Atomic Structure and Properties
Coulombic attraction drives every trend
The nucleus–electron attraction is proportional to the product of charges and inversely proportional to distance (Coulomb's law). Effective nuclear charge (Z_eff) increases left to right across a period because protons are added while electrons enter the same shell; shielding by core electrons stays roughly constant. This single idea explains smaller atomic radii, higher first ionization energies, and more negative electron affinities across a period. Down a group, added shells dominate, so radius increases while ionization energy falls.
Electron configurations follow the Aufbau order, Pauli exclusion, and Hund's rule. Chromium and copper are the classic exceptions because a half-filled or filled d subshell is slightly lower in energy. Photoelectron spectroscopy (PES) is the exam's favorite graph: each peak is a subshell, peak height gives the number of electrons in that subshell, and binding energy increases for electrons closer to the nucleus (1s > 2s > 2p > 3s).
Traps: ionization energy dips at group 13 (removing a p electron from a higher-energy subshell) and at group 16 (removing a paired p electron, which relieves repulsion). Comparing ions requires counting protons: in an isoelectronic series, the most positive ion is smallest. Average atomic mass is a weighted average, so the tallest mass-spectrum peak is not necessarily closest to the average.
Worked example: rank N3−, O2−, F−, Na+, Mg2+, Al3+. All have 10 electrons, so radius falls as nuclear charge rises: N3− > O2− > F− > Na+ > Mg2+ > Al3+.
Rule of thumb: higher Z_eff pulls electrons closer, so radius shrinks and ionization energy rises.