Atomic Structure and the Periodic Table
This domain tests how the quantum-mechanical model of the atom explains the layout of the periodic table and the trends you are asked to predict: electron configuration, ionisation energy, atomic and ionic radius, electronegativity, and the evidence available from mass spectra and successive ionisation energies.
Subatomic structure and isotopes
- Atomic number Z = protons; mass number A = protons + neutrons; isotopes differ only in neutrons, ions only in electrons.
- Relative atomic mass is the abundance-weighted mean of the isotope masses, not the mass number of any single isotope.
- A mass spectrum gives m/z peaks with relative abundances; the molecular ion peak gives Mr, and the M+1 peak (about 1.1% per carbon) indicates the number of carbon atoms.
Electron configuration
- Fill by Aufbau order, obey the Pauli exclusion principle and Hund's rule (maximise unpaired spins within a subshell).
- 4s fills before 3d, but when transition metals ionise the 4s electrons are removed first.
- Cr [Ar]3d5 4s1 and Cu [Ar]3d10 4s1 are the standard exceptions to expect.
- Blocks are named for the subshell receiving the last electron.
Periodic trends
- Across a period, nuclear charge rises while shielding is roughly constant, so effective nuclear charge increases: radius falls, ionisation energy and electronegativity rise.
- Down a group, extra shells increase shielding and distance: radius rises, ionisation energy falls.
- Successive ionisation energies jump sharply at shell boundaries; a large jump after the second value means two outer electrons, so Group 2.
Worked example
Rubidium is 72.2% 85Rb and 27.8% 87Rb. Ar = (85 × 0.722) + (87 × 0.278) = 61.37 + 24.19 = 85.56. Now an element with first ionisation energies 738, 1451 and 7733 kJ/mol: the third value is about five times the second, so two electrons are removed easily and the third comes from an inner shell - a Group 2 element.
Traps that cost marks
- Quoting a periodic-table atomic mass as the mass number of a particular isotope, or forgetting to weight by abundance.
- Defining first ionisation energy without "gaseous atoms" and "one mole".
- Assuming Mg's first ionisation energy exceeds Al's; Al's 3p electron is shielded by the filled 3s, so Al is lower.
- Writing an electron configuration for a transition-metal ion without emptying 4s first.
- Calling Zn a typical transition metal; its common ion is d10, so it lacks the characteristic variable oxidation states and colour.