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− | \(\text{Cl}^-\) |
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
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.
The 4s-before-3d Ionisation 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.
Iron (Fe, Z = 26)
Neutral atom: \([Ar]\,3d^6\,4s^2\)
- 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)
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.
Study this topic on the go
Get active recall flashcards, notes, and topic quizzes in ChemEasy, or build your revision schedule with ChemPlan IB.