Name:     ID: 
 
    Email: 

CHEM 1211 Practice Final Exam Spring 2005  50 points

Multiple Choice
Identify the letter of the choice that best completes the statement or answers the question.
 

 1. 
What is the name of the element with the symbol Pb?
a.
iron
b.
lead
c.
phosphorus
d.
plutonium
e.
rubidium
 

 2. 
Which of the following substances is not an element?
a.
NO
b.
Ca
c.
Na
d.
B
e.
Xe
 

 3. 
A pure substance that is composed of two or more different elements is
a.
a chemical compound.
b.
a multi-element.
c.
an atom.
d.
a heterogeneous mixture.
e.
a homogeneous mixture.
 

 4. 
Which statement concerning NaCl is true?
a.
NaCl has properties similar to sodium metal and chlorine gas.
b.
NaCl is a homogeneous mixture.
c.
NaCl is a heterogeneous mixture.
d.
The percentage of Na in NaCl is dependent on where the sample is obtained.
e.
NaCl is composed of ions, which are electrically charged atoms.
 

 5. 
The molecular model depicts a molecule composed of carbon (black), oxygen (gray), and hydrogen (white. atoms. What is the correct molecular formula?

1211finals05_files/i0060000.jpg
a.
CHO
b.
C6H6O2
c.
C6H7O
d.
C7H6O
e.
C7H6O2
 

 6. 
All of the following are examples of physical properties EXCEPT
a.
the density of neon gas.
b.
the conductivity of copper wire.
c.
the boiling point of water.
d.
the frying of an egg.
e.
the density of mercury metal.
 

 7. 
All of the following are examples of chemical change EXCEPT
a.
the condensation of steam.
b.
the rusting of iron.
c.
the combustion of propane gas.
d.
the tarnishing of silver.
e.
the decomposition of water to hydrogen gas and oxygen gas.
 

 8. 
When 12 copper pennies are submerged in water, the pennies displace 4.13 cm3 of water. If the combined mass of the pennies is 36.93 g, what is the density of copper?
a.
0.745 g/cm3
b.
3.49 g/cm3
c.
8.94 g/cm3
d.
32.8 g/cm3
e.
153 g/cm3
 

 9. 
Thermostats are often set to 72°F. What is this temperature in Celsius?
a.
8°C
b.
22°C
c.
37°C
d.
58°C
e.
72°C
 

 10. 
Which is a correct method for converting Fahrenheit to Celsius?
a.
°C = 1211finals05_files/i0110000.jpg °F + 32
b.
°C = 1211finals05_files/i0110001.jpg°F + 32
c.
°C = 1211finals05_files/i0110002.jpg (°F + 32)
d.
°C = 1211finals05_files/i0110003.jpg (°F - 32)
e.
°C = 1211finals05_files/i0110004.jpg (°F - 32)
 

 11. 
The temperature required to melt NaCl is 801 K. What is this temperature in Celsius?
a.
298°C
b.
327°C
c.
528°C
d.
852°C
e.
1074°C
 

 12. 
According to the kinetic-molecular theory of matter, particles in a liquid
a.
are packed closely together in a regular array.
b.
are close together, but they are not confined to specific positions.
c.
expand to fill their container.
d.
vibrate back and forth about an average position.
e.
move slower as the temperature increases.
 

 13. 
Which of the following is a homogeneous mixture?
a.
italian salad dressing
b.
chocolate chip ice cream
c.
gasoline
d.
a rock, such as granite or marble
e.
a jar of chunky peanut butter
 

 14. 
Which term best describes ethylene glycol, C2H6O2?
a.
chemical compound
b.
solution
c.
homogeneous mixture
d.
heterogeneous mixture
e.
none of the above
 

 15. 
A common wavelength of light emitted from a red laser pointer is 6.50 ´ 102 nm. What is the wavelength in meters?
a.
6.50 ´ 10-9 m
b.
6.50 ´ 10-7 m
c.
6.50 ´ 10-5 m
d.
6.50 ´ 10-3 m
e.
6.50 ´ 100 m
 

 16. 
A rectangular box has dimensions of 20.0 cm ´ 15.0 cm ´ 8.00 cm. Calculate the volume of the box in liters.
a.
2.40 ´ 10-3 L
b.
4.30 ´ 10-3 L
c.
2.40 L
d.
43.0 L
e.
2.40 ´ 103 L
 

 17. 
Which is a correct method of determining the number of liters of gas required to fill an automobile's 15 gallon tank? (1.000 L = 1.057 quarts, 4 quarts = 1 gallon)
a.
15 gallons 1211finals05_files/i0180000.jpg
b.
15 gallons 1211finals05_files/i0180001.jpg
c.
15 gallons 1211finals05_files/i0180002.jpg
d.
15 gallons 1211finals05_files/i0180003.jpg
e.
none of the above
 

 18. 
The mass of a sample weighed on an electronic balance that is sensitive to ±2 mg is 21.7834 g. What is the correct number of significant figures for this measurement?
a.
2
b.
3
c.
4
d.
5
e.
6
 

 19. 
Two electronic balances are tested using a standard weight. The true mass of the standard is 1.0000 g. The results of 5 individual measurements on each balance are recorded below.

