AP Chemistry
Practice Multiple Choice 1 (Calculator was allowed for this exam)
Questions 1-4
a. Heisenberg uncertainty principle
b. Pauli exclusion principle
c. Hund's rule (principle of maximum multiplicity)
d. Shielding effect
e. Wave nature of matter
1. Can be used to predict that a gaseous carbon atom in its
ground state is paramagnetic
2. Explains the experimental phenomenon of electron diffraction
3. Indicates that an atomic orbital can hold no more than two
electrons
4. Predicts that it is impossible to determine simultaneously the
exact position and the exact velocity of an electron
Questions 5-7 refer to the phase diagram below of a pure
substance.
a. Sublimation
b. Condensation
c. Solvation
d. Fusion
e. Freezing
5. If the temperature increases from 10°C to 60°C at a constant
pressure of 0.4 atmosphere, which of the processes occurs?
6. If the temperature decreases from 110°C to 40°C at a
constant pressure of 1.1 atmospheres, which of the
processes occurs?
7. If the pressure increases from 0.5 to 1.5 atmospheres at a
constant temperature of 50°C, which of the processes
occurs?
Questions 8-10 refer to the following diatomic species.
a. Li2
b. B2
c. N2
d. O2
e. F2
8. Has the largest bond-dissociation energy
9. Has a bond order of 2
10. Contains 1 sigma () and 2 pi () bonds
Questions 11-13
a. Pb
b. Ca
c. Zn
d. As
e. Na
11. Utilized as a coating to protect Fe from corrosion
12. Is added to silicon to enhance its properties as a
semiconductor
13. Utilized as a shield from sources of radiation
14. Which of the following is lower for a 1.0-molar aqueous
solution of any solute than it is for pure water?
a. pH
b. Vapor pressure
c. Freezing point
d. Electrical conductivity
e. Absorption of visible light
15. In a molecule in which the central atom exhibits sp3d2 hybrid
orbitals, the electron pairs are directed toward the corners of
a. a tetrahedron
b. a square pyramid
c. a trigonal bipyramid d. a square
e. an octahedron
16. Commercial vinegar was titrated with NaOH solution to
determine the content of acetic acid, HC2H3O2. For 20.0 mL
of the vinegar, 32.0 mL of 0.500-M NaOH solution was
required. What was the concentration of acetic acid in the
vinegar if no other acid was present?
a. 1.60 M
b. 0.800 M
c. 0.640 M
d. 0.600 M
e. 0.400 M
17. Relatively slow rates of chemical reaction are associated with
which of the following?
a. The presence of a catalyst
b. High temperature
c. High concentration of reactants
d. Strong bonds in reactant molecules
e. Low activation energy
18.
2 H2O + 4 MnO4- + 3 CIO2-  4 MnO2 + 3 CIO4- + 4 OHWhich species acts as an oxidizing agent in the reaction
represented above?
a. H2O
b. CIO4c. CIO2d. MnO2
e. MnO419. In which of the following compounds is the mass ratio of
chromium to oxygen closest to 1.6 to 1.0?
a. CrO3
b. CrO2
c. CrO
d. Cr2O
e. Cr2O3
20. _Ag+ + _AsH3 (g) + _OH-  _Ag (s) + _H3AsO3 (aq) + _H2O
When the equation above is balanced with lowest wholenumber coefficients, the coefficient for OH- is
a. 2
b. 4
c. 5
d. 6
e. 7
21. Correct statements about alpha particles include which of the
following?
I. They have a mass number of 4 and a charge of +2.
II. They are more penetrating than beta particles.
III. They are helium nuclei.
a. I only
b. III only
c. I and II
d. I and III
e. II and III
22.
HSO4- + H2O  H3O+ + SO42In the equilibrium represented above, the species that act as
bases include which of the following?
I. HSO4II. H2O
III. SO42a. II only
b. III only
c. I and II
d. I and III
e. II and III
23.
Step 1: Ce4+ + Mn2+  Ce3+ + Mn3+
Step 2: Ce4+ + Mn3+  Ce3+ + Mn4+
Step 3: Mn4+ + TI+  Tl3+ + Mn2+
The proposed steps for a catalyzed reaction between Ce 4+
and TI+ are represented above. The products of the overall
catalyzed reaction are
a. Ce4+ and TI+
b. Ce3+ and Tl3+
c. Ce3+ and Mn3+
d. Ce3+ and Mn4+
e. Tl3+ and Mn2+
24. A sample of 0.010 mole of oxygen gas is confined at 127°C
and 0.80 atm. What would be the pressure of this sample at
27°C and the same volume?
a. 0.10 atm
b. 0.20 atm
c. 0.60 atm
d. 0.80 atm
e. 1.1 atm
25. H2 (g) + ½ O2 (g)  H2O (l)
Ho = x
2 Na (s) + ½ O2 (g)  Na2O (s)
Ho = y
Na (s) + ½ O2 (g) + ½ H2 (g)  NaOH (s)
Ho = z
Based on the information above, what is the standard
enthalpy change for the following reaction?
Na2O (s) + H2O (l)  2 NaOH (s)
a. x + y + z
b. x + y – z
c. x + y - 2z
d. 2z - x - y
e. z - x - y
26. Which of the following actions would be likely to change the
boiling point of a sample of a pure liquid in an open
container?
I. Placing it in a smaller container
II. Increasing the number of moles of the liquid in the
container
III. Moving the container and liquid to a higher altitude
a. I only
b. II only
c. III only
d. II and ill only
e. I, II, and III
27. Which of the following sets of quantum numbers (n, l, ml, ms)
best describes the valence electron of highest energy in a
ground-state gallium atom (atomic number 31)?
a. 4, 0, 0, ½
b. 4, 0, I, ½
c. 4, I, I, ½
d. 4, I, 2, ½
e. 4, 2, 0, ½
28. Given that a solution is 5 % sucrose by mass, what additional
information is necessary to calculate the molarity of the
solution?
I. The density of water
ll. The density of the solution
Ill. The molar mass of sucrose
a. I only
b. II only
c. Ill only
d. I and Ill
e. II and Ill
29. When an aqueous solution of NaOH is added to an aqueous
solution of potassium dichromate, K2Cr2O7, the dichromate
ion is converted to
a. CrO42b. CrO23+
c. Cr
d. Cr2O3 (s)
e. Cr(OH)3 (S)
30. The energy diagram for the reaction X + Y  Z is shown.
The addition of a catalyst to this reaction would cause a
change in which of the indicated energy differences?
a. I only
b. II only
c. Ill only
d. I and II only
e. I, II, and III
31.
H2C2O4 + 2 H2O  2 H3O+ + C2O42Oxalic acid, H2C2O4, is a diprotic acid with K1 = 5 x 10-2 and
K2 = 5 x 10-5. Which of the following is equal to the
equilibrium constant for the reaction represented above?
a. 5 x 10-2
b. 5 x 10-5
c. 2.5 x 10-6
d. 5 x 10-7
-8
e. 2.5 x 10
32. CH3CH2OH boils at 78°C and CH3OCH3 boils at -24°C,
although both compounds have the same composition. This
difference in boiling points may be attributed to a difference in
a. molecular mass
b. density
c. specific heat
d. hydrogen bonding
e. heat of combustion
33. A hydrocarbon gas with an empirical formula CH2 has a
density of 1.88 g/L at 0°C and 1.00 atm. A possible formula
for the hydrocarbon is
a. CH2
b. C2H4
c. C3H6
d. C4H8
e. C5H10
34.
