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Chemistry Final Review
Level 3
Matter & Energy
_____ 1. Which of the following correctly pairs a phase of matter with its description?
A. Solid: Particles have no motion.
B. Liquid: Particles expand to fill any container in which they are placed.
C.
Gas: Particles have higher amounts of energy than when in the liquid phase.
D. Liquid: Particles are more strongly attached to one another than when in the solid phase.
_____ 2. Pure substances include __________.
a. elements only.
c. elements and compounds.
b. compounds and mixtures.
d. elements and mixtures.
_____ 3. The normal boiling point of water is __________.
a. 373 K
b. 173 K
c. 273 K
d. 473 K
_____ 4. The table below shows the physical properties of selected metals.
Physical Properties of Selected Metals
Metal
Molecular
mass (amu)
Melting
Boiling
point (°C) point (°C)
Density
(g/cm3)
Bismuth
209.98
271
1560
9.80
Chromium
52.00
1857
2672
7.20
Polonium
210.05
254
962
9.40
Ruthenium
101.07
2310
3900
12.3
A cube of an unknown metal has a volume of 2.25 cm3 and a mass of 16.2 g. Based on data
in the table above, what is the identity of this metal?
a. bismuth
b. chromium
c. polonium
d. ruthenium
_____ 5. One way that mixtures differ from pure substances is in the methods that can be
used to separate them into their components. Which of the following is a method used to
separate the components of some mixtures?
A. a nuclear reaction
C. a chemical reaction
B. a filtration process
D. an electrolysis process
_____ 6. Which of the following substances is made of particles with the highest average
kinetic energy?
a. Fe (s) at 35C
c. H2O (l) at 30C
b. Br2 (l) at 20C
d. CO2 (g) at 25C
_____ 7. Which of the following describes the separation of the components of a mixture?
A. Water is broken down into hydrogen and oxygen.
B. Salt is isolated from seawater through evaporation.
C. Propane reacts with oxygen to form carbon dioxide and water.
D. Calcium carbonate decomposes to form calcium oxide and carbon dioxide.
_____ 8. Which temperature represents absolute zero?
a. 0 C
b. 0 K
c. 273 K
_____ 9. The graph below compares three states of a substance.
Which of the following choices is the best label for the y-axis?
a. molecular density
c. neutron density
b. molecular motion
d. neutron motion
d. 273 C
_____ 10. A solid cube was put into a cylinder containing four liquids with different densities
as shown below.
The cube fell quickly through layer A, fell slowly through layer B, and stopped upon
reaching layer C. The density of the cube most likely lies between __________.
a. 1.00 and 1.50 g/cm3
c. 3.51 and 6.00 g/cm3
b. 1.51 and 3.50 g/cm3
d. 6.00 and 9.00 g/cm3
Atomic Structure
______ 1. The atomic number of an element indicates which of the following?
A. the number of neutrons in the atom
B. the number of protons in the atom
C. the sum of the neutrons and protons in the atom
D. the sum of the protons and electrons in the atom
_____ 2. Which of the following did scientists learn about the atom from Rutherford’s gold
foil experiment?
A. Atoms combine in simple ratios to form compounds.
B. Electrons travel around the nucleus of an atom in concentric circular paths.
C. The mass of an atom and its positive charge are concentrated in the nucleus.
D. The atomic mass of an atom is equal to the number of protons and neutrons in the
nucleus.
____ 3. In principle energy level 2, what types of orbitals will be present?
A. s
B. s & p
C. s, p, & d
D. s, p, d, & f
4. Deuterium ( H) and protium ( H) are two isotopes of hydrogen. Which of the
following statements best compares a deuterium atom to a protium atom?
A. The deuterium atom has a smaller net charge.
B. The deuterium atom has more electron orbitals.
C. The deuterium atom has a smaller atomic radius.
D. The deuterium atom has more particles in its nucleus.
5. Which of the following describes a particle that contains 36 electrons, 49
neutrons, and 38 protons?
A. an ion with a charge of 2−
B. an ion with a charge of 2+
C. an atom with a mass of 38 amu
D. an atom with a mass of 49 amu
_____ 6. Which of the following elements can form an anion that contains 54 electrons, 74
neutrons, and 53 protons?
B.
C.
D.
A.
_____ 7. Which of the following represents a pair of isotopes?
A.
1H
and 3H
B.
16O2−
and 19F1−
C.
40K
and 40Ca
D.
