AP Chemistry Summer Packet
Complete the following questions, due on the first day of school.
 I encourage you to work with each other.
 The lectures are located online, on nghs-science.wikispaces.com if you would like some additional help. The
questions in this packet are from units 1-3.
 If you get stuck, email, text, or call me (ms.marchilena@gmail.com; Rebekah.marchilena@ngsd.k12.wi.us;
608.417.9574). I will be around for most of the summer and can meet you to help you through anything you
need. I am willing to check your answers at any point to make sure you are doing the problems correctly.
 This will NOT be a participation grade. You will be graded on your answers. There will be no possibility of
redoing the packet, as I will be giving you the answers once you turn in the packet.
 This will count as a test grade (tests count for 50% of your AP Chemistry grade).
 You have nearly 3 months to complete what would have been 4 weeks’ worth of work; do not wait until the last
minute!!
 If you do not complete this packet for the first day of class, please see Mrs. Meoska and find another class to take.
1. Figures a-d represent atomic scale views of different samples of substances:
Figure a
Figure b
Figure c
Figure d
a. Under one set of conditions, the substances in (a) and (b) mix and the result is depicted in (c). Does this
represent a chemical or a physical change?
b. Under a second set of conditions, the same substances mix and the result is depicted in (d). Does this
represent a chemical or a physical change?
c. Under a third set of conditions, the sample depicted in (c) changes into (d). Does this represent a chemical or
a physical change?
d. When the change in part c occurs, does the sample have different chemical properties? Physical properties?
Explain.
2. Define physical property, physical change, chemical property, and chemical change. State which type of
property or change is portrayed in each of the following:
a. Passing an electric current through molten magnesium chloride yields molten magnesium and gaseous
chlorine.
b. A magnet separates a mixture of black iron shavings and white sand.
c. The iron in discarded automobiles slowly forms reddish brown, crumbly rust.
d. Yellow-green chlorine gas attacks silvery sodium to form white crystals of sodium chloride.
3. Which of the following changes can be reversed by changing the temperature? Why is this possible?
a. Dew condensing on a leaf.
b. An egg turning hard when it is boiled.
c. Ice cream melting.
d. A spoonful of batter cooking on a hot griddle.
Why?
4. For each pair, which has the higher potential energy? Kinetic energy?
a. The fuel in your car or the products in its exhaust
b. Wood in a fireplace or the after in the fireplace after the wood burns
c. A sled resting at the top of a hill or a sled sliding down the hill
d. Water above a dam or water falling over the dam
5. How are the key elements of scientific thinking used in the following scenario? While making your breakfast
toast, you notice that it fails to pop out of the toaster. Thinking the spring mechanism is stuck, you notice that the
bread is unchanged. Assuming you forgot to plug in the toaster, you check and find it is plugged in. When you
take the toaster into the dining room and plug it into a different outlet, the toaster works. Returning to the kitchen,
you turn on the switch for the overhead light and nothing happens.
6. When you convert feet to inches, how do you decide which portion of the conversion factor should be in the
numerator and which in the denominator?
7. Write the conversion factors for:
a. in2 to m2
b. km2 to cm2
c. mi/h to m/s
d. lb/ft3 to g/cm3
e. cm/min to in/s
f.
m3 to in3
g. m/s2 to km/h2
h. gal/h to l/min
8. Describe the difference between intensive and extensive properties. Which of the following are intensive?
a. mass
b. density
c. volume
d. melting point
9. For each of the following cases, state whether the density of the object increases, decreases, or remains the same.
a. A sample of chlorine gas is compressed.
b. A lead weight is carried from sea level to the top of a high mountain.
c. A sample of water is frozen.
d. A diamond is submerged in water.
10. The average radius of a molecule of lysozyme, an enzyme in tears, is 1430 pm. What is its radius in nanometers
(nm)?
11. The radius of a barium atom is 2.22x10-10m. What is its radius in angstroms (Å)?
12. A small hole in the wing of a space shuttle requires a 20.7 cm2 patch.
a. What is the patch’s area in square kilometers?
b. If the patching material costs NASA $3.25/in2, what is the cost of the patch?
13. The volume of a certain bacterial cell is 2.56 µm3.
a. What is the volume of the cell in cubic millimeters (mm3)?
b. What is the volume of 105 cells in liters (l)?
14. How many cubic meters of milk are in 1 qt (946.4 ml)? How many liters of milk are in 835 gallons (1 gal = 4 qt)?
15. An empty vial weighs 55.32g.
a. If the vial weighs 185.56 g when filled with liquid mercury (d = 13.53 g/cm3), what is its volume?
b. How much would the vial weigh if it were filled with water (d = 0.997 g/cm3)?
16. Perform the following conversions:
a. 72 °F to °C and K
b. -164 °C to K and °F
c. 0K to °C and °F
17. In the early 20th century, thin metal foils were used to study atomic structure.
a. How many in2 of gold foil with a thickness of 1.6x10-5 in could have been made from 2.0 troy oz (1 troy
oz = 31.1 g; d of gold = 19.3 g/cm3)?
b. If gold cost $20.00/troy oz at that time, how many cm2 of gold foil could have been made for $75.00
worth of gold?
18. A 25.0 g sample of each of three unknown metals is added to 25.0 ml of water in graduated cylinders A, B, and C,
and the three final volumes are depicted in the circles below. Given their densities, identify the metal in each
cylinder: zinc (7.14 g/ml), iron (7.87 g/ml), and nickel (8.91 g/ml).
