Atomic radius & Ionization Energy

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.

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 ~)

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.

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.

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).

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

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.

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

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.