Ionisation Energies and Electron Shielding
Ionisation Energies and Electron Shielding
Ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms (or ions) to form one mole of gaseous positive ions.
First Ionisation Energy
The first ionisation energy is defined as:
X(g) → X⁺(g) + e⁻
It is always endothermic (positive value) because energy must be supplied to overcome the attraction between the nucleus and the electron.
Factors Affecting Ionisation Energy
Four key factors determine ionisation energy:
- Nuclear charge — more protons in the nucleus means a greater attractive force on outer electrons, increasing ionisation energy
- Electron shielding — inner electron shells repel outer electrons, reducing the effective nuclear charge felt by the outermost electron
- Atomic radius — the further an electron is from the nucleus, the weaker the attraction (follows an inverse-square relationship)
- Electron pair repulsion — two electrons sharing the same orbital experience additional repulsion, making one slightly easier to remove
Trends Across a Period
First ionisation energies generally increase across a period (e.g. Period 3, Na → Ar). This is because:
1. Each successive element has one more proton in the nucleus
2. Electrons are added to the same shell, so shielding remains approximately constant
3. The effective nuclear charge increases, pulling electrons closer and binding them more tightly
However, there are two notable drops in the trend across Period 3:
Drop from Mg to Al:
- Mg: 1s² 2s² 2p⁶ 3s²
- Al: 1s² 2s² 2p⁶ 3s² 3p¹
- The 3p electron in aluminium is in a higher energy sub-shell than the 3s electrons in magnesium
- The 3p electron is slightly further from the nucleus and is shielded by the 3s² pair
- Therefore less energy is needed to remove it
Drop from P to S:
- P: 1s² 2s² 2p⁶ 3s² 3p³ (three unpaired 3p electrons)
- S: 1s² 2s² 2p⁶ 3s² 3p⁴ (one 3p orbital now contains a pair)
- The paired electron in sulfur experiences electron pair repulsion within its orbital
- This makes it easier to remove than an unpaired electron in phosphorus
Trends Down a Group
First ionisation energies decrease down a group because:
- Additional electron shells are added, increasing atomic radius
- Shielding increases significantly as more inner shells are present
- Although nuclear charge also increases, the effect of distance and shielding outweighs the extra protons
Successive Ionisation Energies
Successive ionisation energies are the energies to remove the 1st, 2nd, 3rd (etc.) electrons, one at a time, from a gaseous atom.
Each successive ionisation energy is higher than the last because:
- Each electron is removed from an increasingly positive ion
- The remaining electrons are held more tightly by the same nuclear charge
Large jumps in successive ionisation energies indicate that an electron is being removed from a different, inner shell — it is closer to the nucleus and experiences less shielding.
Worked Example
The successive ionisation energies of an element X (in kJ mol⁻¹) are: 578, 1817, 2745, 11578, 14831...
The big jump occurs between the 3rd and 4th ionisation energies. This means:
- The first three electrons are relatively easy to remove (outer shell)
- The fourth electron comes from a much more stable inner shell
- Element X has 3 electrons in its outer shell — it is in Group 3 (e.g. aluminium)
Effective Nuclear Charge
The effective nuclear charge (Z_eff) is the net positive charge experienced by an outer electron after accounting for shielding:
Z_eff = Z − S
Where Z is the actual nuclear charge (number of protons) and S is the shielding constant (approximately equal to the number of inner-shell electrons).
For example, a sodium atom (Z = 11) with 10 inner electrons has Z_eff ≈ 11 − 10 = +1 for its outer 3s electron. A chlorine atom (Z = 17) with 10 inner electrons has Z_eff ≈ 17 − 10 = +7, explaining its much higher ionisation energy.
Evidence for Sub-Shells
The successive ionisation energies of nitrogen (1s² 2s² 2p³) show:
- Electrons 1–3 removed from 2p (relatively lower IE values)
- A small jump to electrons 4–5 from 2s (higher IE, as 2s is lower in energy / closer to nucleus)
- A large jump to electrons 6–7 from 1s (inner shell)
This pattern provides direct evidence for the existence of sub-shells within a principal energy level.
Key Definitions Summary
| Term | Definition |
|---|---|
| First ionisation energy | Energy to remove 1 mol e⁻ from 1 mol gaseous atoms |
| Second ionisation energy | Energy to remove 1 mol e⁻ from 1 mol gaseous 1+ ions |
| Shielding | Repulsion of outer electrons by inner electron shells |
| Effective nuclear charge | Net positive charge felt by outermost electron |
Exam Tips
- Always state that the atom/ion must be in the gaseous state when defining ionisation energy
- When explaining trends, link each factor to the attraction between the nucleus and the outer electron
- For successive IE graphs, count from the right-hand side (the last big jump tells you the shell structure)