Topic 2 covers the structure of atoms and how electron arrangement determines chemical properties. SL focuses on electron configurations and emission spectra. HL extends to subshells, orbitals, and ionisation energy patterns.
Proton: charge +1, mass 1 u, in nucleus. Neutron: charge 0, mass 1 u, in nucleus. Electron: charge −1, mass ≈ 1/1836 u, in orbitals. Atomic number Z = number of protons (defines the element). Mass number A = protons + neutrons. Isotopes: same Z, different A (same element, different number of neutrons).
Electrons fill subshells in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p... Aufbau principle: fill lowest energy first. Pauli exclusion: max 2 electrons per orbital (opposite spins). Hund\'s rule: fill degenerate orbitals singly first. Write configurations: e.g., Na: 1s² 2s² 2p⁶ 3s¹ or [Ne] 3s¹.
Excited electrons fall to lower energy levels, emitting photons: ΔE = hf = hc/λ. Each element has a unique line spectrum — atomic fingerprint. Hydrogen spectrum: Lyman (to n=1, UV), Balmer (to n=2, visible), Paschen (to n=3, IR) series. Lines converge at higher energies, indicating energy levels get closer together.
First IE: energy to remove one mole of electrons from one mole of gaseous atoms. Trends: increases across a period (increasing nuclear charge), decreases down a group (increasing distance and shielding). Successive IEs: large jump when core electrons are removed. IE data provides evidence for electron shells and subshells.
The 4s subshell has lower energy than 3d due to its greater penetration towards the nucleus. However, once 3d is occupied, 3d becomes lower in energy than 4s. This is why transition metals lose 4s electrons first when forming ions (e.g., Fe: [Ar] 3d⁶ 4s² → Fe²⁺: [Ar] 3d⁶, not [Ar] 3d⁴ 4s²).
Book a Trial + Diagnostic session. Get a personalized Learning Path with clear milestones, tutor match, and a plan recommendation — all within 24 hours.
Book Trial + Diagnostic →