Periodic Classification of Elements
Unit 3 tests the modern periodic law and your ability to predict and explain trends in atomic size, ionization enthalpy, electron gain enthalpy and electronegativity.
Core rules
- Properties are periodic functions of atomic number; the table has 7 periods and 18 groups, divided into s, p, d and f blocks.
- Across a period, nuclear charge rises while electrons enter the same shell, so atomic radius falls and ionization enthalpy and electronegativity generally rise.
- Down a group, added shells dominate, so radius increases and ionization enthalpy falls.
- For isoelectronic species, the species with the highest nuclear charge is smallest.
- Anomalies arise from stable configurations: Be and N have higher ionization enthalpies than B and O respectively; F has a less negative electron gain enthalpy than Cl because of its small size.
- Diagonal relationships exist between Li–Mg, Be–Al and B–Si.
Worked example
Arrange N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺ and Al³⁺ in order of decreasing radius. All have 10 electrons, so size depends on nuclear charge: 7, 8, 9, 11, 12 and 13 protons respectively. Greater nuclear charge pulls the same electron set closer, so the order is N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺.
Traps
- Do not compare noble gas radii directly with neighbouring elements; they are measured differently.
- Ionization enthalpy increases across a period but dips at Be→B and N→O.
- Electron gain enthalpy of Cl is more negative than that of F.
- Cations are smaller than their parent atoms; anions are larger.
- Metallic character increases down a group and decreases across a period.