Chemistry 6, 9 AM section, graphic

Atomic Periodic Properties

The periodic variation in electron configurations as one moves sequentially through the Periodic Table from H to ever heavier elements produces a periodic variation in a variety of properties. We have already see the periodic variation in atomic shape, and here we look at three other properties. For more examples, including these three, visit the Dartmouth College Periodic Puzzle in the ChemLab website.

The first such property we will consider is atomic size as measured by the atomic radius, a property discussed in the text starting on page 198 and tabulated in Appendix F along with many other element properties. (And remember: most elements are found under ordinary conditions as diatomic molecules—N2, O2, etc.—polyatomic molecules—P4, S8, etc.—or solid metals. Here we are considering only free atoms, often produced at high temperatures only, and in the gas phase only.) The figure below shows how atomic radius varies qualitatively across the Periodic Table.

The general trend is easy to spot and to understand: heavier atoms (those with large atomic numbers) have more electrons, and increasing the number of electrons means placing electrons into orbitals with ever-increasing principal quantum number—orbitals with ever-increasing size. This clearly explains the trend in increasing size as we go down any one column: H is smaller than Li, etc., ending the first column with the biggest atom for which we have reliable data, Cs. These atoms' sizes are governed by the size of their single highest-energy electron's orbital, which is 1s for H but 6s for Cs.

As we go across a row (look at Li through Ne, for example), the size decreases, but rather slowly: we are adding electrons into orbitals with (for this row) the same principal quantum number. These orbitals are slowly shrinking in size because the nuclear charge is increasing as we go across a row, and increasing nuclear charge means increasing the force attracting electrons to the nucleus, making the orbitals contract.

Next, we look at two energetic properties of atoms, the ionization energy (IE) and the electron affinity (EA). The ionization energy is the energy required to remove the least tightly bound electron from an atom, producing a positive ion and a free electron:

A —> A+ + e

(When we start with a neutral atom, as we have here, we say the energy change is the first ionization energy. If we then remove another electron, as in A+ —> A2+ + e, we call the energy need to do so the second ionization energy, and so forth, until we run out of electrons.) The figure below shows the trend in first ionization energies across the Periodic Table.

Compare this figure to the earlier figure of Atomic Radii and note that in general, small size means large ionization energy. This trend is easy to understand: small atoms have few electrons that are close to the nucleus. The closer an electron is to a nucleus, the more energy it takes to remove that electron from the atom. Now look at the trends going across any one row of the Periodic Table: ionization energies increase. For example, the He ionization energy is greater than the H ionization energy. Both have electrons in the 1s orbital, but the increased nuclear charge of He pulls its two 1s electrons closer to the nucleus, and thus more energy is needed to remove either one of them than is needed for H with only a single nuclear proton. Likewise, as we go from Li to Ne, we are adding electrons to the n = 2 orbitals (2s at first, then 2p). These orbitals are pulled closer to the nucleus as we go from Li to Ne because the nuclear charge is increasing, pulling the electrons closer. As we go down a column (consider H through Cs or He through Rn, for example), the highest energy (outermost spatially) electron or electrons will be found in orbitals of increasing principal quantum number n: n = 1 for H and He, 2 for Li and Ne, etc., through 6 for Cs and Rn. As the principal quantum number of the electron we are trying to remove increases, the energy needed to remove it decreases even though the nuclear charge is increasing as we go down a column. We explain this through the concept of shielding: the electrons in orbitals of smaller principal quantum number shield (or screen) some of the nuclear charge, causing our outer electron of interest to think it is bound to an atom of roughly the same nuclear charge no matter which row it is in.

Finally, we focus on the electron affinity. This is the energy required to remove the outermost electron from an atom that has one extra electron stuck to it:

A —> A + e

In other words, the electron affinity is the ionization energy of the singly charged atomic anion. The trends in electron affinities are shown below.

Note first the elements that fall in the low end of the range. For these elements, the electron affinity is either zero or very close to zero, which means that these elements do not form stable anions. If you compare the locations of these elements to the Periodic Table of spherical elements that we already discussed, you will see that those atoms that are spherical because they have closed shells or subshells also have zero or nearly zero electron affinities. Likewise, those that are spherical because they have half-filled shells or subshells also have zero or very small electron affinities. Why? Consider the rare gases, starting with He. Adding an electron to He means changing the electron configuration from 1s2 to 1s22s1, and the screened He nuclear charge simply isn't strong enough to bind that 2s electron and make a stable He anion. A similar argument holds for all the other rare gases, and for the alkaline earth elements, Be through Ra, plus the elements at the end of the transition metals, Zn, Cd, and Hg, the extra electron would go into a new subshell (as in Be 1s22s2 –> Be 1s22s22p1). Again, the screened nuclear charge cannot bind such an electron. For the half-filled elements such as N (1s22s22p3), the three p electrons, you'll recall, are in different p orbitals, one for each m quantum number value. Adding a fourth to make N forces two electrons into the same p orbital. Since electrons repel and since the nuclear charge doesn't increase, that fourth electron is not bound. N does not exist as a free atomic anion. (For elements below N, you'll notice that the electron affinity is not zero, but it is small. Adding that one extra electron is not quite such a bad idea for these elements because the electron finds itself in a larger p orbital (3p for P, 4p for as, etc.), and these larger orbitals allow the two electrons that are now forced into the same orbital to avoid each other better.)

But for the halogens, F through At, the electron affinity is quite high. These elements are just one electron away from a close shell configuration, and the extra electron can be held by the attractive force of the atom's nucleus. Note as well that Cu, Ag, and Au, which are only one electron away from a closed d subshell, also have high electron affinities.

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