Isotopes and relative masses

Atoms of the same element always have the same number of protons (the atomic number, Z) but can have different numbers of neutrons. Those variants are isotopes. Because chemistry is controlled by electrons, isotopes of an element are chemically identical — they differ only in mass (so in physical properties such as density and rate of diffusion).

• Relative isotopic mass — the mass of an atom of an isotope relative to ¹⁄₁₂ the mass of a carbon-12 atom.
• Relative atomic mass, A_r — the weighted mean mass of the atoms of an element, relative to ¹⁄₁₂ the mass of a carbon-12 atom.
• Relative molecular mass, M_r — the same idea applied to a molecule (use relative formula mass for giant/ionic structures such as NaCl).

Exam habit: relative masses have no units — they are ratios. Quote A_r to the same precision as the data you were given (usually 1 or 2 decimal places).

How a mass spectrometer measures A_r

A time-of-flight (TOF) mass spectrometer separates particles by their mass/charge ratio (m/z):

• Ionisation — electron impact (a high-energy electron knocks one electron out) or electrospray (the sample is protonated). Only ions can be accelerated and deflected by electric/magnetic fields.
• Acceleration — all ions are given the same kinetic energy, so heavier ions travel more slowly.
• Flight tube / detection — the lighter the ion, the shorter its time of flight; the detector records the abundance of each m/z.

For a singly charged ion (z = 1), the m/z value is the relative mass of that ion.

Your turn — weighted mean mass

Your turn — working backwards

Diatomic molecules give a pattern of peaks

When a diatomic molecule of an element is ionised without fragmenting, you see a peak for every possible combination of isotopes — and the peak heights follow the probabilities.

Chlorine gas (³⁵Cl 75%, ³⁷Cl 25%) gives molecular ions at m/z 70, 72 and 74:

• m/z 70 = ³⁵Cl–³⁵Cl → 0.75 × 0.75 = 0.5625
• m/z 72 = ³⁵Cl–³⁷Cl (either way round) → 2 × 0.75 × 0.25 = 0.3750
• m/z 74 = ³⁷Cl–³⁷Cl → 0.25 × 0.25 = 0.0625

Also expect peaks at m/z 35 and 37 from the Cl⁺ fragment, in a 3 : 1 ratio.

Your turn — relative peak heights

Your turn — real spectrometer data

Defining ionisation energies

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.

Ionisation energies are always endothermic (positive): you must do work against the attraction of the nucleus. Three factors decide the size:

• Nuclear charge — more protons pull harder → higher IE.
• Atomic radius — the further the electron, the weaker the attraction (attraction falls off with distance) → lower IE.
• Shielding — inner-shell electrons repel the outer electron and screen it from the nucleus → lower IE.

Successive IEs always rise: each electron is pulled from an increasingly positive ion. A big jump means you have broken into a new (inner) shell — that is how IE data tells you the group number.

Which equation shows the second ionisation energy of calcium?

Your turn — reading IE data

Ionisation energy across a period and down a group

Across Period 3 (Na → Ar) the first ionisation energy generally increases: the nuclear charge rises, the electrons enter the same shell so shielding is roughly constant, and the atomic radius falls. The outer electron is held more tightly.

Down a group (e.g. Li → Cs) the first ionisation energy decreases: each element has an extra full inner shell, so the outer electron is further from the nucleus and better shielded. This outweighs the increase in nuclear charge.

Why does aluminium buck the trend?

And the sulfur dip?

Shells, sub-shells and orbitals

An orbital is a region of space in an atom that can hold up to two electrons — and only if they have opposite spins.

• s orbitals are spherical; there is one per shell → an s sub-shell holds 2 electrons.
• p orbitals are dumb-bell shaped, in three orientations (p_x, p_y, p_z) → a p sub-shell holds 6.
• A d sub-shell has 5 orbitals → 10 electrons; an f sub-shell holds 14.

Quantum shells fill 2, 8, 18, 32 (= 2n²) electrons. Electrons fill the lowest energy sub-shell first, and 4s fills before 3d because 4s is slightly lower in energy.

Writing electronic configurations

Fill in order of increasing energy: 1s 2s 2p 3s 3p 4s 3d 4p. Within a sub-shell, electrons occupy orbitals singly first (Hund's rule) before pairing up.

• Ti: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s² — or [Ar] 3d² 4s²
• Exceptions: Cr = [Ar] 3d⁵ 4s¹ and Cu = [Ar] 3d¹⁰ 4s¹ (a half-full or full d sub-shell is more stable).
• Ions of d-block metals lose 4s electrons first: Fe = [Ar] 3d⁶ 4s² → Fe²⁺ = [Ar] 3d⁶ → Fe³⁺ = [Ar] 3d⁵.

Blocks: the block is named after the sub-shell the outermost electron occupies — s-block (Groups 1 and 2), p-block (Groups 3–0), d-block (the transition series).

d-block ions

s-block, p-block or d-block?

Tap an element, then tap the block its electronic configuration puts it in.

Why ionise the sample?

Periodicity across Period 3

Periodicity = a repeating pattern of properties across the periods, caused by the repeating pattern of electronic configurations.

• Atomic radius decreases across a period: more protons, same shell, same shielding → electrons pulled in.
• First ionisation energy increases across a period (with the Al and S dips).
• Melting temperature across Period 3: rises Na < Mg < Al (metallic bonding gets stronger: more delocalised electrons per atom, smaller ion), peaks sharply at Si (giant covalent — strong covalent bonds must break), then collapses at P₄, S₈, Cl₂ and Ar (simple molecular — only weak London forces between molecules; S₈ > P₄ > Cl₂ because larger molecules have more electrons).

Exam trap: when you explain the drop after silicon, say that intermolecular (London) forces are overcome — not covalent bonds.

Melting temperatures in Period 3

Match the idea to the evidence

Tap an item on the left, then its partner on the right.

The big ideas to know

You've now covered Topic 1: Atomic Structure and the Periodic Table of the Edexcel A-level Chemistry (9CH0) specification. Press Finish to see your score.