Ionisation enthalpy (also called ionisation energy) is the minimum energy required to remove one electron from an isolated gaseous atom in its ground state, to form a gaseous cation. The first ionisation enthalpy removes the first (outermost) electron. Ionisation enthalpy increases across a period (left to right) and decreases down a group. Units: kJ/mol or eV/atom.
Ionisation enthalpy: energy to remove one electron from a gaseous atom in ground state.
Always endothermic: M(g) → M⁺(g) + e⁻
Increases across a period (left to right) due to increasing nuclear charge.
Decreases down a group due to increasing atomic size and shielding.
Anomaly: Be > B (Be has fully filled 2s² subshell — extra stable).
Anomaly: N > O (N has half-filled 2p³ subshell — extra stable).
Highest IE: He (2372 kJ/mol); Lowest IE: Cs (376 kJ/mol).
IE₁ < IE₂ < IE₃ — each successive ionisation energy is higher.
Definition: The minimum energy required to remove the most loosely held electron from a neutral gaseous atom in its ground state:
M(g) → M⁺(g) + e⁻ ΔH = First Ionisation Enthalpy (IE₁) M⁺(g) → M²⁺(g) + e⁻ ΔH = Second Ionisation Enthalpy (IE₂)
Key points: • Always endothermic (energy must be supplied) • Units: kJ/mol or eV/atom • Atom must be gaseous and isolated (no intermolecular forces) • Always: IE₁ < IE₂ < IE₃ (each successive IE is greater) • A very large jump between consecutive IEs indicates a change in electron shell
First ionisation enthalpy values (approximate): • H: 1312 kJ/mol • He: 2372 kJ/mol (highest) • Li: 520 kJ/mol • Na: 496 kJ/mol • F: 1681 kJ/mol • Cl: 1251 kJ/mol
Across a period (left to right): • Ionisation enthalpy generally increases • Reason: nuclear charge increases → electrons are held more tightly • More energy needed to remove an electron • Example (Period 2): Li(520) < Be(900) < B(801) < C(1086) < N(1402) < O(1314) < F(1681) < Ne(2081) kJ/mol
Down a group (top to bottom): • Ionisation enthalpy decreases • Reason: atomic size increases, electron is farther from nucleus, shielding increases → less energy to remove • Example (Group 1): Li(520) > Na(496) > K(419) > Rb(403) > Cs(376) kJ/mol
Highest IE: He (noble gas, completely filled shell) Lowest IE: Cs (largest atom, least nuclear pull on outermost electron)
Two important anomalies in Period 2:
Anomaly 1 — Be > B: • Expected: IE of B > Be (B has more protons) • Actual: IE of Be (900) > IE of B (801) kJ/mol • Reason: Be has configuration 2s² (fully filled subshell — extra stability) B has configuration 2s² 2p¹ — the 2p electron is in a higher energy, more diffuse orbital, shielded by 2s electrons → easier to remove
Anomaly 2 — N > O: • Expected: IE of O > N (O has more protons) • Actual: IE of N (1402) > IE of O (1314) kJ/mol • Reason: N has configuration 2s² 2p³ (half-filled p — extra stability due to exchange energy) O has configuration 2s² 2p⁴ — one p orbital has paired electrons, electron-electron repulsion makes one electron easier to remove
Memory aid for anomalies: Be > B and N > O
Ionisation enthalpy (ionisation energy) is the minimum energy required to remove one electron from an isolated gaseous atom in its ground state: M(g) → M⁺(g) + e⁻. It is always endothermic and measured in kJ/mol or eV/atom.
Ionisation enthalpy generally increases across a period (left to right). This is because the nuclear charge increases while electrons are added to the same shell, so electrons are held more tightly and more energy is needed to remove them.
Ionisation enthalpy decreases down a group. As atomic size increases, the outermost electron is farther from the nucleus and is more shielded by inner electrons. Less energy is needed to remove it.
Beryllium (2s²) has a completely filled 2s subshell, which is extra stable. Boron (2s² 2p¹) has its outermost electron in the 2p orbital, which is higher in energy and shielded by the 2s electrons. This makes the 2p electron of boron easier to remove, so IE(Be) > IE(B).
Nitrogen (2s² 2p³) has a half-filled p subshell, which is extra stable due to exchange energy. Oxygen (2s² 2p⁴) has one p orbital with paired electrons, which experience electron-electron repulsion. This makes it easier to remove one electron from oxygen than from nitrogen, so IE(N) > IE(O).
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