Valence electrons decide how an element reacts
The previous topics showed that valence electrons are the outermost and most loosely held electrons. They are also the ones that change when atoms react. Elements in the same group have the same number of valence electrons in the same kind of subshell: lithium (2s¹), sodium (3s¹) and potassium (4s¹) each have one. That is why the elements of a group behave alike. Sodium and potassium both react fast with water, for the same reason: each gives up one loosely held s electron.
Metals lose electrons; nonmetals gain them
Metals, on the left of the table, have few valence electrons held weakly (low ionization energy). Little energy removes them, so metals form cations. Nonmetals, on the right, have a high effective nuclear charge and a small radius; they attract an extra electron strongly and form anions.
How many electrons? For main-group elements, the ion usually ends up with the same electron configuration as a noble gas, a noble gas configuration (Figure 1).
![Na (1s² 2s² 2p⁶ 3s¹) loses one electron to become the smaller Na⁺, like neon; Cl ([Ne] 3s² 3p⁵) gains one to become the larger Cl⁻, like argon; one of each makes neutral NaCl.](../figures/ion-formation.svg)
| Group | Valence electrons | Usual ion | Examples |
|---|---|---|---|
| 1 | 1 | 1+ | Li⁺, Na⁺, K⁺ |
| 2 | 2 | 2+ | Mg²⁺, Ca²⁺ |
| 13 | 3 | 3+ | Al³⁺ |
| 15 | 5 | 3− | N³⁻, P³⁻ |
| 16 | 6 | 2− | O²⁻, S²⁻ |
| 17 | 7 | 1− | F⁻, Cl⁻, Br⁻ |
Why the charge stops where it does
The reason is energy, not desire. Take aluminum, [Ne] 3s² 3p¹. Its three valence electrons are in shell 3, far from the nucleus and well shielded, so removing them is affordable. The fourth electron would have to come from the 2p core: much closer to the nucleus and shielded only by the 1s pair. Successive ionization energies show the jump: about 2,745 kJ/mol for the third electron, about 11,577 kJ/mol for the fourth. So aluminum forms Al³⁺, not Al⁴⁺.
On the other side, adding electrons to a nonmetal past the noble gas configuration would put them in a new, farther shell with almost no net attraction, so S stops at S²⁻ and Cl at Cl⁻.
Particles with the same number of electrons are called isoelectronic. Na⁺, Mg²⁺, F⁻ and O²⁻ all have 10 electrons, like neon; K⁺, Ca²⁺, Cl⁻ and S²⁻ have 18, like argon.
Worked example. Write the electron configurations of Ca²⁺ and S²⁻.
Ca is [Ar] 4s². Remove the two 4s electrons: Ca²⁺ is [Ar], or 1s² 2s² 2p⁶ 3s² 3p⁶.
S is [Ne] 3s² 3p⁴. Add two electrons to 3p: S²⁻ is [Ne] 3s² 3p⁶, the same as argon.
Both have 18 electrons: they are isoelectronic with argon.
Ionic compounds
An ionic compound is made of cations and anions held together by the attraction of their opposite charges. A sample is electrically neutral, so the ions combine in the ratio that makes the total charge zero. The formula gives that ratio as the smallest whole numbers, with the cation first. The smallest repeating unit, such as one Na⁺ with one Cl⁻, is called a formula unit; ionic compounds do not form separate molecules.
Worked example. Write the formula of the compound of calcium and nitrogen.
Calcium is in group 2: Ca²⁺. Nitrogen is in group 15: N³⁻.
Find the smallest numbers that cancel: 3 × (+2) = +6 and 2 × (−3) = −6.
Formula: Ca₃N₂.
More examples: Mg²⁺ and Cl⁻ give MgCl₂; Al³⁺ and Br⁻ give AlBr₃; K⁺ and O²⁻ give K₂O; Mg²⁺ and O²⁻ give MgO (not Mg₂O₂, which is not the simplest ratio).
| Formula | Charge sum | Correct? |
|---|---|---|
| CaF₂ | +2 + 2(−1) = 0 | yes |
| CaF | +2 − 1 = +1 | no |
| Na₃N | 3(+1) − 3 = 0 | yes |
| MgN | +2 − 3 = −1 | no (Mg₃N₂) |
Ionic compounds in particle diagrams
A particle diagram of a solid ionic compound shows cations and anions in the ratio of the formula. A box of MgCl₂ has twice as many Cl⁻ as Mg²⁺; a box of MgO or NaCl has equal numbers. Counting the ions and adding up their charges is the quickest check: the total in any correct diagram is zero. Unit 2 looks at how these ions are arranged in the solid and how strongly they hold together.