Ionisation Energies

Ionisation energies provide direct experimental evidence for the existence of electron shells and sub-shells. Understanding trends in ionisation energy is essential for explaining periodicity and the properties of elements.

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Definitions

First ionisation energy: The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.

Equation: X(g) → X⁺(g) + e⁻

Second ionisation energy: The energy required to remove one electron from each ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.

Equation: X⁺(g) → X²⁺(g) + e⁻

In general, the nth ionisation energy is always defined for the removal of one electron from the (n−1)+ ion.

Key Point: Ionisation energies are always positive (endothermic) because energy must be supplied to overcome the attraction between the electron and the nucleus. They are measured in kJ mol⁻¹.

Conditions

The definition specifies gaseous atoms because all electrons must be in their normal energy levels (not involved in bonding or intermolecular forces). This ensures a fair comparison between elements.

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Factors Affecting Ionisation Energy

Three main factors determine the size of an ionisation energy:

1. Nuclear Charge

A greater number of protons in the nucleus means a stronger attraction for the outer electrons, so more energy is needed to remove an electron. Increasing nuclear charge increases ionisation energy.

2. Atomic Radius (Distance from Nucleus)

The further an electron is from the nucleus, the weaker the attraction (the force of attraction follows an inverse square law: F ∝ 1/r²). Increasing distance decreases ionisation energy.

3. Electron Shielding

Inner electrons repel outer electrons and reduce the effective nuclear charge experienced by the outer electron. More inner shells means more shielding. Increasing shielding decreases ionisation energy.

Key Definition: The effective nuclear charge is the net positive charge experienced by an outer electron, after accounting for the shielding effect of inner electrons.

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Trend Across a Period (e.g. Period 2: Li to Ne)

The general trend across a period is increasing first ionisation energy. This is because:

• The nuclear charge increases (more protons) as you go across the period.
• Electrons are being added to the same shell, so the shielding remains approximately constant.
• The atomic radius decreases slightly.

Therefore, the effective nuclear charge increases, and the outer electron is held more tightly.

Exceptions Across Period 2

There are two drops in the otherwise increasing trend:

Drop 1: Be (900 kJ mol⁻¹) to B (801 kJ mol⁻¹)

Beryllium: 1s² 2s² Boron: 1s² 2s² 2p¹

The outer electron of boron is in a 2p sub-shell, which is higher in energy and further from the nucleus than the 2s sub-shell of beryllium. The 2p electron is also partially shielded by the 2s² electrons. Therefore, less energy is required to remove it.

Drop 2: N (1402 kJ mol⁻¹) to O (1314 kJ mol⁻¹)

Nitrogen: 1s² 2s² 2p³ — three singly occupied 2p orbitals Oxygen: 1s² 2s² 2p⁴ — one 2p orbital contains a pair of electrons

In oxygen, the paired electrons in the 2p orbital experience electron-electron repulsion within the same orbital. This additional repulsion makes it easier to remove one of the paired electrons, so the ionisation energy of oxygen is lower than that of nitrogen.

First Ionisation Energies Across Period 2

Element	Z	Configuration	1st IE (kJ mol⁻¹)
Li	3	1s² 2s¹	520
Be	4	1s² 2s²	900
B	5	1s² 2s² 2p¹	801
C	6	1s² 2s² 2p²	1086
N	7	1s² 2s² 2p³	1402
O	8	1s² 2s² 2p⁴	1314
F	9	1s² 2s² 2p⁵	1681
Ne	10	1s² 2s² 2p⁶	2081

The Same Pattern Repeats in Period 3

Element	Z	Configuration	1st IE (kJ mol⁻¹)
Na	11	[Ne] 3s¹	496
Mg	12	[Ne] 3s²	738
Al	13	[Ne] 3s² 3p¹	577
Si	14	[Ne] 3s² 3p²	786
P	15	[Ne] 3s² 3p³	1012
S	16	[Ne] 3s² 3p⁴	1000
Cl	17	[Ne] 3s² 3p⁵	1251
Ar	18	[Ne] 3s² 3p⁶	1521