Download end of year review

Chemistry Final Review
Level 4
Matter & Energy
_____ 1. Pure substances include
a. elements only.
b. compounds and mixtures.
c. elements and compounds.
d. elements and mixtures.
_____ 2. Which of the following is NOT TRUE about a mixture? The components of a mixture
a. can be in any ratio.
c. Retain their original identifying properties.
b. can be separated easily.
d. Chemically unite to form one substance.
_____ 3. 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
_____ 4. 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.
_____ 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. Block X and Block Y have the same mass. Both blocks are placed into a container of
pure water. Block X floats in the water, and Block Y sinks to the bottom of the container.
Which of the following statements is an accurate conclusion from this demonstration?
a. Block Y is heavier than Block X.
b. Block Y is less dense than Block X.
c. Block Y has a smaller volume than Block X.
d. Block Y would float if more water were added.
_____ 7. The table below shows the physical properties of selected metals.
Physical Properties of Selected Metals
Metal
Molecular Melting Boiling Density
mass
point
point (g/cm3)
(amu)
(°C)
(°C)
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
_____ 8. 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.
_____ 9. Which of the following best describes exothermic reactions?
a. They always release heat.
c. They never release heat.
b. They always occur spontaneously.
d. They never occur spontaneously.
_____ 10. Which of the following substances is made of particles with the highest average
kinetic energy?
a. Fe (s) at 35 C.
b. Br2 (l) at 20 C
c. H2O (l) at 30 C
d. CO2 (g) at 25 C
_____ 11. 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
_____ 12. The normal boiling point of water is
a. 373 K.
b. 173 K.
c. 273 K.
d. 473 K.
_____ 13. When water freezes, each gram loses an amount of heat equal to its heat of
a. fusion.
b. vaporization.
c. sublimation.
d. reaction.
_____ 14. Which temperature represents absolute zero?
a. 0 K
b. 0 °C
c. 273 K
d. 273 °C
Atomic Structure
______ 1. The number of protons in a neutral atom is always equal to
A. the number of neutrons in the atom
B. the number of electrons in the atom
C. the atomic number of the atom
D. B and C
E. All of the above
____ 2. In principle energy level 2, what sublevels will be present?
A. s
B. s & p
C. s, p, & d
D. s, p, d, & f
_____3. 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.
_____ 4. Which of the following elements can form an anion that contains 54 electrons, 74
neutrons, and 53 protons?
A.
B.
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−
_____ 6. What is the mass of 5.5 moles of H2?
A. 11 g
B. 5.5 g/mol
C. 5.5 g
D. 11 g/mol
_____ 7. How many moles of He are in 7.88 x 1015 particles of He?
A. 4.74 x 1038 mol
B. 6.02 x 1023 mol
C. 1.31 x 10-8 mol
D. 7.64 x 108 mol
_____ 8. Which of the following comparisons correctly describes subatomic particles?
A.
B.
C.
D.
An electron has a negative charge and a mass larger than the mass of a proton.
A neutron has a negative charge and a mass smaller than the mass of a proton.
A neutron has a neutral charge and a mass larger than the mass of an electron.
A proton has a positive charge and a mass smaller than the mass of an electron.
_____9. An element with an atomic number of 26 has how many electrons in the 3d
sublevel?
A. 0
B. 2
C. 6
D. 8
E. 10
_____10. What is the correct orbital diagram for the outer electrons of chromium in the
ground state?
and 32S2
A.
B.
C.
D.
E.
3s: ↑↓ 2d: ↑_ ↑_ ↑_ ↑_ __
4s: ↑↓ 3d: ↑↓ ↑↓ __ __ __
4s: ↑↓ 3d: ↑_ ↑_ ↑_ ↑_ __
4s: ↑_ 3d: ↑_ ↑_ ↑_ ↑_ ↑_
5s: ↑_ 4d: ↑_ ↑_ ↑_ ↑_ ↑_
_____11. An electron ________ energy when it __________ from the _________.
A.
B.
C.
D.
absorbs; falls to the excited state; ground state
emits; jumps to the excited state; ground state
emits; falls to the ground state; excited state
absorbs; jumps to the ground state; excited state
_____12. 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
_____13. 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.
_____14. Which of the following describes a particle that contains 36 electrons, 49 neutrons,
and 38 protons?