19. Underline the significant zeroes in the following numbers:
a. 0.41
b. 0.041
c. 0.0410
d. 4.0100x104
20. Carry out the following calculations, making sure that your answer has the correct number of significant figures:
a. 2.795 m · 3.10 m
6.48 m
b. V = 4πr2 where r = 17.282 cm
3
c. 1.110 cm + 17.3 cm + 108.2 cm + 316 cm
21. Write the following numbers in scientific notation:
a. 131,000.0
b. 0.00047
c. 210,006
d. 2160.5
e. 282.0
f.
0.0380
g. 4270.8
h. 58,200.9
22. Carry out each of the following calculations, paying special attention to significant figures, rounding, and units:
a. (1.84x102g)(44.7m/s2) The term 2 is exact.
2
b.
(1.07x10-4mol/L)2(3.8x10-3mol/L) mol=mole, the SI unit of amount of substance
(8.35x10-5 mol/L)(1.48x10-2 mol/L)3
23. How long is the metal strip shown below? Be sure to answer with the correct number of significant figures.
24. These organic solvents are used to clean compact discs:
Solvent
Density (g/ml) at 20°C
Chloroform
1.492
Diethyl ether
0.714
Ethanol
0.789
Isopropanol
0.785
Toluene
0.867
a. If a 10.00 ml sample of CD cleaner weighs 11.775 g at 20°C, which solvent is most likely present?
b. The chemist analyzing the cleaner calibrates her equipment and finds that the pipet is accurate to ±0.02
ml, and the balance is accurate to ±0.003 g. Is this equipment precise enough to distinguish between
ethanol and isopropanol? Explain.
25. A laboratory instructor gives a sample of amino acid powder to each of four students, I, II, III, and IV, and they
weigh the samples. The true value is 8.72 g. Their results for three trials are:
I: 8.72 g, 8.74 g, 8.70 g
II: 8.56 g, 8.77 g, 8.83 g
III: 8.50 g, 8.48 g, 8.51 g
IV: 8.41 g, 8.72 g, 8.55 g
a. Calculate the average mass from each set of data, and tell which set is most accurate.
b. Precision is a measure of the average deviations of each piece of data from the average value. Which set
of data is the most precise? Is this set also the most accurate? Explain.
c. Which set of data is both the most accurate and most precise?
d. Which set of data is both the least accurate and least precise?
26. To make 2.000 gal of a powdered sports drink, a group of students measure out 2.000 gal of water with 500 ml,
50 ml, and 5 ml graduated cylinders. Show how they could get closest to 2.000 gal of water, using these cylinders
the fewest times.
27. Suppose your dorm room is 11 feet wide by 12 feet long by 8.5 feet high and has an air conditioner that
exchanges air at a rate of 1200 L/min. How long would it take the air conditioner to exchange the air in your
room once?
28. Bromine is used to prepare the pesticide methyl bromide and flame retardants for plastic electronic housings. It is
recovered from seawater, underground brines, and the Dead Sea. The average concentrations of bromine in
seawater (d = 1.024 g/ml) and the Dead Sea (d = 1.22 g/ml) are 0.065 g/L and 0.50 g/L, respectively. What is the
mass ratio of bromine in the Dead Sea to that in seawater?
29. At room temperature (20°C) and pressure, the density of air is 1.189 g/L. An object will float in air if its density
is less than that of air. In a buoyancy experiment with a new plastic, a chemist creates a rigid, thin-walled ball
that weighs 0.12 g and has a volume of 560 cm3.
a. Will the ball float if it is evacuated?
b. Will it float f filled with carbon dioxide (d = 1.830 g/L)?
c. Will it float if filled with hydrogen (d = 0.0899 g/L)?
d. Will it float if filled with oxygen (d = 1.330 g/L)?
e. Will it float if filled with nitrogen (d = 1.165 g/L)?
f.
For any case that will float, how much weight must be added to make the ball sink?
30. Asbestos is a fibrous silicate mineral with remarkably high tensile strength. But it is no longer used because
airborne asbestos particles can cause lung cancer. Grunerite, a type of asbestos, has a tensile strength of 3.5x102
kg/mm2 (thus a strand of grunerite with a 1 mm2 cross sectional area can hold up to 3.5x102 kg). The tensile
strengths of aluminum and steel #5137 are 2.5x104 lb/1n2 and 5.0x104 lb/in2, respectively. Calculate the cross
sectional area (in mm2) of wires made of aluminum and steel that would have the same tensile strength as a fiber
of grunerite with a cross sectional area of 1.0 µm2.
31. Liquid nitrogen is obtained from liquefied air and is used industrially to prepare frozen foods. It boils at 77.36 K.
a. What is the boiling temperature in °C?
b. What is the boiling temperature in °F?
c. At the boiling point, the density of the liquid is 809 g/L and that of the gas is 4.566 g/L. How many liters
of liquid nitrogen are produced when 895.0 L of nitrogen gas is liquefied at 77.36 K?
32. According to the lore of ancient Greece, Archimedes discovered the displacement method of density
determination while bathing, and then used it to find the composition of the king’s crown. If a crown weighing 4
lb 3 oz displaces 186 ml of water, is the crown made of pure gold (d = 19.3 g/cm3)?
33. The Environmental Protection Agency (EPA) proposed a safety standard for microparticles in air: for particles up
to 2.5µm in diameter, the maximum allowable amount is 50. µg/m3. If your room is 10.0 ft x 8.25 ft x 12.5 ft and
just meets the EPA standards, how many of these particles are in your room? How many are in each 0.500 L
breath you take? (Assume the particles are spheres of diameter 2.5 µm and made primarily of soot, a form of
carbon with a density of 2.5 g/cm3.)