16O2−
_____ 8. Which element has the electron configuration 1s22s22p3?
A. boron
B. nitrogen
C. fluorine
D.
phosphorus
and 32S2
_____ 9. Which of the following pieces of evidence best supports Bohr’s idea that electrons
occupy specific energy levels within an atom?
A. Sodium atoms become positive ions when they lose electrons.
B. Each element emits a unique bright-line spectrum when it falls from an excited
state to a ground state.
C. Beryllium atoms bombarded with alpha particles produce beams that are not
influenced by magnetic fields.
D. Each element has physical and chemical properties that are unique to that
element and different from those of other elements.
_____ 10. How many grams are in 7.80 moles of NaCl?
A. 0.134 g
B. 7.44 g
C. 221 g
D. 452 g
_____ 11. What is the molar mass of H2SO4, rounded to the nearest whole number?
A. 100 g
B. 100 g/mol
C. 98 g
D. 98 g/mol
_____ 12. Which of the following comparisons correctly describes subatomic particles?
A. An electron has a negative charge and a mass larger than the mass of a proton.
B. A neutron has a negative charge and a mass smaller than the mass of a proton.
C. A neutron has a neutral charge and a mass larger than the mass of an electron.
D. A proton has a positive charge and a mass smaller than the mass of an electron.
_____ 13. How many moles of He have 7.34 x 1023 particles in it?
A. 0.820 mol
B. 1.22 mol
C. 44.2 mol
22.4 mol
_____ 14. How many neutrons are present in an atom of silver, which has a mass number of
108?
A. 14
B. 47
C. 61
D. 108
_____ 15. The atomic theories of Dalton, Thomson, Rutherford, and Bohr all support which of
the following statements?
A. Atoms are mostly composed of empty space.
B. All matter is composed of tiny, discrete particles called atoms.
C. Electrons orbit the nucleus of an atom at distinct energy levels.
D. Atoms are composed of positively and negatively charged particles.
_____16. When a sample of potassium chloride dissolves in water, it separates into
potassium ions and chloride ions. Which of the following best accounts for the positive
charge of the potassium ions?
A. They have extra mass.
B. They have a large volume.
C. They have fewer electrons than protons.
D. They have a high density of neutrons and protons.
_____17. Do 35 moles of helium or 35 moles of iron have more atoms in it?
A. helium
B. iron
C. they have the same number of atoms
D. not enough information is give to answer this question
Periodicity
_____ 1. Which element is considered malleable?
a. hydrogen
b. gold
c. sulfur
d. radon.
_____ 2. Which element in period 2 has the greatest tendency to gain electrons?
a. fluorine
b. lithium
c. carbon
d. neon
_____ 3. According to Mendeleev, the chemical properties of elements are periodic functions
of their
a. atomic size
b. atomic weight
c. atomic number
_____ 4. A vertical column in the periodic table is known as a(n)
d. isotopic weight
a. octave
b. period
c. group
d. triad
_____ 5. Which of the following elements has characteristics of some metals and also of some
nonmetals?
a. antimony (51Sb)
b. calcium (20Ca)
c. sulfur (16S) d. zinc (30Zn)
_____ 6. According to Moseley, the chemical properties of elements are periodic functions of
their
a. atomic size b. atomic weight
c. atomic number
d. isotopic weight
_____ 7. Which of the following trends in the periodic table should be expected as the atomic
number of the halogens increases from fluorine (F) to iodine (I)?
a. Atomic radius decreases
c. Electronegativity decreases
b. Atomic mass decreases
d. Electron number decreases
_____ 8. Which of the following correctly describes a trend from top to bottom in the group 2
(2A) elements on the periodic table?
a. Ionic radius decreases
c. Ionic charge increases
b. Atomic radius increases
d. Atomic number decreases
_____ 9. Which element will form an ion whose ionic radius is larger than its atomic radius?
a. fluorine
b. potassium
c. lithium
d. magnesium
_____ 10. The most reactive member of the alkali metals is
a. potassium
b. rubidium
c. cesium
d. francium
_____ 11. Which of the following elements has the highest electronegativity?
a. B
b. C
c. O
d. N
_____ 12. Based on its position on the periodic table, which of the following elements is a
nonmetal?
a. potassium
b. vanadium
c. nickel
d. bromine
_____ 13. Which of the following statements describes the elements in family 16 of the
periodic table?
a. They have six valence electrons
b. They are all gases at room temperature
c. They exist commonly as cations in nature
d. They combine easily with elements in family 17
_____ 14. The figure below shows part of the periodic table.
Cu
Ag
Au
Which of the following is an accurate comparison of the atomic number and mass of copper
and gold?
a. Au has a smaller atomic mass and fewer electrons than Cu
b. Au has the same atomic mass as Cu but a greater atomic number
c. Au has the same atomic number as Cu but a much greater atomic mass
d. Au has both a greater atomic number and a greater atomic mass than Cu
_____ 15. The figure below represents the periodic table and the location of four different
elements on the periodic table.