 
Balance A
Balance B
 
0.8888 g
1.3110 g
 
0.9959 g
1.3109 g
 
1.1182 g
1.3111 g
 
1.0033 g
1.3110 g
 
0.9938 g
1.3110 g
average mass =
1.0000 g
1.3110 g

Which statement best describes the results?
a.
A: good precision, good accuracy. B: good precision, good accuracy
b.
A: good precision, good accuracy. B: good precision, poor accuracy
c.
A: poor precision, good accuracy. B: good precision, good accuracy
d.
A: poor precision, good accuracy. B: good precision, poor accuracy
e.
A: poor precision, good accuracy. B: poor precision, poor accuracy
 

 20. 
Express 0.05620 in exponential notation.
a.
5.6 ´ 10-2
b.
5.62 ´ 10-2
c.
5.620 ´ 10-2
d.
5.6 ´ 102
e.
5.62 ´ 102
 

 21. 
All atoms of the same element have the same number of ________ in their nucleus.
a.
neutrons
b.
electrons
c.
protons
d.
neutrons and protons
e.
neutrons, protons, and electrons
 

 22. 
All of the following statements are true EXCEPT
a.
for any neutral element, the number of protons and electrons are equal.
b.
isotopes of an element have the same atomic number.
c.
the mass number is the sum of the number of protons and neutrons.
d.
the atomic number equals the number of protons.
e.
all atoms of a given element have the same mass number.
 

 23. 
How many protons, neutrons, and electrons are in a carbon-13 atom?
a.
6 protons, 6 neutrons, 1 electron
b.
6 protons, 7 neutrons, 6 electrons
c.
7 protons, 6 neutrons, 6 electrons
d.
7 protons, 6 neutrons, 7 electrons
e.
13 protons, 13 neutrons, 13 electrons
 

 24. 
Which two of the ions below have the same number of electrons?

1211finals05_files/i0250000.jpg
a.
1211finals05_files/i0250001.jpg
b.
1211finals05_files/i0250002.jpg
c.
1211finals05_files/i0250003.jpg
d.
1211finals05_files/i0250004.jpg
e.
none of the above
 

 25. 
Which two of the following atoms are isotopes?

1211finals05_files/i0260000.jpg
a.
1211finals05_files/i0260001.jpg
b.
1211finals05_files/i0260002.jpg
c.
1211finals05_files/i0260003.jpg
d.
1211finals05_files/i0260004.jpg
e.
none of the above
 

 26. 
What is the identity of 1211finals05_files/i0270000.jpg?
a.
molybdenum
b.
technetium
c.
americium
d.
copper
e.
iodine
 

 27. 
What is a correct method for calculating the mass of 3.0 ´ 1022 sodium atoms?
a.
3.0 ´ 1022 atoms Na1211finals05_files/i0280000.jpg
b.
3.0 ´ 1022 atoms Na1211finals05_files/i0280001.jpg
c.
3.0 ´ 1022 atoms Na1211finals05_files/i0280002.jpg
d.
3.0 ´ 1022 atoms Na1211finals05_files/i0280003.jpg
e.
6.02 ´ 1023 atoms Na1211finals05_files/i0280004.jpg
 

 28. 
How many elements are in the second period of the periodic table?
a.
2
b.
8
c.
10
d.
18
e.
32
 

 29. 
What element is in the third period in Group 5A?
a.
As
b.
Nb
c.
P
d.
In
e.
Tl
 

 30. 
Which three elements are likely to have similar chemical and physical properties?
a.
nitrogen, oxygen, and neon
b.
sodium, magnesium, and aluminum
c.
calcium, strontium, and barium
d.
nickel, copper, and zinc
e.
uranium, plutonium, and americium
 

 31. 
The smallest unit into which a pure substance such as sugar or water can be divided, while still retaining its composition and chemical properties is a(n) ________.
a.
atom
b.
molecule
c.
isotope
d.
mixture
e.
ion
 

 32. 
Acetylsalicylic acid, commonly known as aspirin, has 9 carbon atoms, 8 hydrogen atoms, and 4 oxygen atoms per molecule. What is the molecular formula of aspirin?
a.
C9H8O4
b.
Ca9H8O2
c.
C9He8O4
d.
C9H8Ox4
e.
C9(HO4)2
 

 33. 
A bromide ion has ________ electrons.
a.
8
b.
34
c.
35
d.
36
e.
37
 

 34. 
Identify the ions present in K2HPO4.
a.
K+ and HPO42-
b.
K22+, H+, and PO43-
c.
K+, H+, P3-, and O2-
d.
K2H3+ and PO43-
e.
K2HPO4 is not ionic.
 

