electrons and the other electrons, and the radius is generally reduced. The inert
gas at the end of the period has a slight increase in radius because of the electron
repulsion in the filled outer principal energy level. For example, lithium’s atomic
radius in Figure 10 is 0.152 nm at the one end of period 2 whereas fluorine has a
radius of only 0.064 nm at the far end of the period. This trend can be seen in
Figure 10 across every period.
TIP
This is the explanation of these trends in periods.TIP
... and in groups.
Atomic Radii in Groups
For a group of elements, the atoms of each successive member have another outer
principal energy level in the electron configuration, and the electrons there are
held less tightly by the nucleus. This is so because of their increased distance
from the nuclear positive charge and the shielding of this positive charge by all
the core electrons. Therefore, the atomic radius increases down a group. For
example, oxygen’s atomic radius in Figure 10 is 0.066 nm at the top of group 16,
whereas polonium has a radius of 0.167 nm at the bottom of the same group. This
trend can be seen in Figure 10 down every group.
Ionic Radius Compared with Atomic Radius
Metals tend to lose electrons in forming positive ions. With this loss of negative
charge, the positive nuclear charge pulls in the remaining electrons closer and
thus reduces the ionic radius below that of the atomic radius.
TIP
Know this relationship between the atomic radius and the ionic radius.Nonmetals tend to gain electrons in forming negative ions. With this added
negative charge, which increases the inner electron repulsion, the ionic radius is
increased beyond the atomic radius. See Figure 10 for relative atomic and ionic
radii values.