A certain element has a ground state electron configuration of 1s22s22p63s23p6. Which
letter in the diagram above represents the position of this element on the periodic table?
a. W
b. X
c. Y
d. Z
_____ 16. Which of the following elements is a member of the halogen family and is located
in period 3?
a. argon (Ar)
b. bromine (Br)
c. chlorine (Cl)
d. sulfur (S)
_____ 17. Which of the following sections of the periodic table contains only metals?
a. group 2
b. group 18
c. period 2
d.
period 6
_____ 18. Which of the following characteristics of an element can be determined precisely
by considering only the element’s specific position on the periodic table?
a. radius of each ion
b. density of the solid
c. boiling point of the liquid
d. number of protons in each atom
Chemical Bonding
_____ 1. Which of the following statements best explains why atoms bond?
A)
B)
C)
D)
Atoms bond to make new substances.
Atoms bond to become less chemically stable.
Atoms bond to change from a liquid to a solid.
Atoms bond to become more chemically stable.
_____ 2. The Lewis dot structure shown below represents an atom of an unknown
metallic element M.
When atoms of this unknown metal react with oxygen, a compound is formed.
Which of the following is the most likely chemical formula of the resulting metal
oxide?
A) MO
B) MO2
C) M2O
D) M2O3
_____ 3. An unknown metal, X, combines with nitrogen to form the compound XN.
Metal X also combines with oxygen to produce the compound X2O3.
Metal X is most likely which of the following elements?
A) 3Li
B) 12Mg
C) 31Ga
D) 50Sn
_____ 4. The table below contains information about an unknown metal.
How many valence electrons does the unknown metal have?
A) 1
B) 3
C) 4
D) 6
_____ 5. Atoms of element A and atoms of element B react to form a compound. In
the reaction, the radius of each atom of element A is decreased.
Which of the following explains this decrease in atomic radius in the reaction?
A)
B)
C)
D)
The atoms of element A lose electrons to atoms of element B.
The atoms of element A gain neutrons from atoms of element B.
Nuclear particles are converted into energy in atoms of element A.
Protons become more densely packed in the nuclei of element A atoms.
_____ 6. The diagram below represents particles of different elements in a crystal.
What type of bond holds these particles together?
A) Covalent
B) Hydrogen
C) Ionic
D) Polar
_____ 7. A student heated a 10 g sample of a compound in an open container. A chemical
reaction occurred. The mass of the sample was measured again and found to be less than
before. Which of the following explains the change in mass of the sample?
A)
B)
C)
D)
The heat caused the compound to become less dense.
The reaction gave off more heat than was added.
Some of the lighter atoms were converted to energy.
One of the reaction products was a gas.
_____ 8. A 1.00 kg sample of water (H2O) contains 0.11 kg of hydrogen (H) and 0.89
kg of oxygen (O). According to the law of definite proportions, how much hydrogen
and oxygen would a 1.5 kg sample of water contain?
A)
B)
C)
D)
0.11 kg H and 0.89 kg O
0.17 kg H and 1.34 kg O
0.22 kg H and 1.78 kg O
1.34 kg H and 0.17 kg O
_____ 9. Which of the following has the same empirical formula as mercury (II)
acetate, Hg2(C2H3O2)2?
A)
B)
C)
D)
Mercury (I) bicarbonate, HgHCO3
Mercury (II) bicarbonate, Hg2(HCO3)2
Mercury (I) acetate, HgC2H3O2
Mercury (I) oxalate, Hg2C2O4
_____ 10. What is the percent mass of oxygen in acetone (C3H6O)?
A) 1.00%
B) 10.3%
C) 27.6%
D) 62.0%
_____ 11. Which of the following properties is not associated with metallic elements?
A) malleability
_____
B) brittleness
C) ductility
12. This substance is held together by metallic bonds
A)
B)
C)
D)
Hydrogen gas, H2
Potassium, K
Aluminum oxide, Al2O3
Bromine, Br2
D) conductivity
_____13. This holds a sample of barium iodide, BaI2, together
A)
B)
C)
D)
Ionic bonding
Metallic bonding
Nonpolar covalent bonding
Polar covalent bonding
_____ 1. Covalent bonds are usually formed by the combination of
A. a metal and a nonmetal
B. a metal and an acid group
C. two nonmetals
D. very active metal and the hydroxide ion.
_____ 2. The illustration below shows two atoms of a fictitious element (M) forming a
diatomic molecule.