X: CH3-CH2-CH2-CH2-CH3
Y: CH3-CH2-CH2-CH2-OH
Z: HO-CH2-CH2-CH2-OH
Based on concepts of polarity and hydrogen bonding, which
of the following sequences correctly lists the compounds
above in the order of their increasing solubility in water?
a. Z < Y < X
b. Y < Z < X
c. Y < X < Z
d. X < Z < Y
e. X < Y < Z
35. For which of the following processes would S have a
negative value?
I. 2 Fe2O3 (s)  4 Fe (s) + 3 O2 (g)
II. Mg2+ + 2 OH-  Mg(OH)2 (s)
III. H2 (g) + C2H4 (g)  C2H6 (g)
a. I only
b. I and II only
c. I and III only
d. II and III only
e. I, II, and III
36.
Zn (s) + Cu2+  Zn2+ + Cu (s)
An electrolytic cell based on the reaction represented above
was constructed from zinc and copper half-cells. The
observed voltage was found to be 1.00 V instead of the
standard cell potential, Eo, of 1.10 V. Which of the following
could correctly account for this observation?
a. The copper electrode was larger than the zinc electrode.
b. The Zn2+ electrolyte was Zn(NO3)2, while the Cu2+
electrolyte was CuSO4.
c. The Zn2+ solution was more concentrated than the Cu2+
solution.
d. The solutions in the half-cells had different volumes.
e. The salt bridge contained KCl as the electrolyte.
37. A 3.0-g sample of an ideal gas at 127°C and 1.0 atm
pressure has a volume of 1.5 L. Which of the following
expressions is correct for the molar mass of the gas? R is
0.08 (L•atm)/(mole•K).
a. (0.08)(400)/(3.0)(1.0)(1.5)
b. (1.0)(1.5)/(3.0)(0.08)(400)
c. (0.08)(1.0)(1.5)/(3.0)(400)
d. (3.0)(0.08)(400)/(1.0)(1.5)
e. (3.0)(0.08)(1.5)/(1.0)(400)
38. Concentrations of colored substances are commonly
measured by means of a spectrophotometer. Which of the
following would ensure that correct values are obtained for
the measured absorbance?
I. There must be enough sample in the tube to cover
the entire light path.
II. The instrument must be periodically reset using a
standard.
III. The solution must be saturated.
a. I only
b. II only
c. I and II only
d. II and III only
e. I, II, and III
39. Samples of F2 gas and Xe gas are mixed in a container of
fixed volume. The initial partial pressure of the F2 gas is 8.0
atm and that of the Xe gas is 1.7 atm. When all of the Xe
gas reacted, forming a solid compound, the pressure of the
unreacted F2 gas was 4.6 atm. The temperature remained
constant. What is the formula of the compound?
a. XeF
b. XeF3
c. XeF4
d. XeF6
e. XeF8
40. The system shown is at equilibrium at 28°C. At this
temperature, the vapor pressure of water is 28 mm Hg.
48.
49.
41.
42.
43.
44.
45.
46.
47.
The partial pressure of O2 (g) in the system is
a. 28 mm Hg
b. 56 mm Hg
c. 133 mm Hg
d. 161 mm Hg
e. 189 mm Hg
A strip of metallic scandium, Sc, is placed in a beaker
containing concentrated nitric acid. A brown gas rapidly
forms, the scandium disappears, and the resulting liquid is
brown-yellow but becomes colorless when warmed. These
observations best support which of the following statements?
a. Nitric acid is a strong acid.
b. In solution scandium nitrate is yellow and scandium
chloride is colorless.
c. Nitric acid reacts with metals to form hydrogen.
d. Scandium reacts with nitric acid to form a brown gas.
e. Scandium and nitric acid react in mole proportions of 1
to 3.
Mass of an empty container
3.0 g
Mass of the container plus the solid
25.0 g
Volume of the solid sample
11.0 cm3
The data above were gathered in order to determine the
density of an unknown solid. The density of the sample
should be reported as
a. 0.5 g/cm3
b. 0.50 g/cm3
3
c. 2.0 g/cm
d. 2.00 g/cm3
e. 2.27 g/cm3
Which of the following pairs of compounds are isomers?
a. CH3–CH2–CH2–CH3 and CH3–CH(CH3)–CH3
b. CH3–CH(CH3)–CH3 and CH3–C(CH3)=CH2
c. CH3–O–CH3 and CH3–CO–CH3
d. CH3–OH and CH3–CH2–OH
e. CH4 and CH2=CH2
Which of the following solutions has the lowest freezing
point?
a. 0.20 m C6H12O6
b. 0.20 m NH4Br
c. 0.20 m ZnSO4
d. 0.20 m KMnO4
e. 0.20 m MgCl2
A sample of an ideal gas is cooled from 50.0oC to 25.0oC in a
sealed container of constant volume. Which of the following
values for the gas will decrease?
I. The average molecular mass of the gas
II. The average distance between the molecules
III. The average speed of the molecules
a. I only
b. II only
c. III only
d. I and III
e. II and III
Which of the following solids dissolves in water to form a
colorless solution?
a. CrCl3
b. FeCl3
c. CoCl2
d. CuCl2
e. ZnCl2
Which of the following has the lowest conductivity?
a. 0.1 M CuSO4
b. 0.1 M KOH
c. 0.1 M BaCl2
d. 0.1 M HF
e. 0.1 M HNO3
50.
51.
52.
53.
54.
55.
PCl3 (g) + Cl2 (g)  PCI5 (g) + energy
Some PCl3 and Cl2 are mixed in a container at 200°C and the
system reaches equilibrium according to the equation above.
Which of the following causes an increase in the number of
moles of PCl5 present at equilibrium?
I. Decreasing the volume of the container
II. Raising the temperature
III. Adding a mole of He gas at constant volume
a. I only
b. II only
c. I and III only
d. II and III only
e. I, II, and III
The isomerization of cyclopropane to propylene is a firstorder process with a half-life of 19 minutes at 500oC. The
time it takes for the partial pressure of cyclopropane to
decrease from 1.0 atm to 0.125 atm at 500oC is closest to
a. 38 minutes
b. 57 minutes
c. 76 minutes
d. 152 minutes
e. 190 minutes
Which of the following acids can be oxidized to form a
stronger acid?
a. H3PO4
b. HNO3
c. H2CO3
d. H3BO3
e. H2SO3
4 HCI (g) + O2 (g)  2 Cl2 (g) + 2 H2O (g)
Equal numbers of moles of HCI and O2 in a closed system
are allowed to reach equilibrium as represented by the
equation above. Which of the following must be true at
equilibrium?
I. [HCI] must be less than [Cl2].
II. [O2] must be greater than [HCI].
IlI. [Cl2] must equal [H2O].
a. I only
b. II only
c. I and III only
d. II and III only
e. I, II, and III
When dilute nitric acid was added to a solution of one of the
following chemicals, a gas was evolved. This gas turned a
drop of limewater, Ca(OH)2, cloudy (a white precipitate). The
chemical was
a. household ammonia, NH3
b. baking soda, NaHCO3
c. table salt, NaCI
d. epsom salts, MgSO4 • 7 H2O
e. bleach, 5% NaOCI
If 87 g of K2SO4 (MM 174 g) is dissolved in enough water to
make 250 mL of solution, what are the concentrations of the
potassium and the sulfate ions?
