Download final-H-2006-07-v2

A physicist, a biologist and a chemist went to the ocean.
The physicist wanted to do research on the waves. He
walked into the ocean and drowned.
The biologist wanted to do research on the flora and
fauna. He too walked into the ocean and drowned.
The chemist shrugged, took out his laboratory notebook,
and wrote, “The physicist and the biologist appear to be
soluble in dilute NaCl.”
Honors Chemistry Final Exam — June 2007
Mr. Bigler & Ms. Manes
Directions:
1. DO NOT OPEN THIS EXAM UNTIL THE PROCTOR INSTRUCTS YOU TO
BEGIN.
2. The exam period ends at 11:30 AM. NO EXTRA TIME WILL BE ALLOWED,
except for students with prior permission. Work accordingly!
3. Print YOUR NAME, your teacher’s name, which class you’re in, and the
TEST NUMBER on your bubble sheet and your open response answer sheet.
There are multiple versions of the test; the test number is necessary in order to
properly grade your exam.
4. Calculators may be used but not shared.
5. You may write on the test booklet, including marking your answer choices.
(This may be helpful if you find yourself accidentally off by one question on the
answer sheet.) However, only the responses on your bubble sheet will be scored.
6. The exam consists of 110 multiple choice questions. A reference sheet with
periodic table and polyatomic ions is included on the last sheet of this exam
packet. You may detach this sheet for your reference.
7. For each question, select the most nearly correct answer and block out the bubble
with a # 2 pencil for the corresponding letter on the answer sheet. Fill in your
choice completely, and completely erase any errors. (Improperly marked or
incompletely erased answers may be marked wrong by the scoring machine.)
8. Do not take too long on any one question. If you get stuck, move on and come
back to it if time permits. Save a few minutes at the end to check your work.
9. When you are done with the exam turn in your answer sheet, scratch paper, and
exam packet to your proctor. However, you may not leave the room until the end
of the test period. Please remain quiet for others who are still testing.
10. May the molementum be with you!
Choose the letter of the correct answer and fill in the corresponding bubble on your
answer sheet. Try your best!
Mole conversions
1. The mass of one mole of neon is
a. 20.18 g
b. 22.4 g
d. 6.02 × 1023 g
c. 4.36 g
2. What is the volume of 3.00 moles of carbon dioxide at STP?
a. 0.134 L
b. 7.47 L
c. 67.2 L
d. 132 L
3. The percentage composition of ammonia (NH3) is
a. 17.8 % N and 82.2 % H
c. 74.9 % N and 25.1 % H
b. 25.1 % N and 74.9 % H
d. 82.2 % N and 17.8 % H
4. A compound was analyzed and found to contain 13.5 g Ca, 10.8 g O, and 0.675 g H.
What is the empirical formula for the compound?
a. HCaO
b. CaOH2
c. Ca2OH
d. Ca(OH)2
5. Which of the following has the greatest mass?
a. 1 mole of H2
b. 2 moles of K
c. 3 moles of Mg
d. 4 moles of He
6. Which of the following represents an empirical formula?
a. C3H6O3
b. C12 H24 O12
c. CH2O
d. C6H12O6
7. The molar mass of a compound is 56 g/mol, and its empirical formula is CH2. What
is its molecular formula?
a. CH2
b. C4H8
c. C6H12
d. C2H4
8. Determine the number of atoms in 10.0 grams of sulfur.
a. 10.0 S atoms
c. 1.88 × 1023 S atoms
b. 320.7 S atoms
d. 6.02 × 1023 S atoms
Redox
9. What is the correct half reaction for aluminum in the following unbalanced reaction:
Al + FeSO4  Al2(SO4)3 + Fe
a. 2 Al + 3 FeSO4  Al2(SO4)3 + 3 Fe
c. 2 Al  2 Al3+ + 6e−
3+
−
b. Al  Al + 3 e
d. Al + FeSO4  Al2(SO4)3 + Fe
10. In question 9 above, what is the reducing agent?
a. aluminum
c. aluminum sulfate
b. iron (II) sulfate
d. iron
11. In the reaction 4 FeO + O2  2 Fe2O3:
a. FeO is oxidized and O2 is reduced.
b. O2 is oxidized and FeO is reduced.
c. Fe is oxidized and O is reduced.
d. O is oxidized and Fe is reduced.
