Shell Filling and the Periodic Table
Shell Filling and the Periodic Table
This table shows electron configurations for the first 20 elements in the periodic table, starting with hydrogen and its single electron and ending with calcium. The Pauli exclusion principle determines the maximum number of electrons allowed in each shell and subshell. But the order in which the shells and subshells are filled is complicated because of the large numbers of interactions between electrons.
Electron Configurations of Elements Hydrogen Through Calcium
| Element | Number of electrons (Z) | Ground state configuration | |||||
|---|---|---|---|---|---|---|---|
| H | 1 | \(1{s}^{1}\) | |||||
| He | 2 | \(1{s}^{2}\) | |||||
| Li | 3 | \(1{s}^{2}\) | \(2{s}^{1}\) | ||||
| Be | 4 | " | \(2{s}^{2}\) | ||||
| B | 5 | " | \(2{s}^{2}\) | \(2{p}^{1}\) | |||
| C | 6 | " | \(2{s}^{2}\) | \(2{p}^{2}\) | |||
| N | 7 | " | \(2{s}^{2}\) | \(2{p}^{3}\) | |||
| O | 8 | " | \(2{s}^{2}\) | \(2{p}^{4}\) | |||
| F | 9 | " | \(2{s}^{2}\) | \(2{p}^{5}\) | |||
| Ne | 10 | " | \(2{s}^{2}\) | \(2{p}^{6}\) | |||
| Na | 11 | " | \(2{s}^{2}\) | \(2{p}^{6}\) | \(3{s}^{1}\) | ||
| Mg | 12 | " | " | " | \(3{s}^{2}\) | ||
| Al | 13 | " | " | " | \(3{s}^{2}\) | \(3{p}^{1}\) | |
| Si | 14 | " | " | " | \(3{s}^{2}\) | \(3{p}^{2}\) | |
| P | 15 | " | " | " | \(3{s}^{2}\) | \(3{p}^{3}\) | |
| S | 16 | " | " | " | \(3{s}^{2}\) | \(3{p}^{4}\) | |
| Cl | 17 | " | " | " | \(3{s}^{2}\) | \(3{p}^{5}\) | |
| Ar | 18 | " | " | " | \(3{s}^{2}\) | \(3{p}^{6}\) | |
| K | 19 | " | " | " | \(3{s}^{2}\) | \(3{p}^{6}\) | \(4{s}^{1}\) |
| Ca | 20 | " | " | " | " | " | \(4{s}^{2}\) |
Examining the above table, you can see that as the number of electrons in an atom increases from 1 in hydrogen to 2 in helium and so on, the lowest-energy shell gets filled first—that is, the \(n=1\) shell fills first, and then the \(n=2\) shell begins to fill. Within a shell, the subshells fill starting with the lowest \(l\), or with the \(s\) subshell, then the \(p\), and so on, usually until all subshells are filled. The first exception to this occurs for potassium, where the \(4s\) subshell begins to fill before any electrons go into the \(3d\) subshell. The next exception is not shown in this table; it occurs for rubidium, where the \(5s\) subshell starts to fill before the \(4d\) subshell. The reason for these exceptions is that \(l=0\) electrons have probability clouds that penetrate closer to the nucleus and, thus, are more tightly bound (lower in energy).
This figure shows the periodic table of the elements, through element 118. Of special interest are elements in the main groups, namely, those in the columns numbered 1, 2, 13, 14, 15, 16, 17, and 18.
Periodic table of the elements (credit: National Institute of Standards and Technology, U.S. Department of Commerce)
The number of electrons in the outermost subshell determines the atom’s chemical properties, since it is these electrons that are farthest from the nucleus and thus interact most with other atoms. If the outermost subshell can accept or give up an electron easily, then the atom will be highly reactive chemically. Each group in the periodic table is characterized by its outermost electron configuration. Perhaps the most familiar is Group 18 (Group VIII), the noble gases (helium, neon, argon, etc.). These gases are all characterized by a filled outer subshell that is particularly stable. This means that they have large ionization energies and do not readily give up an electron. Furthermore, if they were to accept an extra electron, it would be in a significantly higher level and thus loosely bound. Chemical reactions often involve sharing electrons. Noble gases can be forced into unstable chemical compounds only under high pressure and temperature.
Group 17 (Group VII) contains the halogens, such as fluorine, chlorine, iodine and bromine, each of which has one less electron than a neighboring noble gas. Each halogen has 5 \(p\) electrons (a \({p}^{\text{5}}\) configuration), while the \(p\) subshell can hold 6 electrons. This means the halogens have one vacancy in their outermost subshell. They thus readily accept an extra electron (it becomes tightly bound, closing the shell as in noble gases) and are highly reactive chemically. The halogens are also likely to form singly negative ions, such as \({C1}^{-}\), fitting an extra electron into the vacancy in the outer subshell.
In contrast, alkali metals, such as sodium and potassium, all have a single \(s\) electron in their outermost subshell (an \({s}^{1}\) configuration) and are members of Group 1 (Group I). These elements easily give up their extra electron and are thus highly reactive chemically. As you might expect, they also tend to form singly positive ions, such as \({\text{Na}}^{+}\), by losing their loosely bound outermost electron. They are metals (conductors), because the loosely bound outer electron can move freely.
Of course, other groups are also of interest. Carbon, silicon, and germanium, for example, have similar chemistries and are in Group 4 (Group IV). Carbon, in particular, is extraordinary in its ability to form many types of bonds and to be part of long chains, such as inorganic molecules. The large group of what are called transitional elements is characterized by the filling of the \(d\) subshells and crossing of energy levels. Heavier groups, such as the lanthanide series, are more complex—their shells do not fill in simple order. But the groups recognized by chemists such as Mendeleev have an explanation in the substructure of atoms.
PhET Explorations: Stern-Gerlach Experiment
Build an atom out of protons, neutrons, and electrons, and see how the element, charge, and mass change. Then play a game to test your ideas!
This lesson is part of:
Atomic Physics