[K+]
[SO42-]
a. 0.020 M
0.020 M
b. 1.0 M
2.0 M
c. 2.0 M
1.0 M
d. 2.0 M
2.0 M
e. 4.0 M
2.0 M
All of the following statements concerning the characteristics
of the halogens are true EXCEPT:
a. The first ionization energies (potentials) decrease as the
atomic numbers of the halogens increase.
b. Fluorine is the best oxidizing agent.
c. Fluorine atoms have the smallest radii.
d. Iodine liberates free bromine from a solution of bromide
ion.
e. Fluorine is the most electronegative of the halogens.
What volume of 0.150-M HCI is required to neutralize 25.0
mL of 0.120-M Ba(OH)2?
a. 20.0 mL
b. 30.0 mL
c. 40.0 mL
d. 60.0 mL
e. 80.0 mL
56. It is suggested that SO2 (MM 64 g), which contributes to acid
rain, could be removed from a stream of waste gases by
bubbling the gases through 0.25-M KOH, thereby producing
K2SO3. What is the maximum mass of SO2 that could be
removed by 1,000. L of the KOH solution?
a. 4.0 kg
b. 8.0 kg
c. 16 kg
d. 20. kg
e. 40. kg
57. Molecules that have planar configurations include which of
the following?
I. BCl3
ll. CHCl3
Ill. NCl3
a. I only
b. Ill only
c. I and II only
d. II and Ill only
e. I, ll, and Ill
58.
N2 (g) + 3 H2 (g)  2 NH3 (g)
The reaction indicated above is thermodynamically
spontaneous at 298 K, but becomes nonspontaneous at
higher temperatures. Which of the following is true at 298 K?
a. G, H, and S are all positive.
b. G, H, and S are all negative.
c. G and H are negative, but S is positive.
d. G and S are negative, but H is positive.
e. G and H are positive, but S is negative.
59. When a 1.25-g sample of limestone was dissolved in acid,
0.44 g of CO2 was generated. If the rock contained no
carbonate other than CaCO3, what was the percent of
CaCO3 by mass in the limestone?
a. 35%
b. 44%
c. 67%
d. 80%
e. 100%
60.
I2 (g) + 3 Cl2 (g)  2 ICl3 (g)
According to the data in the table below, what is the value of
Ho for the reaction represented above?
Bond
Average Bond Energy (kilojoules/mole)
I–I
150
CI–CI
240
I–Cl
210
a. - 870 kJ
b. - 390 kJ
c. +180 kJ
d. + 450 kJ
e. +1,260kJ
61. A 1-M solution of which of the following salts has the highest
pH?
a. NaNO3
b. Na2CO3
c. NaCI
d. NaHSO4
e. Na2SO4
62. The electron-dot structure (Lewis structure) for which of the
following molecules would have two unshared pairs of
electrons on the central atom?
a. H2S
b. NH3
c. CH4
d. HCN
e. CO2
63. Which of the following expressions is correct for the
maximum mass of copper, in grams, that could be plated out
by electrolyzing aqueous CuCl2 for 16 hours at a constant
current of 3.0 A?
(1 F = 96,500 C)
a. (16)(3,600)(3.0)(63.55)(2)/(96,500)
b. (16)(3,600)(3.0)(63.55)/(96,500)(2)
c. (16)(3,600)(3.0)(63.55)/(96,500)
d. (16)(60)(3.0)(96,500)(2)/(63.55)
e. (16)(60)(3.0)(96,500)/(63.55)(2)
64. At 25°C, a sample of NH3 (MM 17 g) effuses at the rate of
0.050 mole/minute. Under the same conditions, which of the
following gases effuses at approximately one-half that rate?
a. O2 (MM 32 g)
b. He (MM 4.0 g)
c. CO2 (MM 44 g)
d. Cl2 (MM 71 g)
e. CH4 (MM 16 g)
65. Barium sulfate is LEAST soluble in a 0.01-M solution of
which of the following?
a. Al2(SO4)3
b. (NH4)2SO4
c. Na2SO4
d. NH3
e. BaCl2
66. What is the pH of a 1.0 x l0-2-M solution of HCN ? (For HCN,
Ka = 4.0 x 10-10)
a. 10
b. Between 7 and 10
c. 7
d. Between 4 and 7
e. 4
67. Substances X and Y that were in a solution were separated
in the laboratory using the technique of fractional
crystallization. This fractional crystallization is possible
because substances X and Y have different
a. boiling points
b. melting points
c. densities
d. crystal colors
e. solubilities
68. Which of the following molecules has a dipole moment of
zero?
a. C6H6 (benzene)
b. NO
c. SO2
d. NH3
e. H2S
69. Correct procedures for a titration include which of the
following?
I. Draining a pipet by touching the tip to the side of the
container used for the titration
II. Rinsing the buret with distilled water just before
filling it with the liquid to be titrated
III. Swirling the solution frequently during the titration
a. I only
b. II only
c. I and III only
d. II and III only
e. I, II, and III
70. To determine the molar mass of a solid monoprotic acid, a
student titrated a weighed sample of the acid with
standardized aqueous NaOH. Which of the following could
explain why the student obtained a molar mass that was too
large?
I. Failure to rinse all acid from the weighing paper into
the titration vessel
II. Addition of more water than was needed to dissolve
the acid
III. Addition of some base beyond the equivalence point
a. I only
b. III only
c. I and II only
d. II and III only
e. I, II, and III
71.
_Fe(OH)2 + _O2 + _H2O  _Fe(OH)3
If 1 mole of O2 oxidizes Fe(OH)2 according to the reaction
represented above, how many moles of Fe(OH)3 can be
formed?
a. 2
b. 3
c. 4
d. 5
e. 6
72. The nuclide 24996Cm is radioactive and decays by the loss of
one beta (-) particle. The product nuclide is
a. 24594Pu
b. 24595Am
c. 24896Cm
d. 25096Cm
e. 24997Bk
73.
2 SO2 (g) + O2 (g)  2 SO3 (g)
When 0.40 mole of SO2 and 0.60 mole of O2 are placed in an
evacuated 1.00-L flask, the reaction represented above
occurs. After the reactants and the product reach equilibrium
and the initial temperature is restored, the flask is found to
contain 0.30 mole of SO3. Based on these results, the
expression for the equilibrium constant, Kc, of the reaction is
a. (0.30)2/(0.45)(0.10)2
b. (0.30)2/(0.60)(0.40)2
c. (2 x 0.30)/(0.45)(2 x 0.10)
d. (0.30)/(0.45)(0.10)
e. (0.30)/(0.60)(0.40)
74. A solution of calcium hypochlorite, a common additive to
swimming-pool water, is
a. basic because of the hydrolysis of the OCI- ion
b. basic because Ca(OH)2 is a weak and insoluble base
c. neutral if the concentration is kept below 0.1 molar
d. acidic because of the hydrolysis of the Ca2+ ions
e. acidic because the acid HOCI is formed
75. A direct-current power supply of low voltage (less than 10 V)
has lost the markings that indicate which output terminal is
positive and which is negative. A chemist suggests that the
power supply terminals be connected to a pair of platinum
electrodes that dip into 0.1-M KI solution. Which of the
following correctly identifies the polarities of the power supply
terminals?
a. A gas will be evolved only at the positive electrode.
b. A gas will be evolved only at the negative electrode.
c. A brown color will appear in the solution near the
negative electrode.
d. A metal will be deposited on the positive electrode.
e. None of the methods above will identify the polarities of
the power supply terminals.