 35. 
Identify the ions in (NH4)2SO4.
a.
N3-, H+, S2-, and O2-
b.
NH22+ and H2SO4
c.
NH4+ and SO42-
d.
NH42+ and SO42-
e.
NH4+ and SO4-
 

 36. 
Which atom is most likely to form a -2 ion?
a.
Zn
b.
S
c.
F
d.
P
e.
Ca
 

 37. 
Which atom is most likely to form a +2 ion?
a.
K
b.
C
c.
Kr
d.
Ba
e.
Se
 

 38. 
All of the following formulas are correct EXCEPT
a.
Ba(NO3)2
b.
KClO4
c.
Na3N
d.
Al2(SO4)3
e.
Ca2HPO4
 

 39. 
For a nonmetal in Group 6A of the periodic table, the most common monatomic ion will have a charge of ________.
a.
-2
b.
-1
c.
0
d.
+2
e.
+6
 

 40. 
Which formula represents the binary compound formed by strontium ions and phosphate ions?
a.
Sr2(PO4)3
b.
SrPO4
c.
Sr2P3
d.
Sr3(PO4)2
e.
SrP
 

 41. 
What is the correct formula for a binary compound that contains magnesium and bromine?
a.
MgBr
b.
Mg2Br2
c.
Mg2Br
d.
MgBr2
e.
Mg2Br3
 

 42. 
What are the integer values for x and y, respectively, for Alx(CO3)y?
a.
1 and 2
b.
2 and 3
c.
1 and 3
d.
3 and 1
e.
3 and 2
 

 43. 
What is the correct formula for potassium nitrate?
a.
KN
b.
K3N
c.
KNO
d.
KNO2
e.
KNO3
 

 44. 
What is the correct formula for magnesium carbonate?
a.
Mg2C
b.
MgCO3
c.
Mg2CO3
d.
Mg(CO3)2
e.
Mg3(CO3)2
 

 45. 
What is the correct formula for aluminum sulfide?
a.
AlSO3
b.
Al2(SO4)3
c.
AlS
d.
Al3(SO3)2
e.
Al2S3
 

 46. 
What is the correct name for NH4NO3?
a.
ammonium nitrate
b.
nitrogen tetrahydronitrogen trioxide
c.
nitrogen tetrahydronitrate
d.
tetrahydronitrogen nitrate
e.
dinitrogen hydrogenate
 

 47. 
What is the correct name for CoBr2?
a.
cobalt(II) dibromate
b.
cobalt(II) dibromide
c.
cobalt bromine
d.
monocobalt dibromate
e.
cobalt(II) bromide
 

 48. 
What is the common name for NH3?
a.
mononitrogen trihydrogen
b.
hydrazine
c.
nitrogen
d.
ammonia
e.
trihydrogen nitride
 

 49. 
What is the common name for N2O4?
a.
nitrogen tetraoxide
b.
di(nitrogen dioxide)
c.
dinitrogen oxide
d.
dinitroxide
e.
dinitrogen tetraoxide
 

 50. 
The empirical formula of a hydrocarbon with a molar mass of 78.11 g/mol is CH. What is the molecular formula?
a.
C6H6
b.
C5H28
c.
C5H2O
d.
C2H4
e.
C8H18
 

 51. 
What is the molar mass of nitroglycerine, C3H5(ONO2)3?
a.
41.07 g/mol
b.
227.1 g/mol
c.
103.1 g/mol
d.
165.1 g/mol
e.
204 g/mole
 

 52. 
What is the mass percent of iron in iron(II) oxalate, FeC2O4?
a.
14.29%
b.
61.18%
c.
32.07%
d.
38.82%
e.
81.17%
 

 53. 
A molecule is found to contain 47.35% C, 10.60% H, and 42.05% O. What is the empirical formula for this molecule?
a.
C2H6O
b.
C2H6O2
c.
C3H8O2
d.
C3H6O3
e.
C4H6O
 

 54. 
Sulfur dioxide may be prepared by the reaction of sulfur with oxygen gas according to the chemical equation below.

__ S8(s) + __ O2(g) ® __ SO2(g)

What are the respective coefficients when the equation is balanced with the smallest whole numbers?
a.
8, 8, 8
b.
2, 16, 8
c.
1, 8, 8
d.
2, 16, 16
e.
1, 2, 1
 

 55. 
When ethanol undergoes complete combustion, the products are carbon dioxide and water.

__ C2H5OH(ª) + __ O2(g) ® __ CO2(g) + __ H2O(g)

What are the respective coefficients when the equation is balanced with the smallest whole numbers?
a.
2, 7, 4, 6
b.
1, 3, 2, 3
c.
2, 2, 1, 4
d.
1, 2, 3, 2
e.
2, 4, 6, 4
 

 56. 
The products of the complete combustion of a hydrocarbon are carbon dioxide and water. Write a balanced chemical equation for the combustion of butane, C4H10.
a.
2 C4H10(g) + 13 O2(g) ® 8 CO2(g) + 10 H2O(g)
b.
C4H10(g) ® 4 C(s) + 4 H2(g)
c.
C4H10(g) + 13 O2(g) ® 4 CO2(g) + 5 H2O(g)
d.
C4H10(g) + 9 O2(g) ® 4 CO2(g) + 5 H2O(g)
e.
None of the above are correctly balanced.
 