12. What is the oxidation number of chlorine in the perchlorate ion, ClO4−?
a. +7
b. +4
c. +1
d. −1
Chemical Equations
13. In a chemical reaction, the total mass of the products compared to the total mass of
the reactants is
a. always equal
c. always larger
b. always smaller
d. sometimes smaller, sometimes larger
14. What happens to the bonds between atoms in a substance that undergoes a chemical
reaction?
a. All single covalent bonds become double covalent bonds.
b. Some existing bonds will weaken while others will strengthen.
c. Some existing bonds are broken and atoms are rearranged with the formation of
new bonds.
d. Bonds between atoms are unaffected during a chemical reaction.
15. In the balanced equation below:
SiO2 (s)
+
4 HF (aq)
the symbol (l) after H2O stands for
a. liters
b. liquid

SiF4 (g)
+
2 H2O (l)
c. length
d. left over
16. An atom’s ability to undergo chemical reactions is determined by its
a. protons
c. neutrons
b. innermost electrons
d. outermost electrons
17. What are the reactants in the following chemical equation:
Zn + CuSO4  ZnSO4 + Cu
a. zinc and copper
b. zinc and copper (II) sulfate
c. zinc sulfate and copper
d. only zinc
18. Which of the following chemical reactions is a decomposition reaction?
a. BaCO3  BaO +CO2
c. 3 Br2 + 2 FeI3  2 FeBr3 + 2 I2
b. 2 Ca + O2  2 CaO
d. MgCl2 + H2SO4  MgSO4 + 2 HCl
19. Copper in the compound CuSO4 can be isolated in the following reaction with iron:
Fe + CuSO4  FeSO4 + Cu
What type of reaction is shown above?
a. decomposition
b. synthesis
c. single displacement
d. double displacement
20. Aluminum reacts vigorously and exothermically with copper (II) chloride. Which of
the following is the correct balanced equation for this reaction?
a. Al + CuCl2  AlCl3 + Cu
b. Al + 3 CuCl2  2 AlCl3 + Cu
c. 2 Al + 3 CuCl2  2 AlCl3 + 3 Cu
d. 3 Al + 2 CuCl2  3 AlCl3 + 2 Cu
21. Li2SO4 and BaCl2 react. What are the correct formulas for the products?
a. Li2Ba and SO4Cl2
c. LiCl2 and Ba2SO4
b. Li2Cl2 and BaSO4
d. LiCl and BaSO4
22. Magnesium and bromine react. What is/are the correct formula(s) for the product(s)?
a. MgBr
b. MgBr2
c. MgBr3
d. Mg2Br3
e. Mg +Br2
23. When the following equation is balanced, what is the coefficient of H2SO4?
____ Al + ____ H2SO4  ____ Al2(SO4)3 + _____ H2
a. 1
b. 2
c. 3
d. 4
e. 5 or greater
24. When the following equation is balanced, what is the coefficient of C6H6?
_____ C6H6 + _____ O2  ______ H2O + _____ CO2
a. 1
b. 2
c. 3
d. 4
e. 5 or greater
25. When the following equation is balanced, what is the coefficient of Ca3(PO4)2?
_____ Na3PO4 + _____ Ca(NO3)2  _____ Ca3(PO4)2 + _____ NaNO3
a. 1
b. 2
c. 3
d. 4
e. 5 or greater
26. Predict the products for this reaction: potassium phosphate + barium chloride 
a. Ba(PO4)3 + KCl
c. BaPO4 + KCl2
e. Ba3PO4 + KCl
b. Ba2(PO4)3 + KCl
d. Ba3(PO4)2 + KCl
Stoichiometry
27. According to the chemical equation shown below, 4.0 moles of H2O2 (l) can
decompose, producing ____ moles of O2 (g).
2 H2O2 (l)
a. 0.50 mole
b. 2.0 moles

2 H2O (l)
+
c. 3.0 moles
d. 4.0 moles
O2 (g)
e. 5.0 moles
28. How many liters of hydrogen react with 2.00 mol nitrogen at STP in the following
reaction?