Practice Multiple Choice 2 (Calculator was not allowed for this exam)
Questions 1-4 refer to the following types of energy.
Questions 17-18 refer to the following elements.
a. Activation energy
a. Lithium
b. Nickel
b. Free energy
c. Bromine
d. Uranium
c. Ionization energy
e. Fluorine
d. Kinetic energy
17. Is a gas in its standard state at 298 K
e. Lattice energy
18. Reacts with water to form a strong base
1. The energy required to convert a ground-state atom in the
19. Which of the following best describes the role of the spark
gas phase to a gaseous positive ion
from the spark plug in an automobile engine?
2. The energy change that occurs in the conversion of an ionic
a. The spark decreases the energy of activation for the
solid to widely separated gaseous ions
slow step.
3. The energy in a chemical or physical change that is available
b. The spark increases the concentration of the volatile
to do useful work
reactant.
4. The energy required to form the transition state in a chemical
c. The spark supplies some of the energy of activation for
reaction
the combustion reaction.
Questions 5-8 refer to atoms for which the occupied atomic
d. The spark provides a more favorable activated complex
orbitals are shown below.
for the combustion reaction.
e. The spark provides the heat of vaporization for the
a. 1s 2s_
volatile hydrocarbon.
b. 1s2s
20. What mass of Au is produced when 0.0500 mol of Au2S3 is
c. 1s2s2p_ _ 
reduced completely with excess H2?
d. 1s2s2p  
a. 9.85 g
b. 19.7 g
e. [Ar] 4s3d _ _ _
c. 24.5 g
d. 39.4 g
5. Represents an atom that is chemically unreactive
e.
48.9
g
6. Represents an atom in an excited state
21. When a solution of sodium chloride is vaporized in a flame,
7. Represents an atom that has four valence electrons
the color of the flame is
8. Represents an atom of a transition metal
a. blue
b. yellow
Questions 9-12 refer to aqueous solutions containing 1:1 mole
c. green
d. violet
ratios of the following pairs of substances. Assume all
e.
white
concentrations are 1 M.
22. Of the following reactions, which involves the largest
a. NH3 and NH4CI
decrease in entropy?
b. H3PO4 and NaH2PO4
a. CaCO3 (s)  CaO (s) + CO2 (g)
c. HCI and NaCI
b. 2 CO (g) + O2 (g)  2 CO2 (g)
d. NaOH and NH3
c. Pb(NO3)2 (s) + 2 KI (s)  PbI2 (s) + 2 KNO3 (s)
e. NH3 and HC2H302 (acetic acid)
d. C3H8 (g) + 5 5O2 (g)  3 CO2 (g) + 4 H2O (g)
9. The solution with the lowest pH
e. 4 La (s) + 3 O2 (g)  2 La2O3 (s)
10. The most nearly neutral solution
11. A buffer at a pH > 8
23. A hot-air balloon rises. Which of the following is the best
12. A buffer at a pH < 6
explanation for this observation?
Questions 13-16 refer to the following descriptions of bonding in
a. The pressure on the walls of the balloon increases with
different types of solids.
increasing temperature.
a. Lattice of positive and negative ions held together by
b. The difference in temperature between the air inside and
electrostatic forces
outside the balloon produces convection currents.
b. Closely packed lattice with delocalized electrons
c. The cooler air outside the balloon pushes in on the walls
throughout
of the balloon.
c. Strong single covalent bonds with weak intermolecular
d. The rate of diffusion of cooler air is less than that of
forces
warmer air.
e. The air density inside the balloon is less than that of the
d. Strong multiple covalent bonds (including 1 bonds) with
surrounding air.
weak intermolecular forces
e. Macromolecules held together with strong polar bonds
13. Cesium chloride, CsCI (s)
14. Gold, Au (s)
15. Carbon dioxide, CO2 (s)
16. Methane, CH4 (s)
24. The safest and most effective emergency procedure to treat
an acid splash on skin is to do which of the following
immediately?
a. Dry the affected area with paper towels
b. Sprinkle the affected area with powdered Na2SO4 (s)
c. Flush the affected area with water and then with a dilute
NaOH solution
d. Flush the affected area with water and then with a dilute
NaHCO3 solution
e. Flush the affected area with water and then with a dilute
vinegar solution
25. The cooling curve for a pure substance as it changes from a
liquid to a solid is shown.
26.
27.
28.
29.
30.
The solid and the liquid coexist at
a. point Q only
b. point R only
c. all points on the curve between Q and S
d. all points on the curve between R and T
e. no point on the curve
_ C10H12O4S (s) + _ O2 (g)  _ CO2 (g) + _ SO2 (g) + _ H2O (g)
When the equation above is balanced and all coefficients are
reduced to their lowest whole-number terms, the coefficient
for O2 (g) is
a. 6
b. 7
c. 12
d. 14
e. 28
Appropriate uses of a visible-light spectrophotometer include
which of the following?
I. Determining the concentration of a solution of
Cu(NO3)2
II. Measuring the conductivity of a solution of KMnO4
III. Determining which ions are present in a solution
that may contain Na+, Mg2+, Al3+
a. I only
b. II only
c. III only
d. I and II only
e. I and III only
The melting point of MgO is higher than that of NaF.
Explanations for this observation include which of the
following?
I. Mg2+ is more positively charged than Na+.
II. O2- is more negatively charged than F-.
III. The O2- ion is smaller than the F- ion.
a. II only
b. I and II only
c. I and III only
d. II and III only
e. I, II, and III
O
II
CH3–C–CH2–CH3
The organic compound represented above is an example of
a. an organic acid
b. an alcohol
c. an ether
d. an aldehyde
e. a ketone
H2Se (g) + 4 O2F2 (g)  SeF6 (g) + 2 HF (g) + 4 O2 (g)
Which of the following is true regarding the reaction
represented above?
a. Oxidation number of O does not change.
b. Oxidation number of H changes from -1 to + 1.
c. Oxidation number of F changes from + 1 to -1.
d. Oxidation number of Se changes from -2 to +6.
e. It is a disproportionation reaction for F.
31. If the temperature of an aqueous solution of NaCI is
increased from 20°C to 90°C, which of the following
statements is true?
a. The density of the solution remains unchanged.
b. The molarity of the solution remains unchanged.
c. The molality of the solution remains unchanged.
d. The mole fraction of solute decreases.
e. The mole fraction of solute increases.
32. Types of hybridization exhibited by the C atoms in propene,
CH3CHCH2 include which of the following?
I. sp
II. sp2
III. sp3
a. I only
b. III only
c. I and II only
d. II and III only
e. I, II, and III
33. A 1.0 L sample of an aqueous solution contains 0.10 mol of
NaCI and 0.10 mol of CaCl2. What is the minimum number
of moles of AgNO3 that must be added to the solution in
order to precipitate all of the CI- as AgCI (s)? (Assume that
AgCI is insoluble.)
a. 0.10 mol
b. 0.20 mol
c. 0.30 mol
d. 0.40 mol
e. 0.60 mol
Questions 34-35 refer to an electrolytic cell that involves the
following half-reaction.
AIF63- + 3 e-  Al + 6 F34. Which of the following occurs in the reaction?
a. AIF63- is reduced at the cathode.
b. Al is oxidized at the anode.
c. Aluminum is converted from the -3 oxidation state to the
0 oxidation state.
d. F- acts as a reducing agent.
e. F- is reduced at the cathode.
