Two of the key factors in the chemistry of an atom are its size and the charge on its core (the nucleus plus the electrons of the preceding noble gas). The distance of the outer-shell electrons from the nucleus, and the charge on the core, jointly determine the pull that these electrons feel, and hence determine the electronegativity, or electron-holding power, of the atom.

In Chapters 3 and 6 we introduced several measures of atomic size: metallic and covalent radii, ionic radii, and van der Waals radii. Metallic and covalent radii form a smooth, continuous series across the periodic table because they both describe atoms held together by shared electrons. In metals the electrons are shared between many atoms, and it is natural to think of the metal ions and the bonding metallic electrons separately. In contrast, one seldom regards an molecule as consisting of two ions held together by four bonding electrons, but the distinction is more conventional than real.

Ionic radii are larger than covalent or metallic radii for negative ions, which have picked up more electrons, and smaller for positive ions, which have lost them. Van der Waals radii are important mainly for nonmetals, and are large because they represent packing of atoms without electron sharing.

Metallic and covalent radii for all the representative elements that form bonds are shown at the bottom of the page. Two trends are obvious: shrinkage from left to right across one row because of the pull of the increased nuclear charge, and expansion from top to bottom because of the higher quantum number of the outer shell being filled.

The transition metals add little to the trend and need not be shown, but their absence can be detected by the extra decrease in radii between Groups IIA and IIIA: Ca and Ga, Sr and In, Ba and Tl.