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.