N2 + 3 H2  2 NH3
a. 3.00 L
b. 22.4 L
c. 67.2 L
d. 134 L
29. Percent yield is the quantity of a product actually produced compared to the quantity
a. of a product theoretically possible
c. usually produced on average
b. of a limiting reactant
d. of a reactant in excess
30. What are the values of standard temperature and pressure (STP)?
a. 0°C and 1 atm
c. 0°C and 22.4 atm
b. 20°C and 1 atm
d. 20°C and 22.4 atm
31. A piece of copper with a mass of 5.00 g is placed in a solution of silver nitrate
containing excess AgNO3. The silver metal produced has a mass of 15.2 g. What is
the percent yield for this reaction?
Cu + 2 AgNO3  Cu (NO3)2 + 2 Ag
a. 32.9 %
b. 89.4 %
c. 3.04 %
d. 14.0 %
32. How many molecules of chlorine gas can be produced when 63.0 g of HCl react with
excess O2?
4 HCl
−24
a. 1.45 × 10
+

O2
b. 5.27 × 10
23
2 Cl2
c. 6.83 × 10
2 H2 O
26
d. 2.77 × 1027
33. The limiting reactant in a completed chemical reaction will be the substance
a. left over
b. used up
c. with the greatest mass before the reaction begins
d. with the least mass before the reaction begins
Use the figure 11-4 below to answer questions 34 and 35.
34. Which drawing in Figure 11-4 above shows the products of a reaction between
hydrogen and chlorine?
a. a
b. b
c. c
d. d
35. Which molecule is in excess in the reaction of hydrogen and chlorine pictured by the
drawings in Figure 11-4?
b. Cl2
c. H2
d. None
a. HCl
36. If you add 0.10 moles of HCl to 0.10 moles of Mg, which one of the following is
true?
2 HCl + Mg  MgCl2 + H2
a. Mg will be a limiting reactant.
b. Mg will remain in excess.
c. HCl will remain in excess.
d. none of the above is correct.
37. A balanced chemical reaction is shown below.
C5H12 + 8 O2  5 CO2 + 6 H2O
If 3 moles of C5H12 are reacted completely, how many moles of water are formed?
a. 3
b. 6
c. 12
d. 18
Questions 38 and 39 refer to the following chemical reaction, in which a 2.00 g sample of
ammonia is mixed with 4.00 g of oxygen.
4 NH3 (g) + 5 O2 (g)  4 NO (g) + 6 H2O (g)
38. Which is the limiting reactant?
a. NH3
b. O2
c. NO
d. H2O
39. How much NO would be produced?
a. 3.00 g
b. 3.53 g
c. 3.75 g
d. 14.1 g
States of Matter
40. Compared to a liquid, a gas is more compressible because the particles
a. move more slowly
c. are more attracted to each other
b. have greater space between them
d. have more frequent collisions
41. The state of matter at which a substance exists at room temperature depends on
a. the volume of the particles
b. the number of particles
c. the strength of the attractive forces between particles
d. the number of collisions between particles
Figure 14-2 represents the kinetic energies of the molecules in two systems, one at
temperature T1 and the other at temperature T2.
42. How do the two temperatures represented by curves T1 and T2 in Figure 14-2
compare?
a. T1 is lower than T2
c. T1 equals T2
b. T1 is higher than T2
d. Not enough information to tell
43. Substances with particles experiencing strong intermolecular forces
a. have high boiling and high melting points
b. have high boiling and low melting points
c. have low boiling and high melting points
d. have low boiling and low melting points
e. are usually highly reactive
Use Figure 14-4 to answer questions 44 and 45.
44. In Figure 14-4, where do the three phases exist in equilibrium?
a. only at the origin
c. on any solid line
b. only at the triple point
d. between any two solid lines
45. In Figure 14-4, what phase would be present at 15 oC and 20 mmHg?
a. solid
b. liquid
c. gas
d. vapor
46. Which of the following compounds has the highest melting and boiling points?
a. Ar
b. HBr
c. CaCl2
d. PCl3
47. Particles in nonpolar liquids are attracted to each other because of
a. covalent bonding
d. dipole-dipole forces
b. ionic bonding
e. metallic bonding
c. dispersion forces
48. A glass of ice water at 0°C absorbs heat until all the ice melts. Even though heat is
being absorbed, the temperature does not change. Why is this true?
a. The energy is being used to break the intermolecular attractions between the
water molecules in the ice.
b. The energy is being used to increase the kinetic energy of the water molecules
in the liquid state.
c. The energy is being used to increase the kinetic energy of the water molecules
in the solid state.
d. The energy is being used to increase the kinetic energy of the water molecules
in both liquid and solid states.
e. None of the above.
