium but similar in appearance, was discovered before americium in 1944. The separation of americium from curium took so long (abut a year) and was so painstaking that researchers working on them first called these elements pandemonium (meaning demons or hell in Greek) and delirium (madness), respectively. A 7-minute RSC podcast describes the discovery of americium (and how smoke detectors work as well as americium's connection to the nuclear bomb).

Seaborg was the first to suggest that the f energy shell was filling up before the d shell. Before this, the actinides were thought to be filling up a fourth d-block row. He published his theory despite warnings from his colleagues that this radical departure from the current formulation of the periodic table would ruin his career. In 1951, he received the Nobel prize in Chemistry not for placing the actinide/lanthanide series as a separate group called the f block, but for discovering or co-discovering ten radioactive elements, some of which are actinides and several of which are man-made, including seaborgium (106) (not an actinide) and more than 100 atomic isotopes as well.

How the F-block Works Differently Than the D-block

The reason why these elements act this way is because the energy difference between the 4f and 5d (and between the 5f and 6d) subshells is very small to begin with. After lanthanum, for example, the energy of the 4f sub shell actually falls below that of the 5d sub shell and that's why it's filled first. The evidence for this came from the unique emission line spectra of these elements, revealing unexpected energy differences as electrons shift between these energy shells.

Another electron arrangement unique to the lanthanides series is that the 5s and 5p orbitals penetrate into the 4f-subshell. This has the effect that 4f-subshell electrons are not shielded from the increasing positive nuclear charge and that means that the attractive force makes the radii of their orbitals contract so they nudge closer to the nucleus. This means that the radii (physical size) of the lanthanide atoms actually decreases as atomic number increases. For example, the atomic radius of cerium (second lightest lanthanide, Z=58) is 2.70 angstroms across and the atomic radius of lutetium (heaviest lanthanide, Z=71) is just 2.25 angstroms. As proton number increases and radius decreases, the ionization energy (the energy you need to pull an electron into a higher energy shell or the energy released as it returns to ground state) increases, as you might expect. Increasing ionization energy means that outermost electrons are held more and more tightly to the atom. Density, hardness and melting points are high to start with in the lanthanides and they increase even further as you go left to right in the series as a consequence of this. This makes the lanthanides easier to separate from non-lanthanides, but it makes it more difficult to separate lanthanides from each other. This, and their chemical similarity to one another, is why the rare earths are notoriously hard to purify (we will explore these fascinating rare earths in Periodic Table Part 5: Rare Earths - A Story of Discovery, Demand and International Intrigue.

The unusual progression toward smaller atomic size is called the lanthanide contraction. It is similar to a contraction that occurs in the d-block elements (transition metals), called the scandide contraction, but the mechanisms behind the two contractions are slightly different: The scandide contraction is caused by poor shielding of the nuclear charge by electrons in the d-subshell, whereas the lanthanide contraction is caused by poor nuclear shielding by electrons in the f-subshell.

Next up: History of the Periodic Table Part 5.

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