34. Molecular scenes A and B depict changes in matter at the atomic level:
a. Which show(s) a physical change?
b. Which show(s) a chemical change?
c. Which result(s) in different physical
properties?
d. Which result(s) in different chemical
properties?
e. Which result(s) in a change in state?
35. Which of the following are pure substances? Explain.
a. Calcium chloride, used to melt ice on roads, consists of 2 elements, calcium and chlorine, in a fixed mass
ratio.
b. Sulfur consists of sulfur atoms combined into octatomic molecules.
c. Baking powder, a leavening agent, contains 26% to 30% sodium hydrogen carbonate and 30% to 35%
calcium dihydrogen phosphate by mass.
d. Cytosine, a component of DNA, consists of H, C, N, and O atoms bonded in a specific arrangement.
36. Each molecular-scale scene below represents a mixture. Describe each one in terms of the number of elements
and/or compounds present.
37. Identify the mass law that each of the following observations demonstrates, and explain your reasoning.
a. A sample of potassium chloride from Chile contains the same percent by mass of potassium as one from
Poland.
b. A flashbulb contains magnesium and oxygen before use and magnesium oxide afterward, but its mass
does not change.
c. Arsenic and oxygen form one compound that is 65.2 mass % arsenic and another that is 75.8 mass %
arsenic.
38. State the mass law(s) demonstrated by the following experimental results, and explain your reasoning:
a. A student heats 1.00g of a blue compound and obtains 0.64 g of a white compound and 0.36g of a
colorless gas.
b. A student heats 3.25g of the same blue compound and obtains 2.08g of a white compound and 1.17g of a
colorless gas.
39. Fluorite, a mineral of calcium, is a compound of the metal with fluorine. Analysis shows that a 2.76 g sample of
fluorite contains 1.42g of calcium. Calculate the following:
a. The mass of fluorine in the sample.
b. Mass fractions of calcium and fluorine in fluorite.
c. Mass percents of calcium and fluorine in fluorite.
40. A compound of copper and sulfur contains 88.39g of metal and 44.61g of nonmetal. How many grams of copper
are in 5264kg of compound? How many grams of sulfur?
41. A compound of iodine and cesium contains 63.94g of metal and 61.06g of nonmetal. How many grams of cesium
are in 38.77g of compound? How many grams of iodine?
42. Show, with calculations, how the following data illustrate the law of multiple proportions:
Compound I: 47.5 mass % sulfur and 52.5 mass % chlorine
Compound II: 31.1 mass % sulfur and 68.9 mass % chlorine
43. Use Dalton’s theory to explain why potassium nitrate from India or Italy has the same mass percents of K, N, and
O.
44. Choose the correct answer. The difference between the mass number of an isotope and its atomic number is:
a. Directly related to the identity of the element
b. The number of electrons
c. The number of neutrons
d. The number of isotopes
45. Argon has three naturally occurring isotopes, 36Ar, 38Ar, and 40Ar. What is the mass number of each? How many
protons, neutrons, and electrons are present in each?
46. Do both members of the following pairs have the same number of protons? Neutrons? Electrons?
a. 168O and 178O
b.
40
and 4119K
c.
60
and 6028Ni
18Ar
27Co
d. Which pair(s) consist(s) of atoms with the same Z value? N value? A value?
47. Write the AZX notation or each atomic depiction:
48. Gallium has two naturally occurring isotopes, 69Ga (isotopic mass 68.9256 amu, abundance 60.11%) and 71Ga
(isotopic mass 70.9247 amu, abundance 39.89%). Calculate the atomic mass of gallium.
49. Magnesium has three naturally occurring isotopes, 24Mg (isotopic mass 23.9850 amu, abundance 78.99%), 25Mg
(isotopic mass 24.9858 amu, abundance 10.00%), and 26Mg (isotopic mass 25.9826 amu, abundance 11.01%).
Calculate the atomic mass of magnesium.
50. Copper has two naturally occurring isotopes, 63Cu (isotopic mass 62.9396 amu) and 65Cu (isotopic mass 64.9278
amu). If copper has an atomic mass of 63.546 amu, what is the abundance of each isotope?
51. Correct each of the following statements:
a. In the modern periodic table, the elements are arranged in order of increasing atomic mass.
b. Elements in a period have similar chemical properties.
c. Elements can be classified as either metalloids or nonmetals.
52. Give the name, atomic symbol, and group number of the element with the following Z value, and classify it as a
metal, nonmetal, or metalloid:
a. Z = 32
b. Z = 15
c. Z = 2
d. Z = 3
e. Z = 42
f. Z = 33
g. Z = 20
h. Z = 35
i. Z = 19
j. Z = 13
53. Given that the ions in LiF and MgO are of similar size, which compound has stronger ionic bonding? Use
Coulomb’s law in your explanation.
54. Describe the formation of solid magnesium choride (MgCl2) from large numbers of magnesium and chlorine
atoms.
55. Does potassium nitrate (KNO3) incorporate ionic bonding, covalent bonding, or both? Explain.
k
56. What monatomic ions do potassium (Z=19) and iodine (Z=53) form?
57. What monatomic ions do barium (Z=56) and selenium (Z=34) form?
58. An ionic compound forms when lithium (Z=3) reacts with oxygen (Z=8). If a sample of the compound contains
8.4 x 1021 lithium ions, how many oxide ions does it contain?
59. The ionic compound forms when calcium (Z=20) reacts with iodine (Z=53). If a sample of the compound
contains 7.4 X 1021 calcium ions, how many iodide ions does it contain?