49. A process that releases heat is a(n) __________ process.
a. exothermic
b. endothermic
c. polythermic
d. ectothermic
50. What phase(s) is/are present at point B in
the heating curve to the right?
a. solid
b. solid & liquid
c. liquid
d. liquid & gas
e. gas
Gases
51. The temperature at which the motion of particles ceases is _________.
a. –273 K
b. 0 K
c. 0°C
d. 273°C
52. Which of the following occurs when a rigid container of gas is heated?
a. The pressure inside the container increases.
b. The pressure inside the container decreases.
c. The pressure inside the container stays the same.
d. The pressure inside the container changes the composition of the gas.
53. When the amount and volume of a gas are kept constant, if the pressure increases, the
temperature of the gas
a. is unchanged
b. increases
c. decreases
54. A gas at a pressure of 608 mmHg is held in a container with a volume of 545 cm3.
The volume of the container is then increased to 1065 cm3 without change in
temperature. What is the new pressure of the gas?
a. 1188 mmHg
b. 311 mmHg
c. 955 mmHg
55. The temperature inside of a car on a summer’s day is 42°C. What is that temperature
expressed in Kelvin?
a. 231 K
b. 315 K
c. 215 K
d. –131 K
56. How many moles of a gas at 100°C does it take to fill a 1.00 L flask to a pressure of
1.50 atm? {R = 0.0821 (L × atm) / (mol × K) }
a. 0.0490 mol
b. 0.183 mol
c. 5.47 mo
d. 20.4 mol
57. Which gas effuses faster at the same pressure and temperature, CO2 or CO?
a. CO
c. they effuse at the same rate
b. CO2
d. not enough information
58. If the pressure of the gas in the open mercury manometer at the right is 815 torr, what
is the atmospheric pressure?
a.
b.
c.
d.
e.
880 torr
750 torr
880 atm
750 atm
There is not enough information to determine the
atmospheric pressure.
59. Two samples of gas represented below have the same volume, temperature, and
pressure.
Based on this information, these two samples of gas must also have the same
a. chemical reactivity
c. mass
b. density
d. number of molecules
60. You are given a mixture of gases consisting of only nitrogen, hydrogen and helium.
The pressure of the nitrogen is 33.4 kPa and the pressure of hydrogen is 58.1 kPa. If
the total pressure of the mixture is 125.4 kPa, then the pressure of the helium is:
a. 33.9 kPa
b. 216.9 kPa
c. 94.6 kPa
d. 52.3 kPa
61. The increase in pressure of a gas with a decrease in volume is the result of
a. increased numbers of collisions
c. decreased force on the container
b. decreased number of collisions
d. none of the above
62. A fixed sample of gas occupies 40.0 mL at –123°C. What volume does the sample
occupy at 27°C, assuming constant pressure?
a. 182 mL
b. 8.80 mL
c. 80.0 mL
d. 20.0 mL
63. An inflated ball
a. weighs more than when it is deflated
b. weights less than when it is deflated
c. weighs the same when it is deflated
64. A real gas deviates from ideal behavior at
a. high temperature and low pressure
b. low temperature and high pressure
c. low pressure and low temperature
d. high pressure and high temperature
65. Four different gases are all observed to have the same temperature. Which of the
following conclusions is supported by this observation?
a. All four gases must have the same mass.
b. All four gases must have the same pressure.
c. All four gases must have equal numbers of particles.
d. All four gases must have equal average kinetic energies.
Reaction rates
66. A catalyst
a. is not used up in a reaction
b. is one of the reactants in a single-replacement reaction
c. is a solid product in a reaction
d. is the product of a combustion reaction
67. Activation energy is the energy required to form the
a. original bonds.
c. reaction complex.
b. activated complex.
d. reactant molecules.
68. Which of the following factors does NOT affect the speed of a chemical reaction?
a. the size of the beaker in which the reaction proceeds.
b. temperature.
c. nature of the reactants.
d. concentration of the reactants.
e. an inhibitor
69. In general, conditions which increase the number of collisions taking place between
reactant molecules have no effect on the rate of a reaction.
a. true
b. false
70. Which will react faster with water, a lump of calcium, or finely powdered calcium?
a. The lump of calcium will react faster.
b. The finely powdered calcium will react faster.
c. Both the lump of calcium and the powdered calcium will react at the same rate.
