Name:
3 Dec. 2007
Chemical Principles I
Third Midterm
Constants:
c = 3.00 x 108 m/s ; h = 6.63 x 10-34 J-s
PART A1 Circle the letter for the correct answer. (1/2 pt each)
1. What two measurable aspects of electromagnetic radiation always equal the speed of
light when they are multiplied together?
a. Intensity and frequency.
b. Wavelength and amplitude.
c. Wavelength and frequency.
d. Amplitude and frequency.
2. What phenomenon is defined by electrons being ejected from a metal surface when it
is bombarded by radiation above a certain, threshold frequency?
a. The wave - particle duality.
b. The photon electron effect.
c. The Stefan - Boltzmann Law.
d. The photoelectric effect.
3.
a.
b.
c.
d.
What is a wave function?
A mathematical function that varies with time.
A mathematical function that varies with position.
A semi-quantitative, mathematical function that varies with position.
A semi-quantitative, mathematical function that varies with amplitude.
4.
a.
b.
c.
d.
What is defined by the principle quantum number, n?
The energy levels of a hydrogen atom.
The energy bursts of a hydrogen atom.
The direction of the orbitals of an atom.
The orientation of the orbitals of an atom.
5.
a.
b.
c.
d.
What does the orbital angular momentum quantum number specify?
The size of an orbital.
The shape of an orbital.
The orientation of an orbital.
The number of electrons in an orbital.
6.
a.
b.
c.
d.
How is electron density distributed at a nodal plane?
There is a large amount of electron density there.
There is only a small amount of electron density there.
There is no electron density there.
There is a concentration of electron density there.
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7.
a.
b.
c.
d.
What does Hund’s Rule state?
That no two electrons in an atom can have the same four quantum numbers.
That no two atoms can have the same Z.
When electrons half fill the orbitals in a shell, the electrons must have the same spin.
If more than one orbital in a subshell is available, electrons will half fill each orbital
prior to pairing in an orbital.
8. How is the exact radius of an atom measured?
a. The exact radius cannot be measured.
b. Through combining the atom with several other atoms and averaging the bond
distances.
c. Through averaging the atomic radii of the two atoms adjacent on the periodic table to
the atom in question.
d. Through the use of a van der Waal’s micrometer.
9.
a.
b.
c.
d.
Why do ionization energies typically increase down a group?
Because there are more electron shells as you move down a group.
Because the valence electrons are more polarizable.
Because the electron shells are farther from the nucleus.
They do not; rather they decrease down a group.
10. A diagonal relationship involves similarity in properties between what elements?
a. Between diagonal neighbors of any groups of the periodic table.
b. Between diagonal neighbors of the main groups of the periodic table.
c. Between diagonal neighbors only in the transition metals of the periodic table.
d. Between diagonal neighbors solely in the non-metal groups of the periodic table.
11. What chemical properties do all s-block metals share?
a. They all form acidic oxides.
b. They are all non-reactive elements.
c. They are all reactive elements.
d. They all do not form oxides.
12. Why do transition metals of the same period have very different properties?
a. Because of the differences in the number of outer shell d - electrons.
b. Because of the differences in atomic radii.
c. Because of the differences in inner shell d - electrons.
d. They do not; rather, their properties are usually similar.
14. Which of the following is a characteristic of the transition metals?
a. The ability to have negative oxidation numbers, because of the gain of electrons.
b. The ability to exist in more than one oxidation state.
c. The ability of each transition metal to exist as a solid or liquid at room temperature.
d. There are no special characteristics of the transition metals.
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15. Which is the best description of a p orbital?
a. A region where the electron will be found in two lobes.
b. A region where the + and – charge of the electron are shown.
c. A region of high probability of finding an electron.
d. A mathematical function.
