Problem Set
Periodic Table Review
Complete all questions on your own sheet(s) of paper.
1. Draw a periodic table and indicate the trends in terms of atomic radii, ionic radii, ionization energy, and
metallic character. Be sure to show the trends down the table as well as across the table.
Atomic radii decreases 

Ionic radii decreases 

Ionization energy increases 

Metallic character decreases 

2. State the Periodic Law.
If the elements are arranged according to increasing atomic mass (number), a pattern can be seen in which
the properties repeat on a regular basis.
3. How do metal atoms and metal ions compare in terms of atomic (ionic) radius? Explain why using the
factors that affect atomic radius.
Metal atoms are larger than metal ions because ions have lost electrons and their original valence shell!
4. Nearly all elements react with fluorine. Write the formulas of the compounds that form with the
elements of the 3 period when reacted with fluorine.
NaF MgF2 AlF3 SiF4 PF3 SF2 ClF
5. State the meaning of the terms: nuclear chargesum of total positive charge in the nucleus, shielding
effectthe weakening of the attractive forces between the nucleus and valence electrons (inner electrons
shield outer electrons from the positive charge of the nucleus), electron-electron repulsionselectrons in
each shell repel each other (all negatively charged) causing the electrons to spread out from each other as
much as possible.
6. What two statements can be made about the structure of the periodic table as it relates to electron
structure of atoms?
The group number indicates how many valence electrons an atom has while the period number tells us how
many shells to expect for the atom.
7. What is the group name (chemical family) that the following elements are located in:
a) Ba
alkaline earths
b) Fe transition metals
c) Kr noble gas
d) Cl
halogens
e) Pu actinide series
f) Rb
8. Explain the trend of the atomic radius as you move from left to right on the periodic table. Provide
reasons why this trend exists.
The atomic radii decreases since the larger nuclear charge tends to attract valence electrons in more and
more as we move across the period. This causes the atoms to be smaller. The shielding effect of the inner
electrons also becomes less so outer electrons “see” more of the nuclear charge.
9. a) What does ionization energy refer to? The amount of energy required to remove an electron from the
valence shell of a gaseous atom.
b) Write an equation to show the first ionization of a chlorine atomCl (g) + I.E.  Cl+ (g) + e10. Explain which is larger and why it is:
a) Ca and Br atoms
Ca is bigger. Br has greater nuclear charge.
b) Br and I atoms
Iodine atom is larger since it has more shells.
c) As ion and Se ion
As ion has greater electron-electron repulsions
d) Sr ion and Sr atom
Sr atom, Sr ion has lost valence electrons and a valence shell.
11. Explain why magnesium has a 1st ionization energy of 418 kJ/mol and 2nd ionization energy of 789
kJ/mol while its 3rd ionization energy is 4067 kJ/mol. The third electron is now being pull out of an
inner shell that is closer to the nucleus and this requires much more energy than the electrons coming out
of the valence shell.
12. a) Name the element with the largest atomic radius. Francium
b) Name the element with the largest ionization energy (excluding noble gases). Fluorine
13. How do non-metal atoms and non-metal ions compare in terms of atomic (ionic) radius? Explain why
using the factors that affect atomic radius. Nonmetal ions are larger than their respective atoms since
more electrons are added into the valence shells and there are greater electron-electron repulsions
(electrons move out further to accommodate the additional electrons).
14. Predict some of the chemical properties and physical properties for each element: francium, bromine,
and neon.
Francium - soft metal, low density, low melting point, low boiling point, reactive with water
Bromine – brightly colored, high density, low melting point
Neon – unreactive, gas, low melting and boiling points, complete valence shell
15. Which element would be more metallic: Beryllium or Calcium? Why? Calcium – according to metallic
character trends (not testing this).
16. If shown a graph of ionization energy vs. atomic number of the first 54 elements, what would you expect
its shape to be? Sketch it and explain how you arrived at the your answer. Saw tooth - review
http://mmsphyschem.com/ionEng.htm
17. Where would one expect to find the noble gases on this graph of Ionization Energy vs. Atomic Number?
Top of the peaks
18. Oxygen and sulphur are very different elements in that one is a colourless gas and the other is a yellow
crystalline solid. Why are they both in Group VIA? Chemically, they both produce 2- ions and have
same number of electrons in valence shell. Make similar kinds of compounds.
19. What does the term isoelectronic mean? The same electron configuration.
20. Why would it be so much more difficult to remove the a second electron (called the second ionization
energy (IE 2) from an ion than the first (IE 1)? Since there are less electron-electron repulsions the
electron is actually held a little closer to the nucleus therefore requiring more energy to remove it.
21. Explain why the element fluorine is the most reactive element on the periodic table. Highest
electronegativity (and electron affinity).
22. Explain why the ionization energy will decrease as you go down a family? The valence electrons are
easier to remove since they are much further away from the nucleus.
23. The radius of P is much smaller than the ionic radius of P3-. Why is this? There are greater electronelectron repulsions found in the phosphide ion. This causes electrons to spread out a bit more to
accommodate the additional electron.