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Atomic structure is the foundation of chemistry: everything from bonding to the periodic table to reaction energetics builds on it. The good news is that it’s a topic with clear rules and predictable exam questions. This guide organises it into seven units, gives the core facts for each, includes checkpoint questions with answers, and links to a detailed article for every idea. Work through it in order, or jump to the units you find hardest.
Unit 1: The three subatomic particles
Core facts
| Particle | Relative charge | Relative mass | Location |
|---|---|---|---|
| proton | +1 | 1 | nucleus |
| neutron | 0 | 1 | nucleus |
| electron | −1 | about 1/1836 | around the nucleus, in shells/orbitals |
- Atoms are neutral: number of electrons = number of protons.
- The nucleus is tiny (about 1/100,000 of the atom’s diameter) but holds almost all the mass.
Read more: what is an atom?, the proton, the neutron, the electron, how small is an atom?
Checkpoint
- Which particle has almost no mass?
- Why is an atom neutral overall?
Unit 2: Atomic number, mass number and isotopes
Core facts
- Atomic number (Z) = number of protons; it defines the element.
- Mass number (A) = protons + neutrons.
- Neutrons = A − Z.
- Isotopes have the same Z but different A (different numbers of neutrons). They have the same chemical properties but different masses.
- In nuclide notation, ᴬ_Z X: mass number top left, atomic number bottom left.
Read more: atomic number explained, mass number explained, isotope notation
Checkpoint 3. How many protons, neutrons and electrons are in ⁵⁶₂₆Fe? 4. What’s the difference between chlorine-35 and chlorine-37?
Unit 3: Relative atomic mass
Core facts
- Relative atomic mass (Aᵣ) is the weighted average mass of an element’s isotopes compared with 1/12 of the mass of a carbon-12 atom.
- Aᵣ = Σ(isotope mass × abundance) ÷ Σ(abundance)
- Mass spectrometers measure isotope masses and abundances.
Read more: relative atomic mass calculations, finding isotope abundances, mass spectrometry and isotopes
Checkpoint 5. Boron is 20% boron-10 and 80% boron-11. Calculate its Aᵣ.
Unit 4: The history of atomic models
Core facts
| Scientist | Date | Contribution |
|---|---|---|
| Dalton | 1803–1808 | atoms as indivisible particles; elements have characteristic atoms |
| Thomson | 1897 / 1904 | discovered the electron; plum pudding model |
| Rutherford | 1911 | gold foil experiment; nuclear model |
| Bohr | 1913 | electrons in fixed energy levels |
| Chadwick | 1932 | discovered the neutron |
| Schrödinger and others | 1926 onwards | quantum mechanical model with orbitals |
For each, know the evidence and why the model changed. The gold foil experiment is the most frequently examined: most alpha particles passed straight through (atom mostly empty space), a few were deflected strongly (tiny dense positive nucleus).
Read more: Dalton’s atomic theory, Thomson’s plum pudding model, Rutherford’s gold foil experiment, the Bohr model, Chadwick and the neutron, the quantum mechanical model
Checkpoint 6. What did Rutherford conclude from alpha particles bouncing back from gold foil?
Unit 5: Electron arrangement
Core facts (introductory level)
- Electrons occupy shells: 2, 8, 8, 2 for the first 20 elements.
- Period number = number of occupied shells; group (1–2, 13–18) relates to the number of outer electrons.
Core facts (advanced level)
- Shells contain subshells (s, p, d, f) holding 2, 6, 10 and 14 electrons.
- Filling order: 1s 2s 2p 3s 3p 4s 3d 4p…
- Rules: Aufbau (lowest energy first), Pauli (two electrons per orbital, opposite spins), Hund (fill equal orbitals singly first).
- Exceptions: Cr [Ar] 3d⁵ 4s¹ and Cu [Ar] 3d¹⁰ 4s¹.
Read more: electron shells explained, electron configuration rules, noble gas shorthand, orbital diagrams, electron configuration exceptions
Checkpoint 7. Write the shell arrangement of chlorine and the full configuration of iron.
Unit 6: Ions
Core facts
- Metals lose outer electrons to form positive ions; non-metals gain electrons to form negative ions.
- Ions often have noble gas configurations: Na⁺ and F⁻ both have [Ne].
- Transition metals lose 4s electrons before 3d: Fe²⁺ is [Ar] 3d⁶.
- In an isoelectronic series, more protons means a smaller ion.
Read more: what is an ion?, electron configurations of ions, isoelectronic species
Checkpoint 8. Write the configuration of Fe³⁺.
Unit 7: Spectra and ionisation energy
Core facts
- Electrons absorb or emit photons when they move between energy levels: ΔE = hf.
- Line spectra prove that energy levels are quantised; each element has a unique spectrum.
- Hydrogen’s visible lines (Balmer series) come from electrons falling to n = 2.
- First ionisation energy increases across a period and decreases down a group; dips occur at group 13 and group 16.
- Successive ionisation energies show a big jump when an electron is removed from a new, inner shell, revealing the group.
Read more: photons and energy levels, the hydrogen emission spectrum, flame test colours, ionization energy trend, effective nuclear charge
Checkpoint 9. Why do atoms give line spectra rather than continuous spectra? 10. An element’s successive ionisation energies (kJ/mol) are 738, 1451, 7733, 10543. Which group is it in?
Checkpoint answers
- The electron.
- It has equal numbers of protons (+) and electrons (−).
- 26 protons, 30 neutrons, 26 electrons.
- Chlorine-37 has two more neutrons (20 instead of 18); same protons and electrons.
- (10 × 20 + 11 × 80) ÷ 100 = 10.8.
- The atom is mostly empty space, with its positive charge and most of its mass in a tiny, dense nucleus.
- Cl: 2, 8, 7. Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ (or [Ar] 3d⁶ 4s²).
- [Ar] 3d⁵.
- Electrons can only have certain energies, so only certain energy differences, and therefore certain photon wavelengths, are possible.
- The big jump is between the 2nd and 3rd ionisation energies, so it has 2 outer electrons: group 2 (it’s magnesium).
A one-week revision plan
| Day | Focus | Practice |
|---|---|---|
| 1 | Units 1–2 | atomic structure practice questions |
| 2 | Unit 3 | isotope and Aᵣ calculations |
| 3 | Unit 4 | explain each model change in two sentences |
| 4 | Unit 5 | electron configuration practice |
| 5 | Unit 6 | ion configurations, isoelectronic ranking |
| 6 | Unit 7 | spectra, ionisation energy graphs |
| 7 | mixed review | atomic structure misconceptions; redo weakest checkpoints |
Exam technique tips
- Show working in calculations (Aᵣ, neutrons), and give Aᵣ without units.
- Link evidence to conclusion in history questions (“most passed through → mostly empty space”).
- Use precise vocabulary: shell, subshell, orbital; atom, ion, isotope.
- Check electron counts by adding superscripts and comparing with Z (minus the charge).
- Explain trends using nuclear charge, distance and shielding, not “atoms want”.
- Draw diagrams where they help: a labelled atom, a shell diagram or an energy-level diagram often earns marks faster than words.
Key takeaways
- Protons define the element; neutrons give isotopes; electrons determine chemistry and form ions.
- Aᵣ is a weighted average of isotope masses.
- Atomic models changed as new evidence appeared; know the evidence for each change.
- Electron configurations follow Aufbau, Pauli and Hund, with known exceptions.
- Spectra and ionisation energies provide the evidence for energy levels and shells.
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