PART A2 Circle the letter for the correct answer. (1/2 pt each)
1. The ionic model of binding is best utilized to describe what type of atoms joining
together?
a. Two metallic elements.
b. Two non-metallic elements.
c. A metallic and a non-metallic element, especially a metallic p-block element.
d. A non-metallic and a metallic element, especially a metallic s-block element.
2. Why are ionic structures typically brittle, and why do they have high melting points?
a. Because they are stacked in irregular, non-crystalline structures displaying large,
coulombic interactions.
b. Because they are stacked in regular, crystalline structures displaying large, coulombic
interactions.
c. Because they are stacked in regular, crystalline structures displaying small, ionic
interactions
d. None of the above.
3. Which elements, when forming molecules, do not reach an octet in their Lewis
structures?
a. He, Li, and F.
b. H, Li, and Mg.
c. He, Li, and Be.
d. H, Li, and Be.
4.
a.
b.
c.
d.
What property do all covalent bonds have in common?
The sharing of pairs of electrons.
The loss of one or more outer shell electrons.
The gain of one or more outer shell electrons.
The loss or gain of one or more outer shell electrons.
5.
a.
b.
c.
d.
What electron arrangement must a compound have to be considered a radical?
One or more unshared electrons pairs within the compound.
One or more metallic elements with an unpaired electron in the molecule.
At least one unpaired electron somewhere in the molecule.
At least one atom in a molecule with more than eight electrons.
6. How are most actual bonds defined in terms of the ionic and covalent models?
a. As mostly ionic.
b. As mostly covalent.
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c. As somewhere between purely covalent and purely ionic bonds.
d. As none of the above.
7. How is ionic character in a bond determined using electronegativities?
a. The greater the electronegativity difference between the atoms in a bond, the more
ionic character the bond exhibits.
b. The greater the electronegativity difference between the atoms in a bond, the less
ionic character the bond exhibits.
c. The less the electronegativity difference between the atoms in a bond, the more ionic
character the bond exhibits.
d. There is no direct correlation between electronegativity and ionic bond character.
8.
a.
b.
c.
d.
For a cation to have a high polarizing power, what requirements must be met?
It must be either large or highly charged.
It must be either small or of low charge.
It must be both large and highly charged.
It must be both small and highly charged.
9.
a.
b.
c.
d.
Dissociation energy is a measure of what?
The energy added to a bond to force its electrons to separate.
The energy required to separate bonded atoms.
The energy required to cause new bonds to form in a molecule.
None of the above.
10. Why are multiple bonds longer than single bonds?
a. Because of additional bonding electrons.
b. Because more electrons in the multiple bond require more physical space.
c. Because more electrons in the bond push the bonding nuclei farther apart.
d. They are not longer; rather, they are shorter because of additional bonding electrons.
PART A3 Circle the letter for the correct answer. (1/2 pt each)
1. What does the VSEPR model dictate about areas in a molecule with high electron
density?
a. That they will repel each other.
b. That they will attract each other.
c. That they will repel each other if they are electrons in a bond.
d. That they will attract each other only if they are electrons in a lone pair.
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2. How do lone pairs of electron density influence the shape of a molecule?
a. They occupy minimal space and repel other electrons, thus influencing only slightly
the overall molecular shape.
b. They occupy space and repel other electrons, thus influencing overall molecular
shape.
c. They occupy space and attract other electrons, thus influencing overall molecular
shape
d. They do not actually influence the shape.
3.
a.
b.
c.
d.
What type of molecule has a non-zero dipole moment?
An ionic compound.
A covalent molecule.
A polar molecule.
A molecule with polar bonds.
4.
a.
b.
c.
d.
Where is pi-bond electron density located?
Along the bond axis of two atoms.
Above and below the bond axis between two atoms.
Above and below the bond axis as well as on the bond axis between two atoms.
On either end of the bond axis between two atoms.
5. What does a hybridized orbital denote?
a. A blending of two atomic orbitals to fit the experimentally determined shape of a
molecule.
b. A blending of the s, p, or d character of two different atoms.
c. Any mixing of s, p, or d orbital character.
d. A mixing of s, p, and/or d orbital character to fit the experimentally determined shape
of a molecule.
