Properties of atoms change systematically across the periodic table because nuclear attraction, electron shielding and the distance of occupied shells change in regular ways.

Across a period from left to right, nuclear charge increases while added valence electrons remain in roughly the same shell. The resulting increase in effective attraction generally pulls the electron cloud inward, so atomic radius decreases.

Down a group, electrons occupy shells farther from the nucleus and are more strongly shielded by inner electrons, so atomic radius increases.

The first ionization energy is the energy required to remove the least tightly bound electron from a gaseous atom. It generally increases across a period and decreases down a group: smaller atoms with stronger effective nuclear attraction hold their outer electrons more tightly.

Electron affinity describes the energy change when a gaseous atom gains an electron. Its detailed pattern has more exceptions than atomic radius or ionization energy, but atoms near the right side of a period often gain electrons more favorably than those near the left.

These are trends, not exact laws. Subshell structure and electron-electron interactions create local exceptions. Their value is predictive: periodic position gives a first explanation for why some atoms lose electrons easily, others attract additional electron density, and chemical behavior recurs from one period to the next.