35. A steady current of 10 amperes is passed through an
aluminum-production cell for 15 minutes. Which of the
following is the correct expression for calculating the number
of grams of aluminum produced? (1 faraday = 96,500
coulombs)
a. (10)(15)(96,500)/(27)(60) g
b. (10)(15)(27)/(60)(96,500) g
c. (10)(15)(60)(27)/(96,500)(3) g
d. (96,500)(27)/(10)(15)(60)(3) g
e. (27)(3)/(96,500)(10)(15)(60) g
36. The initial-rate data in the table were obtained for the
reaction represented below.
Initial Rate of
Initial
Initial [O2] Formation of
Exp.
[NO]
(mol L-1)
NO2
(mol L-1)
(mol L-1 s-1)
1
0.10
0.10
2.5 x 10-4
2
0.20
0.10
5.0 x 10-4
3
0.20
0.40
8.0 x 10-3
What is the experimental rate law for the reaction?
2 NO (g) + O2 (g)  NO2 (g)
a. Rate = k[NO][O2]
b. Rate = k[NO][O2]2
c. Rate = k[NO]2[O2]
d. Rate = k[NO]2[O2]2
e. Rate = k[NO][O2]-1
37. The ionization energies for element X are listed in the table.
Ionization Energies for element X (kJ mol-1)
First
Second
Third
Fourth
Fifth
580
1,815
2,740
11,600
14,800
On the basis of the data, element X is most likely to be
a. Na
b. Mg
c. Al
d. Si
e. P
38. A molecule or an ion is classified as a Lewis acid if it
a. accepts a proton from water
b. accepts a pair of electrons to form a bond
c. donates a pair of electrons to form a bond
d. donates a proton to water
e. has resonance Lewis electron-dot structures
39. The phase diagram for a pure substance is shown.
40.
41.
42.
43.
44.
45.
Which point on the diagram corresponds to the equilibrium
between the solid and liquid phases at the normal melting
point?
a. A
b. B
c. C
d. D
e. E
Of the following molecules, which has the largest dipole
moment?
a. CO
b. CO2
c. O2
d. HF
e. F2
2 SO3 (g)  2 SO2 (g) + O2 (g)
After the equilibrium represented above is established, some
pure O2 (g) is injected into the reaction vessel at constant
temperature. After equilibrium is reestablished, which of the
following has a lower value compared to its value at the
original equilibrium?
a. Keq for the reaction
b. The total pressure in the reaction vessel
c. The amount of SO3 (g) in the reaction vessel
d. The amount of O2 (g) in the reaction vessel
e. The amount of SO2 (g) in the reaction vessel
_Li3N (s) + _H2O (I)  _Li+ (aq) + _OH- (aq) + _NH3 (g)
When the equation above is balanced and all coefficients
reduced to lowest terms, the coefficient for OH- (aq) is
a. 1
b. 2
c. 3
d. 4
e. 6
A sample of 61.8 g of H3BO3, a weak acid, is dissolved in
1,000 g of water to make a 1.0-molal solution. Which of the
following would be the best procedure to determine the
molarity of the solution? (Assume no additional information is
available.)
a. Titration of the solution with standard acid
b. Measurement of the pH with a pH meter
c. Determination of the boiling point of the solution
d. Measurement of the total volume of the solution
e. Measurement of the specific heat of the solution
A rigid metal tank contains oxygen gas. Which of the
following applies to the gas in the tank when additional
oxygen is added at constant temperature?
a. the volume of the gas increases.
b. The pressure of the gas decreases.
c. The average speed of the gas molecules remains the
same.
d. The total number of gas molecules remains the same.
e. The average distance between the gas molecules
increases.
What is the H+ (aq) concentration in 0.05 M HCN (aq)?
(The Ka for HCN is 5.0 x 10-10.)
a. 2.5 x 10-11 M
b. 2.5 x 10-10 M
c. 5.0 X 10-10 M
d. 5.0 x 10-6 M
-4
e. 5.0 x 10 M
46. Which of the following occurs when excess concentrated NH3
(aq) is mixed thoroughly with 0.1 M Cu(NO3)2 (aq)?
a. A dark red precipitate forms and settles out.
b. Separate layers of immiscible liquids form with a blue
layer on top.
c. The color of the solution turns from light blue to dark
blue.
d. Bubbles of ammonia gas form.
e. The pH of the solution decreases.
47. When hafnium metal is heated in an atmosphere of chlorine
gas, the product of the reaction is found to contain 62.2
percent Hf by mass and 37.4 percent Cl by mass. What is
the empirical formula for this compound?
a. HfCI
b. HfCl2
c. HfCl3
d. HfCl4
e. Hf2Cl3
48. If 87.5 percent of a sample of pure 131I decays in 24 days,
what is the half-life of 131I?
a. 6 days
b. 8 days
c. 12 days
d. 14 days
e. 21 days
49. Which of the following techniques is most appropriate for the
recovery of solid KNO3 from an aqueous solution of KNO3?
a. Paper chromatography
b. Filtration
c. Titration
d. Electrolysis
e. Evaporation to dryness
50. In the periodic table, as the atomic number increases from 11
to 17, what happens to the atomic radius?
a. It remains constant.
b. It increases only.
c. It increases, then decreases.
d. It decreases only.
e. It decreases, then increases.
51. Which of the following is a correct interpretation of the results
of Rutherford's experiments in which gold atoms were
bombarded with alpha particles?
a. Atoms have equal numbers of positive and negative
charges.
b. Electrons in atoms are arranged in shells.
c. Neutrons are at the center of an atom.
d. Neutrons and protons in atoms have nearly equal mass.
e. The positive charge of an atom is concentrated in a
small region.
52. Under which of the following sets of conditions could the
most O2 (g) be dissolved in H2O (I)?
Pressure of O2 (g)
Temperature
Above H2O (I)
of H2O (I)
(atm)
(oC)
a.
5.0
80
b.
5.0
20
c.
1.0
80
d.
1.0
20
e.
0.5
20
53.
W (g) + X (g)  Y (g) + Z (g)
Gases W and X react in a closed, rigid vessel to form gases
Y and Z according to the equation above. The initial
pressure of W is 1.20 atm and that of X is 1.60 atm. No Y or
Z is initially present. The experiment is carried out at
constant temperature. What is the partial pressure of Z when
the partial pressure of W has decreased to 1.0 atm?
a. 0.20 atm
b. 0.40 atm
c. 1.0 atm
d. 1.2 atm
e. 1.4 atm
54.
55.
56.
57.
58.
59.
60.
61.
