A
Chemistry 1251
Final Exam (3 hours)
Dec. 10, 2010
Page Points Score Grader Page Points Score Grader
MC
75
8
27
9
23
6
23
10
20
7
32
Total
200
First Name
Last Name
Signature
SID# 800-
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Please read each question carefully and answer it completely and clearly. Most of the factual information
you need is contained in the text of the problem. Answer the questions you find easy first; questions which
appear difficult at first usually require a revision of your thought process. Always keep in mind the
intended examination topics if a problem appears too hard.
The quiz should have 10 pages for a total of 41 questions. Individual point values for free response are
given in the corner of each answer space. Multiple choice questions are worth 3 points each. You must
show your work to receive any credit for problems involving calculations. All calculated answers and
work must include units, and should be rounded to the appropriate number of significant figures!
88.91
1
A
Multiple Choice Instructions: Select the SINGLE BEST answer of the given choices, and cleanly bubble
your answer into the scantron form. Each question is worth 3 points (except questions 1-3).
1. The exam version, listed in the upper right hand corner, is…
(A) Version A
(B) Version B
(C) Version C
2. Did you take chemistry in high school?
(A) Yes; I earned an A.
(D) Yes; I earned a D or F.
(B) Yes;I earned a B.
(C) Yes; I earned a C.
(E) No, I did not take chemistry in high school.
3. How long ago did you finish your high school chemistry course?
(A) between Fall 2010  Spring 2011
(C) between Fall 2006  Spring 2009
(E) I did not take a high school chemistry class.
(B) between Fall 2009 – Spring 2010
(D) before Fall 2006
4. In the early 1800s, John Dalton established 4 postulates which became the basis for modern atomic
theory. Which of Dalton’s postulates of atomic theory listed below is most completely invalidated by the
subsequent discovery of isotopes?
(A) All matter is composed of small, indivisible particles called atoms.
(B) All atoms of a given element are identical in mass and properties.
(C) Compounds are formed by a combination of two or more atoms in definite arrangements in the
ratio of small whole numbers.
(D) Atoms are not created, destroyed or converted into other kinds of atoms during chemical reactions.
They are simply rearranged into new compounds.
5. Why will two H atoms react to form H2?
(A) The atoms stay together because they have higher energy, which is stored in the bond.
(B) The atoms stay together because they are ions and ions attract each other.
(C) The atoms stay together because the compound is lower in energy.
(D) The atoms stay together because they emit a line spectrum.
6. A student trying to identify an unknown solid hypothesizes that the solid is sodium chloride because the
solid is white. Which of the following results would NOT support this hypothesis?
(A) The solid dissolves in water.
(B) An aqueous solution of the material conducts electricity.
(C) When an aqueous solution of the solid is mixed with AgNO3(aq), a precipitate forms.
(D) The solid combusts to form CO2 (g) and H2O (g) in a very exothermic process.
7. One of the crucial nutrients for plant growth is nitrogen. Considering only the percent composition of
the following compounds, which would deliver the largest amount of nitrogen to increase the nitrogen
content of soil if 1.00 gram of each compound were added to the soil?
(A) NH3
(B) NH4NO3
(C) (NH4)2SO4
2
(D)
(NH2)2CO
A
8. The relative sizes of rubidium and bromine ATOMS are shown to the right:
Which of the following illustrations best illustrates the relatives sizes of Rb+ and Br –?
9. Rutherford’s gold foil experiment involved positively charged alpha particles interacting with the atoms
in a thin film of gold. As a result of this experiment, Rutherford…
(A) concluded that the mass of the proton and the neutron were very close and the mass of the electron
was negligible.
(B) calculated the charge and the mass of the electron and found that the charge of the electron was
negative.
(C) calculated the mass-charge ratio of the electron and found that the charge of the electron was negative.
(D) found that the nucleus was the small, dense and positively charged center of the atom.
(E) discovered that the cathode rays consisted of alpha, beta and gamma rays, for which he found the
charges to be positive, negative and neutral, respectively.
10. What is the correct formula for cobalt(III) oxide?
(A) CoO3
(B) Co3O
(C) Co2O3
(D) CO3
(E) Co3O2
11. A compound has the empirical formula C4H4O and a molar mass of 136 g/mol. What is the molecular
formula of the compound?
(A) C8H8O
(B) C8H8O2
(C) C4H4O
(D) C9H12O
(E) C5H6O2
12. GeF3H can be produced from the reaction: GeH4(s) + 3 GeF4 (s)  4 GeF3H. How many moles of GeF4
are needed to produce 8.50 moles of GeF3H?
(A) 3.24 moles
(B) 5.56 moles
(C) 6.48 moles
(D) 2.78 moles
(E) 6.38 moles
13. While trying to make a potential anti-tumor drug, researchers were able to generate 1.44 g of
compound. The percent yield from the reaction was 34.9 %. What was the theoretical yield of the
reaction?