What type of bonding occurs between these two atoms?
A. covalent
B. ionic
C. nuclear
D. polar
_____ 3. The chemical formula for ammonia is NH3. Which of the following is the correct
Lewis electron dot structure for ammonia?
A.
C.
B.
D.
_____ 4. Which of the following Lewis dot structures represents the compound methane
(CH4)?
A.
B.
C.
D.
_____ 5. Which of the following statements explains why the bond in hydrogen chloride
(HCl) is polar covalent?
A. The atomic mass of chlorine is greater than that of hydrogen.
B. The electronegativity of chlorine is greater than that of hydrogen.
C. The diameter of a chlorine atom is greater than that of a hydrogen atom.
D. The number of valence electrons in a chlorine atom is greater than that in a hydrogen
atom
_____ 6. Which is an example of a non-polar molecule that contains polar covalent bonds?
A. CCl4
B. N2
C. H2S
D. NH3
_____7. Two compounds that contain the elements carbon and chlorine are carbon
tetrachloride (CCl4) and chloroform (CHCl3). Which of the following statements describes
the geometry around carbon in these two compounds?
A. CCl4 and CHCl3 have bent geometries.
B. CCl4 and CHCl3 have tetrahedral geometries.
C. CCl4 has linear geometry and CHCl3 has bent geometry.
D. CCl4 has tetrahedral geometry and CHCl3 has trigonal planar geometry.
_____8. Two elements in a molecule have the same electronegativity values. Which of the
following most likely holds the elements together and why?
A. an ionic bond, because electrons transfer from one element to the other
B. a nonpolar covalent bond, because the elements share electrons equally
C. a polar covalent bond, because the elements do not share electrons equally
D. an intermolecular force, because the elements do not form a chemical bond
_____9. What is the empirical formula for C4Br2F8?
A. CBrF
B. C2BrF4
C. C2BrF6
D. C8Br8F8
_____10. A leaf gently floats on a pond. Which of the following statements best explains why
the leaf stays on top of the water?
A. The leaf has nonpolar covalent bonds between its atoms.
B. The density of the leaf is greater than the density of the water.
C. The water molecules are held tightly together by hydrogen bonding.
D. The hydrogen and oxygen atoms in the water are chemically bonded.
_____11. Which of the following elements does not form a diatomic molecule?
A. Oxygen
B. Nickel
C. Bromine
D. Hydrogen
_____12. The Lewis dot structure of a compound is shown below.
Which of the following elements does X represent in the structure?
A.
carbon (C)
C.
oxygen (O)
B.
nitrogen (N)
D.
fluorine (F)
_____13. Which of the following could be the molecular formula of a compound that has the
empirical formula CHO? Its molar mass is 116.1 g/mol.
A. CHO
B. C2H2O2
C. C3H3O3
D. C4H4O4
_____14. What is holding the following compounds near each other?
A. A metallic bond
B. An intermolecular force
C. An intramolecular bond
D. none of the above.
Chemical Reaction & Stoichiometry
_____ 1. Which of the following chemical equations is balanced correctly?
A.
B.
C.
D.
_____ 2. Which of the following chemical reactions is a decomposition reaction?
A. BaCO3
→ BaO + CO2
B. 2 Ca + O2
→
2 CaO
C. 3 Br2 + 2 FeI3
→
2 FeBr3 + 3I2
D. MgCl2 + H2SO4
→ MgSO4 + 2HCl
_____ 3. An unbalanced chemical equation is shown below.
H3BO3
B2O3 + H2O
What are the coefficients of the balanced equation?
A. 2:1:3
C. 3:1:2
B. 2:2:3
D. 3:2:2
_____ 4. Potassium carbonate (K2CO3) is an important component of fertilizer. The partially
balanced equation for the reaction of 6 moles of potassium hydroxide (KOH) and 3 moles of
carbon dioxide (CO2) to produce potassium carbonate and water is given below.
6KOH + 3CO2 → K2CO3 + 3H2O
When this equation is balanced, what is the coefficient for potassium carbonate?
