Atomic Structure: Orbitals and Ionisation Energy

This guide covers subtopics 1.3, Electrons, energy levels and atomic orbitals, and 1.4, Ionisation energy, from Topic 1 of Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Both are AS Level content, assessed from the first year of the course.

Before studying this

This resource assumes you already know, from IGCSE or O Level, that atoms are made of protons, neutrons and electrons, that isotopes are atoms of the same element with different numbers of neutrons, and that electrons occupy shells — written as a simple list like 2,8,8,2. If any of that is unfamiliar, start with Atomic Structure first.

It also assumes Topic 1’s first two subtopics, covered in Atomic Structure: Particles, Radius and Isotopes — subatomic particle behaviour in an electric field, and atomic/ionic radius trends explained by nuclear charge and shielding. This page begins where that one ends, moving from the shell-level picture to sub-shells and orbitals.

Where this fits in 9701

9701 replaces “shells” as the whole picture with shells, sub-shells and orbitals — a finer structure that explains things the IGCSE shell model can’t, such as why the third shell doesn’t simply fill before the fourth begins. This is also where you meet ionisation energy properly for the first time: not just “how reactive is this element,” but as measurable, tabulated data you’re expected to interpret.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — AS Level, Topic 1

Both subtopics are examined for elements hydrogen to krypton only, and every atom or ion described is assumed to be in the ground state.

1.3 Electrons, energy levels and atomic orbitals — the terms shell, sub-shell, orbital, principal quantum number (n) and ground state; the number of orbitals and electron capacity of s, p and d sub-shells; the order of increasing energy of sub-shells within the first three shells plus 4s and 4p; full electronic configurations by shell, sub-shell and orbital; explaining configurations in terms of electron energy and inter-electron repulsion; determining electronic configuration from proton number and charge, in full or shorthand notation; electrons-in-boxes notation; the shapes of s and p orbitals; the definition of a free radical.

1.4 Ionisation energy — defining and using first ionisation energy; writing equations for first, second and subsequent ionisation energies; trends in ionisation energy across a period and down a group, and their explanation; variation in successive ionisation energies of one element; the origin of ionisation energy in nuclear attraction; explaining ionisation energy in terms of nuclear charge, atomic/ionic radius, shielding, and spin-pair repulsion; deducing electronic configuration and Periodic Table position from successive ionisation energy data.

There is no Core/Extended tiering at 9701 — every outcome above is examinable for every candidate.

Sub-shells and orbitals

Each shell (n = 1, 2, 3…) is divided into sub-shells, labelled s, p, d and f, and each sub-shell is made of one or more orbitals — regions of space where an electron is likely to be found. Each orbital holds a maximum of two electrons.

Sub-shell	Number of orbitals	Maximum electrons
s	1	2
p	3	6
d	5	10

Sub-shells fill in order of increasing energy, and that order is not simply 1s, 2s, 2p, 3s, 3p, 3d, 4s — the 4s sub-shell is lower in energy than 3d and fills first:

1s  2s  2p  3s  3p  4s  3d  4p

This is why potassium (proton number 19) has the configuration 1s²2s²2p⁶3s²3p⁶4s¹ rather than filling 3d first, and why the shorthand form for iron is written [Ar] 3d⁶4s² — the 4s electrons are still shown as filled first, even though (once present) 3d electrons are slightly lower in energy and are removed before 4s electrons when a transition metal ion forms.

Electrons-in-boxes and Hund’s rule

Electrons-in-boxes notation shows each orbital as a box and each electron as an arrow, with arrow direction representing spin. Electrons occupy separate orbitals within a sub-shell, spins unpaired, before any orbital is doubly occupied — this minimises the repulsion between electrons of like charge, which is exactly the “inter-electron repulsion” reasoning the syllabus asks you to explain configurations with. For fluorine’s 2p⁵: three 2p orbitals hold arrows pointing up in each, then a second arrow (opposite spin) doubles up in only one of them.