60. Consider a mixture of 10 billion O2 molecules and 10 billion H2 molecules. In what way is this mixture similar to
a sample containing 10 billion hydrogen peroxide (H2O2) molecules? In what way is it different?
61. Write an empirical formula for each of the following:
a. Hydrazine, a rocket fuel, molecular formula N2H4
b. Glucose, sugar, molecular formula C6H12O6
c. ethylene glycol, car antifreeze, molecular formula C2H6O2
d. peroxodisulfuric acid, a compound wised to make bleaching agents, molecular formula H2S2O8
62. Give the name and formula of the compound formed from the following elements:
a. sodium and nitrogen
b. oxygen and strontium
c. aluminum and chlorine
d. cesium and bromine
e. sulfur and barium
f.
calcium and fluorine
63. Give the name and formula of the compound formed from the following elements:
a. 12Ψ and 9Θ
b.
30Ψ
and 16Θ
c.
17Ψ
and 38Θ
d.
37Ψ
and 35Θ
e.
8Ψ
f.
20Ψ
and 13Θ
and 53Θ
64. Give the systematic names for the formulas or the formulas for the names:
a. Tin(IV) chloride
b. FeBr3
c. Cuprous bromide
d. Mn2O3
e. Na2HPO4
f.
Potassium carbonate dehydrate
g. NaNO2
h. Ammonium perchlorate
65. Correct each of the following formulas:
a. Barium oxide is BaO2.
b. Iron(II) nitrate is Fe(NO3)3.
c. Magnesium sulfide is MnSO3.
66. Correct each of the following names:
a. CuI is cobalt(II) iodide.
b. Fe(HSO4)3 is iron(II) sulfate.
c. MgCr2O7 is magnesium dichromium heptaoxide.
67. Give the number of atoms of the specified element in a formula unit of each of the following compounds, and
calculate the molecular (formula) mass:
a. Oxygen in aluminum sulfate, Al2(SO4)3.
b. Hydrogen in ammonium hydrogen phosphate, (NH4)2HPO4.
c. Oxygen in the mineral azurite, Cu3(OH)2(CO3)2.
d. Hydrogen in ammonium benzoate, C6H5COONH4.
e. Nitrogen in hydrazinium sulfate, N2H6SO4.
f.
Oxygen in the mineral leadhillite, PB4SO4(CO3)2(OH)2.
68. Write the formula of each compound, and determine its molecular (formula) mass:
a. Ammonium sulfate
b. Sodium dihydrogen phosphate
c. Potassium bicarbonate
d. Sodium dichromate
e. Ammonium perchlorate
f.
Magnesium nitrate trihydrate
69. What is the difference between a homogeneous and a heterogeneous mixture?
70. Classify each of the following as a compound, a homogeneous mixture, or a heterogeneous mixture:
a. Distilled water
b. Gasoline
c. Beach sand
d. Wine
e. Air
f.
Orange juice
g. Vegetable soup
h. Cement
i.
Calcium sulfate
j.
Tea
71. Scenes A-I depict various types of matter on the atomic scale. Choose the correct scene(s) for each of the
following:
a. A mixture that fills its container
b. A substance that cannot be broken down into simpler ones
c. An element with a very high resistance to flow
d. A homogenous mixture
e. An element that conforms to the walls of its container and displays a surface
f.
A gas consisting of diatomic particles
g. A gas that can be broken down into simpler substances
h. A substance with a 2:1 ratio of its component atoms
i.
Matter that can be separated into its component substances by physical means
j.
A heterogeneous mixture
k. Matter that obeys the law of definite proportions
72. Nitrogen forms more oxides than any other element. The percents by mass of N in three differen nitrogen oxides
are (I) 46.69%, (II) 36.85%, and (III) 25.94%.
a. Determine the empirical formula of each compound.
b. How many grams of oxygen per 1.00 g of nitrogen are in each compound?
73. Dinitrogen monoxide (N2O; nitrous oxide) is a greenhouse gas that enters the atmosphere principally from natural
fertilizer breakdown. Some studies have shown that the isotope ratios of 15N to 14N and of 18O to 16O in N2O
depend on the source, which can thus be determined by measuring the relative abundance of molecular masses in
a sample of N2O.
a. What different molecular masses are possible for N2O?
b. The percent abundance of 14N is 99.6%, and that of 16O is 99.8%. Which molecular mass of N2O is least
common, and which is most common?
74. The scenes below represent a mixture of two monatomic gases undergoing a reaction when heated. Which mass
law(s) is (are) illustrated by this change?
75. When barium reacts with sulfur to form barium sulfide (BaS), each Ba atom reacts with an S atom. If 2.50 cm3 of
Ba reacts with 1.75 cm3 of S, are there enough Ba atoms to react with the S atoms? (dBa = 3.51 g/cm3; dS = 2.07
g/cm3)
76. The two isotopes of potassium with significant abundance in nature are 39K (isotopic mass 38.9637 amu,
93.258%) and 41K (isotopic mass 40.9618 amu, 6.730%). Fluorine has only one naturally occurring isotope, 19F
(isotopic mass 18.9984 amu). Calculate the formula mass of potassium fluoride.
77. Dimercaprol (HSCH2CHSHCH2OH) is a complexing agent developed during WWI and an antidote to arsenicbased poison gas and used today to treat heavy-metal poisoning. Such an agent binds and removes the toxic
element from the body.
a. If each molecule binds one arsenic (As) atom, how many atoms of As could be removed by 250 mg of
dimercaprol?
b. If one molecule binds one metal atom, calculate the mass % of each of the following metals in a metaldimercaprol complex:
i. Mercury
ii. Thallium
iii. Chromium
78. The anticancer drug Platinol (cicplatin), Pt(NH3)2Cl2, reacts with the cancer cell’s DNA and interferes with its
growth.
a. What is the mass % of platinum in Platinol?
b. If Pt costs $19/g, how many grams of Platinol can be made for $1.00 million (assume the cost of Pt
determines the cost of the drug)?