A. 2
C. 6
B. 3
D. 9
_____ 5. Aluminum reacts vigorously and exothermically with copper(II) chloride. Which of
the following is the balanced equation for this reaction?
A.
Al + CuCl2 → AlCl3 + Cu
B.
Al + 3CuCl2 → 2AlCl3 + Cu
C. 2Al + 3CuCl2 → 2AlCl3 + 3Cu
D. 3Al + 2CuCl2 → 3AlCl3 + 2Cu
_____ 6. Which of the following represents a double displacement reaction?
A. ABC → AB + C
B. A + B → AB
C. AB + CD → AD + CB
D. A + BC → AC + B
_____ 7. The balanced equation below shows the reaction used to make calcium sulfate
(CaSO4), an ingredient in plaster.
In an experiment, 0.500 mol CaCO3 reacted with excess sulfuric acid (H2SO4). The reaction
produced 0.425 mol CaSO4. What was the percent yield for the reaction?
A. 42.5%
C. 73.5%
B. 50.0%
D. 85.0%
_____ 8. When pure N2O5 is heated under certain conditions, O2 and NO2 are produced. What
type of reaction is this?
A. combustion
C. double displacement
B. decomposition
D. synthesis (combination)
_____ 9. Calcium combines with boron, as represented by the chemical equation below.
Ca + 6B  CaB6
What is the minimum amount of calcium, in grams, that could completely react with 54.0 g
of boron?
A. 9.0 g
C. 87 g
B. 33 g
D. 240 g
_____10. Which of the following diagrams represents a single displacement (replacement)
reaction?
A.
B.
C.
D.
_____11. A chemical reaction involving substances A and B stops when B is completely
used. B is the
A. excess reactant.
C. limiting reactant.
D. excess product.
B. limiting product.
_____12. The following reaction can be identified as a ____________ reaction:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
A. Synthesis
B. Double-Replacement
C. Combustion
D. Decomposition
_____13. The figure below represents a reaction.
What type of reaction is shown?
A. synthesis
B. decomposition
C. single displacement
D. double displacement
Nuclear Chemistry
_____ 1. Uranium forms thorium and helium, as shown in the equation below.
Which of the following does this equation represent?
A. decomposition reaction
C. radioactive decay
B. physical change
D. synthesis reaction
_____ 2. Which of the following statements applies to a nuclear fission reaction?
A. The reaction has no commercial applications.
B. The reaction takes place only at very high temperatures.
C. The reaction produces only short-lived radioactive waste.
D. The reaction releases large amounts of energy when nuclei split apart.
_____ 3. Gold-198 has a half-life of approximately 3 days. If a 100 g sample of gold-198
decays for 9 days, approximately how much gold-198 remains in the sample?
A. 13 g
C. 33 g
B. 25 g
D. 50 g
_____ 4. Which of the following statements accurately describes alpha particles in terms of
charge and mass?
A. Alpha particles are positively charged & less massive than beta particles.
B. Alpha particles are negatively charged & less massive than beta particles.
C. Alpha particles are positively charged & more massive than beta particles.
D. Alpha particles are negatively charged & more massive than beta particles.
_____ 5. The atomic number of an element indicates which of the following?
A. the number of neutrons in the atom
B. the number of protons in the atom
C. the sum of the neutrons and protons in the atom
D. the sum of the protons and electrons in the atom
_____ 6. The equation below shows the radioactive decay of thorium (Th).
Which of the following particles is released in this reaction?
A.
B.
C.
D.
_____ 7. The final elements produced by radioactive decay differ from the original
radioactive elements because the nuclei of the final elements are always
A. more stable.
B. increased in mass.
C. half as radioactive.
D. positively charged.
_____ 8. A radioactive source emits a beam containing alpha, beta, and gamma radiation. The
beam passes between two charged plates before striking a detection screen. One plate is
negatively charged and the other plate is positively charged, as shown in the diagram
below.
Which of the following tables indicates the location where each type of radiation will most
likely strike the detection screen after passing between the charged plates?
A.
C.
B.
D.
_____ 9. The positions of copper (Cu) and carbon (C) are identified on the periodic table below.
When carbon-14 decays, it emits a beta particle to produce nitrogen-14, as shown below.
When copper-67 undergoes beta decay, which of the following isotopes is produced?
A.
copper-66
B.
copper-68
C.
nickel-67
D.
zinc-67
_____ 10. An equation is shown below.
Which kind of reaction does the equation represent?
