1.4Unit 1. Structure, Bonding and Introduction to Organic Chemistry

Electrons & Ions — IAL Edexcel Modular (Chemistry) Revision Notes & Topical Questions

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1.4 Electrons & Ions — IAL EDEXCEL MODULAR (CHEMISTRY) notes: Define first and successive ionisation energies, describe the shapes of orbitals, write electronic configurations of atoms and ions, and link configuration to chemical properties and the blocks of the Periodic Table.Fi…

First ionisation energy
Second ionisation energy
Later ionisation energies
Always endothermic
Nuclear charge
Distance from the nucleus

Section 1 of 11Ionisation energy

1First ionisation energy

First ionisation energy the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions

EquationX(g) → X⁺(g) + e^-, for example Na(g) → Na⁺(g) + e^- (+496 kJ mol^-1).
Units are kJ mol^-1, measured under standard conditions.

2Second ionisation energy

Second ionisation energy the energy required to remove one mole of electrons from one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions

EquationX⁺(g) → X²⁺(g) + e^-.

3Later ionisation energies

Later ionisation energies the nth IE removes electrons from gaseous (n−1)+ ions
EquationX⁽ⁿ⁻¹⁾⁺(g) → Xⁿ⁺(g) + e^-, for example Al²⁺(g) → Al³⁺(g) + e^-.

4Always endothermic

Always endothermic energy must be supplied to overcome the attraction between the electron and the nucleus
All ionisation energies are therefore positive.

Section 2 of 11Factors affecting ionisation energy

1Nuclear charge

Nuclear charge more protons attract the electrons more strongly, so ionisation energy increases

2Distance from the nucleus

Distance from the nucleus attraction falls off rapidly with distance, so outer electrons further away are removed more easily

3Shielding

Shielding more inner shells and sub-shells shield the outer electron from the nucleus, so it is removed more easily

4Spin-pair repulsion

Spin-pair repulsion two electrons in one orbital repel each other, so a paired electron is removed more easily

5Down a group the first IE decreases

Down a group the first IE decreases larger atomic radius and more shielding outweigh the extra nuclear charge
Group 1 metals have low ionisation energies and the noble gases have very high ones.

Section 3 of 11Successive ionisation energies

Successive ionisation energies

Successive IEs always increase

Successive IEs always increase each electron is removed from an increasingly positive ion

Large jumps show a change of shell

Large jumps show a change of shell the next electron comes from a shell closer to the nucleus
Smaller jumps within a shell show a change of sub-shell, which supports the sub-shell model.

Deducing the group

Deducing the group the position of the first large jump gives the number of outer electrons
A huge jump after the first IE places an element in group 1, as in sodium (1s² 2s² 2p⁶ 3s¹).
A huge jump after the second IE places an element in group 2, as in magnesium.
A huge jump after the third IE places an element in group 3, as in aluminium.

Use of the data

Use of the data confirms the electronic configuration and the number of outer-shell electrons

Section 4 of 11Shells, sub-shells and orbitals

Shells
- numbered n = 1, 2, 3 and so on, with the energy rising as n increases
- Shell n holds up to 2n² electrons.

Sub-shell	Orbitals	Maximum electrons
s	1	2
p	3	6
d	5	10
f	7	14

Orbital

Orbital a region of space that holds up to two electrons of opposite spin

This is the Pauli exclusion principle.

Occurrence

Occurrence the first shell has only a 1s orbital, p orbitals begin at n = 2 and d orbitals at n = 3

Degenerate orbitals

Degenerate orbitals orbitals in the same sub-shell, such as 2p_x, 2p_y and 2p_z, have equal energy

Section 5 of 11Shapes of orbitals

s orbitals

s orbitals spherical
They get larger as the shell number increases, so 3s is bigger than 1s.

p orbitals

p orbitals dumbbell-shaped with two lobes
The three p orbitals lie along the x, y and z axes at right angles to each other.
The lobes become larger and longer as the shell number increases.

d orbitals

d orbitals their shapes are not required

Section 6 of 11Filling orbitals

Energy order

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

Aufbau principle

Aufbau principle sub-shells are filled in order of increasing energy
The 4s sub-shell fills before the 3d sub-shell.

