1) (5 pts) Arrange the following atoms in order of decreasing first ionization energy (IE1): Ar, Cl, Cs, K
See section 9-4 of the text.
smallest IE1 
______
______
______
______
 largest IE1
2) (5 pts) Answer the following questions about magnetic properties, and explain your answer. Use orbital
diagrams to support your answers if appropriate. No hint!
a) Must all atoms with an odd atomic number be paramagnetic?
____________________________________________________________________________________
b) Must all atoms with an even atomic number be diamagnetic?
____________________________________________________________________________________
3) (5 pts) Match each of the lettered items on the left with one appropriate numbered item on the right. All the
numbered items should be used, and some more than once. (Hint: Identify the elements from their atomic
numbers. You can answer this without actual numbers by using electron configurations and periodic
trends.) End of chapter problem 45 from chapter 9.
a) Z = 8
Number = ______
(1) diamagnetic
b) Z = 20
Number = ______
(2) two unpaired p electrons
c) Z = 32
Number = ______
(3) first ionization energy lower than that of Ca
but greater than that of Cs
d) Z = 38
Number = ______
(4) more negative electron affinity than elements on
on either side of it in the same period
e) Z = 48
Number = ______
f) Z = 53
Number = ______
4) (5 pts) For each of the following pairs, indicate the atom that has the larger size:
See section 9-3 of the text.
(a) Te or Br
(d) N or O
__________
__________
(b) K or Ca
(e) O or P
__________
__________
(c) Ca or Cs __________
(f) Al or Au __________
5) (5 pts) Answer the following questions. See sections 8-10, 8-11, and 9-3 of the text.
a) Why the 4s subshell is lower in energy than the 3d subshell at the beginning of the 4th period?
____________________________________________________________________________________
b) Use the concept of shielding to briefly explain the similar properties of the alkaline earth metals.
____________________________________________________________________________________
6) (5 points) Match each molecular geometry (VSEPR shape) from the left with its name on the right.
Not all the names are used. T-shaped and See saw are not used.
(A)
______ Bent
(B)
______ Linear
(C)
______ Octahedral
(D)
______ Square planar
(E)
______ Square pyramidal
(F)
______ Tetrahedral
______ Trigonal bipyramidal
(G)
______ Trigonal planar
(H)
______ Trigonal pyramidal
______ T-shaped
(I)
______ Seesaw
7) (5 points) Draw the Lewis structure for the nitronium cation, ONO+. Draw all possible resonance
structures. Determine the formal charges of the atoms in each structure, and circle the best Lewis structure.
VSEPR shape is linear.
8) (5 points) Calculate the enthalpy change (∆H) for the following reaction using bond dissociation energies.
The multiple bonds are shown; all other bonds are single. Add up all reactant bond energies and subtract all
product bond energies.
2 NH3(g) + 3 O=O(g) + 2 CH4(g) →
2 H–C≡N(g) + 6 H2O(g)
9) (5 points) First ionization energies usually increase as we go from left to right across a period, as shown
below. Explain why the ionization energy decreases from nitrogen to oxygen. The ionization energies are
shown in kJ/mol: See section 9-4 of the text.
B = 800
C = 1086
N = 1402
O = 1314
F = 1681
Ne = 2088
10) (5 points) Arrange the following bonds in order of increasing polarity: See section 10-3 of the text.
Al–Br
Al–Cl
C–Br
Al–F
Cl–Cl
11) (5 points) Answer these questions for the F2+ cation.
Fill the valence electrons into the molecular orbitals in the middle of the diagram. See annotated 2015 notes.
a. Fill in the following molecular orbital diagram. You do not have to fill in the atomic orbitals (such as 2s
and 2px) at the left and right of the figure; only fill in the molecular orbitals in the center.
b. The Bond Order = __________
c. The F2+ cation is a (stable OR unstable) species ____________________
d. The ion is (paramagnetic OR diamagnetic) ____________________
e. The bond length in the ion is (shorter OR longer) than in molecular F2 ___________________
12) (5 points) Describe the bonding schemes for the σ and π bonds shown in the following diagram. You can
draw an arrow to a bond, and then label the arrow with the bonding scheme (bond type and orbital overlap).
If there are several bonds with identical schemes, you may connect multiple arrows to one bonding scheme.
The C and O atoms use hybridized orbitals for σ bonding.
H
H
C
H
C
H
C
..
.O.
H
H
13) (5 points) Draw the Lewis structure for TeF4, and predict the VSEPR geometry. If the molecule is polar,
draw a molecular dipole arrow over the Lewis structure. It’s like SF4 from lab (S and Te in same family!).
14) (5 points) Ammonia (NH3) has a greater dipole moment than nitrogen trifluoride, NF3. This means that
ammonia is the more polar molecule. Explain. (Electronegativities: H=2.1, N=3.0, F=4.0)
15) (10 pts) Copy the appropriate molecular orbital energy level scheme from the reference page and fill it in
for the ion O22+. Predict the bond order and magnetism of this species.