 57. 
Aluminum reacts with oxygen to produce aluminum oxide.

4 Al(s) + 3 O2(g) ® 2 Al2O3(s)

If 3.0 moles of Al reacts with excess O2, how many moles of Al2O3 can be formed?
a.
1.5 mol
b.
2.0 mol
c.
2.7 mol
d.
3.0 mol
e.
4.5 mol
 

 58. 
What mass of carbon is needed to react completely with 23.14 grams of SiO2 according to the following equation?

SiO2(s) + 3 C(s) ® SiC(s) + 2 CO(g)
a.
1.16 g
b.
4.62 g
c.
13.9 g
d.
38.6 g
e.
116 g
 

 59. 
What is a correct method for determining how many grams of oxygen will react with 1.00 gram of propane?

C3H8(g) + 5 O2(g) ® 3 CO2(g) + 4 H2O(g)
a.
1.00g C3H81211finals05_files/i0600000.jpg=
b.
1.00g C3H81211finals05_files/i0600001.jpg=
c.
1.00g C3H81211finals05_files/i0600002.jpg=
d.
1.00g C3H81211finals05_files/i0600003.jpg=
e.
1.00g C3H81211finals05_files/i0600004.jpg=
 

 60. 
What is a correct method for determining how many moles of magnesium oxide will be formed from the reaction of 5.00 g magnesium with excess oxygen?

2 Mg(s) + O2(g) ® 2 MgO(s)
a.
5.00g Mg1211finals05_files/i0610000.jpg
b.
5.00g Mg1211finals05_files/i0610001.jpg
c.
5.00g Mg1211finals05_files/i0610002.jpg
d.
5.00g Mg1211finals05_files/i0610003.jpg
e.
none of the above
 

 61. 
Which one of the following solutions will have the highest electrical conductivity?
a.
0.010 M KCl
b.
0.010 M CaI2
c.
0.010 M MgSO4
d.
0.010 M Al(NO3)3
e.
0.010 M Na2SO4
 

 62. 
Which statement about the reaction below is correct?

K2SO4 + Ba(NO3)2 ® BaSO4 + 2 KNO3
a.
BaSO4 will precipitate.
b.
KNO3 will precipitate.
c.
Both BaSO4 and KNO3 will precipitate.
d.
Neither BaSO4 nor KNO3 will precipitate.
e.
No reaction will occur because K2SO4 is insoluble.
 

 63. 
Which of the following compounds is a weak acid?
a.
HCl
b.
HF
c.
HBr
d.
HNO3
e.
HClO4
 

 64. 
What is the net ionic equation for the reaction below?

AgNO3(aq) + KBr(aq) ® AgBr(s) + KNO3(aq)
a.
K+(aq) + NO3-(aq) ® KNO3(s)
b.
AgNO3(aq) + KBr(aq) ® AgBr(s)
c.
K+(aq) + NO3-(aq) ® KNO3(aq)
d.
AgNO3(aq) + KBr(aq) ® AgBr(s) + KNO3(aq)
e.
Ag+(aq) + Br-(aq) ® AgBr(s)
 

 65. 
What is the net ionic equation for the reaction of sodium hydroxide with iron(III) nitrate?
a.
3 Na+(aq) + Fe3+(aq) ® Na3Fe(s)
b.
NaOH(aq) + FeNO3(aq) ® FeOH(s) + NaNO3(aq)
c.
Fe3+(aq) + 3 NO3-(aq) ® Fe(NO3)3(s)
d.
Fe3+(aq) + 3 OH-(aq) ® Fe(OH)3(s)
e.
Na+(aq) + NO3-(aq) ® NaNO3(s)
 

 66. 
What is the net ionic equation for the reaction of potassium hydroxide and hydrochloric acid?
a.
H+(aq) + KOH(aq) ® H2O(ª) + K+(aq)
b.
K+(aq) + Cl-(aq) ® KCl(aq)
c.
HCl(aq) + KOH(aq) ® H2O(ª)
d.
H+(aq) + OH-(aq) ® H2O(ª)
e.
KOH(aq) + H2O(ª) ® H+(aq) + K(OH)2(s)
 

 67. 
What is the oxidation number of manganese in KMnO4?
a.
-2
b.
0
c.
+3
d.
+5
e.
+7
 

 68. 
What is the oxidation number of phosphorus in CaHPO4?
a.
-3
b.
-1
c.
+1
d.
+3
e.
+5
 

 69. 
If 1.928 g KNO3 is dissolved in enough water to make 250.0 mL of solution, what is the molarity of potassium nitrate?
a.
6.912 ´ 10-4 M
b.
4.767 ´ 10-3 M
c.
7.627 ´ 10-2 M
d.
1.297 ´ 10-1 M
e.
7.712 M
 

 70. 
Which of the following directions correctly describe the preparation of 0.500 L of 0.150 M NaOH from a 6.00 M stock solution?
a.
Dilute 0.200 L of 6.00 M NaOH to a volume of 0.500 L.
b.
Dilute 12.5 mL of 6.00 M NaOH to a volume of 0.500 L.
c.
Combine 0.200 L of 6.00 M NaOH with 0.500 L of water.
d.
Dilute 475 mL of 6.00 M NaOH to a volume of 0.500 L.
e.
Combine 12.5 mL of 6.00 M NaOH with 0.500 L of water.
 