A. alpha decay
B. beta decay
C. nuclear fission
D. nuclear fusion
_____ 11. The three main types of nuclear radiation are alpha, beta, and gamma. Which of the
following lists these types of radiation from highest penetrating power to lowest
penetrating power?
A. alpha, gamma, beta
C. beta, gamma, alpha
B. beta, alpha, gamma
D. gamma, beta, alpha
_____ 12. Which of the following is an example of nuclear fusion?
A. Hydrogen-1 and hydrogen-2 combine to form helium-3.
B. Polonium-210 decays into lead-206 and an alpha particle.
C. Carbon-14 breaks down into a beta particle and nitrogen-14.
D. Uranium-235 and a neutron produce barium-141, krypton-92, and three neutrons.
States of Matter
_____ 1. 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.
_____ 2. 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.
_____ 3. The 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.
B. density.
C. mass.
D. number of molecules
_____ 4. The air inside a beach ball is at a temperature of 25°C and a pressure of
1.0 atm. If the ball contains 0.85 mol of air, what is its volume?
A. 1.7 L
C. 21 L
B. 6.1 L
D. 27 L
_____ 5. Oxygen (O2) and nitrogen (N2) molecules are contained in a flask, which is
separated from a second flask by a closed valve as shown below. The second
flask, of equal volume, is a vacuum.
The valve separating the two flasks is opened. Which of the following diagrams
represents the most likely arrangement of molecules after the valve is opened?
A.
B.
C.
D.
_____ 6. Assuming pressure is held constant, which of the following graphs shows how the
volume of an ideal gas changes with temperature?
A.
C.
B.
D.
_____ 7. Which of the following is not true of a sample of gas as it is heated in a rigid,
closed container?
A. The pressure of the molecules increases.
B. The average speed of the molecules increases.
C. The average distance between molecules increases.
D. The number of collisions between molecules increases.
_____ 8. Which of the following gas particles would effuse most quickly if all are held at
the same temperature and pressure?
A. NH3
B. He
C. O2
D. H2
_____ 9. What is the volume of 1 mole of hydrogen gas (H2) at standard temperature
and pressure?
A. 1.0 L
B. 2.0 L
C. 22.4 L
D. 44.8 L
_____ 10. Which of the following gases behaves most like an ideal gas at STP?
A. C3H8
B. NH3
C. He
D. HCl
_____ 11. A cylinder of gas particles is shown below.
The cylinder is fitted with a moveable piston that can be raised and lowered. Which of the
following would result in an increase in the pressure of the gas below the piston?
A. Increasing the volume of the cylinder
B. Removing some of the gas from the cylinder
C. Decreasing the volume of the cylinder
D. Decreasing the pressure outside the cylinder
_____ 12. The picture below shows a gas at standard conditions in a container with a
moveable piston.
According to Charles’ Law, what will happen to the piston when the gas is heated?
A. The piston will move up because the gas particles get larger.
B. There will be no change because the temperature change will not affect the system.
C. The piston will move up because the gas particles will move faster and get farther apart.
D. The piston will move down because the gas particles will move more slowly and get
closer together.
_____ 13. The pressure exerted by a gas is due to the
A. chemical nature of the container
B. diameter of the gas molecules
C. color of the gas
D. collisions of the gas molecules with the walls of the container
_____ 14. An assumption of the kinetic molecular theory is that the particles of a gas
experience _____________ collisions and have a(n) _____________ volume.
A. elastic... significant
B. inelastic... significant
C. elastic... insignificant
D. inelastic... insignificant
_____ 15. Real gases behave most like ideal gases under conditions of
A. high temperature and high pressure
B. high temperature and low pressure
C. low temperature and high pressure
D. low temperature and low pressure
_____ 16. A sample of nitrogen (N2) gas in a 10.0 L container has a pressure of 1.0 atm at 297
K. Assuming ideal gas behavior, what will the pressure be if the same amount of
nitrogen gas is put into a 5.0 L container at 297 K?
A. 0.40 atm
C. 2.0 atm
B. 0.50 atm
D. 2.5 atm
_____ 17. The illustrations below represent the expansion of a gas in a cylinder of an engine.
The piston moves as the gas volume changes.
What could have been done to the gas in the cylinder to bring about this change in volume?
A. Half of the molecules were released.
B. The Kelvin temperature was doubled.
C. The condensation rate for the gas was doubled.
D. The amount of heat in the gas was reduced by one half.
_____ 18. The four tanks shown in the diagram below contain compressed nitrogen gas. The
temperature of the gas is the same in each tank.