A. An ion with a charge of 2B. 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
_____15. In going from 1s22s22p63s23p64s1 to 1s22s22p63s23p54s2, an electron would
A.
B.
C.
D.
E.
absorb energy
emit energy
relax to the ground state
bind to another atom
undergo no change in energy
_____16. The total number of electrons that can be accommodated in the fourth principal
energy level is ______.
A. 2
B. 8
C. 18
D. 32
E. 50
_____ 17. In a neutral atom of argon-40, the number of protons
A. equals the number of electrons
B. equals the number of neutrons
C. is less than the number of electrons
D. is greater than the number of neutrons
_____18. What is the energy (in Joules) of a photon that has a frequency of 4.00 x 1010 Hz?
A. 1.99 x 10-25 J
B. 2.65 x 10-23 J
C. 7.50 x 10-3 J
D. 1.20 x 1019 J
E. 6.02 x 1023 J
_____19. 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.
1.
2.
STATEMENT I
An element with an atomic
number of X and a mass number of
N has X-N neutrons
BECAUSE
T
F
All atoms of an element are
identical.
T
STATEMENT II
Elements have more neutrons than
protons
T
F
All atoms of the same element have the
same number of protons.
BECAUSE
F
T
F
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
d. isotopic weight
_____ 4. 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)
_____ 5. 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
_____ 6. 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
_____ 7. Which element will form an ion whose ionic radius is larger than its atomic radius?
a. fluorine
b. potassium
c. lithium
d. magnesium
_____ 8. The most reactive member of the alkali metals is
a. potassium
b. rubidium
c. cesium
d. francium
_____ 9. The dominant factor in determining the variation in size of successive atoms in a
group is the
a. increase in nuclear charge
c. decrease in the number of electron shells.
b. addition of an energy level
d. increase in the number of neutrons
_____ 10. Which ion has the largest ionic radius?
a. F-
b. Cl-
c. Br-
d. I-
_____ 11. Which element has the greatest first ionization energy?
a. aluminum
b. calcium
c. phosphorous
d. sodium
_____ 12. Which of the following elements has the highest electronegativity?
a. B
b. C
c. O
d. N
_____ 13. Which of the following is the same for both hydrogen and potassium?
a. atomic mass
c. number of valence electrons
b. total mass of neutrons
d. number of filled energy levels
_____ 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. 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
_____ 16. 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
_____ 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
c. boiling point of the liquid
b. density of the solid
d. number of protons in each atom
_____ 19. Which of the following elements has the lowest electronegativity?
a. Cesium
d. Barium
b. Strontium
e. Potassium
c. Calcium
_____20. Which element has the greatest ionization energy?
a. Chlorine
d. Phosphorus
b. Oxygen
e. Fluorine
c. Sulfur
_____21. Which of the following elements is most electronegative?
a. S
d. Mg
b. Cl
e. P
c. Na
23.
STATEMENT I
On the periodic chart, atomic
radius increases from left to right
STATEMENT II
The number of protons is increasing.
24.
25.
T
F
Sulfur chemically resembles
oxygen
T
F
The ionization energy generally
increases as you move from left to
right across the periodic table
T
26.
BECAUSE
BECAUSE
BECAUSE
F
The fluoride ion has a larger
radius than the fluorine atom.
T
F
T
F
They are in the same period.
BECAUSE
T
F
Effective nuclear charge increases as you
move from left to right across the
periodic table
T
F
The fluoride ion has 8 electrons and 9
protons
T
F
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 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
_____ 3. 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.
_____ 4. 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
_____ 5. 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.
_____ 6. 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
_____ 7. Which of the following has the same empirical formula as dimercury (II)
acetate, Hg2(C2H3O2)2?
A)
B)
C)
D)
Mercury (I) bicarbonate, HgHCO3
Dimercury (II) bicarbonate, Hg2(HCO3)2
Mercury (I) acetate, HgC2H3O2
Mercury (I) oxalate, Hg2C2O4
_____ 8. Which of the following properties is not associated with metallic elements?
A) malleability
B) brittleness
C) ductility
D) conductivity
_____ 9. What is the percent composition by weight of Ag in Ag(NH3)2+?
A) 4
B) 20
C) 76
D) 80
E) 96
_____10. Which of the following compounds contains the greatest percentage of oxygen by
weight?