79. Choose the box color(s) in the periodic table below that match(es) the following:
a. Four elements that are nonmetals.
b. Two elements that are metals.
c. Three elements that are gases at room temperature.
d. Three elements that are solid at room temperature.
e. One pair of elements likely to form a covalent compound.
f.
Another pair of elements likely to form a covalent compound.
g. One pair of elements likely to form an ionic compound with the formula MX.
h. Another pair of elements likely to form an ionic compound with the formula MX.
i.
Two elements likely to form an ionic compound with the formula M2X.
j.
Two elements likely to form and ionic compound with the formula MX2.
k. An element that forms no compounds.
l.
A pair of elements whose compounds exhibit the law of multiple proportions.
80. A rock is 5.0% by mass fayalite (Fe2SiO4), 7.0% by mass forsterite (Mg2SiO4), and the remainder silicon dioxide.
What is the mass percent of each element in the rock?
81. Which of the following steps in an overall process involve(s) a physical change and which involve(s) a chemical
change?
82. The atomic mass of Cl is 35.45 amu, and the atomic mass of Al is 26.98 amu. What are the masses in grams of 3
mol of Cl and 2 mol of Al atoms?
83. Why might the expression “1 mol of chlorine” be confusing? What change would remove any uncertainty? For
what other elements might a similar confusion exist?
84. What advantage is there to using a counting unit (the mole) in chemistry rather than a mass unit?
85. Calculate the molar mass of each of the following:
a. (NH4)3PO4
b. CH2Cl2
c. CuSO4  5H2O
d. BrF2
86. Calculate the molar mass of each of the following:
a. N2O4
b. C8H10
c. MgSO4  7H2O
d. Ca(C2H3O2)2
87. Calculate each of the following quantities:
a. Mass in kilograms of 4.6  1021 molecules of NO2
b. Moles of Cl atoms in 0.0615g of C2H4Cl2
c. Number of H- ions in 5.82g of SrH2
88. Calculate the following quantities:
a. Total number of ions in 38.1g of SrF2
b. Mass in kilograms of 3.58 mol of CuCl2  2H2O
c. Mass in milligrams of 2.88  1022 formula units of Bi(NO3)3  5H2O
89. Calculate each of the following quantities:
a. Mass in grams of 8.42 mol of chromium (III) sulfate decahydrate.
b. Mass in grams of 1.83  1024 molecules of dichlorine heptaoxide.
c. Number of moles and formula units in 6.2 g of lithium sulfate.
d. Number of lithium ions, sulfate ions, S atoms, and O atoms in the mass of compound in part c.
90. Calculate each of the following:
a. Mass % of I in strontium periodate.
b. Mass % of Mn in potassium permanganate.
91. Propane is widely used in liquid form as a fuel for barbeque grills and camp stoves. For 85.5 g of propane,
calculate:
a. The moles of compound.
b. The grams of carbon.
92. Hemoglobin, a protein in red blood, carries O2 from the lungs to the body’s cells. Iron (as ferrous ion, Fe2+)
makes up 0.33 mass% of hemoglobin. If the molar mass of hemoglobin is 6.8  104 g/mol, how many Fe2+ ions
are in one molecule?
93. What is the empirical formula and the empirical formula mass for each of the following compounds?
a. C4H8
b. C3H6O3
c. P4O10
d. Ga2(SO4)3
e. Al2Br6
94. What is the molecular formula of each compound?
a. Empirical formula CH2 (MW = 42.08 g/mol)
b. Empirical formula C3H6O2 (MW = 32.05 g/mol)
c. Empirical formula HgCl (MW = 472.1 g/mol)
d. Empirical formula C7H4O2 (MW = 240.20 g/mol)
95. Determine the empirical formulas of the following compounds:
a. 0.039 mol of iron atoms combined with 0.052 mol of oxygen atoms.
b. 0.903 g of phosphorus combined with 6.99 g of bromine.
c. A hydrocarbon with 79.9 mass % carbon.
96. A 0.370 mol sample of a metal oxide (M2O3) weighs 55.4g.
a. How many moles of O are in the sample?
b. How many moles of M are in the sample?
c. What element is represented by the symbol M?
97. In the process of balancing the equation Al + Cl2  AlCl3
Student I writes: Al + Cl2  AlCl2
Student II writes: Al + Cl2 + Cl  AlCl3
Student III writes 2Al + 3Cl2  2AlCl3
Is the approach of Student I valid? Student II? Student III? Explain.
98. Write the balanced equations for each of the following reactions:
a. ___Cu(NO3)2 (aq) + ___KOH (aq)  Cu(OH)2 (s) + ___KNO3 (aq)
b. ___BCl3 (g) + ___H2O (l)  ___H3BO3 (s) + ___HCl (g)
c. ___CaSiO3 (s) + ___HF (g)  ___SiF4 (g) + ___CaF2 (s) + ___H2O (l)
d. ___(CN)2 (g) + ___H2O (l)  ___H2C2O4 (aq) + ___NH3 (g)
99. Convert the following into balanced equations:
a. When lead(II) nitrate solution is added to potassium iodide solution, solid lead(II) iodide forms and
potassium nitrate solution remains.
b. Liquid disilicon hexachloride reacts with water to from solid silicon dioxide, hydrogen chloride gas, and
hydrogen gas.
c. When nitrogen dioxide gas is bubbled into water, a solution of nitric acid forms and gaseous nitrogen
monoxide is released.
