1. Atomic radius
&
Ionization Energy

2. Atomic radii increase going down the
groups.
Electrons are added to successive energy
levels (shielding effect). Each period
represents another energy level. Atomic
radii increase as electrons are added to
successive energy levels.

3. 2. Atomic radii decrease across periods. As
one goes across a period each element has
an additional proton in the nucleus that
pulls the electrons closer. (avg. energy ~)

4. Na or (Rb)
(Sr) or Cd
(Ba) or Hg
(Hf) or Pt
(Cs) or Li
A sodium atom, Na, would be larger than a
sodium ion, Na1+ , which has lost an electron
and only has 2 energy levels (vs. 3).
A chloride ion, Cl1- , is larger because the extra
electron weakens the pull of the nucleus and
results in shell expansion.

5. 4. Ionization energy decreases going down
within a group. Energy is applied against
the pull of the nucleus. Electrons are
further from the nucleus and the shielding
effect increases as one goes down the
groups, so electrons are easier to remove.

6. 5. Across periods ionization energy increases
because more protons are added to the
nucleus (w/out significant change in the
average energy of the electrons).

7. 6. Metals have lower ionization energy going
down the group. They react more readily
because it takes less energy to remove
electrons from them.

8. 7. The inert gasses have the highest ionization
energies. This is because they are smaller &
have filled s & p sublevels, which is a stable
electron arrangement. Therefore they don’t
form ions nor tend to react.

9. Hydrogen has a very high ionization
energy because its one electron is close
to the nucleus with no other levels to
shield it.

10. Similarities between the graphs include a
similar shape of peaks and valleys.
Differences include that the peaks of one
are the valleys of the other.