Nature of Matter and Atomic Structure
Core concepts
Matter is classified by composition (elements, compounds, mixtures) and by particle arrangement (atoms, ions, molecules). Candidates must move fluently between macroscopic descriptions and subatomic models: protons define the element (atomic number, Z), neutrons set the isotope, and electrons in shells/subshells govern chemical behaviour.
Understand the rules of electron configuration: Aufbau (fill lowest energy first), Pauli (two electrons per orbital, opposite spins), and Hund (maximise unpaired spins within a subshell). Know trends in atomic radius, ionisation energy, electron affinity, and electronegativity across periods and down groups, and the reasons behind anomalies (e.g. the stability of half-filled and fully filled subshells).
Traps to watch
- Confusing mass number with relative atomic mass (an average weighted by isotopic abundance).
- Forgetting that ions change electron count, not proton count.
- Mixing up a physical change (state, shape) with a chemical change (new substance).
- Assuming ionisation energy always rises smoothly across a period; there are dips at group-to-group boundaries.
- Treating isotopes as chemically different - they differ in mass and nuclear stability, not in bonding behaviour.
Worked example
Chlorine has isotopes of mass 35 (75%) and 37 (25%). Relative atomic mass = (35 × 0.75) + (37 × 0.25) = 26.25 + 9.25 = 35.5. Rule of thumb: weight each isotope by its fractional abundance, then sum; the answer must lie between the lightest and heaviest isotope masses.