 71. 
Potassium hydrogen phthalate (KHP) is a weak acid that is used to standardize sodium hydroxide according to the net ionic equation below.

HC8H4O4-(aq) + OH-(aq) ® H2O(ª) + C8H4O42-(aq)

If 1.02 g KHP (molar mass = 204.2 g/mol) is titrated with 28.34 mL of NaOH, what is the concentration of NaOH?
a.
0.03536 M
b.
0.1385 M
c.
0.2004 M
d.
0.2176 M
e.
0.2713 M
 

 72. 
All of the following statements are true EXCEPT
a.
In an endothermic process heat is transferred from the surroundings to the system.
b.
The greater the specific heat of an object, the more thermal energy it can store.
c.
The SI unit of specific heat capacity is joules per gram per kelvin.
d.
Heat is transferred from the system to the surroundings in an exothermic process.
e.
The temperature of a system is a state function.
 

 73. 
Specific heat capacity is
a.
the quantity of heat required to melt 1.00 g of a substance.
b.
the mass of a substance 1.00 J of energy will heat by 1.00 K.
c.
the mass of a substance 1.00 cal of energy will heat by 1.00 K.
d.
the temperature change undergone when 1.00 g of a substance absorbs 1.00 cal.
e.
the quantity of heat needed to change 1.00 g of a substance by 1.00 K.
 

 74. 
If 34.8 J is required to change the temperature of 10.0 g of mercury by 25 K, what is the specific heat of mercury?
a.
0.139 J/g·K
b.
0.338 J/g·K
c.
0.718 J/g·K
d.
0.870 J/g·K
e.
1.93 J/g·K
 

 75. 
Calculate the amount of heat required to change 35.0 g ice at -25.0ºC to steam at 125ºC. (Heat of fusion = 333 J/g; heat of vaporization = 2260 J/g; specific heats: ice = 2.09 J/g·K, water = 4.18 J/g·K, steam = 1.84 J/g·K)
a.
22.0 kJ
b.
90.9 kJ
c.
109 kJ
d.
276 kJ
e.
3290 kJ
 

 76. 
10.0 g of ice at 0.00ºC is mixed with 50.0 g of water at 32.0ºC. What is the final temperature of the mixture? (Heat of fusion = 333 J/g; specific heats: ice = 2.09 J/g·K, water = 4.184 J/g·K)
a.
-4.59ºC
b.
0.00ºC
c.
4.59ºC
d.
13.4ºC
e.
23.8ºC
 

 77. 
Calculate DE of a gas for a process in which the gas absorbs 42 J of heat and does 14 J of work on the surroundings (i.e. the gas expands)?
a.
-56 J
b.
-28 J
c.
+28 J
d.
+42 J
e.
+56 J
 

 78. 
For a particular process q = 25 kJ and w = -15 kJ. What conclusions may be drawn for this process?
a.
DE = 40 kJ
b.
DE = -40 kJ
c.
This is a product favored reaction.
d.
Work is done by the system on the surroundings.
e.
Both answer b and d are correct.
 

 79. 
One statement of the first law of thermodynamics is that
a.
the amount of work done on a system is independent of pathway.
b.
the total energy flow in or out of a system is equal to the sum of the heat absorbed and the work done on the system.
c.
the heat flow in or out of a system is independent of pathway.
d.
the total work done on a system must equal the heat absorbed by the system.
e.
in any chemical process the sum of the heat flow and the work must equal zero.
 

 80. 
Calculate the standard enthalpy of formation of carbon monoxide,

C(s) + 1/2 O2(g) ® CO(g), given the enthalpies of the reactions below.

C(s) + O2(g) ® CO2(g)
DH = -393.5 kJ
2 CO(g) + O2(g) ® 2 CO2(g)
DH = -566.0 kJ
a.
-959.6 kJ
b.
-421.6 kJºC
c.
-172.5 kJ
d.
-110.5 kJ
e.
172.5 kJ
 

 81. 
Calculate the enthalpy for the formation of calcium carbonate from calcium oxide and carbon dioxide,

CaO(s) + CO2(g) ® CaCO3(s)

given the enthalpies of the reactions below.