(A) 0.503 g
(B) 0.937 g
(C) 2.21 g
3
(D) 4.13 g
(E) 5.57 g
A
14. The following diagram represents the molecular view of the reaction that occurs between H2(g) (lighter
color) and O2(g) (darker color) to produce H2O(g) under a given set of conditions. Which of the
following statements is correct regarding the diagram?
Insert Fig. 3.15
(A) O2(g) is the limiting reactant.
(B) The amount of water vapor produced is determined by the hydrogen gas present.
(C) For every 2 moles of hydrogen gas consumed, 1 mole of H2O(g) is produced.
(D) There will be hydrogen leftover after the reaction is complete.
(E) 10 g of H2 and 7 g of O2 will react to produce 10 g of H2O and 2 g of O2.
15. What is the concentration, in M K2SO4, of a solution made by dissolving 8.71 g K2SO4 in 250.0 mL of
water? (The molar mass of K2SO4 is 174.27 g/mol.)
(A) 2.00104 M
(B) 0.0348 M
(C) 0.0500 M
(D) 0.200 M
(E) 34.8 M
Matching. Match the description shown on the left with the proper quantum number shown on the right.
(Answers may be used more than once.)
16. Differentiates electrons in a magnetic field
(A) Principal Quantum Number (n)
17. Orientation of an orbital in space
(B) Angular Momentum Quantum Number (l)
18. Shape of an orbital
(C) Magnetic Quantum Number (ml)
19. Energy of an orbital
(D) Spin Quantum Number (ms)
20. Consider a ground state carbon atom. Which of the following statements is TRUE?
(A)
(B)
(C)
(D)
(E)
Carbon has 2 unpaired electrons.
Carbon has 4 valence electrons in p orbitals.
The 2s electrons are the highest energy electrons in carbon.
There are 2 empty p orbitals.
The valence electrons are shielded from the nucleus by the net effects of 4 electrons.
21. Which ion is isoelectronic with argon?
(A) Na+
(B) Ne
(C) Ca2+
4
(D) Se2
(E) Br
A
22. Consider the chemical equation: X+(g)  X2+(g) + e¯. Which of the following is FALSE?
(A) This represents the second ionization energy of X.
(B) This ionization energy is higher than that of X(g)  X+(g) + e¯.
(C) This is an exothermic reaction.
(D) If X is an alkali metal, IE2 will be much greater than IE1.
(E) Outer electrons in X2+ are held to the nucleus more tightly than outer electrons in X+.
Consider the structure of TNT (trinitrotoluene), shown to the right. Use this structure to answer the next
three (3) questions.
23. How many pi (π) bonds are in this structure?
(A) 0
(B) 3
(C) 6
(D) 12
(E) 21
24. What are the approximate O-N-O bond angles?
(A) 90°
(B) 109.5°
(C) 120°
(D) 180°
25. How many electron domains are there for the carbon
indicated by the arrow?
(A) 2
(B) 3
(C) 4
(D) 5
(E) 6
26. Which of the following statements is TRUE regarding the relative lattice energies or LiF and KI?
(A) LiF has a higher lattice energy because fluorine is more electronegative than iodine.
(B) LiF has a higher lattice energy because the Li+/F− distance is shorter than the K+/I− distance.
(C) KI has a higher lattice energy because the K+/I− distance is longer than the Li+/F− distance.
(D) KI has a higher lattice energy because KI has higher magnitude charges than LiF.
(E) KI has a higher lattice energy because there is a KI triple bond, but only an Li-F single bond.
27. Rank the C-O bond lengths from shortest to longest for: CO, CH2O, CO32−
(A) CO < CH2O < CO32−
(D) CH2O < CO32− < CO
(B) CO < CO32− < CH2O
(E) CO32− < CH2O < CO
(C) CH2O < CO < CO32−
28. In which of these structures is the dipole arrow for the overall molecular polarity shown correctly?
(A) ICl3
(B) BCl3
(C) CH3Cl
5
(D) NH3
(E) HCN
A
Free Response Instructions: Solve the following problems in the boxes provided below each question.
Be sure to show all your work and units, and round your answers to the appropriate significant figures.
29. 10.12 grams of solid (NH4)2SO4 are dissolved in 100.0 grams of water, and the temperature of the
solution goes from 25.00 °C to 23.97 °C.
(a) Write the balanced, chemical reaction showing how solid (NH4)2SO4 dissociates in water.
3
(b) Calculate the value of qrxn, in J; report your answer to 3 significant figures. Assume that ssoln = sH2O.
5
(c) Calculate ΔHrxn for the dissolution of (NH4)2SO4, in kJ/mol (NH4)2SO4.