Diagram like that in question 11 would be given.
Bond Order = __________
Magnetism = ____________________
16) (5 pts) Fill in the following blanks No hints! :-(
a) In an sp3d hybridization scheme, lone pairs are always placed in the __________________ position.
b) Molecular orbital theory replaces multiple resonance structures by using _______________ orbitals.
c) Atoms and ions that have the same number of electrons are said to be ____________________.
d) Electrons fill the 4s orbital before the 3d because the 4s ____________________ inside the 3d.
e) The _____________________ states that two electrons in the same orbital must have opposite spins.
f) An atom with all its electrons in the lowest energy orbitals is said to be in the _________________.
g) Similar properties recur periodically when elements are arranged according to increasing ________.
h) When Main Group elements form ions, they achieve the electron configuration of ______________.
i) Ionization energies decrease as atomic radii ____________________.
17) (5 pts) Explain the following observations: See section 9-4 of the text.
a) Generally, ionization energies is to increase from left to right across a period, but the first ionization
energy (I1) decreases from 1012 kJ/mol for P to 999.6 kJ/mol for S.
b) The ionization energy of Al increases dramatically from I3 to I4 (see table).
18) (5 pts) Place the following ions in order of increasing radius, from smallest to largest.
See section 9-3 of the text. (Isoelectronic!)
Ca2+
Cl−
smallest 
Ga3+
______
P3−
K+
______
S2−
______
______
______
______
 largest
19) (5 pts) Use bond energies to estimate ∆Hrxn for the following reaction See section 10-9 of the text.
H
H
+
C C
2
H
H
H
2 H Cl
+
O O
H
H
+
C C
2
2
O
H
H
Cl
20) (5 pts) Draw the Lewis structure of the carbonate anion, CO32−, including resonance structures.
VSEPR shape is trigonal planar.
21) (10 pts) Draw the Lewis structure of SbF52−. Label each atom with its Formal Charge. Predict the shape of
this species and the hybridization of the central atom. You do have to show a promotion and hybridization
scheme for the central atom.
VSEPR shape is square pyramidal. (This ion has the same number of valance electrons as the BrF5
molecule from lab!)
22) (5 pts) The following figure is the Lewis structure of the amino acid alanine. Write the bonding schemes
showing bond type and orbital overlap, such as σ:C(sp3)−H(1s), for the C and O atoms labeled with an
asterisk (*) in this molecule. Count BONDED ATOMS and LONE PAIRS on the two starred atoms.
H
H
H
..O..
N
..
C
C*
H
C
H
..
O*
..
H
H
33. (10 pts) An incomplete Lewis structure of piperine, the active compound in white and black pepper, is
shown below. The structures containing the letters a and f are rings made of six C atoms. It might help you
to write in carbon atom labels at the vertices (points) of each hexagon.
All first and second period elements, so no octet expansion.
(a) Complete the Lewis structure by showing all lone pairs of electrons (write them in above). Use dots!
(b) How many carbon atoms are sp, sp2, sp3, sp3d, and sp3d2 hybridized in piperine?
sp: ________
sp2: ________
sp3: ________
sp3d: ________
sp3d2: ________
(c) What is the hybridization of the nitrogen atom in piperine? ________
(d) Give values for the bond angles marked a through f above. You can neatly write the values of the bond
angles next to the letters above.
23) (5 pts) Draw the molecular orbital scheme for the diatomic species CO2−. Fill-in the diagram with
electrons, calculate the bond order, and predict whether CO2− is diamagnetic or paramagnetic. (Hint: The
element with the higher atomic number has priority.)
Diagram like that in question 11 would be given. Then fill the valence electrons into the molecular orbitals
in the middle of the diagram. See annotated 2015 notes. There are 12 valence electrons to fill in:
C = 4, O = 6, and 2 for the 2− charge.
Bond Order: __________________________________
Diamagnetic or Paramagnetic (circle one)
24) (5 points) Use the molecular orbital energy-level diagram for the neutral hydroxyl radical (HO) below to
answer the questions that follow. I love this question! But it is too much for our test! :-(
(a) The number of bonding electrons
_____________
(b) The number of antibonding electrons _____________
(c) The number of non-bonding electrons _____________
(d) The Bond Order =
_____________
(e) The magnetism (di or para) =
_____________
27. Draw the molecular orbital energy-level diagram for the F2 molecule in the space below. Be sure to label
the molecular orbitals, and to place electrons in the orbitals. Note that the energy of σ2p < π2p for F2.
Diagram like that in question 11 would be given.
Is F2 diamagnetic or paramagnetic?
F atom
F2 molecule
2p ↑↓ ↑↓ ↑_
2s
↑↓
F atom
↑↓ ↑↓ ↑_ 2p
↑↓
Be careful counting valence electrons! F2 has 14, while F2+ (questions 11) has 13.
2s