amazon.co.ukA* revision for atomic structure: isotopes, time-of-flight mass spectrometry, electron configuration and ionisation energy trends, mapped to AQA, Edexcel, OCR, WJEC and CCEA specifications with exam-style questions.
Atoms consist of a nucleus of protons and neutrons surrounded by electrons in orbitals. Protons and neutrons each have a relative mass of approximately 1 (the neutron is very slightly heavier at 1.0087), while the electron has a relative mass of 1/1836 — effectively all the mass sits in the nucleus. Protons carry a charge of +1, electrons carry −1 and neutrons are neutral, so an atom is neutral because the number of protons equals the number of electrons.
| Particle | Relative mass | Relative charge |
|---|---|---|
| Proton | 1 | +1 |
| Neutron | 1 | 0 |
| Electron | 1/1836 | −1 |
The atomic (proton) number Z defines the element and fixes its position in the periodic table; the mass number A is protons plus neutrons. Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons — same Z, different A. Isotopes have identical chemical properties (same electron configuration, so same bonding and reactions) but different physical properties such as mass, density and diffusion rate. Radioactive isotopes additionally have unstable nuclei.
Relative atomic mass (Ar) is defined as the weighted mean mass of an atom of an element compared with 1/12 of the mass of one atom of carbon-12. Relative isotopic mass is the mass of one isotope compared with the same standard — for light isotopes this equals the mass number. Relative molecular mass (Mr) is the sum of the relative atomic masses in a molecule; relative formula mass is used for ionic compounds.
Mass spectrometry of elements: a time-of-flight (TOF) instrument works in four stages. (1) Ionisation: in electron impact (EI) the vaporised sample is bombarded with high-energy electrons which knock one electron off each atom to give M+(g) ions; in electrospray ionisation the sample is dissolved and protonated to MH+, a gentler method suited to fragile biological molecules. (2) Acceleration: an electric field accelerates all ions to the same kinetic energy. (3) Ion drift: ions enter a field-free tube where lighter ions travel faster — time of flight is inversely proportional to √(mass), since KE = ½mv². (4) Detection: ions strike a detector producing a current proportional to their abundance. Exam calculations often ask you to find a time of flight ratio or the Ar: take each isotopic mass, multiply by its relative abundance (signal area), total these and divide by total abundance.
Electrons occupy orbitals — regions of space that can hold a maximum of two electrons with opposite spins. An s orbital is spherical; there are three p orbitals (px, py, pz) shaped like dumbbells at right angles; there are five d orbitals and seven f orbitals. Sub-shells group orbitals of the same type: s holds 2, p holds 6, d holds 10 electrons. Sub-shell energies increase in the order 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p, so the 4s orbital fills before 3d.
Two rules govern filling: the Pauli exclusion principle (no orbital holds more than two electrons, and they must have opposite spin) and Hund's rule (within a sub-shell electrons occupy orbitals singly with parallel spins before pairing up, minimising repulsion). Two exceptions must be known: chromium is [Ar]3d54s1 and copper is [Ar]3d104s1, because a half-filled or filled d sub-shell is more stable than the expected arrangement.
For ions, remember the rule: 4s electrons are removed first. Although 4s fills before 3d, once the 3d sub-shell is occupied it drops below 4s in energy, so the 4s orbital is emptied first during ionisation. Fe (Z=26): 1s²2s²2p⁶3s²3p⁶3d⁶4s²; Fe²⁺: 1s²2s²2p⁶3s²3p⁶3d⁶; Fe³⁺: 1s²2s²2p⁶3s²3p⁶3d⁵ (a stable half-filled d sub-shell, which is why Fe³⁺ is more stable than Fe²⁺). Other examples: Cl⁻ is 1s²2s²2p⁶3s²3p⁶ (isoelectronic with Ar); Ca²⁺ is [Ar]; Na⁺ is [Ne].
Writing configurations from a noble gas core saves time: Al = [Ne]3s²3p¹. Exam tip: when asked for the full electron configuration of a transition metal ion, always write the 3d before 4s in the final answer even though 4s filled first — most boards penalise 3d written after 4s for ions.
First ionisation energy is the enthalpy change when one mole of gaseous 1+ ions forms from one mole of gaseous atoms: Na(g) → Na⁺(g) + e⁻, ΔH = +496 kJ/mol. It is always endothermic — energy must be supplied to overcome the attraction between the nucleus and the electron. Successive ionisation energies measure removal of electrons one after another from the same atom, and each is larger than the last because the same positive charge attracts fewer electrons more strongly and the ion is increasingly positive.
