Electronic Configuration and Ions
Electron Configuration
Electrons are arranged in energy levels (shells) around the nucleus. Each shell can hold a limited number of electrons:
| Shell | Maximum electrons |
|---|---|
| 1st (closest to nucleus) | 2 |
| 2nd | 8 |
| 3rd | 8 |
| 4th | 18 (but at GCSE, fills to 8 first) |
Electrons fill the lowest available energy level first. This is sometimes called the aufbau principle.
Writing Electron Configurations
The electron configuration is written as numbers separated by dots or commas, showing how many electrons are in each shell.
| Element | Atomic number | Configuration |
|---|---|---|
| Hydrogen | 1 | 1 |
| Carbon | 6 | 2.4 |
| Sodium | 11 | 2.8.1 |
| Chlorine | 17 | 2.8.7 |
| Calcium | 20 | 2.8.8.2 |
Electron Configuration and the Periodic Table
The periodic table is arranged so that:
- The group number tells you the number of electrons in the outer shell (for Groups 1-7)
- The period number tells you the number of occupied electron shells
For example, potassium is in Group 1, Period 4, so its configuration is 2.8.8.1.
This arrangement explains why elements in the same group have similar chemical properties — they have the same number of outer shell electrons.
Ions
An ion is an atom (or group of atoms) that has gained or lost electrons and therefore carries an electrical charge.
How Ions Form
Atoms form ions to achieve a stable electron configuration, usually that of the nearest noble gas (a full outer shell).
Metal atoms lose electrons to form positive ions (cations):
- Sodium (2.8.1) → Na⁺ (2.8) — loses 1 electron
- Magnesium (2.8.2) → Mg²⁺ (2.8) — loses 2 electrons
- Aluminium (2.8.3) → Al³⁺ (2.8) — loses 3 electrons
Non-metal atoms gain electrons to form negative ions (anions):
- Chlorine (2.8.7) → Cl⁻ (2.8.8) — gains 1 electron
- Oxygen (2.6) → O²⁻ (2.8) — gains 2 electrons
- Nitrogen (2.5) → N³⁻ (2.8) — gains 3 electrons
Key Rules
- The number of protons does not change when an ion forms — only electrons change
- A positive ion has fewer electrons than protons
- A negative ion has more electrons than protons
Predicting Ion Charges from the Periodic Table
| Group | Electrons in outer shell | Ion charge | Example |
|---|---|---|---|
| 1 | 1 | 1+ | Na⁺, K⁺, Li⁺ |
| 2 | 2 | 2+ | Mg²⁺, Ca²⁺ |
| 3 | 3 | 3+ | Al³⁺ |
| 5 | 5 | 3− | N³⁻ |
| 6 | 6 | 2− | O²⁻, S²⁻ |
| 7 | 7 | 1− | Cl⁻, Br⁻, F⁻ |
Group 0 elements do not normally form ions because they already have full outer shells.
Ions and Ionic Bonding
When a metal reacts with a non-metal, electrons are transferred from the metal atom to the non-metal atom. Both atoms become ions. The opposite charges attract each other, forming an ionic bond.
For example, in sodium chloride:
Na (2.8.1) → Na⁺ (2.8) + e⁻
Cl (2.8.7) + e⁻ → Cl⁻ (2.8.8)
The Na⁺ and Cl⁻ ions are held together by strong electrostatic attraction.
Exam Tips
- Always state the full electron configuration of the ion, not just the charge
- When drawing dot-and-cross diagrams for ions, show the square brackets and charge
- Remember that the charge on a transition metal ion can vary (e.g. Fe²⁺ and Fe³⁺) and must be given in the question
- Group 0 elements have full outer shells, which is why they are unreactive