Orbital shapes

An s orbital is spherical, centred on the nucleus. A p orbital is shaped like two lobes on either side of the nucleus (sometimes described as a figure-of-eight or dumbbell), and each p sub-shell has three such orbitals oriented along different axes.

Free radicals

A free radical is a species with one or more unpaired electrons — a concept you’ll meet again when organic mechanisms introduce free-radical substitution.

Ionisation energy

The first ionisation energy of an element is the energy needed to remove one electron from each atom in a mole of gaseous atoms, forming a mole of gaseous 1+ ions:

X(g) → X+(g) + e-

Second and subsequent ionisation energies remove a further electron each time, always starting from the ion already formed:

X+(g) → X2+(g) + e-

Trends, and what explains them

Ionisation energy decreases down a group: atomic radius increases, and the extra filled shells between the nucleus and the outer electron increase shielding, both of which weaken the nuclear attraction on the outer electron — outweighing the fact that nuclear charge is also increasing.

Ionisation energy increases across a period: nuclear charge increases while shielding from inner shells stays roughly constant, so the outer electron is held more tightly. Two small dips interrupt this general rise: the first electron removed from a p sub-shell requires less energy than the last electron of the preceding s sub-shell (a new sub-shell of slightly higher energy has begun), and an electron from a doubly-occupied p orbital requires less energy than one from a singly-occupied p orbital, because of spin-pair repulsion between the two electrons sharing that orbital.

Successive ionisation energies as evidence

Because each successive ionisation energy is larger than the last — removing an electron from an increasingly positive ion takes more energy — and because a large jump between two successive values marks the boundary between two different shells, a full set of successive ionisation energies can be used to work out an element’s electron configuration, and from that, its group in the Periodic Table.

Worked example. Magnesium’s first three ionisation energies are approximately IE₁ = 738 kJ/mol, IE₂ = 1451 kJ/mol, IE₃ = 7733 kJ/mol. The jump from IE₂ to IE₃ is more than four times larger than the jump from IE₁ to IE₂ — because the third electron removed comes from a full inner shell, much closer to the nucleus and far less shielded. This large jump after the second ionisation energy is exactly the evidence that places magnesium in Group 2: two outer-shell electrons, then a large jump into the next shell in.

Common mistakes

• Writing 3d before 4s in a full electron configuration. The filling order places 4s first (1s²2s²2p⁶3s²3p⁶4s²3d¹⁰…), even though 4s electrons are removed before 3d electrons when transition metal ions form — both facts are correct and examinable, and are not contradictions.
• Explaining the increase in ionisation energy across a period using only “more protons.” Full marks need nuclear charge and the fact that shielding does not increase to match it.
• Forgetting the small dips at Group 13 and Group 16. A strictly monotonic increase across a period is the wrong picture — s-to-p sub-shell change and spin-pair repulsion both interrupt it.
• Using “energy released” language for ionisation energy. It is always energy required — ionisation energies are endothermic by definition.
• Mixing up which jump in a successive-ionisation-energy graph identifies the group. The jump comes after the number of electrons equal to the group’s outer-shell count have been removed, not before.

Quick revision checklist

• Shells, sub-shells, orbitals, and the number of orbitals/electrons in s, p, d
• The 4s-before-3d filling order, and writing full or shorthand configurations
• Electrons-in-boxes notation and why electrons occupy orbitals singly first
• s and p orbital shapes; the definition of a free radical
• First ionisation energy: definition and equation
• Explaining the down-a-group and across-a-period trends, including the two dips
• Reading successive ionisation energy data to deduce configuration and group

Related resources

• Atomic Structure — the IGCSE/O Level shell model this resource builds on
• Atomic Structure: Particles, Radius and Isotopes — subtopics 1.1–1.2, the rest of Topic 1
• Atoms, Molecules and Stoichiometry at AS — the next AS topic
• Chemical Bonding: Shapes and Intermolecular Forces — electron arrangement applied to bonding
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.