100.
Chromium(III) oxide reacts with hydrogen sulfide (H2S) gas to form chromium(III) sulfide and water:
Cr2O3 (s) + 3H2S (g)  Cr2S3 (s) + 3H2O (l)
To produce 421 g of Cr2S3,
a. How many moles of Cr2O3 are required?
b. How many grams of Cr2O3 are required?
101.
acid.
Calculate the mass of each product formed when 174g of silver sulfide reacts with excess hydrochloric
Ag2S (s) + HCl (aq)  AgCl (s) + H2S (g) (unbalanced)
102.
Elemental sulfur occurs as octatomic molecules, S8. What mass of fluorine gas is needed to react
completely with 17.8g of sulfur to form sulfur hexafluoride?
103.
Metal hydrides react with water to form hydrogen gas and the metal hydroxide. For example:
SrH2 (s) + 2H2O (l)  Sr(OH)2 (s) + 2H2 (g)
You wish to calculate the mass of hydrogen gas that can be prepared from 5.70g of SrH2 and 4.75g of H2O.
a. How many moles of H2 can form from the given mass of SrH2?
b. How many moles of H2 can form from the given mass of H2O?
c. Which is the limiting reagent?
d. How many grams of H2 can form?
104.
Calculate the maximum numbers of moles and grams of H2S that can form when 158g of aluminum
sulfide reacts with 131g of water:
Al2S3 + H2O  Al(OH)3 + H2S (unbalanced)
What mass of the excess reactant remains?
105.
A mixture of 0.0375g of hydrogen and 0.0185 mol of oxygen in a closed container is sparked to initiate a
reaction. How many grams of water can form? Which reactant is in excess and how many grams of it will
remain after the reaction?
106.
Calcium nitrate and ammonium fluoride react to form calcium fluoride, dinitrogen monoxide, and water.
What mass of each substance is present after 16.8g of calcium nitrate and 17.50g of ammonium fluoride react
completely?
107.
Two successive reactions, D  E and E  F have yields of 48% and 73% respectively. What is the
overall percent yield for conversion of D  F.
108.
What is the percent yield of a reaction in which 200.0 grams of phosphorus trichloride reacts with excess
water to form 128g of HCl and aqueous phosphorus acid (H3PO3)?
109.
When 56.6 g of calcium and 30.5 g of nitrogen gas undergo a reaction that has a 93.0% yield, what mass
of calcium nitride forms?
110.
Compressed butane gas is used as a liquid fuel in disposable cigarette lighters and lightweight camping
stoves. Supposed a lighter contains 5.50 ml of butane (d = 0.579 g/ml).
a. What mass of oxygen is needed to burn the butane completely?
b. How many moles of H2O form when all the butane burns?
c. How many total molecules of gas form when the butane burns completely?
111.
Sodium borohydride (NaBH4) is used industrially in many organic synthesis. One way to prepare it is by
reacting sodium hydride with gaseous diborane (B2H6). Assuming an 88.5% yield, how many grams of NaBH4
can be prepared by reacting 7.98g of sodium hydride and 8.16 g of diborane?
112.
Calculate each of the following quantities:
a. Volume in milliliters of 2.26M potassium hydroxide that contains 8.42g of solute.
b. Number of Cu2+ ions in 52 L of 2.3M copper(II) chloride.
c. Molarity of 275ml of solution containing 135mmol of glucose.
113.
A sample of concentrated nitric acid has a density of 1.41g/ml and contains 70.0% HNO3 by mass.
a. What mass of HNO3 is present per liter of solution?
b. What is the molarity of the solution?
114.
115.
How many milliliters of 0.083M HCl are needed to react with 16.2g of CaCO3?
2HCl + CaCO3  CaCl2 + CO2 + H2O
Muriatic acid, an industrial grad of concentrated (11.7M) HCl, is used to clean masonry and cement.
a. Write instructions for diluting the concentrated acid to make 3.0 gallons of 3.5M acid for routine use (1
gal = 4 qt; 1 qt = 0.946L).
b. How many milliliters of the muriatic acid solution contain 9.66g of HCl?
116.
Narceine is a narcotic in opium. It crystallizes from water solution as a hydrate that contains 10.8 mass%
water. If the molar mass of narceine hydrate is 499.52 g/mol, determine the x in narceine  xH2O.
117.
During studies of the reaction below:
2N2H4 (l) + N2O4 (l)  3N2 (g) + 4H2O (g)
a chemical engineer measured a less-than-expected yield of N2 and discovered that the following side reaction
occurs:
N2H4 (l) + 2N2O4 (l)  6NO (g) + 2H2O (g)
In one experiment, 10.0 g of NO formed when 100.0g of each reactant was used. What is the highest percent
yield of N2 that can be expected to form?
118.
Seawater is approximately 4.0% by mass dissolved ions. About 85% of the mass of the dissolved ions is
from NaCl.
a. Calculate the mass percent of NaCl in seawater.
b. Calculate the mass percent of Na+ ions and Cl- ions in seawater.
c. Calculate the molarity of NaCl in seawater at 15C (d of seawater at 15C = 1.025 g/ml).
119.
In each pair, which quantity is larger, or are they equal?
a. Entities: 0.4 mol of O3 or 0.4 mol of O atoms
b. Grams: 0.4 mol of O3 or 0.4 mol of O atoms
c. Moles: 4.0 g of N2O4 or 3.3g of SO2
d. Grams: 0.6 mol of C2H4 or 0.6 mol of F2
e. Total ions: 2.3 mol of sodium chlorate or 2.2 mol of magnesium chloride
f.