6.
a.
b.
c.
d.
What is required for sp3d hybridization to be possible?
There must be d orbitals present on one atom within the molecule.
There must be d orbitals present on the central atom of the molecule.
There must be at least four bonds to the central atom of the molecule.
There must be four or more bonds to the central atom of the molecule.
7.
a.
b.
c.
d.
How are pi-bonds best described by the valence bond theory?
As unhybridized bonds.
As hybridized bonds.
As covalent bonds.
As polar covalent bonds.
8.
a.
b.
c.
d.
What character does the pi-bonding in a double bond impart to a molecule?
It makes the molecule resistant to chemical reactivity.
It makes the double bond in the molecule resistant to lengthening or shortening.
It makes the double bond in the molecule resistant to twisting.
It makes the molecule resistant to twisting.
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9.
a.
b.
c.
d.
For a compound to be paramagnetic, what must it possess?
A lone pair of electrons.
A double bond.
A double or triple bond.
Unpaired electrons.
10. How are bonding molecular orbitals formed?
a. By the constructive interference of two molecular orbitals.
b. By the constructive interference of two atomic orbitals.
c. By the destructive interference of two atomic orbitals.
d. By the partial overlap of two or more atomic orbitals.
11. How many electrons can be accommodated in an electron orbital, according to the
Pauli Exclusion Principle?
a. 1
b. 2
c. 3
d. 4
12. What feature do all heteronuclear diatomic molecules share?
a. A single or double bond.
b. A non-polar bond.
c. An ionic bond.
d. A polar bond.
13. How does the delocalization of electrons in a molecule affect bonding in it, according
to molecular orbital theory?
a. It spreads the electron density over specific areas where bonding will occur.
b. It allows increased electron density at areas of lowered electron density.
c. Delocalization spreads bonding effects over the entire molecule.
d. Delocalization destabilizes the entire molecule.
PART B Write answers on exam. Show calculations.
1) What is an orbital? (5 pts)
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2) What is the difference in energy for a hydrogen atom with its electron in a 2s
orbital and in a 2p orbital? (2 pts)
BONUS: What causes the difference in energy between electrons in a 2s orbital and in a
2p orbital in atoms?
3) The physical property that demonstrates the wave-like behavior of light is called
_________________________. (3 pts)
4) Calculate the energy of a photon that is released when the electron in hydrogen
moves from n = 6 to n = 2. (RB = 2.18 x 10-18 J) (3 pts)
5) Fill in the numerical value of quantum numbers n and l corresponding to the
following orbital designations: (2 pts)
Orbital
n
l
3p
2s
4f
5d
6) Fill in the maximum number of electrons that occupy each of the subshells in the
following table: (2 pts)
3d
4s
2p
5d
Max number e-
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7) Write the electron configurations for the following atoms or ions, using the
appropriate noble-gas core abbreviations: (4 pts)
Electron configuration
Cs
Cu
Ni2+
Se2-
8) Fill in the quantum number(s) that govern the following properties of orbitals: (4
pts)
Quantum number (or numbers)
shape
energy
spin
spatial orientation
9) Arrange the following atoms or ions in order of increasing size: (5 pts)
Lowest, middle, highest
Ca, Mg, Be
Ga, Br, Ge
Se2-, Te2-, Se
Co2+, Fe2+, Fe3+
Ca, Ti4+, Sc3+
10) For each of the following pairs, circle the element with the higher metallic
character. (4 pts)
Li or Be
Li or Na
Sn or P
Al or B
11) Predict which of the following oxides are ionic or molecular by putting the letter I
under the ionic and M under the molecular. (3 pts)
SO2, MgO, Li2O, P2O5, N2O, XeO3
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12) Write a balanced equation for the reaction that occurs in the following cases: (2
pts each)
a) Potassium metal burns in an atmosphere of chlorine gas:
b) Strontium oxide is added to water:
c) Sodium metal is added to molten sulfur:
d) Iron(II) oxide reacts with phosphoric acid:
e) Sulfur trioxide reacts with water:
f) Carbon dioxide reacts with aqueous sodium hydroxide.