71. In order for molecules to react they must
a. collide with the right amount of energy
b. collide with the proper orientation
c. collide with the walls of their container
d. both a & b
e. all of the above
Solutions
Use Figure 15-1 to answer questions 72 and 73.
72. According to Figure 15-1, if 75 grams of KCl dissolve in 100 grams of water at 90°C,
the solution is
a. supersaturated
c. unsaturated
b. dilute
d. saturated
73. Use Figure 15-1 to determine how many grams of KNO3 you would dissolve in
100 grams of water at 40°C to make a saturated solution.
a. 25 grams
c. 75 grams
b. 67 grams
d. less than 50 grams
74. In ocean water, NaCl (salt) would be considered the
.
a. solvent
b. solute
c. solution
d. colloid
e. suspension.
75. Solution concentration expressed as moles of solute/kilogram of solvent is known as
a. molarity
c. mole fraction
b. molality
d. millimolarity
76. Molarity can be described as
a. a measure of concentration
b. moles of solute per liter of solution
c. a constant ratio between solute and solvent
d. being represented by the symbol M
e. all of the above
77. Mole fraction is
a. moles of reactant compared to moles product
b. moles of acid compared to moles of base
c. moles of substance compared to total moles
78. A solution which, upon mixing for a long period of time, still contains undissolved
particles is most likely a(n) _____ solution.
a. supersaturated
b. unsaturated
c. saturated
79. How many moles of sugar are present in 4.0 liters of a 0.25 M solution?
a. 0.042 mole
c. 24 moles
b. 0.15 mole
d. a, b, and c are all incorrect
80. Which of the following is NOT an important factor influencing solubility?
a. chemical nature of solute
c. chemical nature of solvent
b. temperature
d. all of the above are important
81. A substance dissolves in a polar solvent. The substance is
a. polar
d. either polar or ionic
b. nonpolar
e. either nonpolar or ionic
c. ionic
82. When a solute is dissolved in a solvent, the boiling point of the resulting solution is
________ the boiling point of the pure solvent.
a. greater than
b. lower than
c. the same as
d. either greater than or lower than, depending on the nature of the solute
83. How does temperature affect the solubility of a gas?
a. the higher the temperature, the greater the solubility
b. the higher the temperature, the lower the solubility
c. there is no relationship between the temperature and the solubility
84. Many laboratory preparations of solutions call for stirring the solvent while adding
the solute. Which of the following is always an effect of this procedure?
a. It decreases the reactivity of the solute.
b. It decreases the solubility of the solute.
c. It brings the solute and solvent rapidly into contact.
d. It produces a double displacement reaction.
85. A data table and two prepared beakers are shown below.
Solid KNO3 was added to each beaker. Each beaker was stirred at the same rate until
all of the solid dissolved. The table shows the solubilities of KNO3 at different
temperatures. How will the rates of dissolving compare?
a. KNO3 will dissolve faster in Beaker B because of increased surface area.
b. KNO3 will dissolve faster in Beaker A because the water molecules are farther
apart.
c. KNO3 will dissolve faster in Beaker B because the overall kinetic energy is
increased.
d. KNO3 will dissolve at the same rate in Beaker A and Beaker B because the
concentrations are the same.
86. How many mL of 3.25 M NaCl would be needed to make 15.5 mL of a 0.500 M NaCl
solution?
a. 0.105 mL
b. 0.419 mL
c. 2.38 mL
d. 9.54 mL
87. What is the boiling point of a 0.222 m solution of K3PO4 in water? (Kb for water is
0.512°C/m)
a. 0.114°C
c. 100.114°C
e. 100.455°C
b. 0.455°C
d. 100.227°C
Equilibrium
88. At equilibrium,
a. the forward reaction rate is less than the reverse reaction rate.
b. the forward reaction rate is greater than the reverse reaction rate.
c. the forward reaction rate is equal to the reverse reaction rate.
d. no reactions take place.
−8
89. In an equilibrium reaction with a Keq of 1 × 10 , the
a. reactants are favored.
c. the products are favored.
b. the reaction is spontaneous.
d. the reaction is exothermic.
90. According to LeChâtelier's Principle, if conditions of a system at equilibrium are
changed, the system will
a. shift in the direction that tends to reduce that change in conditions.
b. shift to minimize the amount of reactants.
c. shift to minimize the amount of products.
d. shift so that the products equal the reactants.
