Ionisation energies orbitals and electronic configuration

Key Definition The first ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous 1+ ions. The second and third ionisation energies remove the next electron from one mole of gaseous 1+ ions and 2+ ions respectively. All ionisation energies are endothermic (positive) because energy must be supplied to overcome the electrostatic attraction between the electron and the nucleus.

State symbols and the gaseous state are essential in defining and writing equations for these processes.

1st IE: X(g) → X⁺(g) + e⁻ 2nd IE: X⁺(g) → X²⁺(g) + e⁻ 3rd IE: X²⁺(g) → X³⁺(g) + e⁻

• Nuclear charge
• Shielding by inner electrons
• The distance of the electron from the nucleus
• The sub-shell from which the electron is removed

Greater nuclear charge raises ionisation energy, whereas greater shielding, greater distance from the nucleus, and removal from a higher-energy sub-shell all lower it.

Successive ionisation energies provide direct evidence for quantum shells through the large jumps that mark transitions between shells, and the count of electrons removed before the first big jump reveals the group of the element.

First ionisation energies across a period reveal sub-shell structure through the characteristic dips at Mg→Al and P→S.

• The Mg→Al dip occurs because aluminium's outermost electron is removed from a 3p orbital, which is higher in energy and more shielded than the 3s orbital from which magnesium loses its electron.
• The P→S dip occurs because sulfur's 3p subshell forces two electrons to pair in one 3p orbital, and the resulting electron–electron repulsion makes one electron easier to remove.

Key Definition An orbital is a region around the nucleus that can hold up to two electrons with opposite spins. s orbitals are spherical, while p orbitals are dumb-bell shaped and lie along three mutually perpendicular axes (px, py, pz).

Within a sub-shell, orbitals fill singly with parallel spins before any electrons pair up (Hund's rule).

Electronic configurations from hydrogen to krypton follow the filling order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, with anomalies at chromium ([Ar]3d⁵4s¹) and copper ([Ar]3d¹⁰4s¹) due to the extra stability of half-filled and fully-filled 3d sub-shells.

For d-block ions, the 4s electrons are removed before the 3d electrons, even though 4s fills first.

Electronic configuration governs chemical properties because outer-shell electrons determine bonding and reactivity, and elements with the same outer configuration belong to the same group with similar chemistry.

The Periodic Table divides into s, p and d blocks reflecting the sub-shell of the highest-energy electron, with sub-shell capacities of 2, 6 and 10 respectively.

Connecting these ideas, ionisation energy patterns serve as the experimental backbone for the entire model of shells, sub-shells and orbitals.

Examiners reward

• Precise definitions
• Full state symbols
• Structured comparisons referencing all relevant factors
• Accurate s, p, d notation, including the correct order of removal when forming ions