Electrons & Ions

Revise Electrons & Ions for Chemistry 1 (IAL) WCH11 (AS Level) — revision notes and instant AI marking.

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Edexcel IAL Chemistry · Unit 1

Electrons & Ions

Big idea: Every atom's electrons sit in an ordered, energy-ranked "address system" (shells → subshells → orbitals), and pulling one electron out of that system — ionisation — always costs energy, with the exact amount telling you precisely where that electron used to live.

Summary — What This Chapter Covers

Ionisation energy (IE) is the energy needed to remove one mole of electrons from one mole of gaseous atoms/ions — it's always endothermic (energy IN).
Successive ionisation energies (1st, 2nd, 3rd...) get progressively bigger, with big jumps revealing shell changes and small jumps revealing subshell changes.
Four factors control the size of an ionisation energy: nuclear charge, distance from nucleus, shielding, and spin-pair repulsion.
Orbitals are regions where electrons are found — s orbitals are spherical, p orbitals are dumbbell-shaped, each orbital holds a max of 2 electrons (opposite spins).
Subshells (s, p, d, f) contain a fixed number of orbitals: s=1, p=3, d=5, f=7 — giving max electron capacities of 2, 6, 10, and 14.
Energy order for filling: shells increase with principal quantum number n; subshells go s < p < d < f — except 4s fills before 3d.
Hund's Rule — electrons fill separate orbitals in a subshell singly (same spin) before pairing up.
Pauli Exclusion Principle — a single orbital can only ever hold 2 electrons, and they must have opposite spins.
Electron configurations can be written in full (1s² 2s² 2p⁶...) or shorthand using the nearest preceding noble gas (e.g. [Ar] 4s¹).
Cr and Cu are exceptions — they "steal" one 4s electron to make the 3d subshell half-full or full, because that's more energetically stable.
The Periodic Table's shape (s-block, p-block, d-block, f-block) directly reflects where each element's highest-energy electron sits.

1. Ionisation Energy

1.1What Ionisation Energy Actually Means

Imagine an electron as a tiny magnet trapped in orbit around a positively charged nucleus, held there purely by electrostatic attraction. To rip that electron away completely — so it becomes a free, gaseous electron with no more pull from the nucleus — you have to do work against that attraction. That "work" is the ionisation energy.

Formally: the ionisation energy (IE) of an element is the energy required to remove one mole of electrons from one mole of gaseous atoms of that element, to form one mole of gaseous 1+ ions. It's measured under standard conditions (298 K, 101 kPa) and given in kJ mol⁻¹.

Notice three "one mole" conditions packed into that definition — this precision matters in exam answers. Also notice the state symbols: everything must be gaseous. If atoms were in a solid or liquid, neighbouring particles would interfere with the measurement, so chemists always define IE for isolated gas-phase particles.

First Ionisation Energy — General Equation

X (g) → X⁺ (g) + e⁻

Worked example — calcium: Ca (g) → Ca⁺ (g) + e⁻ IE₁ = +590 kJ mol⁻¹

Why is IE always positive (endothermic)?

You are pulling a negative electron away from a positive nucleus — that's fighting an attractive force, so it always takes energy in. There is no such thing as a negative first ionisation energy for a neutral atom.