91. What is the equilibrium expression for the following hypothetical reaction:
A+2B
c. K eq  [A][B] 2 [C]
[A][B] 2
[C]
[C]

[A][B] 2
a. K eq 
b. K eq
C
d. K eq 
[A]  [B] 2
[C]
e. K eq 
92. A reaction's equilibrium constant will change if the
a. temperature
b. pressure
c. concentration of one of the reactants or one of the products
d. none of these affects Keq.
[C]
[A]  [B] 2
is changed.
93. The concentrations of solids and pure liquids are not included in equilibrium
expressions.
a. true
b. false
Questions 94–96 refer to a hypothetical reversible reaction in which reactant Y is a bright
yellow color, reactant C is colorless, and the product B is a bright blue color
C (g) + Y (g)
3 B (g) + heat
The reaction is slightly exothermic, with Keq = 1.50. The system is initially at
equilibrium, with a green color. (Note that green is a roughly equal mixture of blue and
yellow.)
94. Adding a large quantity of C favors a
a. blue color
b. yellow color
c. green color
95. Heating the system favors a
a. blue color
b. yellow color
c. green color
96. Increasing the pressure on the system favors a
a. blue color
b. yellow color
c. green color
97. Ammonia is made according to the following equation:
N2 (g) + H2 (g)
2 NH3 (g) + heat
In the Haber Process, NH3 is continuously removed as the above reaction occurs.
What effect does this have on the equilibrium?
a. the reaction would shift to the right
b. the reaction would shift to the left
c. there would be no change
98. Consider the following equilibrium equation:
H2 (g) + I2 (g)
2HI (g)
At 448°C in a 2.0 L closed system, 2.0 × 10−2 mol HI, 1.0 × 10−2 mol of H2, and
3.0 × 10−2 mol of I2 are present. The equilibrium constant for the reaction is 51.
Which of the following is true?
a. The reaction is at equilibrium.
b. The reaction is not at equilibrium. The reaction must proceed to the right to
reach equilibrium.
c. The reaction is not at equilibrium. The reaction will proceed to the left to reach
equilibrium.
d. Not enough information is provided to answer the question.
Acid-Base
99. 25.0 mL of a HNO3 solution of unknown concentration is titrated with 0.50 M NaOH.
75.0 mL of the sodium hydroxide solution are required to reach the equivalence point.
The original concentration of the HNO3 solution was:
a. 0.25 M
b. 0.50 M
c. 0.75 M
d. 1.5 M
e. 3.0 M
100.
What is the pH of a solution that has an H+ concentration of 0.0001 M?
a. 1
b. 2
c. 3
d. 4
e. 11
101.
What is the pOH of a solution that has a pH of 6?
a. 1
b. 4
c. 6
d. 8
e. 10
102.
Substances that change color when added to acids or bases are called
a. indicators
b. electrolytes
c. insulators
103.
KOH is a(n)
a. acid
b. base
c. salt
d. indicator
104. Which of the following is the best answer for the products of the reaction between
an acid and a base?
a. acid + salt
c. base + salt
e. water + salt
b. salt
d. water
105. Distilled water is boiled to remove all impurities. What is true of the
concentrations of [H3O+] and [OH−] in this water?
a. they are the same
b. they are about equal to the concentration of water molecules
c. They are extremely small but still significant
d. Answers a and c are correct
106.
For basic solutions, the concentration of hydroxide ions, [OH−], is
a. greater than 1 × 10 –7 M
b. 1 × 10 7 M
c. equal to the hydrogen ion concentration
d. zero
107.
What is the [OH−] in a solution with a pH of 3.21?
a. 1.6 × 10−11 M
b. 6.3 × 10−4 M
c. 1.6 × 103 M
d. 6.3 × 104 M
Which of the following substances is amphoteric?
a. H2SO4
b. NH4+
c. HCO3−
d. HBr
108.
109. If calcium hydroxide and sulfuric acid were mixed equally, the resulting solution
would be
a. acidic
b. basic
c. neutral
110.
What is the conjugate base of HSO4−?
a. H2SO4
b. HSO4+
c. SO42−
d. OH−
If you have not yet turned in your book, please do so after you hand in your exam.
Congratulations! You have finished the final exam for Honors
Chemistry. Thank you for the wonderful year!