Three factors control ionisation energy. Nuclear charge: more protons pull electrons harder. Distance: the further the outer electron is from the nucleus, the weaker the attraction. Shielding: inner shells repel outer electrons and screen them from the nuclear charge.
Trends: ionisation energy increases across a period — nuclear charge rises while shielding stays roughly constant because electrons enter the same outer shell, so Zeff increases. Small dips break the trend and are standard A*/explain questions: aluminium (3s²3p¹) has a lower first IE than magnesium (3s²) because the highest-energy electron is in the higher-energy 3p sub-shell, which is easier to remove; sulfur (3p⁴) has a lower first IE than phosphorus (3p³) because sulfur's paired 3p electrons repel each other, so one is easier to remove. Ionisation energy decreases down a group — each element has an extra full inner shell, so atomic radius and shielding both increase and outweigh the rise in nuclear charge.
Graphs of log(successive IE) against electron number provide the classic evidence for shells: within one shell the increase is gradual, but a large jump occurs when the next electron comes from a new, inner shell because it is much closer to the nucleus and far less shielded. Counting the number of electrons before the first big jump tells you the number of outer-shell electrons, hence the group. For example, an element with a huge jump after the fourth ionisation energy is in Group 4.
A mass spectrum of an element shows one line per isotope: the position on the m/z axis gives the isotopic mass and the height gives the relative abundance. From this you calculate Ar. Worked example: chlorine gives peaks at m/z 35 (75.0%) and m/z 37 (25.0%), so Ar = (35 × 75.0 + 37 × 25.0) / 100 = 35.5.
Two diatomic patterns are examined constantly. Chlorine (³⁵Cl 75%, ³⁷Cl 25%) gives molecular-ion peaks for Cl₂⁺ at m/z 70, 72 and 74 in the ratio 9:6:1 (0.75² : 2 × 0.75 × 0.25 : 0.25²), while any fragment containing one Cl atom (e.g. CH₃Cl⁺) appears as a doublet at m/z 49/51 or 35/37 in a 3:1 ratio. Bromine (⁷⁹Br 50%, ⁸¹Br 50%) gives 1:1 doublets, and Br₂⁺ peaks at m/z 158, 160, 160, 162 in the ratio 1:2:1. These ratios let you deduce how many halogen atoms a fragment contains — a favourite AS-level question.
TOF calculation style: two ions of the same charge given the same kinetic energy — the heavier ion flies more slowly. If ion A (mass mA) takes time tA and ion B takes tB, then tB/tA = √(mB/mA). Rearrange to find the unknown mass, then identify the isotope. Always work in kg for mass (divide by 1000 from gmol⁻¹ × 10³ to match SI joules) unless the question asks for a ratio.
Mass spectra and ionisation energy graphs reward precise method writing. For TOF calculations, state the principle first — same charge, same kinetic energy, so velocity depends on mass — then apply tB/tA = √(mB/mA). For Ar calculations from abundance data, show the weighted-mean expression with all terms before evaluating; the expression itself carries method marks even if arithmetic slips. When asked to identify an element from successive ionisation energies, count the electrons before the large jump, state the number of outer-shell electrons, then name the group — three separate reasoning steps, three marks.
For ionisation energy trend questions, the mark scheme expects the three factors named explicitly: nuclear charge (number of protons), distance of the outer electron from the nucleus, and shielding by inner shells. 'More protons attract more strongly' alone scores one mark; adding 'same shell so shielding unchanged, so effective nuclear charge increases' scores the rest. For discontinuity explanations, name the sub-shells of both elements compared — Mg 3s² vs Al 3s²3p¹, or P 3p³ vs S 3p⁴ — before giving the reason (higher-energy sub-shell, or paired-electron repulsion).
Electron configuration answers: write full configurations in the order asked (AQA asks ascending order 1s first; for transition metal ions, 3d before 4s in the final answer), use superscript notation correctly, and double-check the electron count against the charge — an Fe²⁺ answer with 24 electrons is a silent error that costs a mark. Common traps: writing 4s² in a transition-metal ion's configuration; forgetting the Cr/Cu exceptions; treating isotopes as chemically different; giving ionisation energy as exothermic. Practise writing the causal chain — extra inner shell → more shielding and greater distance → weaker attraction → lower IE — because chained explanations are what separate A* scripts from A/B.