Which of the tanks contains the greatest number of gas particles?
A. tank 1
B. tank 2
C. tank 3
D. tank 4
_____ 19. Which is not an assumption of the kinetic-molecular theory?
A.
B.
C.
D.
Matter is composed of tiny particles.
The particles of matter are in continual motion.
The particles of a gas are far relatively apart.
When individual gas particles collide, the collision is inelastic.
_____ 20. How many liters of O2 are used if 35 g of H2O are produced according to the
following equation:
2H2(g) + O2(g)  2H2O(g)
A. 2.18 L
B. 392 L
C. 21.8 L
D. 39.2 L
______1. A crystal of table salt (NaCl) is dissolved in water. Which of the following
statements explains why the dissolved salt does not recrystallize as long as the
temperature and the amount of water stay constant?
A. Na+ and Cl– ions lose their charges in the water.
B. Water molecules surround the Na+ and Cl– ions.
C. Na+ and Cl– ions leave the water through vaporization.
D. Water molecules chemically react with the Na+ and Cl– ions.
______2. A pharmacist mixes together 20 g of crystals of compound A and 10 g of crystals of
compound B. The mixture is then dissolved in 120 mL of water to make cough syrup. The
mixture will most likely dissolve fastest under which of the following sets of conditions?
A.
C.
B.
D.
______3. Which of the following solutes can most likely dissolve in the solvent CF4?
A. NaBr
B. NaCO3
C. Cl2
D. HCl
______4. When stirred in 30°C water, 5 g of powdered potassium bromide, KBr, dissolves
faster than 5 g of large crystals of potassium bromide. Which of the following best explains
why the powdered KBr dissolves faster?
A. Powdered potassium bromide exposes more surface area to water molecules than
large crystals of potassium bromide.
B. Potassium ions and bromide ions in the powder are smaller than potassium ions
and bromide ions in the large crystals.
C. Fewer potassium ions and bromide ions have been separated from each other in the
powder than in the crystals.
D. Powdered potassium bromide is less dense than large crystals of potassium
bromide.
______5. Which of the following will be electrically conductive?
I. Salt dissolved in water
II. Pure water
III. Pure solid salt
A. I only
B. I and II only
C. I and III only
D. I, II, and II
______6. A mug of hot chocolate (hot chocolate mix and water) has a powdery residue on the
bottom of the mug after being thoroughly mixed. The hot chocolate solution is most likely
A. saturated
C. supersaturated
B. unsaturated
D. none of the above
______7. How many grams of KCl are dissolved in 2.00 L of a 0.200 M solution of KCl?
A. 0.400 g
C. 29.8 g
B. 14.9 g
D. 400 g
______8. Which of the following solutions has the highest concentration of solute?
A. 1.0 mol solute in 200 mL solvent
B. 2.0 mol solute in 500 mL solvent
C. 3.0 mol solute in 1 L solvent
D. 4.0 mol solute in 1.5 L solvent
______9. A person left a bottle of distilled water and a bottle of a sugary drink outside
overnight. In the morning, one liquid was frozen but the other was not. Which liquid was
frozen and why did it freeze?
A. The sugary drink froze because solutions are more dense than pure substances.
B. The distilled water froze because pure substances are more dense than solutions.
C. The sugary drink froze because solutions have a higher freezing point than pure
substances.
D. The distilled water froze because pure substances have a higher freezing point than
solutions.
______10. The diagram below shows gas inside a sealed container before and after force is
applied to the container’s movable piston. The temperature inside the container remains
the same after the force is applied.
Applying force to the piston results in compression of the gas particles and an increase in
gas pressure. Which of the following statements best describes the change in gas particles
after compression?
A.
B.
C.
D.
The kinetic energy of the gas particles increases.
The kinetic energy of the gas particles decreases.
The velocity with which the gas particles hit the container wall increases.
The frequency with which the gas particles hit the container wall increases.
______11. What is the volume of 0.25 mol of ammonia gas (NH3) at 1.00 atm and 273 K?
A. 0.33 L
B. 5.6 L
C. 95 L
D. 1530 L
______12. A picture of a balloon is shown below.
If the temperature of this balloon were to decrease suddenly, how would the balloon
change?