2 Ca(s) + O2(g) ® 2 CaO(s)
DH = -1270.2 kJ
C(s) + O2(g) ® CO2(g)
DH = -393.5 kJ
2 Ca(s) + 2 C(s) + 3 O2(g) ® 2 CaCO3(s)
DH = -2413.8 kJ
a.
-4077.3 kJ
b.
-2235.5 kJ
c.
-750.1 kJ
d.
-350.2 kJ
e.
-178.3 kJ
 

 82. 
Determine the heat of reaction for the oxidation of iron,

4 Fe(s) + 3 O2(g) ® 2 Fe2O3(s)

given the enthalpies of the reactions below.

2 Fe(s) + 6 H2O(ª) ® 2 Fe(OH)3(s) + 3 H2(g)
DH = 321.8 kJ
2 H2(g) + O2(g) ® 2 H2O(ª)
DH = -571.7 kJ
Fe2O3(s) + 3 H2O(ª) ® 2 Fe(OH)3(s)
DH = 288.6 kJ
a.
-1648.7 kJ
b.
-636.9 kJ
c.
-505.3 kJ
d.
387.0 kJ
e.
+1447.1 kJ
 

 83. 
Calculate the molar enthalpy of combustion of C3H6(g),

C3H6(g) + 9/2 O2(g) ® 3 CO2(g) + 3 H2O(ª)

using standard enthalpies of formation.

molecule
DHfº (kJ)
C3H6(g)
+53.3
CO2(g)
-393.5
H2O(ª)
-285.8
a.
-2091.2 kJ
b.
-1984.6 kJ
c.
-187.8 kJ
d.
-62.6 kJ
e.
+732.3 kJ
 

 84. 
What type of orbital is designated n = 4, ª = 2, mª = +1?
a.
4s
b.
4p
c.
4d
d.
2f
e.
none
 

 85. 
What type of orbital is designated n = 2, ª = 0, mª = 0?
a.
2s
b.
2p
c.
2d
d.
2f
e.
none
 

 86. 
All of the following sets of quantum numbers are allowed EXCEPT
a.
n = 6,ª = 4, mª = +2
b.
n = 3, ª = 2, mª = -1
c.
n = 4, ª = 1, mª = 0
d.
n = 1, ª = 0, mª = 0
e.
n = 2, ª = 3, mª = +3
 

 87. 
What is the maximum number of orbitals when n = 6 and ª = 2?
a.
1
b.
2
c.
3
d.
5
e.
10
 

 88. 
The n = _____ shell is the lowest that may contain an f-orbital.
a.
1
b.
2
c.
3
d.
4
e.
5
 

 89. 
Which of the following diagrams represents a d-orbital?

1211finals05_files/i0900000.jpg
a.
(I) only
b.
(II) only
c.
(III) only
d.
(IV) only
e.
(I) and (IV)
 

 90. 
Which of the following diagrams represent a p-orbital?

1211finals05_files/i0910000.jpg
a.
(I) only
b.
(II) only
c.
(III) only
d.
(IV) only
e.
(I) and (II)
 

 91. 
The Pauli exclusion principle states that
a.
electrons can have either positive or negative _ spins.
b.
no two electrons in an atom can have the same 4 quantum numbers.
c.
two electrons can share the same orbital if they have the same spin.
d.
no two electrons in an atom can have the same spin.
e.
atoms with one or more unpaired electrons are paramagnetic.
 

 92. 
What is the maximum number of electrons that can exist in the shell n = 2?
a.
2
b.
8
c.
18
d.
32
e.
50
 

 93. 
Which +3 ion has the electron configuration [Ar]3d3?
a.
Fe
b.
Nb
c.
Cr
d.
Mo
e.
Sc
 

 94. 
If the electron configuration of an element is [Ar]3d104s24p4, what is the charge on the monoatomic anion of the element?
a.
-4
b.
-2
c.
+2
d.
+4
e.
+6
 

 95. 
What is the electron configuration for Cu+?
a.
[Ar]
b.
[Ar]3d8
c.
[Ar]4s23d8
d.
[Ar]3d10
e.
[Ar]4s13d9
 

 96. 
What is the electron configuration for an iodine atom?
a.
[Kr]4d105s25p5
b.
[Kr]4f145d106s26p5
c.
[Kr]5d106s26p5
d.
[Xe]5p-1
e.
[Kr]4d104f145d5
 

 97. 
Which of the following atoms and ions have the same electron configuration: I-, Pb4+, Xe, Ba2+ and Sn2+?
a.
I-, Xe, and Ba2+
b.
I-, Pb4+, Xe, and Ba2+
c.
Pb4+ and Sn2+
d.
Pb4+ and Ba2+
e.
none of the above
 

 98. 
What element has the following electron configuration?

1211finals05_files/i0990000.jpg
a.
Br
b.
Ag
c.
Ga
d.
Kr
e.
I
 

 99. 
What -1 ion the following electron configuration?

1211finals05_files/i1000000.jpg
a.
Na-
b.
Ar
c.
In-
d.
K-
e.
Cl-
 

 100. 
Which type of elements have no affinity for electrons?
a.
transition metals
b.
main group metals
c.
noble gases
d.
main group nonmetals
e.
semiconductors
 

