Download Chemistry Final Review- Level 2

Chemistry Final Review- Level 2
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
B. a filtration process
C. a chemical reaction
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
____
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.
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
_____ 4. Which of the following elements can form an anion that contains 54 electrons, 74
neutrons, and 53 protons?
B.
A.
C.
D.
_____ 5. Which of the following represents a pair of isotopes?
A.
1H
and 3H
B.
16O2−
and 19F1−
C.
40K
and 40Ca
D.
16O2−
and 32S2
_____ 6. Which element has the electron configuration 1s22s22p3?
A. boron
B. nitrogen
C. fluorine
D.
phosphorus
_____ 7. How many grams are in 7.80 moles of NaCl?
A. 0.134 g
B. 7.44 g
C. 221 g
D. 452 g
_____ 8. 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.
_____ 9. 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.
_____10. 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.
Periodicity
_____ 1. Which element is considered a metal?
a. hydrogen
b. gold
c. sulfur
d. radon.
_____ 2. Which element in period 2 has the greatest electronegativity?
a. fluorine
b. lithium
c. carbon
d. neon
_____ 3. Mendeleev organized the periodic table by:
a. atomic size b. atomic weight
c. atomic number
d. isotopic weight
_____ 4. A vertical column in the periodic table is known as a(n)
a. octave
b. period
c. group
d. triad
_____ 5. Which of the following elements is a metalloid?
a. antimony (51Sb)
b. calcium (20Ca)
c. sulfur (16S) d. zinc (30Zn)
_____ 6. If one electron was removed from a sodium atom, the new particle could be
represented as
a. Na+
b. Na-
c. Na+3
d. Na-3
_____ 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 element will form an ion whose ionic radius is larger than its atomic radius?
a. fluorine
b. potassium
c. lithium
d. magnesium
_____ 9. The most reactive family of metals is the:
a. inner transition
b. alkaline earth
c. alkali
d. transition
_____ 10. Atoms increase in size down a group because:
a. nuclear charge increases c. a new shell is removed
b. a new shell is added
d. neutrons increase
_____ 11. Based on its position on the periodic table, which of the following elements is a
nonmetal?
a. potassium
b. vanadium
c. nickel
_____ 12. The figure below shows part of the periodic table.
d. bromine
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
_____ 13. The bar graph below represents four elements and their relative atomic numbers.
What would be the most likely positioning of these unknown elements in the periodic
table?
a.
b.
c.
d.
_____ 14. 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
_____ 15. Which of the following sections of the periodic table contains only metals?
a. group 2
b. group 18
c. period 2
d. period 6
Chemical Bonding
____1. Which of the following 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. What is the percent mass of oxygen in acetone (C3H6O)?
A) 1.00%
_____
B) 10.3%
C) 27.6%
D) 62.0%
10. This substance is held together by metallic bonds
A)
B)
C)
D)
Hydrogen gas, H2
Potassium, K
Aluminum oxide, Al2O3
Bromine, Br2
_____ 1. Covalent bonds are usually formed by the combination of
A. a metal and a nonmetal
B. a metal and a 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. nonpolar
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. The chemical structure of formaldehyde is shown below.
What is the geometry around the carbon atom?
A. bent
B. linear
C. tetrahedral
D. trigonal planar
_____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. Which of the following statements best explains why ice floats on water?
A. Water has a higher specific heat than ice.
B. Ice has the same molecular mass as water.
C. Heat is absorbed when water changes from the solid state to the liquid state.
D. Hydrogen bonding causes water to be less dense in the solid state than in the liquid
state.
_____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)
B. nitrogen (N)
C. oxygen (O)
D. fluorine (F)
_____13. 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 Reactions & 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
B. 2:2:3
C. 3:1:2
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
B. 3
C. 6
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
_____ 8. When pure N2O5 is heated under certain conditions, O2 and NO2 are produced. What
type of reaction is this?
A. combustion
B. decomposition
C. double displacement
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
B. 33 g
C. 87 g
D. 240 g
_____10. Which of the following diagrams represents a single displacement (replacement)
reaction?
A.
B.
C.
D.
_____12. The following reaction can be identified as a ____________ reaction:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
A.
B.
C.
D.
Synthesis
Double-Replacement
Combustion
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
B. physical change
C. radioactive decay
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
B. 25 g
C. 33 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
C. nuclear fission
B. beta decay
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.
Gases
_____ 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
B. 6.1 L
C. 21 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.
C.
B.
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. 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
_____ 9. 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
_____ 10. 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.
_____ 11. 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
_____ 12. 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
_____ 13. 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.
_____ 14. 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
Solutions
______ 1. Which statement below best describes what happens when sodium chloride, NaCl,
is dissolved in water?
A. The NaCl separates into Na+ and Cl– ions.
B. The NaCl separates into uncharged Na and Cl.
C. The NaCl reacts with water to form NaH and HCl.
D. The NaCl reacts with water to form NaOH and Cl2.
______2. 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
______3. 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
______4. 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.
______5. 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.
______6. Which of the following samples of sugar will dissolve fastest in a pitcher of
lemonade?
A. 5 g of cubed sugar in 5°C lemonade
B. 5 g of cubed sugar in 20°C lemonade
C. 5 g of granulated sugar in 5°C lemonade
D. 5 g of granulated sugar in 20°C lemonade
______7. The table below gives information about four aqueous solutions of sodium
nitrate (NaNO3).
In which beaker will an additional 10 g of sodium nitrate (NaNO3) dissolve at the slowest
rate?
A. 1
B. 2
C. 3
D. 4
______8. A solution that contains less solute than it can hold at a given temperature is
A. disassociated.
C. supersaturated.
B. saturated.
D. unsaturated.
______9. 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.
______10. 5 g of sugar are poured into 25 mL of water and stirred. In this example, sugar is
the
A. Solution
C. Solvent
B. Solute
D. Mixture
Acids & Bases
_____ 1. 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
B. 3
C. 7
D. 10
_____ 2. 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
B. hydroxide ions
C. magnesium ions
D. oxide ions
_____ 3. 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
_____ 4. 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+
_____ 5. 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.
_____ 6. 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.
_____ 7. Which of the following substances has the highest concentration of hydrogen ions in
solution?
A. bleach – pH 13
B. water – pH 7
C. tomato juice – pH 4
D. vinegar – pH 3
_____ 8. 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
B. creek
C. pool
D. faucet
_____ 9. 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.
_____ 10. 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-