Hund's rule

Hund's rule orbitals of equal energy are filled singly with parallel spins before any pairing
Three electrons in a p sub-shell occupy p_x, p_y and p_z singly, so nitrogen has three unpaired 2p electrons.

Electrons-in-boxes

Electrons-in-boxes each box is an orbital and each electron is an arrow, with opposite arrows for opposite spins
The boxes are arranged in order of increasing energy.

Section 7 of 11Writing electronic configurations

Find the number of electrons from the atomic number and adjust for any charge.

Fill the sub-shells in order from 1s, putting at most two electrons in each orbital.

Fill degenerate orbitals singly first.

Write the sub-shells in numerical order, with 3d before 4s.

For positive ions, remove electrons from the outermost shell, and from 4s before 3d.

Full configuration

Full configuration lists every sub-shell from 1s
Potassium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹.
Germanium is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p².

Shorthand configuration

Shorthand configuration the core is replaced by the preceding noble gas in brackets
Potassium is [Ar] 4s¹.

Exceptions

Exceptions chromium is [Ar] 3d⁵ 4s¹ and copper is [Ar] 3d¹⁰ 4s¹
A half-full or full 3d sub-shell is energetically stable.
Iron is [Ar] 3d⁶ 4s² and Fe³⁺ is [Ar] 3d⁵.

Section 8 of 11Blocks of the Periodic Table and chemical properties

Blocks of the Periodic Table and chemical properties

Blocks

Blocks an element's block is the sub-shell that holds its outermost electron
s-block elements are in groups 1 and 2, p-block elements are in groups 3 to 0, and d-block elements are the transition metals.

Periods and groups

Periods and groups elements in a period have the same number of shells and elements in a group have the same number of outer electrons

Reactions

Reactions atoms lose or gain electrons to reach a stable noble gas configuration
Mg²⁺ and O^2- both have the configuration [Ne].

Noble gases

Noble gases full outer sub-shells and high ionisation energies make them very unreactive

Section 9 of 11More from the specification

Subshell shielding and atomic energy levels

Subshell shielding and atomic energy levels successive ionisation energies increase sharply when electrons are extracted from inner, less shielded quantum energy levels.

Subshells

Each principal shell is split into s, p, d and f subshells, holding up to 2, 6, 10 and 14 electrons respectively.

Quantum numbers

The principal quantum number n labels the shell and its energy, and each orbital holds a maximum of two electrons with opposite spin.

Electronic structure

Electrons fill the lowest energy subshell first, so calcium is 1s2 2s2 2p6 3s2 3p6 4s2, and the 4s fills before 3d.

Valence electrons

The outer-shell electrons determine bonding, and are the ones lost in forming ions, so calcium forms Ca²⁺.

Ionisation energy

The energy needed to remove one mole of electrons from one mole of gaseous atoms, an endothermic process;
- successive values show the shell structure.

Subshell

A set of orbitals of the same type within a shell: s, p, d or f.

Quantum

A packet of energy.

Valence electron

An electron in the outermost shell of an atom.

Energy level

A main shell of fixed energy labelled by principal quantum number n.

Section 10 of 11Key terms

Term	Meaning
First ionisation energy	The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Orbital	A region of space around the nucleus that holds up to two electrons of opposite spin.
Sub-shell	A group of orbitals of the same type and energy within a shell.
Principal quantum number	The number n of the main energy level (shell) an electron occupies.
Degenerate orbitals	Orbitals of equal energy in the same sub-shell.
Hund's rule	Orbitals of equal energy are filled singly before electrons pair up.
Shielding	Repulsion of outer electrons by inner electrons, reducing the nuclear attraction.

Section 11 of 11Exam technique

Writing an ionisation energy as exothermic — all ionisation energies are endothermic and positive

Forgetting state symbols in ionisation equations — the atoms and ions must be gaseous

Removing 3d electrons before 4s when forming transition metal ions — 4s electrons are lost first

Giving chromium as [Ar] 3d⁴ 4s² or copper as [Ar] 3d⁹ 4s² — they are [Ar] 3d⁵ 4s¹ and [Ar] 3d¹⁰ 4s¹

Pairing electrons in a p sub-shell before each orbital has one — Hund's rule requires single occupancy first

Last updated: 15 October 2026