2 NO (g) + O2 (g)  2 NO2 (g)
H < 0
Which of the following changes alone would cause a
decrease in the value of Keq for the reaction represented
above?
a. Decreasing the temperature
b. Increasing the temperature
c. Decreasing the volume of the reaction vessel
d. Increasing the volume of the reaction vessel
e. Adding a catalyst
10 HI + 2 KMnO4 + 3 H2SO4  5 I2 + 2 MnSO4 + K2SO4 + 8 H2O
According to the balanced equation above, how many moles
of HI would be necessary to produce 2.5 mol of I2, starting
with 4.0 mol of KMnO4 and 3.0 mol of H2SO4?
a. 20.
b. 10.
c. 8.0
d. 5.0
e. 2.5
A yellow precipitate forms when 1 M Nal (aq) is added to a 1 M
solution of which following ion?
a. Pb2+ (aq)
b. Zn2+ (aq)
c. CrO42- (aq)
d. SO42- (aq)
e. OH (aq)
M (s) + 3 Ag+ (aq)  3 Ag (s) + M3+ (aq)
Eo = + 2.46 V
Ag+ (aq) + e-  Ag (s)
Eo = + 0.80 V
According to the information above, what is the standard
reduction potential for the half-reaction
M3+ (aq) + 3 e-  M (s)?
a. –1.66 V
b. –0.06 V
c. 0.06 V
d. 1.66 V
e. 3.26 V
On a mountaintop, it is observed that water boils at 90°C, not
at 100°C as at sea level. This occurs because on the
mountaintop the
a. equilibrium water vapor pressure is higher due to the
higher atmospheric pressure
b. equilibrium water vapor pressure is lower due to the
higher atmospheric pressure
c. equilibrium water vapor pressure equals the atmospheric
pressure at a lower temperature
d. water molecules have a higher average kinetic energy
due to the lower atmospheric pressure
e. water contains a greater concentration of dissolved
gases
A 40.0 mL sample of 0.25 M KOH is added to 60.0 mL of
0.15 M Ba(OH)2. What is the molar concentration of OH- (aq)
in the resulting solution? (Assume that the volumes are
additive.)
a. 0.10 M
b. 0.19 M
c. 0.28 M
d. 0.40 M
e. 0.55 M
NH4NO3 (s)  N2O (g) + 2 H2O (g)
A 0.03 mol sample of NH4NO3 (s) is placed in 1 L evacuated
flask, which is then sealed and heated. The NH4NO3 (s)
decomposes completely according to the balanced equation
above. The total pressure in the flask measured at 400 K is
closest to which of the following? (R, is 0.082 L atm mol-1 K-1.)
a. 3 atm
b. 1 atm
c. 0.5 atm
d. 0.1 atm
e. 0.03 atm
C2H4 (g) + 3 O2 (g)  2 CO2 (g) + 2 H2O (g)
For the reaction of ethylene represented above, H is -1,323
kJ. What is the value of H if the combustion produced liquid
water H2O (I), rather than water vapor H2O (g)?
(H for the phase change H2O (g)  H2O (l) is -44 kJ mol-1.)
a. -1,235 kJ
b. -1,279 kJ
c. -1,323 kJ
d. -1,367 kJ
e. -1,411 kJ
HC2H3O2 (aq) + CN- (aq)  HCN (aq) + C2H3O2- (aq)
The reaction represented above has an equilibrium constant
equal to 3.7 x 104. Which of the following can be concluded
from this information?
a. CN- (aq) is a stronger base than C2H3O2- (aq).
b. HCN (aq) is a stronger acid than HC2H3O2 (aq).
c. The conjugate base of CN- (aq) is C2H3O2- (aq).
d. The equilibrium constant will increase with an increase in
temperature.
e. The pH of a solution containing equimolar amounts of
CN- (aq) and HC2H3O2 (aq) is 7.0.
63. The graph shows the results of a study of the reaction of X
with a large excess of Y to yield Z. The concentrations of X
and Y were measured over a period of time.
62.
64.
65.
66.
67.
According to the results, which of the following can be
concluded about the rate law for the reaction under the
conditions studied?
a. It is zero order in [X].
b. It is first order in [X].
c. It is second order in [X].
d. It is first order in [Y].
e. The overall order of the reaction is 2.
Equal numbers of moles of He (g), Ar (g), and Ne (g) are placed
in a glass vessel at room temperature. If the vessel has a
pinhole-sized leak, which of the following will be true
regarding the relative values of the partial pressures of the
gases remaining in the vessel after some of the gas mixture
has effused?
a. PHe < PNe < PAr
b. PHe < PAr < PNe
c. PNe < PAr < PHe
d. PAr < PHe < PNe
e. PHe = PAr = PNe
Which of the following compounds is NOT appreciably
soluble in water but is soluble in dilute hydrochloric acid?
a. Mg(OH)2 (s)
b. (NH4)2CO3 (s)
c. CuSO4 (s)
d. (NH4)2SO4 (s)
e. Sr(NO3)2 (s)
When solid ammonium chloride, NH4Cl (s), is added to water
at 25oC, it dissolves and the temperature of the solution
decreases. Which of the following is true for the values of H
and S for the dissolving process?
H
S
a.
Positive
Positive
b.
Positive
Negative
c.
Positive
Equal to zero
d.
Negative
Positive
e.
Negative
Negative
What is the molar solubility in water of Ag2CrO4?
(The Ksp for Ag2CrO4 is 8 x 10-12.)
a. 8 x 10-12 M
b. 2 x 10-12 M
c. (4 x 10-12)½ M
d. (4 x 10-12)⅓ M
e. (2 x 10-12)⅓ M
68. In which of the following processes are covalent bonds
broken?
a. I2 (s)  I2 (g)
b. CO2 (s)  CO2 (g)
c. NaCl (s)  NaCl (l)
d. C(diamond)  C (g)
e. Fe (s)  Fe (l)
69. What is the final concentration of barium ions, [Ba 2+], in
solution when 100. mL of 0.10 M BaCI2 (aq) is mixed with 100.
mL of 0.050 M H2SO4 (aq)?
a. 0.00 M
b. 0.012 M
c. 0.025 M
d. 0.075 M
e. 0.10 M
70. When 100 mL of 1.0 M Na3PO4 is mixed with 100 mL of 1.0
M AgNO3, a yellow precipitate forms and [Ag+] becomes
negligibly small. Which of the following is a correct listing of
the ions remaining in solution in order of increasing
concentration?
a. [PO43-] < [NO3-] < [Na+]
b. [PO43-] < [Na+] < [NO3-]
c. [NO3-] < [PO43-] < [Na+]
d. [Na+] < [NO3-] < [PO43-]
e. [Na+] < [PO43-] < [NO3-]
71. In a qualitative analysis for the presence of Pb 2+, Fe2+, and
Cu2+ ions in aqueous solution, which of the following will
allow the separation of Pb2+ from the other ions at room
temperature?
a. Adding dilute Na2S (aq) solution
b. Adding dilute HCI (aq) solution
c. Adding dilute NaOH (aq) solution
d. Adding dilute NH3 (aq) solution
e. Adding dilute HNO3 (aq) solution
72. After completing an experiment to determine gravimetrically
the percentage of water in a hydrate, a student reported a
value of 38 percent. The correct value for the percentage of
water in the hydrate is 51 percent. Which of the following is
the most likely explanation for this difference?
a. Strong initial heating caused some of the hydrate sample
to spatter out of the crucible.
b. The dehydrated sample absorbed moisture after heating.
c. The amount of the hydrate sample used was too small.
d. The crucible was not heated to constant mass before
use.
e. Excess heating caused the dehydrated sample to
decompose.
73. The volume of distilled water that should be added to 10.0
mL of 6.00 M HCI (aq) in order to prepare a 0.500 M HCI (aq)
solution is approximately
a. 50.0 mL
b. 60.0 mL
c. 100. mL
d. 110. mL
e. 120. mL
74. Which of the following gases deviates most from ideal
behavior?
a. SO2
b. Ne
c. CH4
d. N2
e. H2
75. Which of the following pairs of liquids forms the solution that
is most ideal (most closely follows Raoult's law)?
a. C8H18 (l) and H2O (l)
b. CH3CH2CH2OH (I) and H2O (I)
c. CH3CH2CH2OH (I) and C8H18 (I)
d. C6H14 (l) and C8H18 (l)
e. H2SO4 (l) and H2O (I)
Practice Free Response 1
1.
