Study Guide: Periodic Table and Periodic Trends (Honors)
1. Differentiate between Mendeleev’s periodic table and Moseley’s periodic table.
Mendeleev put the elements in order by atomic mass, but he noticed that in several
places, the atomic masses did not always increase sequentially. Moseley discovered
that ordering by atomic number was the correct way to describe the order.
2. State the periodic law.
The properties of the elements are periodic functions of their atomic numbers.
3. Vertical columns on the periodic table are called groups or families and horizontal rows are
called periods.
4. Elements with similar chemical properties are placed in the same group.
5. Complete the table below.
Electron Config.
[Xe] 6s24f145d106p6
[Xe] 6s1
[Ar] 4s13d5
[He] 2s2
[Kr] 5s1 4d10
[He] 2s2 2p5
Block
Period
p
s
d
s
d
p
6
6
4
2
5
2
Group/Type
Element
18 / Noble Gas
1 / Alkali Metal
6 / Trans. Metal
2 / Alk. Earth Metal
11 / Trans. Metal
17 / Halogen
Radon
Cesium
Chromium
Beryllium
Silver
Fluorine
6. Determine whether the following statements are describing alkali metals (A), alkaline earth
metals (AE), transition metals (TM), halogens (H), noble gases (NG):
a) most reactive non-metals H
b) s2p6 valence configuration NG
c) second most reactive metals AE
d) s1 valence configuration A
e) exceptions found to e- configurations TM
f) soft metals A
g) d orbitals being filled TM
7. Would you expect Li and S to have different chemical properties? Why? Yes; Li is an alkali
metal with 1 valence e-, and S is a nonmetal with 6 valence e-.
8. Explain why a magnesium atom is smaller than: a sodium atom; a calcium atom. Mg is
smaller than Na because the greater nuclear charge in Mg pulls the e- cloud in
tighter (periodic trend). Mg is smaller than Ca because it has fewer energy
levels (group trend).
9. Predict the size of the astatine atom compared to the sulfur atom. At is larger than S,
because At is 3 periods down from S, but only 1 group over. The trends compete,
but the the sizeable increase in energy levels overcomes the slight difference in
nuclear charge.
10. Would you expect at Cl- ion to be larger or smaller than an Mg2+ ion. Explain. A Cl- ion is
larger than a Mg2+ ion, because the chloride ion has 3 energy levels, matching
the e- configuration of argon, while the magnesium ion has only 2 (matching the
e- configuration of neon).
11. Explain why the sulfide ion (S2-) is larger than the chloride ion (Cl-). Both of these ions
match the e- configuration of argon, but S2- is larger because the nuclear charge
is less, resulting in less “pulling in” of the e- cloud.
12. Compare the ionization energy of sodium to that of potassium. Ionization energy
decreases down a group, so the IE of sodium is greater than that of potassium.
13. Explain the difference in ionization energy between lithium and beryllium. Beryllium has a
greater IE than lithium because Be’s greater nuclear charge holds the electrons
tighter, resulting in increased IE required to remove an e-.
14. Will the electronegativity of barium be larger or smaller than that of strontium? smaller
15. Will a magnesium have a tendency to form a cation or an anion? cation (loses 2 e-)
16. Arrange oxygen, fluorine, and sulfur in order of increasing electronegativity. S, O, F
17. The most metallic element on the periodic table is francium; the least metallic element is
fluorine.
18. As you go across a period, the electron affinity decreases (becomes a larger negative
number) . Explain why. Nuclear charge increases across a period, leading to a
stronger attraction for electrons.
19. When an electron is added to a neutral atom, energy is released (exo, negative). This is
electron affinity. When an electron is removed from a neutral atom, energy is absorbed
(endo, positive). This is ionization energy.
 A+ + e-, energy is a reactant and ionization
20. In an ionization energy equation, A 
energy is an endothermic process.
 A-, energy is a product and electron affinity
21. In an electron affinity equation, A + e- 
is an exothermic process.
22. Write an equation for the formation of F- from F. (energy = - 327.8 kJ/mol) Does this
represent ionization energy or electron affinity?
F (g) + e -  F – (g) + 327.8 kJ electron affinity
23. For Group 3, list the element symbol for
a) most metallic Ac
b) least metallic Sc
c) largest atomic radius Ac
d) smallest atomic radius Sc
e) largest ionic radius Ac
f) largest ionization energy Sc
g) largest electronegativity Sc
h) smallest electron affinity Sc
24. For Period 3, list the element symbol for
a) most metallic Na
b) least metallic Cl
c) largest atomic radius Na
d) smallest atomic radius Ar
e) largest ionic radius P
f) largest IE Ar
g) largest electronegativity Cl
h) smallest electron affinity Cl