Introduction
Atoms are electrically neutral — they have equal numbers of protons and electrons. But atoms don't always stay neutral! When atoms gain or lose electrons, they become ions — charged particles. Understanding why and how this happens is key to predicting the formulas of compounds and explaining many chemical reactions, including those of group 1 metals and group 17 non-metals (Talbot, p.233).
Why do atoms form ions?
Atoms "want" to achieve a stable electron arrangement, usually one that matches the nearest noble gas (a full outer shell). This is sometimes called the octet rule (8 electrons in the outer shell), though helium's stable arrangement is just 2 electrons.
- Metals tend to lose electrons from their outer shell, because it is easier to remove a few electrons than to gain many. Losing electrons leaves more protons than electrons, so the particle becomes positively charged — a cation.
- Non-metals tend to gain electrons to complete their outer shell. Gaining electrons gives more electrons than protons, producing a negative charge — an anion.
Predicting ion charge from the Periodic Table
The group number of an element gives a strong clue about the ion it usually forms, because it tells us the number of valence (outer-shell) electrons (Talbot, p.132).
| Group | Valence electrons | Tendency | Typical ion charge |
|---|---|---|---|
| 1 | 1 | lose 1 e⁻ | +1 |
| 2 | 2 | lose 2 e⁻ | +2 |
| 13 | 3 | lose 3 e⁻ | +3 |
| 15 | 5 | gain 3 e⁻ | −3 |
| 16 | 6 | gain 2 e⁻ | −2 |
| 17 | 7 | gain 1 e⁻ | −1 |
| 18 | 8 (full) | neither | 0 |
Worked Example 1 — Sodium forming Na⁺
Sodium's electron configuration is 1s² 2s² 2p⁶ 3s¹ (2,8,1). It has 1 valence electron. Losing it gives:
Na → Na⁺ + e⁻
The resulting ion has configuration 1s² 2s² 2p⁶ (2,8), identical to neon — a stable noble gas arrangement (Talbot, p.132).
Worked Example 2 — Sulfur forming S²⁻
Sulfur has configuration 1s² 2s² 2p⁶ 3s² 3p⁴ (2,8,6), with 6 valence electrons. Gaining 2 electrons completes the outer shell:
S + 2e⁻ → S²⁻
Giving configuration 2,8,8 (1s² 2s² 2p⁶ 3s² 3p⁶), matching argon (Talbot, p.132).
Worked Example 3 — Aluminium forming Al³⁺
Aluminium (2,8,3) loses its 3 outer electrons:
Al → Al³⁺ + 3e⁻
Resulting configuration: 2,8 — same as neon.
A special case: transition metals
Transition elements (d-block) don't follow a simple group-number rule. Most form more than one stable cation, because electrons can be removed from different sublevels with similar energy. For example:
- Copper forms Cu⁺ (copper(I)) and Cu²⁺ (copper(II))
- Iron forms Fe²⁺ (iron(II)) and Fe³⁺ (iron(III))
Roman numerals are used in naming to show which oxidation state (charge) is present (Talbot, p.132). Zinc is an exception among transition metals — it only forms one ion, Zn²⁺ (Talbot, p.263).
Key takeaways
- Atoms form ions by losing or gaining electrons to reach a stable, noble-gas-like electron arrangement.
- Metals lose electrons → cations (positive); non-metals gain electrons → anions (negative).
- The group number usually predicts the charge of a simple ion for main-group elements.
- Transition metals often form multiple stable ions (e.g., Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺), shown using Roman numerals; zinc is an exception, forming only Zn²⁺.
- Writing the electron configuration of an ion helps confirm it matches a noble gas arrangement.