Molecules: 1.0 g of H2O or 1.0 g of H2O2
g. Na+ ions: 0.500L of 0.500M NaBr or 0.0146kg of NaCl
h. Mass: 6.02  1023 atoms of 235U or 6.02  1023 atoms of 238U
120.
Assuming that the volumes are additive, what is the concentration of KBr in a solution prepared by
mixing 0.200L of 0.053M KBr with 0.550L of 0.078M KBr?
121.
Hydrocarbon mixtures are used as fuels.
a. How many grams of CO2 (g) are produced by the combustion of 200.0 g of a mixture that is 25.0% CH4
and 75.0% C3H8 by mass?
b. A 252 g gaseous mixture of CH4 and C3H8 burns in excess 02, and 748 g of CO2 gas is collected. What is
the mass % of CH4 in the initial mixture?
122.
Nitrogen (N), phosphorus (P), and potassium (K) are the main nutrients in plant fertilizers. By
convention, the numbers on a label refer to the mass percents of N2, P2O5, and K2O, in that order. Calculate the
N/P/K ratio of a 30/10/10 fertilizer in terms of moles of each element and express it as X/Y/1.0.
123.
When carbon containing compounds are burned in a limited amount of air, some CO (g) as well as CO2
(g) is produced. A gaseous product mixture is 35.0 mass % CO and 65.0 mass % CO2. What is the mass % C in
the mixture?
124.
Hemoglobin is 6.0 % heme (C34H32FeN4O4) by mass. To remove the heme, hemoglobin is treated with
acetic acid and NaCl to form hemin (C34H32N4O4FeCl). At a crime scene, a blood sample contains 0.65 g of
hemoglobin.
a. How many grams of heme are in the sample?
b. How many moles of heme?
c. How many grams of Fe?
d. How many grams of hemin could be formed for a forensic chemist to measure?
125.
When zinc is heated with sulfur, a violent reaction occurs, and zinc sulfide forms:
Zn+S8  ZnS (unbalanced)
Some of the reactants also combine with oxygen in air to form zinc oxide and sulfur dioxide. When 83.2 g of Zn
reacts with 52.4 g of S8, 104.4 g of ZnS forms. What is the percent yield of ZnS? If all the remaining reactants
combine with oxygen, how many grams of each of the two oxides form?
The following questions are straight from past AP Chemistry exams.
1. How will the time it takes to hard-boil an egg compare at higher altitudes?
a. It will take longer because of the lower boiling temperature of the water.
b. It will take longer because of the higher vapor pressure of the water.
c. It will take less time because of the higher boiling temperature of the water.
d. It will take less time because of the lower vapor pressure of the water.
e. It will take the same amount of time, regardless of the altitude.
2. At normal atmospheric pressure and a temperature of 60 C, which phase(s) of H2O can exist?
a. Ice and water
b. Ice and water vapor
c. Water only
d. Water vapor only
e. Ice only
3. The normal boiling point for the substance represented in the phase diagram above is
a. 33 C
b. 35 C
c. 44 C
d. 50 C
e. 88 C
Questions 4-6 refer to the phase diagram of a pure substance, shown below:
a.
b.
c.
d.
e.
Freezing
Melting
Sublimation
Condensation
Vaporization
4. If the temperature increased from 40 C to 60 C at a pressure of 1.5 atm, which process is occurring?
5. If the temperature increases from 20 C to 60 C at a pressure of 0.5 atm, which process is occurring?
6. If the pressure increases from 0.5 atm to 1.0 atm at 60 C, which process is occurring?
7. Determine the molecular formula of a compound that contains 40.0% C, 6.71% H, 53.29% O and has a molecular
mass of 60.05g.
a. C2H4O2
b. CH2O
c. C2H3O4
d. C2H2O4
e. C4H8O4
8. What was the key finding in Rutherford’s famous “gold foil” experiment?
a. Electrons orbit the nucleus in concentric rings.
b. All neutrons are located in a central nucleus.
c. Most of the mass of an atom is located in a central, dense core.
d. Atoms are composed of positively and negatively charged particles.
e. Alpha particles are attracted to a negatively charged plate.
9. What is the oxidation state of Cr in the compound K2Cr2O7
a. +2
b. +3
c. +5
d. +6
e. +7
10. The correct name for the compound [Co(NH3)6]Cl3 is
a. Cobalt (III) hexamine chloride
b. Hexamminecobalt chloride (III)
c. Hexamminecobalt (III) chloride
d. Hexamminecobalt trichloride
e. Hexamminecobalt (VI) chloride
11. Which of the following represents a pair of isotopes of the same element?
Atomic Number
Mass Number
a.
I.
92
208
II.
92
211
b.
I.
45
106
II.
92
206
c.
I.
92
238
II.
89
235
d.
I.
92
233
II.
89
233
e.
I.
7
15
II.
8
16
12. What is the percentage composition of Mg in the compound Mg3(PO4)2?
a. 21.92%
b. 23.57%
c. 27.74%
d. 32.32%
e. 48.70%
13. Determine the name of [Cu(NH3)4]SO4.
a. Tetraamminecuprate (II) sulfate
b. Tetracuprateammine (IV) sulfate
c. Tetrasulfonoammine (IV) cuprate
d. Amminecuprate (IV) sulfate
e. Cupric ammonium sulfate
14. Which of the following pairs of atoms represents an isotope?
Atomic Number
Mass Number
a.
I.
8
18
II.
9
18
b. I.
8
9
II.
18
18
c.