A. Its mass would increase.
B. Its mass would decrease.
C. Its volume would increase.
D. Its volume would decrease.
______13. At the same temperature and pressure, which gas would effuse most quickly?
A. Oxygen (O2)
B. Neon (Ne)
C. Chlorine (Cl2)
D. Xenon (Xe)
______14. What is one difference between a real gas and an ideal gas?
A. Real gases expand to fill their containers, while ideal gases remain compressed.
B. Real gases particles have some attractive forces, while ideal gas particles have no
attractive forces.
C. Real gas particles have no volume, while ideal gas particles do take up space.
D. Real gases have low density, while ideal gases are very dense.
Acids & Bases
_____ 1. Which of the following characteristics allows blood to resist changes in pH?
A. acidity
C. buffering capacity
B. basicity
D. clotting factors
_____ 2. Calcium hydroxide, Ca(OH)2, is used as a soil conditioner in home gardens. When
mixed with water, it releases hydroxide ions. Which of the following is the most likely pH
for a solution of calcium hydroxide and water?
A. 1
C. 7
B. 3
D. 10
_____ 3. The compound Mg(OH)2 is classified as an Arrhenius base because, when the
compound dissolves in water, there is an increase in the concentration of which of the
following ions?
A. hydrogen ions
C. magnesium ions
B. hydroxide ions
D. oxide ions
_____ 4. The pH of four different solutions of common materials is measured. Which of the
following lists the solutions in order from most acidic to most basic?
A. battery acid, lemon juice, blood,laundry detergent
B. lemon juice, battery acid, blood,laundry detergent
C. laundry detergent, blood, lemonjuice, battery acid
D. battery acid, blood, laundry detergent, lemon juice
_____ 5. In the reaction of hydrobromic acid (HBr) and ammonia (NH3), ammonia acts as a
Brønsted base. Which of the following ions is formed?
A. N+
C. NH2–
B. NH2–
D. NH4+
_____ 6. The formula for carbonic acid is H2CO3, and the formula for hydrogen carbonate is
HCO3— Together they form a buffer that is found in blood. Which of the following reactions
represents what happens when excess base enters the
bloodstream?
A. HCO3–(aq) + H3O+(aq) → H2CO3(aq) + H2O(l)
B.
H2CO3(aq) + OH–(aq) → HCO3–(aq) + H2O(l)
C.
HCO3–(aq) + H2O(l) → H2CO3(aq) + OH–(aq)
D.
H2CO3(aq) + H2O(l) → HCO3–(aq) + H3O+(aq)
_____ 7. Which of the following statements explains how a buffer maintains pH when small
amounts of a strong base are added?
A. The salt in the buffer absorbs the base.
B. The water in the buffer dilutes the base.
C. The weak acid of the buffer neutralizes the base.
D. The weak base of the buffer neutralizes the base.
_____ 8. A chemical equation representing the reaction of water (HOH) and ammonia (NH3)
is shown below.
Which of the following statements best explains the chemical action of the reactants in this
equation?
A. Both water and ammonia are acting as acids.
B. Both water and ammonia are acting as bases.
C. Water is acting as an acid, and ammonia is acting as a base.
D. Water is acting as a base, and ammonia is acting as an acid.
_____ 9. The equation below shows ammonia dissolving in water.
Why is water considered an acid when ammonia is dissolved in it?
A. Water acts as a proton donor.
B. Water acts as a proton acceptor.
C. Water contains hydrogen atoms.
D. Water has a 2:1 ratio of hydrogen to oxygen.
_____ 10. Which of the following substances has the highest concentration of hydrogen ions
in solution?
A. bleach – pH 13
C. tomato juice – pH 4
B. water – pH 7
D. vinegar – pH 3
_____ 11. The table below contains data for water samples from four sources.
Nancy analyzed water samples from several sources: rainfall, a nearby creek, a swimming
pool, and her kitchen faucet. She recorded her data in the table. Which sample was most
acidic?
A. rain
C. pool
B. creek
D. faucet
_____ 12. Sodium hydroxide (NaOH) is a strong base. The dissociation of NaOH in an
aqueous solution is given below.
NaOH(aq) → Na+(aq) + OH−(aq)
According to the Arrhenius theory, why is sodium hydroxide a base?
A. NaOH is a neutralizer.
B. NaOH is a proton acceptor.
C. NaOH is a hydroxide ion donor.
D. NaOH is an electron pair provider.
_____ 13. The equation below represents the reaction of hydrogen iodide with water.
HI + H2O → H3O+ + I−
Which reactant in this equation acts as a Brønsted base?
A. HI
B. H2O
C. H3O+
D. I-