Electron Configurations of Atoms and Ions

1From 2,8,1 to 1s² 2s² 2p⁶ 3s¹

In S1 you wrote the electron arrangement of sodium as 2,8,1: two electrons in the first shell, eight in the second, one in the third. That shell picture works well up to calcium, then it breaks. Potassium is 2,8,8,1, not 2,8,9, even though the third shell can hold 18 electrons. The reason is the subshells. The 4s subshell has slightly lower energy than 3d, so it fills first. Writing configurations subshell by subshell explains this, and it explains the shape of the periodic table, the charges of ions and much of the chemistry in the rest of this course.

2Three rules for filling orbitals

1. Aufbau principle ("building up"): electrons occupy the lowest-energy orbitals available first. The order of subshell energies, for neutral atoms filling up, is It follows the $(n+l)(n\; +\; l)$ rule: the subshell with the lower value of $n+ln\; +\; l$ fills first; if two have the same $n+ln\; +\; l$, the one with lower $nn$ fills first. 4s ($4+0=44\; +\; 0\; =\; 4$) fills before 3d ($3+2=53\; +\; 2\; =\; 5$). Figure 4.1 turns the rule into a diagram.
2. Pauli exclusion principle: each orbital holds at most two electrons, with opposite spins (↑↓).
3. Hund's rule: when electrons enter orbitals of equal energy (the three p, or five d), they occupy them singly, with parallel spins, before any pairing. Electrons repel each other, so they spread out.

An orbital box diagram draws each orbital as a box and each electron as an arrow. Figure 4.2 shows Hund's rule at work in the 2p subshell.

3Shorthand: the noble-gas core

The inner electrons of an atom are arranged like those of the previous noble gas. Write that noble gas in square brackets and add only the outer electrons: sodium is $1s^{2}2s^{2}2p^{6}3s^{1}1s^2\backslash ,2s^2\backslash ,2p^6\backslash ,3s^1$, or [Ne] $3s^{1}3s^1$.

Table 4.1: Ground-state configurations of some elements

Once 3d holds electrons, it is customary to write it before 4s, grouping each shell together. Chromium and copper break the pattern because a half-filled ($d^{5}d^5$) or completely filled ($d^{10}d^\{10\}$) d subshell is especially stable: moving one 4s electron into 3d lowers the total energy.

4Configurations of ions

• Negative ions: add the extra electrons to the next available orbitals, following the same rules.
• Positive ions: remove electrons from the outermost shell, the one with the highest $nn$, first. For main-group elements this is simply the last subshell filled. For the d-block, 4s electrons are removed before 3d electrons: once 3d is occupied, 4s is the outermost subshell and is higher in energy.

Species with the same number of electrons and the same configuration are isoelectronic. For example ${\mathrm{N}}^{3-}\backslash mathrm\{N^\{3-\}\}$, ${\mathrm{O}}^{2-}\backslash mathrm\{O^\{2-\}\}$, ${\mathrm{F}}^{-}\backslash mathrm\{F^\{-\}\}$, Ne, $\mathrm{N}{\mathrm{a}}^{+}\backslash mathrm\{Na^\{+\}\}$, $\mathrm{M}{\mathrm{g}}^{2+}\backslash mathrm\{Mg^\{2+\}\}$ and $\mathrm{A}{\mathrm{l}}^{3+}\backslash mathrm\{Al^\{3+\}\}$ all have 10 electrons, $1s^{2}2s^{2}2p^{6}1s^2\backslash ,2s^2\backslash ,2p^6$.