4
30. PSCl3 (thiophosphoryl chloride) is a compound that is used to make pesticides.
(a) Draw the Lewis structure of PSCl3 that obeys the
(b) Draw the Lewis structure of PSCl3 that has
octet rule. Be sure to label non-zero formal charges.
minimized formal charges.
3
2
31. Consider the hybrid orbitals in the molecule ethylene, C2H4, which is used to make plastic bags.
(a) Identify the orbital overlaps that lead to the
sigma (σ) bonds between…
C
&
H
C
&
(b) Draw in the atomic orbitals used to create the
carbon- carbon pi (π) bond.
C
4
2
6
A
32. Consider the reaction: 3 NO2(g) + H2O(l)  2 HNO3(aq) + NO(g)
(a) Name each of the reactants and products in this reaction.
NO2
H2O
HNO3
NO
4
(b) Identify the oxidation state for the nitrogen identified in each of the compounds, including +/- signs:
NO2
HNO3
NO
6
(c) Draw one Lewis structure of HNO3. (Hint: H is connected to O.) Be sure to include any non-zero formal
charges on individual atoms.
(d) Can HNO3 exhibit resonance?
Resonance?
YES
NO
1
circle one
3
(e) If you have 3.00 grams of NO2 and 1.00 grams of H2O, what is the limiting reactant? Be sure to explicitly
show any mole ratios and any comparisons you use.
8
33. Consider hydrocyanic acid, HCN, which has the following VSEPR and Lewis structure:
(a) In the CN triple bond, there are
(b) To the right, draw the orbital energy
diagrams for the valence electrons
for C and for N. Be sure to clearly label
each set of orbitals with the hybrid or
atomic orbital name, and clearly show
the number of orbitals and electrons.
(c) Circle the electrons that correspond to
the lone pair on N.
(d) Put a box around the electrons that
participate in  bonding.
 bond(s) and
E
 bond(s). (Enter a number.)
2
E
C in HCN
7
N in HCN
8
A
34. 35.2 mL of H2SO4 (a diprotic acid) was titrated 46.5 mL of 0.100 M NaOH to reach the equivalence
point. What was the original concentration of the sulfuric acid?
8
35. Consider the green line in the hydrogen emission spectrum, which has a wavelength of 487 nm.
(a) What is the frequency, in Hz, associated with this wavelength?
3
(b) Another hydrogen emission line occurs at 657 nm. Does this correspond to a photon of higher or lower
energy than a 487 nm photon? Justify your answer with an equation, explanation, and/or calculation.
3
(c) The highest energy visible emission line is emitted when an electron moves from n = 6 to n = 2.
Calculate ΔE, the change in energy of this Bohr model atom, in J. Pay careful attention to the sign.
4
36. Dichloromethane, CH2Cl2, has a density of 1.33 g/mL and a molar mass of 84.93 g/mol. How many
Cl atoms are found in 7.00 mL of CH2Cl2?
9
8
A
37. Consider the combustion of 1 mole of methane gas, CH4(g): CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g).
(a) Draw an accurate 3-D structure for CH4,
using lines, dashes, and wedges to show
orientation.
Show the position of all
nonbonding pairs of electrons.
(b) What is the hybridization of the central atom?
2
(c) What is the molecular geometry of the molecule?
2
(d) Is the molecule polar or non-polar?
3
2
(e) Calculate ΔH°rxn for the combustion of 1 mole of methane from the ΔH°f values below. As a reminder,
the overall equation is: CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g).
ΔH°f (CH4) = −74.8 kJ/mol
ΔH°f(CO2) = −393.5 kJ/mol
ΔH°f(H2O) = −241.8 kJ/mol
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(f) Calculate ΔH°rxn for the combustion of 1 mole of methane from the AVERAGE bond enthalpies below. As
a reminder, the overall equation is: CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g).
O−H
460. kJ/mol
O−O
142 kJ/mol
O=O
499 kJ/mol
C−O
351 kJ/mol
C=O
799 kJ/mol
CO
1070. kJ/mol
C−H
413 kJ/mol
6
38. The exact mass of 198Os is 188.958 amu. Why doesn’t this match the mass listed in the periodic table?
2
9
A
39. Na3PO4(aq) reacts with Fe(NO3)2(aq), and a solid precipitate forms.
(a) Give the systematic chemical name for the reactants.
Na3PO4
Fe(NO3)2
2
(b) Write the balanced, molecular equation, including states of matter.
3
(c) Write the ionic equation, including states of matter.
3
(d) Write the net ionic equation, including states of matter.
1
40. (a) Draw the orbital diagram, including electrons, for a neutral sulfur atom.
(b) Circle the valence electrons in your orbital diagram.
(c) Identify the quantum numbers for the last
electron added to your orbital diagram.
n
l
ml
E
ms
2
5
41. Complete the table shown.
Nuclide
Representation
Protons
Neutrons
33
42
Electrons
Net Charge
198 Os
10
3−
4