Ions are formed when atoms gain or lose electrons to achieve a more stable electron configuration. Typically, this means achieving the nearest noble-gas configuration. The IB demands precise terminology throughout.
Cations (+)
Metal atoms lose valence electrons to form positively charged ions. The atom becomes smaller as an entire electron shell is typically lost.
Anions (−)
Non-metal atoms gain electrons into their valence shell to form negatively charged ions. The ion is larger due to increased electron-electron repulsion.
Predicting Main-Group Ion Charges
For s-block and p-block elements, the group number provides a highly reliable prediction:
| Group | Process | Ion Charge | Example |
|---|---|---|---|
| Group 1 | Loses 1 electron | 1+ | \(\text{Na}^+\) |
| Group 2 | Loses 2 electrons | 2+ | \(\text{Mg}^{2+}\) |
| Group 13 | Loses 3 electrons | 3+ | \(\text{Al}^{3+}\) |
| Group 15 | Gains 3 electrons | 3− | \(\text{N}^{3-}\) |
| Group 16 | Gains 2 electrons | 2− | \(\text{O}^{2-}\) |
| Group 17 | Gains 1 electron | 1− | \(\text{Cl}^-\) |
📐 Worked Example. Sodium & Chlorine
Na: Configuration \(1s^2\,2s^2\,2p^6\,3s^1\). Loses its single 3s electron → \(\text{Na}^+\) with configuration \(1s^2\,2s^2\,2p^6\) (isoelectronic with neon).
Cl: Configuration \(1s^2\,2s^2\,2p^6\,3s^2\,3p^5\). Gains one electron → \(\text{Cl}^-\) with configuration \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\) (isoelectronic with argon).
Transition Metals & Variable Charges
📘 IB Definition. Transition Metal
A transition metal is an element that possesses an incomplete d-subshell in its neutral atomic state, or one that can form at least one stable cation with an incomplete d-subshell.
Under this strict IB definition, scandium (Sc) and zinc (Zn) are excluded from the first-row transition metals because neither forms an ion with a partially filled d-subshell.
⚠️ Critical Rule. The 4s-before-3d Rule
While filling electron configurations (Aufbau), the 4s orbital fills before the 3d. But when ionizing transition metals, the 4s electrons are always removed first. They are the outermost, highest-energy electrons once the atom is fully populated.
Example: Iron (\(\text{Fe}\), Z = 26)
Neutral: \([Ar]\,3d^6\,4s^2\)
- Fe²⁺: Remove 2 × 4s electrons → \([Ar]\,3d^6\)
- Fe³⁺: Remove 2 × 4s + 1 × 3d electron → \([Ar]\,3d^5\). Particularly stable due to half-filled d-subshell symmetry
🟣 HL Extension. Why Variable Oxidation States?
HL students must explain variable oxidation states using successive ionization energies. In transition metals, the 4s and 3d subshells are very close in energy:
- Main-group metals show a catastrophic energy jump when removing core electrons, which strictly limits them to one charge (e.g. K is only ever 1+)
- Transition metals show a gradual, continuous increase across 3d electrons. The energy cost is recoverable through lattice enthalpy or hydration enthalpy
This is why manganese can exhibit oxidation states from +2 all the way to +7.
Ion Size vs Atom Size
Cations are SMALLER
Metals lose their outermost electron shell entirely. The remaining electrons are pulled closer to the nucleus by a now-dominant nuclear charge. There are more protons than electrons, increasing the effective nuclear charge.
Anions are LARGER
Non-metals gain electrons into the same shell. More electrons means greater electron-electron repulsion, which pushes the electron cloud outward. The nuclear charge is now shared across more electrons, reducing its pull per electron.