2.
3.
HF(aq) + H2O(l)  H3O+(aq) + F-(aq)
Hydrofluoric acid, HF(aq), dissociates in water as represented by the equation above.
a. Write the equilibrium-constant expression for the dissociation of HF(aq) in water.
b. Calculate the molar concentration of H3O+ in a 0.40 M HF(aq) solution.
HF(aq) reacts with NaOH(aq) according to the reaction represented below.
HF(aq) + OH-(aq)  H2O(l) + F-(aq)
A volume of 15 mL of 0.40 M NaOH(aq) is added to 25 mL of 0.40 M HF(aq) solution. Assume the volumes are additive.
c. Calculate the number of moles of HF(aq) remaining in the solution.
d. Calculate the molar concentration of F-(aq) in the solution.
e. Calculate the pH of the solution
N2(g) + 3 F2(g)  2 NF3(g)
Ho298 = -264 kJ mol-1; So298 = -278 J K-1 mol-1
The following questions relate to the synthesis reaction represented by the chemical equation in the box above.
a. Calculate the value of the standard free energy change, Go298, for the reaction.
b. Determine the temperature at which the equilibrium constant, Keq, for the reaction is equal to 1.00. (Assume that Ho and So
are independent of temperature.)
c. Calculate the standard enthalpy change, Ho, that occurs when a 0.256 mol sample of NF3(g) is formed from N2(g) and F2(g)
at 1.00 atm and 298 K.
d. How many bonds are formed when two molecules of NF3 are produced according to the equation the box above?
e. Use both the information in the box above and the table of average bond enthalpies below to calculate the average enthalpy of
the F – F bond.
Bond
NN
N–F
F–F
Average Bond Enthalpy (kJ mol-1)
946
272
?
An external direct-current power supply is connected to two platinum electrodes immersed in a beaker containing 1.0 M CuSO4(aq)
at 25oC, as shown in the diagram above. As the cell operates, copper metal is deposited onto one electrode and O 2(g) is produced
at the other electrode. The two reduction half-reactions for the overall reaction that occurs in the cell are shown in the table below
Eo (V)
Half-Reaction
+
O2(g) + 4 H (aq) + 4 e  2 H2O(l)
+1.23
Cu2+(aq) + 2 e-  Cu(s)
+0.34
a. On the diagram, indicate the direction of electron flow in the wire.
b. Write a balanced net ionic equation for the electrolysis reaction that occurs in the cell.
c. Predict the algebraic sign of Go for the reaction. Justify your prediction.
d. Calculate the value of Go for the reaction.
An electric current of 1.50 amps passes through the cell for 40.0 minutes.
e. Calculate the mass, in grams, of the Cu(s) that is deposited on the electrode.
f. Calculate the dry volume, in liters, measured at 25oC and 1.16 atm, of the O2(g) that is produced.
4.
5.
6.
For each of the following three reactions, in part (i) write a balanced equation for the reaction and in part (ii) answer the question
about the reaction. In part (i), coefficients should be in terms of the lowest whole numbers. Assume that solutions are aqueous
unless otherwise indicated. Represent substances in solutions as ions if the substances are extensively ionized. Omit formulas for
any ions or molecules that re unchanged by the reaction.
a. A solution of sodium hydroxide is added to a solution of lead(II) nitrate.
i. Balanced equation:
ii. If 1.0 L volumes of 1.0 M solutions of sodium hydroxide and Lead(II) nitrate are mixed together, how many moles of
product(s) will be produced? Assume the reaction goes to completion.
b. Excess nitric acid is added to solid calcium carbonate.
i. Balanced equation:
ii. Briefly explain why statues made of marble (calcium carbonate) displayed outdoors in urban areas are deteriorating.
c. A solution containing silver(I) ion (an oxidizing agent) is mixed with a solution containing iron(II) ion (a reducing agent).
i. Balanced equation:
ii. If the contents of the reaction mixture described above are filtered, what substance(s), if any, would remain on the filter
paper?
5 Fe2+(aq) + MnO4-(aq) + 8 H+(aq)  5 Fe3+(aq) + Mn2+(aq) + 4 H2O(l)
The mass percent of iron in a soluble iron(II) compound is measured using a titration based on the balanced equation above.
a. What is the oxidation number of manganese in the permanganate ion, MnO 4-(aq)?
b. Identify the reducing agent in the reaction represented above.
The mass of a sample of iron(II) compound is carefully measured before the sample is dissolved in distilled water. The resulting
solution is acidified with H2SO4(aq). The solution is then titrated with MnO4-(aq) until the end point is reached.
c. Describe the color change that occurs in the flask when the end point of the titration has been reached. Explain why the color
of the solution changes at the end point.
d. Let the variables g, M, and V be defined as follows:
g = the mass, in grams, of the sample of the iron(II) compound
M = the molarity of the MnO4-(aq) used as the titrant
V = the volume, in liters, of MnO4-(aq) added to reach the end point
In terms of these variables, the number of moles of MnO4-(aq) added to reach the end point of the titration is expressed as M x
V. Using the variables defined above, the molar mass of iron (55.85 g mol-1), and the coefficients in the balanced chemical
equation, write the expression for each of the following quantities.
(i) The number of moles of iron in the sample
(ii) The mass of iron in the sample, in grams
(iii) The mass percent of iron in the compound
e. What effect will adding too much titrant have on the experimentally determined value of the mass percent of iron in the
compound? Justify your answer.
Answer the following questions, which pertain to binary compounds.
a. Draw a complete Lewis electron-dot diagram for the IF3 molecule.
b. On the basis of the Lewis electron-dot diagram that you drew in part a, predict the molecular geometry of the IF 3 molecule.
c. In the SO2 molecule, both of the bonds between sulfur and oxygen have the same length. Explain this observation, supporting
your explanation by drawing a Lewis electron-dot diagram (or diagrams) for the SO2 molecule
d. On the basis of your Lewis electron-dot diagram(s) in part c, identify the hybridization of the sulfur atom in the SO2 molecule.
The reaction between SO2(g) and O2(g) to form SO3(g) is represented below.
2 SO2(g) + O2(g)  2 SO3(g)
The reaction is exothermic. The reaction is slow at 25oC; however, a catalyst will cause the reaction to proceed faster.
e. Draw the complete potential-energy diagram for both the catalyzed and uncatalyzed reactions. Clearly label the curve that
represents the catalyzed reaction.
f. Predict how the ratio of the equilibrium pressures, PSO2/PSO3, would change when the temperature of the uncatalyzed reaction
mixture is increased. Justify your prediction.
g. How would the presence of a catalyst affect the change in the ratio described in part f? Explain.
Practice Free Response 2
1.
C(s) + CO2(g)  2 CO(g)
Solid carbon and carbon dioxide gas at 1,160 K were placed in a rigid 2.00 L container, and the reaction represented above
occurred. As the reaction proceeded, the total pressure in the container was monitored. When equilibrium was reached, there was
still some C(s) remaining in the container. Results are recorded in the table below.
