Atomic structure

Build H2 Chemistry atomic models using orbitals and electron configurations, and interpret first and successive ionisation-energy evidence.

GCE A-Level H2 Chemistry 2027
5 lessons

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Before you begin

Build an atomic model from particles, shells, subshells and orbitals. Then use ionisation-energy evidence to test how the electrons are arranged.

Be comfortable with: atoms, elements and isotopes, electron configuration and the periodic table.

Learning goals

Subatomic Particles, Isotopes, Ions
Atomic Orbitals: Energies and Shapes
Orbitals and Electron Configuration
First Ionisation Energy Trends
Successive Ionisation Energies

Syllabus statements covered

identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses
deduce the behaviour of beams of protons, neutrons and electrons in an electric field
describe the distribution of mass and charges within an atom
deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (and charge)
describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number
distinguish between isotopes on the basis of different numbers of neutrons present
describe the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals
describe the shapes of s, p and d orbitals (see also Section 13) [knowledge of wave functions is not required]
state the electronic configuration of atoms and ions given the proton number (and charge)
explain the factors influencing the ionisation energies of elements (see the Data Booklet) (see also Section 5)
deduce the electronic configurations of elements from successive ionisation energy data
interpret successive ionisation energy data of an element in terms of the position of that element within the Periodic Table

How the ideas connect

When Fe forms Fe²⁺, both electrons leave 4s. The six 3d electrons stay. Boxes show orbital occupancy, not position or an energy scale; the [Ar] core is omitted.

Count the particles
Use proton number and nucleon number. Isotopes change neutron count; forming an ion changes electron count.
Step 1
Arrange the electrons
Fill orbitals using energy, spin and pairing rules. Distinguish filling an atom from removing electrons to form an ion.
Step 2
Interpret the evidence
Explain ionisation-energy trends using nuclear charge, distance and shielding. A large successive jump reveals a more strongly bound shell.
Step 3

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Topic reference

Quick reference

Particles in an ion

Protons = Z; neutrons = A-Z.

For charge + n, electrons = Z-n. For charge -n, electrons = Z + n.

Common pitfall

Forming an ion changes electrons, not protons. Isotopes have the same proton number but different neutron numbers.

Read particles, isotopes and ions →

Remove 4s first

Fe: [Ar]3d⁶ 4s²

Fe²⁺: [Ar]3d⁶

Write the neutral atom’s configuration first, then remove its 4s electrons before its 3d electrons.

Common pitfall

The neutral-atom filling order is not the order for removing electrons from a d-block atom.

Explore the iron-ion simulation →

Find the large jump

Count the electrons removed before the first large successive ionisation-energy jump.

For a main-group element, this reveals the number of outer-shell electrons. Ionisation energy is commonly in kJ mol⁻¹.

Common pitfall

Count removals before the jump, not the label of the ionisation energy after it.

Read the evidence explanation →

Complete quick-reference table

Question	First move	What to show
Particles in an ion	protons = Z; neutrons = A-Z	electrons = Z-n for Xⁿ⁺, Z + n for Xⁿ⁻
A beam in an electric field	the sign of the charge gives the direction	the size of the deflection depends on charge ÷ mass
Capacity of a subshell	s has 1 orbital, p has 3, d has 5	two electrons per orbital: 2, 6 and 10
Configuration of an atom	fill 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p	Cr is [Ar]3d⁵ 4s¹ and Cu is [Ar]3d¹⁰ 4s¹
Configuration of a d-block ion	write the atom’s configuration first	remove 4s electrons before 3d
A trend in IE₁	nuclear charge, distance and shielding	how the attraction for the outer electron changes
A dip in IE₁	Group 2 to 13, or Group 15 to 16?	a new p subshell, or a paired p electron
Successive ionisation energies	find the first large jump	for main-group elements, count the outer-shell electrons removed before it

Core knowledge to remember

Definitions and core knowledge

Proton number, Z: the number of protons in the nucleus. Nucleon number, A: the total number of protons and neutrons.
Isotopes: atoms of the same element with the same number of protons but different numbers of neutrons.
Orbital: a region of space around the nucleus with a high probability of finding an electron; it holds at most two electrons, with opposite spins.
Electrons fill the lowest-energy subshell first, and occupy the orbitals of a subshell singly before pairing.
An s orbital is spherical; a p orbital has two lobes. Four d orbitals have four lobes; d_z² has two lobes along the z axis and a ring around the centre.
First ionisation energy: the energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1 + ions, X(g) → X⁺(g) + e⁻.

Common mistakes

More mistakes to check

Changing the number of protons when an ion forms.
Defining isotopes without both parts: the same number of protons and different numbers of neutrons.
Drawing all five d orbitals with four lobes.
Removing 3d electrons before 4s when a transition metal forms an ion.
Leaving “gaseous” or “one mole” out of the definition of ionisation energy.
Saying shielding increases greatly across a period.
Reading the group from the label of the ionisation energy after the jump.