I.
8
18
II.
18
36
d. I.
9
18
II.
9
19
e.
I.
6
12
II.
12
18
In questions 16 & 17, select the scientists primarily responsible for the findings in each question.
a.
b.
c.
d.
e.
Dalton
Einstein
Heisenberg
Thomson
Rutherford
15. Made important discoveries about the properties of cathode rays.
16. Proposed the existence of a nucleus.
17. In the compound Mn2O7, what is the oxidation number of manganese?
a. +2
b. +3
c. +5
d. +7
e. +8
18. In the molecule NaAsO3, what is the oxidation number of As?
a. 1
b. 2
c. 3
d. 4
e. 5
19. How many grams are in a 6.94 mol sample of sodium hydroxide?
a. 40.0 g
b. 278 g
c. 169 g
d. 131 g
e. 34.2 g
20. A 0.4647 g sample of a compound containing only carbon, hydrogen, and oxygen was burned in excess pure
oxygen to yield 0.8635 g of CO2 and 0.1767 g of H2O. What is the empirical formula of the compound?
a. CHO
b. C2H2O
c. C3H3O2
d. C6H3O2
e. C3H6O2
21. . . .NH3 (g) + . . .O2 (g)  . . .NO (g) + . . .H2O
If the equation above is balanced, the coefficient before the O2 will be
a. 1
b. 2
c. 3
d. 5
e. 8
22. ___As(OH)3(s) + ___H2SO4(aq)  ___As2(SO4)3(aq) + ___H2O(l)
What is the coefficient for water when the equation above is balanced?
a. 1
b. 2
c. 4
d. 6
e. 12
23. How many grams of silver nitrate (AgNO3) are required to produce 44.0 g of aluminum nitrate (Al(NO3)3)?
6 AgNO3 + Al2(SO4)3  3 Ag2SO4 + 2 Al(NO3)3
a.
b.
c.
d.
e.
105.3 g
132.0 g
169.9 g
213.0 g
264.0 g
24. Si (s) + 2 Cl2 (g)  SiCl4 (l)
If 3.84 mol Cl2 react with an excess of Si, how many moles of SiCl4 will be produced?
a. 0.96 mol
b. 1.92 mol
c. 3.84 mol
d. 4.00 mol
e. 5.73 mol
25. Ms + 2H2O  Mg(OH)2 + H2
What mass of hydrogen can be produced by reaction of 4.73 g of magnesium with 1.83 g of water?
a. 0.103 g
b. 0.0162 g
c. 0.0485 g
d. 0.219 g
e. 3.20 g
26. When 2.00 g of ethylene (C2H4) burns in oxygen to give carbon dioxide and water, how many grams of CO2 are
formed?
a. 200 g
b. 255 g
c. 314 g
d. 400 g
e. 6.28 g
27. 3Mg + N2  Mg3N2
What mass of magnesium nitride can be made from reaction of 1.22 g of magnesium with excess nitrogen?
a. 1.69 g
b. 15.2 g
c. 5.07 g
d. 5.02 g
e. 0.592 g
28. Ca(OH)2(aq) + 2 HNO3(aq)  ca(NO3)2(aq) + 2 H2O(l)
How many grams of Ca(NO3)2 can be produced by reacting excess HNO3 with 7.40 g of Ca(OH)2?
a. 10.2 g
b. 16.4 g
c. 32.8 g
d. 8.22 g
e. 7.40 g
29. 4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)
In the above reaction, 3.10 g of NH3 reacts with 2.50 g of O2. What is the theoretical yield of NO?
a. 1.88 g
b. 5.46 g
c. 8.20 g
d. 24.0 g
e. 120 g
30. A lab instructor is preparing 5.00 liters of a 0.100 M Pb(NO3)2 (molecular weight = 331) solution. She should
weigh out
a. 165.5 g of Pb(NO3)2 and add 5.00 kg of H2O.
b. 165.5 g of Pb(NO3)2 and add H2O until the solution has a volume of 5.00 liters.
c. 33.1 g of Pb(NO3)2 and add H2O until the solution has a volume of 5.00 liters.
d. 331 g of Pb(NO3)2 and add 5.00 liters of H2O.
e. 165.5 g of Pb(NO3)2 and add 5.00 liters of H2O.
31. What volume of 18.0 M sulfuric acid must be used to prepare 15.5 L of 0.195 M H2SO4?
a. 168 ml
b. 0.336 L
c. 92.3 ml
d. 226 ml
e. 125 ml
32. You are supposed to mix 250 ml of a 0.1 M solution of Pb(NO3)2 solution (molar mass = 331.2 g). You would
need to mix ____ of Pb(NO3)2 with enough water to make 250 ml solution.
a. 331.2 g
b. 33.12 g
c. 8.28 g
d. 3.312 g
e. 0.828 g
33. The volume of distilled water that should be added to 10.0 ml of 12.0 M HCl(aq) in order to prepare a 1.00 M
HCl(aq) solution is approximately
a. 50.0 ml
b. 60.0 ml
c. 100 ml
d. 110 ml
e. 120 ml
34. Answer each of the following 4 parts using appropriate chemical principles.
(a) When ice at the freezing temperature is compressed, it liquefies, while CO2, when compressed, solidifies.
(b) Rock salt mixed with ice is used as a freezing mixture in home ice cream makers. The slushy salt-ice mixture
remains fluid at temperatures well below 0 C.
(c) Helium balloons deflate more rapidly than the same size balloons filled with air.
(d) Ice cubes can appear cloudy because of dissolved gases that are present at freezing. To make clear ice cubes, it is
preferable to use hot water over cold water.