 101. 
Which group of the periodic table elements forms only +1 ions?
a.
group 1A
b.
group 2A
c.
group 7B
d.
group 7A
e.
group 8A
 

 102. 
Which of the following ions is least likely to be formed?
a.
Al3+
b.
Cu+
c.
Na+
d.
Ti4+
e.
Sr3+
 

 103. 
What is the expected number of valence electrons for a group 3A element?
a.
0
b.
3
c.
5
d.
6
e.
10
 

 104. 
Which of the following elements is most likely to form a molecule that exceeds the octet rule?
a.
Ne
b.
C
c.
O
d.
P
e.
Be
 

 105. 
Which of the following combinations is most likely to produce ionic bonds?
a.
O and H
b.
Al and S
c.
C and N
d.
N and O
e.
S and Cl
 

 106. 
When both of the electrons in a molecular bond originate from the same atom, the bond is called a(n)
a.
ionic bond.
b.
free radical bond.
c.
coordinate covalent bond.
d.
Lewis dot structure.
e.
double bond.
 

 107. 
What is the total number of valence electrons in a carbon tetrachloride molecule?
a.
24
b.
30
c.
32
d.
36
e.
48
 

 108. 
Which of the following is a correct Lewis structure for SO2?

1211finals05_files/i1090000.jpg
a.
1
b.
2
c.
3
d.
4
e.
5
 

 109. 
H3PO3 is a diprotic acid (i.e. it has two acid functions). Which of the following Lewis structures is most likely correct for H3PO3?

1211finals05_files/i1100000.jpg
a.
1
b.
2
c.
3
d.
4
e.
5
 

 110. 
Which of the following is a correct Lewis structure for sulfate ion?

1211finals05_files/i1110000.jpg
a.
1
b.
2
c.
3
d.
4
e.
5
 

 111. 
Which of the following is a possible Lewis structures for C2H6O?

1211finals05_files/i1120000.jpg
a.
1
b.
2
c.
3
d.
1 and 2
e.
1 and 3
 

 112. 
Electronegativity is a measure of
a.
the charge on an electron.
b.
a molecule's polarity.
c.
the charge on an atom.
d.
the number of extra electrons on an anion.
e.
an atom's ability to attract electrons to itself.
 

 113. 
Predict which of the following compounds will have the bond that is most polar.
a.
NH3
b.
CF4
c.
H2O
d.
HF
e.
HI
 

 114. 
When heated, azomethane decomposes into nitrogen gas and methane gas.

CH3N=NCH3(g) ® N2(g) + C2H6(g)


Bond
Bond Enthalpy (kJ/mol)
 

Bond
Bond Enthalpy (kJ/mol)
C-H
413
 
N-N
163
C-N
305
 
N=N
418
C-C
346
 
NºN
945

Using average bond enthalpies, calculate the enthalpy of reaction.
a.
-609 kJ/mol
b.
-583 kJ/mol
c.
-462 kJ/mol
d.
-263 kJ/mol
e.
-197 kJ/mol
 

 115. 
Use VSEPR theory to predict the electron pair geometry and the molecular geometry of SO2.
a.
e- pair geometry = trigonal planar, molecular geometry = bent
b.
e- pair geometry = trigonal planar, molecular geometry = linear
c.
e- pair geometry = tetrahedral, molecular geometry = bent
d.
e- pair geometry = tetrahedral, molecular geometry = trigonal planar
e.
e- pair geometry = tetrahedral, molecular geometry = linear
 

 116. 
Use VSEPR theory to predict the molecular geometry of SF2.
a.
bent
b.
linear
c.
trigonal pyramidal
d.
tetrahedral
e.
octahedral
 

 117. 
Use VSEPR theory to predict the molecular geometry HCN.
a.
bent
b.
linear
c.
trigonal planar
d.
tetrahedral
e.
octahedral
 

 118. 
Use VSEPR theory to predict the molecular geometry of BCl3.
a.
bent
b.
trigonal pyramidal
c.
trigonal planar
d.
tetrahedral
e.
t-shaped
 

 119. 
What are the bond angles in SCN-?
a.
90°
b.
109°
c.
120°
d.
180°
e.
90° and 109°
 

 120. 
What are the bond angles in SiH4?
a.
90°
b.
109°
c.
120°
d.
180°
e.
90° and 109°
 

 121. 
How many sigma (s) bonds and pi (p) bonds are in the following molecule?

1211finals05_files/i1220000.jpg
a.
seven s and two p
b.
six s and two p
c.
eleven s and zero p
d.
nine s and two p
e.
two s and nine p
 

 122. 
How many sigma (s) bonds and pi (p) bonds are in acetone?

1211finals05_files/i1230000.jpg
a.
eight s and one p
b.
six s and one p
c.
nine s and one p
d.
one s and nine p
e.
one s and eight p
 