Time (hours)
0.0
2.0
4.0
6.0
8.0
10.0
Total Pressure of Gases in Container (atm)
5.00
6.26
7.09
7.75
8.37
8.37
a. Write the expression for the equilibrium constant, Kp, for the reaction.
b. Calculate the number of moles of CO2(g) initially placed in the container. (Assume that the volume of the solid carbon is
negligible.)
c. For the reaction mixture at equilibrium at 1,160 K, the partial pressure of the CO2(g) is 1.63 atm. Calculate
(1) the partial pressure of CO(g), and
(2) the value of the equilibrium constant, Kp.
d. If a suitable solid catalyst were placed in the reaction vessel, would the final total pressure of the gases at equilibrium be
greater than, less than, or equal to the final total pressure of the gases at equilibrium without the catalyst? Justify your
answer. (Assume that the volume of the solid catalyst is negligible.)
In another experiment involving the same reaction, a rigid 2.00 L container initially contains 10.0 g of C(s), plus CO(g) and CO2(g),
each at a partial pressure of 2.00 arm at 1,160 K.
e. Predict whether the partial pressure of CO2(g) will increase, decrease, or remain the same as this system approaches
equilibrium. Justify your prediction with a calculation.
2.
3.
Answer the following questions relating to gravimetric analysis. In the first of two experiments, a student is assigned the task of
determining the number of moles of water in one mole of MgCl 2•n H2O. The student collects the data shown in the following table.
Mass of empty container 22.347 g
Initial mass of sample and container 25.825 g
Mass of sample and container after first heating 23.982 g
Mass of sample and container after second heating 23.976 g
Mass of sample and container after third heating 23.977 g
a. Explain why the student can correctly conclude that the hydrate was heated a sufficient number of times in the experiment.
b. Use the data above to
(1) calculate the total number of moles of water lost when the sample was heated, and
(2) determine the formula of the hydrated compound.
c. A different student heats the hydrate in an uncovered crucible, and some of the solid spatters out of the crucible. This
spattering will have what effect on the calculated mass of the water lost by the hydrate? Justify your answer.
In the second experiment, a student is given 2.94 g of a mixture containing anhydrous MgCl 2 and KNO3. To determine the
percentage by mass of MgCl2 in the mixture, the student uses excess AgNO3(aq) to precipitate the chloride ion as AgCI(s).
d. Starting with the 2.94 g sample of the mixture dissolved in water, briefly describe the steps necessary to quantitatively
determine the mass of the AgCI precipitate.
e. The student determines the mass of the AgCI precipitate to be 5.48 g. On the basis of this information, calculate each of the
following.
(1) The number of moles of MgCl2 in the original mixture
(2) The percent by mass of MgCl2 in the original mixture
Answer the following questions related to chemical reactions involving nitrogen monoxide, NO(g). The reaction between solid
copper and nitric acid to form copper(II) ion, nitrogen monoxide gas, and water is represented by the following equation.
3 Cu(s) + 2 NO3-(aq) + 8 H+(aq)  3 Cu2+(aq) + 2 NO(g) + 4 H2O(l)
Eo = +0.62 V
a. Using the information above and in the table below, calculate the standard reduction potential, E o, for the reduction of NO3- in
acidic solution.
Standard Reduction Potential, Eo
Half-Reaction
2+
Cu (aq) + 2 e  Cu(s)
+0.34 V
NO3-(aq) + 4 H+(aq) + 3 e-  NO(g) + 2 H2O(l)
?
b. Calculate the value of the standard free energy change, Go, for the overall reaction between solid copper and nitric acid.
c. Predict whether the value of the standard entropy change, So, for the overall reaction is greater than 0, less than 0, or equal
to 0. Justify your prediction.
Nitrogen monoxide gas, a product of the reaction above, can react with oxygen to produce nitrogen dioxide gas, as represented
below.
2 NO(g) + O2(g)  2 NO2(g)
A rate study of the reaction yielded the data recorded in the table below.
Experiment
Initial Concentration of NO
Initial Concentration of O2
Initial Rate of Formation of NO2
(mol/L)
(mol/L)
(mol/L•s)
1
0.0200
0.0300
8.52 x 10-2
2
0.0200
0.0900
2.56 x 10-1
3
0.0600
0.0300
7.67 x 10-1
d. Determine the order of the reaction with respect to each of the following reactants. Give details of your reasoning, clearly
explaining or showing how you arrived at your answers.
(1) NO
(2) O2
e. Write the expression for the rate law for the reaction as determined from the experimental data.
f. Determine the value of the rate constant for the reaction, clearly indicating the units.
4.
5.
6.
For each of the following three reactions, in part (1) write a balanced equation for the reaction and in part (2) answer the question
about the reaction. In part (1), coefficients should be in terms of lowest whole numbers. Assume that solutions are aqueous
unless otherwise indicated. Represent substances in solutions as ions if the substances are extensively ionized. Omit formulas for
any ions or molecules that are unchanged by the reaction.
a. Aqueous sodium hydroxide is added to a saturated solution of aluminum hydroxide, forming a complex ion.
(1) Balanced equation:
(2) If the resulting mixture is acidified, would the concentration of the complex ion increase, decrease, or remain the same?
Explain.
b. Hydrogen chloride gas is oxidized by oxygen gas.
(1) Balanced equation:
(2) If three moles of hydrogen chloride gas and three moles of oxygen gas react as completely as possible, which reactant, if
any, is present in excess? Justify your answer.
c. Solid potassium oxide is added to water.
(1) Balanced equation:
(2) If a few drops of phenolphthalein are added to the resulting solution, what would be observed? Explain.
Using principles of atomic and molecular structure and the information in the table below, answer the following questions about
atomic fluorine, oxygen, and xenon, as well as some of their compounds.
Atom
F
O
Xe
First Ionization Energy (kJ/mol)
1,681.0
1,313.9
?
a. Write the equation for the ionization of atomic fluorine that requires 1,681.0 kJ/mol.
b. Account for the fact that the first ionization energy of atomic fluorine is greater than that of atomic oxygen. (You must discuss
both atoms in your response.)
c. Predict whether the first ionization energy of atomic xenon is greater than, less than, or equal to the first ionization energy of
atomic fluorine. Justify your prediction.
d. Xenon can react with oxygen and fluorine to form compounds such as XeO3 and XeF4. Draw the complete Lewis electron-dot
diagram for each of the molecules.
e. On the basis of the Lewis electron-dot diagrams you drew for part (d), predict the following:
(1) The geometric shape of the XeO3 molecule
(2) The hybridization of the valence orbitals of xenon in XeF 4
f. Predict whether the XeO3 molecule is polar or nonpolar. Justify your prediction.
Answer the following questions by using principles of molecular structure and intermolecular forces.
a. Structures of the pyridine molecule and the benzene molecule are shown below. Pyridine is soluble in water, whereas
benzene is not soluble in water. Account for the difference in solubility. You must discuss both of the substances in your
answer.
Pyridine
Benzene
b.
Structures of the dimethyl ether molecule and the ethanol molecule are shown below. The normal boiling point of dimethyl
ether is 250 K, whereas the normal boiling point of ethanol is 351 K. Account for the difference in boiling points. You must
discuss both of the substances in your answer.
Dimethyl Ether
Ethanol
c.
SO2 melts at 201 K, whereas SiO2 melts at 1,883 K. Account for the difference in melting points. You must discuss both of
the substances in your answer.
The normal boiling point of Cl2(l) (238 K) is higher than the normal boiling point of HCl(l) (188 K). Account for the difference in
normal boiling points based on the types of intermolecular forces in the substances. You must discuss both of the substances
in your answer.
d.