13) Write an equation (defining your symbols) that accounts for the large increase in
lattice energy in the series of isoelectronic substances KF < CaO < ScN (2 pts)
14) Circle the most electronegative atom in each of the following sets: (4 pts)
(a) P, S, As, Se;
(b) Be, B, C, Si;
(c) Zn, Ga, Ge, As; (d) Na, Mg, K, Ca
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15) Sketch the 1s, 2s, and 3s orbitals on the following graphs (3pts)
BONUS (1 pt) Show the value of each possible node.
1
1
1



0
0
0
5
x
1s = e -| x |
0
0
0
5
5
x
x
2s = (5 – x) e -| x/2 |
3s = (x2-6x +5) e -| x/3 |
16) A 1s orbital placed 2 units to the left on the x axis has the form, e -| x +2 | and
another placed 2 units to the right has the form, e -| x - 2|. Sketch the  MO on the
graph on the left and the * MO on the right, and write the mathematical function
for each (4 pts).


1
1
x
x
-5
5
-1
 = ___________________________
5
-5
-1
* = _________________________
BONUS (3 pts) HeH+ is a molecular ion found in space. Sketch its  MO with its nuclei
at +/- 2 and write its mathematical function, knowing that the 1s orbital for different
atoms is e -| Zx |
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17) Draw any “good” Lewis structure for each and predict the electron-domain name,
molecular structure, and hybrid orbital set used by the central atom: (15 pts)
BONUS ( 1pt each) When there are resonance structures involving formal charges,
show any one “good” Lewis structure with formal charges and any one possible
“best” Lewis structure with formal charges (and label which is good and best).
Lewis structure
electron-domain
molecular structure
hybrid orbital
+
H3O
SCN
-
CS2
BrO3
-
-
ICl4
BF3
CF4
-
NO2
SF2
-
AlCl4
KrF2
-
HSO3
SO3
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2-
SO3
ICl2-
18) Use each of the Molecular Orbital diagrams to fill in the number of electrons of
+
the three species NO, NO , NO and then calculate the bond order and give the
word to describe the magnetic property. (6 pts)
a) NO
z
x* y *
px py pz
px py pz
x y
E
z
s
s
s
s
N
Bond Order: ___________
NO
O
Magnetism: _______________________
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b) NO
z
x*y*
px py pz
px py pz
x y
E
z
s
s
s
s
N
NO
O
Bond Order: ___________
Magnetism: _______________________
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+
c) NO
z
x*y*
px py pz
px py pz
x y
E
z
s
s
s
s
N
NO+
O
Bond Order: ___________
Magnetism: _______________________
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BONUS Questions
1) Pose and answer a question from chapters 1-3 that you thought would be on
the exam but wasn’t. (up to 2 pts)
2) White blood cells generate the superoxide ion O2 to destroy other cells. Use
any of the above MO energy diagrams to describe the bonding and magnetic
properties of the superoxide ion (work need not be shown for credit). (2 pts)
3) Using bond enthalpies in the table, estimate the enthalpy of formation for
NCl3 (Calculations must be shown to receive credit: (3 pts)
Bond
Bond enthalpy (kJ/mol)
N - Cl
200
Cl - Cl
242
N≡N
941
Cl
2
Cl
N
N
N
+
3
Cl
Cl
Cl
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4) Give a reason why fulminate CNO is explosive while its isomer, cyanate,
NCO is not. (3 pts)
BONUS questions generated from the peanut gallery (as of last mid-night):
1) According to your textbook authors, who were the first two scientists to propose
the periodic table? (1 pt)
2)
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