 123. 
In order to form a set of sp hybrid orbitals, how many pure atomic orbitals must be mixed?
a.
one s, one p
b.
two s, one p
c.
two s, two p
d.
one s, two p
e.
zero s, two p
 

 124. 
In order to form a set of sp3d hybrid orbitals, how many pure atomic orbitals must be mixed?
a.
one s, one p, and one d
b.
one s, three p, and one d
c.
two s, one p, and two d
d.
two s, six p, and two d
e.
none of the above
 

 125. 
What is the maximum number of hybridized orbitals that can be formed by a nitrogen atom?
a.
1
b.
2
c.
3
d.
4
e.
6
 

 126. 
What is the maximum number of hybridized orbitals that can be formed by xenon?
a.
0
b.
2
c.
4
d.
5
e.
6
 

 127. 
In which of the following molecules does the carbon atom have sp hybridization: HCN, CH4, CO2, and CH2O?
a.
CH4 only
b.
CH4 and CH2O
c.
HCN and CH2O
d.
HCN and CO2
e.
HCN, CO2, and CH2O
 

 128. 
In which of the following molecules or ions does the central atom have sp2 hybridization: NH2-, H2O, BH3, SO2?
a.
NH2- and BH3
b.
H2O and SO2
c.
H2O, BH3, and SO2
d.
NH2-, H2O, and SO2
e.
BH3 and SO2
 

 129. 
What is the hybridization of each oxygen atom in O2?
a.
sp
b.
sp2
c.
sp3
d.
sp3d
e.
sp3d2
 

 130. 
What is the hybridization of the sulfur atom in sulfate ion, SO42-?
a.
sp
b.
sp2
c.
sp3
d.
sp3d
e.
sp3d2
 

 131. 
What is the molecular geometry around an atom that is sp3 hybridized and has two lone pairs of electrons?
a.
bent
b.
linear
c.
trigonal pyramidal
d.
trigonal planar
e.
trigonal bipyramidal
 

 132. 
What is the molecular geometry around an atom that is sp3 hybridized, has three sigma bonds, no pi bonds, and one lone pair?
a.
bent
b.
linear
c.
trigonal pyramidal
d.
trigonal planar
e.
tetrahedral
 

 133. 
What is the molecular geometry around an atom that is sp hybridized, has two sigma bonds, two pi bonds, and no lone pairs?
a.
bent
b.
linear
c.
trigonal planar
d.
tetrahedral
e.
octahedral
 

 134. 
Which of the following hybridized atoms is not possible?
a.
an sp hybridized oxygen atom
b.
an sp3 hybridized nitrogen atom
c.
an sp2 hybridized carbon atom
d.
an sp2 hybridized boron atom
e.
an sp3d2 hybridized fluorine atom
 

 135. 
At constant temperature, 10.0 L of N2 at 0.983 atm is compressed to 2.88 L. What is the final pressure of N2?
a.
0.283 atm
b.
0.293 atm
c.
2.98 atm
d.
3.41 atm
e.
28.3 atm
 

 136. 
Avogadro's law states that equal volumes of gases under the same conditions of temperature and pressure have equal ________.
a.
masses
b.
numbers of molecules
c.
molar masses
d.
densities
e.
velocities
 

 137. 
Which of the following relationships are true for gases?

1. The volume of a gas is directly proportional to its pressure in mm Hg.
2. The pressure of a gas in inversely proportional to its temperature in kelvin.
3. The moles of a gas are directly proportional to the gas constant R.
a.
1 only
b.
2 only
c.
3 only
d.
2 and 3
e.
none are true
 

 138. 
Ammonia gas is synthesized according to the balanced equation below.

N2(g) + 3 H2(g) ® 2 NH3(g)

If 2.50 L N2 react with 7.00 L H2, what is the theoretical yield (in liters) of NH3? Assume that the volumes of reactants and products are measured at the same temperature and pressure.
a.
2.50 L
b.
4.67 L
c.
5.00 L
d.
7.00 L
e.
10.5 L
 

 139. 
Which of the following are postulates of kinetic-molecular theory of gases?

1. The distance between gas molecules is large in comparison to their size.
2. The velocity of a gas molecule is inversely proportional to its temperature.
3. Gas molecules are in constant, random motion.
4. At a given temperature, all gases have the same average kinetic energy.
a.
1 and 4
b.
1, 2, and 4
c.
1, 3, and 4
d.
2 and 3
e.
3 and 4
 

 140. 
Non-ideal behavior for a gas is most likely to be observed under conditions of
a.
high temperature and high pressure.
b.
low temperature and high pressure.
c.
low temperature and low pressure.
d.
standard temperature and pressure.
e.
high temperature and low pressure.
 

 141. 
One way in which real gases differ from ideal gases is that the molecules of a real gas
a.
have no kinetic energy.
b.
occupy no volume.
c.
are attracted to each other.
d.
have positive and negative spins.
e.
are always polar.
 



 
Submit          Reset Help