Atomic Structure and Bonding (Units 1 and 2) establishes the foundation — from subatomic particles through electron configuration to all types of chemical bonding and intermolecular forces.
Subatomic particles: proton (p, +1, mass 1), neutron (n, 0, mass 1), electron (e, −1, mass 1/1836). Atomic number Z = protons. Mass number A = protons + neutrons. Isotopes: same Z, different A. Mass spectrometer: ionise → accelerate → deflect → detect. Relative atomic mass from mass spectrum. Electron configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s²... Fill in order: 1s, 2s, 2p, 3s, 3p, 4s, 3d. Exceptions: Cr [Ar] 3d⁵ 4s¹, Cu [Ar] 3d¹⁰ 4s¹. First ionisation energy: energy to remove one mole of electrons from one mole of gaseous atoms. Trends: increases across period (more protons, same shielding), decreases down group (more shells, greater shielding). Dips at Group 3 and 6 explained by sub-shell structure.
Ionic: metal transfers electrons to non-metal → ions → electrostatic attraction in giant lattice. High MP, conduct when molten/dissolved. Covalent: shared electron pairs. Dot-cross diagrams. Dative (coordinate): both electrons from one atom (e.g., NH₄⁺, CO). VSEPR: electron pairs (bonding and lone) repel → shapes. Linear (2 bp, 180°), trigonal planar (3 bp, 120°), tetrahedral (4 bp, 109.5°), trigonal bipyramidal (5 bp), octahedral (6 bp). Lone pairs repel more → e.g., pyramidal (107°), bent (104.5°). Electronegativity: ability to attract bonding electrons. Polar bond: unequal sharing. Polar molecule depends on symmetry. Intermolecular forces: London dispersion (all molecules, temporary dipoles), permanent dipole-dipole, hydrogen bonding (N-H, O-H, F-H with lone pair). H-bonding explains high BP of water, ice density.
Nitrogen has electron configuration 1s² 2s² 2p³ — three electrons in 2p, one in each orbital (maximally stable half-filled sub-shell). Oxygen has 2p⁴ — one 2p orbital has a pair of electrons. The paired electrons in oxygen\'s 2p orbital experience inter-electron repulsion, making it easier to remove one of them. Despite oxygen having more protons (more nuclear charge), this repulsion effect dominates. This creates a dip in the ionisation energy trend across Period 2: it generally increases from Li to Ne, but with dips from Be to B (2s→2p, where 2p is higher energy and easier to remove) and from N to O (half-filled stability). This pattern is key evidence for the sub-shell model of the atom.
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