Name________KEY___________________________
Unit 4 Study Guide
1.
Write the complete electron configuration, orbital diagram, and report the number of lone pairs on a separate
sheet of paper for each of the following:
a. Fluorine
d. Phosphorus
b. Germanium
e. Potassium
c. Lead
2. Write the abbreviated electron configuration for each of the following:
a. Tin
c. berkelium
b. Iridium
d. Thallium
3. In your own words explain each of the following: Hund’s rule, Aufbau Principle, and Pauli Exclusion Principle.
Illustrate the rule.
4. What trend in atomic radius do you see as you go down a group on the periodic table? What causes this trend?
Atomic Radius increases as you go down a group on the periodic table. This occurs because you are adding
energy levels and the valence electrons have greater shielding from the pull of the nucleus.
5. What trend in atomic radius do you see as you go across a period on the periodic table? What causes this
trend?
Atomic Radius decreases as you go across a period on the periodic table. This occurs because the electrons are
being added to the same energy level (so the shielding remains constant) but you are adding protons to the
nucleus and thus there is greater nuclear charge to “pull” on the electrons. This greater pull “pulls” the electrons
in closer.
6. Circle the atom in each pair that has the largest atomic radius. (YOU MUST BE ABLE TO EXPLAIN WHY!)
a. Al or B (more E-levels thus more shielding)
d. O or F (less nuclear charge)
b. Na or Al (less nuclear charge)
e. Br or Cl (more E-levels thus more shielding)
c. S or O (more E-levels thus more shielding)
f. Mg or Ca (more E-levels thus more
shielding)
Name________KEY___________________________
7. Define ionization energy.
Energy required to remove an electron. 1st I.E is the energy required to move the first electron, 2nd IE is the
energy required to remove the second electron…and so on.
8. What trend in ionization energy do you see as you go down a group on the periodic table? What causes this
trend?
Ionization Energy decreases as you go down a group on the periodic table. This occurs because the valence
electrons being removed are further away and shielded from the nucleus so the nuclear charge cannot “hold on
to them” as tightly so it is easier to take them away.
9. What trend in ionization energy do you see as you go across a period on the periodic table? What causes this
trend?
Ionization Energy increases as you go across a period on the perioidic table. This occurs because the valence
electrons are at the same distance from the nucleus “holding on to them” and have equal shielding. However,
there are more protons in the nucleus resulting in a greater nuclear charge to “hold on to the electrons”
stronger. This stronger “hold” make is harder to remove the electron.
10. Circle the atom in each pair that has the greater ionization energy. (YOU MUST BE ABLE TO EXPLAIN WHY)
a. Li or Be (greater nuclear charge=stronger hold)
d. P or Ar (greater nuclear charge=stronger hold)
b. Ca or Ba (closer to the nucleus/less shielding)
e. Cl or Si (greater nuclear charge=stronger hold)
c. Na or K (closer to the nucleus/less shielding)
f. Li or K (closer to the nucleus/less shielding)
11. Define electronegativity.
Ability to attract electrons
12. What trend in electronegativity do you see as you go down a group on the periodic table? What causes this
trend?
Electronegativity decreases as you go down a group on the periodic table. This occurs because the nucleus is
shielded from the valence shell.
13. What trend in electronegativity do you see as you go across a period on the periodic table? What causes this
trend?
Electronegativity increases as you go across a period on the periodic table. This occurs because there is equal
shielding and greater nuclear charge to pull and attract electrons. Also, as you go across the periodic table the
number of valence electrons in the outer energy level increases and is getting closer to having a full octet.
14. Circle the atom in each pair that has the greater electronegativity. (YOU MUST BE ABLE TO EXPLAIN WHY)
a. Ca or Ga (greater nuclear charge/more valence e-)
e. Ba or Sr (less nuclear shielding)
b. Br or As Ga (greater nuclear charge/more valence e )
f. Cl or S (greater nuclear charge/more valence e-)
c. Li or O (greater nuclear charge/more valence e-)
g. O or S (less nuclear shielding)
d.
15. For each of the following sets of atoms, rank the atoms from smallest to largest atomic radius.
a. Li, C, F F,C,Li
d. C, N, Al N, C, Al
b. Li, Na, K Li, Na, K
e. Al+3, O-2, Ne Al+3, Ne, O-2
c. Ge, P, O O, P, Ge
Name________KEY___________________________
16. For each of the following sets of atoms, rank the atoms from lowest to highest ionization energy.
a. Mg, Si, S Mg, Si, S
d. Ba, Cu, Ne Ba, Cu, Ne
b. Mg, Ca, Ba Ba, Ca, Mg
e. Si, P, H Si, P, H
c. F, Cl, Br Br, Cl,
17. For each of the following sets of atoms, rank the atoms from lowest to highest electronegativity.
a. Li, C, N Li, C, N
d. K, Mg, P K, Mg, P
b. C, O, Ne Ne, C, O
e. S, F, He He, S, F
c. Si, P, O Si, P, O