AP Exam Questions
Explain each of the following observations using principles
of atomic structure and/or bonding
a) Potassium has a lower
first-ionization energy than lithium
b) The ionic radius of N3–
is larger than that of O2–.
c) A calcium atom is larger
than a zinc atom.
d) Boron has a lower
first-ionization energy than beryllium.
Explain each of the following observations using principles
of atomic structure and/or bonding
a) Potassium has a lower first-ionization energy
than lithium
K (1s2)(2s2,2p6)(3s2,3p6)
(3d0) (4s1) Z*
= 19 – (10 · 1.00 + 8 · 0.85) = 2.2
Li (1s2)(2s1) Z*
= 3 - (2 · 0.35) = 2.3
Explanation not just in the difference in effective
nuclear charge. As one can there is not much. The primary
reason is related to the energy of a single s
electron in the n = 4 level and the n = 2
level. The electron in the n = 4 level is higher
in energy compared to an electron in n = 2 level
and therefore is easier to remove.
b) The ionic radius of N3–
is larger than that of O2–.
N3– (1s2)(2s2,2p6)
Z* = 7 – (2 · 0.85 + 7 · 0.35) = 2.85
O2– (1s2)(2s2,2p6)
Z* = 8 – (2 · 0.85 + 7 · 0.35) = 3.85
The electrons in N3– experience a
smaller effective nuclear charge and are less bound to
the nucleus compared to the electrons in oxygen. The
result is the electrons in nitrogen will take up a occupy
volume.
c) A calcium atom is larger than a zinc atom.
Ca (1s2)(2s2,2p6)(3s2,3p6)
(3d0) (4s2) Z*
= 20 – (10 · 1.00 + 8 · 0.85 + 1 · 0.35) = 2.85
Zn (1s2)(2s2,2p6)(3s2,3p6)
(3d10) (4s2) Z*
= 30 – (10 · 1.00 + 18 · 0.85 + 1 · 0.35) = 4.35
The electrons in the 4s level experience a
greater effective nuclear charge in Zn than in Ca. The
electrons in Zn are therefore attracted to the nucleus to
a larger extent than the electrons in calcium. So the
atomic radius is smaller. In general the Z*
increases more slowly across a d block compared to
a p block. The reason is the electrons being added
in the dblock are inside the ns electrons,
while for the p block the electrons are not inside
the ns electrons.
d) Boron has a lower first-ionization energy than
beryllium.
Be (1s2)(2s2)
Z* = 4 – (2 · 0.85 + 1 · 0.35) = 1.95
B (1s2)(2s2,2p1)
Z* = 8 – (2 · 0.85 + 2 · 0.35) = 2.6
We would expect from Z* arguments that it
should require more energy to remove the first electron
in boron compared to beryllium. However, the electron
that is removed from boron in in the 2p level
which is higher in energy compared to the 2s
level, therefore it is easier to remove. Recall that Z*
for the Group IA metals (after Li) are effectively the
same value, yet we all know the first ionization energy
for the valence electron in Cs is easier to remove than
the valence electron in sodium because the electron is
located in a higher energy level.
Question #2
The diagram shows the first ionization energies for the
elements Li to Ne. Briefly, explain each of the following in
terms of atomic structure.
a) In general, there is an increase in the first
ionization energy from Li to Ne.
b) The first ionization energy of B is lower than that of
Be.
c) The first ionization energy of O is lower than that of
N.
d) Predict how the first ionization energy of Na compares
to those of Li and of Ne. Explain.
Question #3
Account for each of the following in terms of principles of
atomic structure.
a) The second ionization energy of sodium is about three
times greater than the second ionization energy of magnesium.
16. Which of the following are reasonable values for the
first four ionization energies for Al?
|
1st
|
2nd
|
3rd
|
4th
|
A)
|
496 kJ
|
4562 kJ
|
6912 kJ
|
9543 kJ
|
B)
|
578 kJ
|
1817 kJ
|
2744 kJ
|
11,577 kJ
|
C)
|
738 kJ
|
1451 kJ
|
7733 kJ
|
10,540 kJ
|
D)
|
657 kJ
|
1269 kJ
|
2136 kJ
|
2752 kJ
|
9. Which of the following are reasonable values for the first
four ionization energies for Mg?
|
1st
|
2nd
|
3rd
|
4th
|
A)
|
496 kJ
|
4562 kJ
|
6912 kJ
|
9543 kJ
|
B)
|
578 kJ
|
1817 kJ
|
2744 kJ
|
11,577 kJ
|
C)
|
738 kJ
|
1451 kJ
|
7733 kJ
|
10,540 kJ
|
D)
|
657 kJ
|
1269 kJ
|
2136 kJ
|
2752 kJ
|