A) C3H6O5Cl B) C3H6O2
C) C5H10O5
D) C4H8O3
E) They are all equal
_____11. What is the mass of nitrogen in a 50.0 g sample of sodium nitrite (NaNO2)?
A) 20.2 g
B) 16.4 g
C) 10.1 g
D) 8.23 g
E) 23.4 g
_____12. This substance is held together by metallic bonds
A) Hydrogen gas, H2
B) Carbon monoxide, CO
C) Potassium, K
D) Aluminum oxide, Al2O3
E) Bromine, Br
_____13. This holds a sample of barium iodide, BaI2, together
A) Hydrogen bonding
B) Ionic bonding
C) Metallic bonding
D) Nonpolar covalent bonding
E) Polar covalent bonding
_____14. Element X has an electron configuration of 1s22s22p63s2. Element X will most likely
form oxides with the formula
A) X2O
B) X2O3
C) XO
D) XO2
_____15. An oxide of arsenic contains 65.2% arsenic by weight. What is its simplest
formula?
A) AsO
B) As2O3
C) AsO2
D) As2O5
E) As2O
_____16. A certain mass of sulfur required 16 grams of oxygen to be converted into sulfur
dioxide, SO2. If this same mass of sulfur were to be converted into sulfur trioxide, SO3, the
mass of oxygen required would be
A) 4.0 g
B) 8.0 g
C) 12 gD) 24 g
E) 32 g
For questions 17-19, refer to the following formulas:
Answer Choice
A
B
C
D
E
Formula
AB
AB2
A2B
AB3
A2B3
Which of the above represents the formula for the most common compound of A and B,
where A and B represent given pairs of elements or polyatomic ions as indicated below?
Question
#
17.
18.
19.
28.
29.
30.
Na and Cl will form an ionic
bond
T
F
The most important factor in
determining the chemical
properties of an element is
the number of electrons in
the outermost shell
T
F
Magnesium fluoride is an
example of ionic bonding
T
F
A
B
Be
F
NH4+
Cl
Al
O
BECAUSE
BECAUSE
BECAUSE
Cl donates an electron to Na
T
F
The number of electrons in the
outer shell determines the
bonding characteristics of that
element
T
F
Magnesium and fluorine have
the same electronegativity
T
F
_____ 1. 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
_____ 2. The chemical formula for ammonia is NH3. Which of the following is the correct
Lewis electron dot structure for ammonia?
A.
C.
B.
D.
_____ 3. 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
_____ 4. When elements combine to form compounds:
A. only the outermost electrons of the atoms are involved
B. all the electrons of the atoms are involved
C. the protons and electrons are involved
D. only the protons are involved.
_____ 5. Which is an example of a non-polar molecule that contains polar covalent bonds?
A. CCl4
B. N2
C. H2S
D. NH3
_____6. 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.
_____7. Palmitic acid, a component of most animal fats, has the molecular formula
CH3(CH2)14COOH. Which of the following is the empirical formula for palmitic acid?
A. CHO
B. C3H6O2
C. C8H16O
D. C16H32O2
_____8. Which molecule is incorrectly matched with the molecular geometry?
Molecule Molecular geometry
A.
CO2
bent
B.
CH4
tetrahedral
C.
SO3
trigonal planar
D.
SiCl4
tetrahedral
E.
PH3
trigonal pyramidal
The following choices are for Questions 9-12
A. H2
B. O2
C. N2
D. CO2
E. NH3
_____9. Which molecule is polar?
_____10. Which molecule contains a triple bond?
_____11. Which molecule has no unshared electron pairs?
_____12. Which molecule has trigonal pyramidal molecular geometry?
_____13. 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.
_____14. The Lewis dot structure of a compound is shown below.
Which of the following elements does X represent in the structure?
A.
B.
C.
D.
carbon (C)
nitrogen (N)
oxygen (O)
fluorine (F)
_____15. Which of the following molecules contains both ionic and covalent bonds?
A. C6H14
B. MgCl2
C. (NH4)2SO4
D. H2O
E. C2H4
_____16. Which of the following elements does not form a diatomic molecule?
A. Oxygen
B. Nickel
C. Bromine D. Hydrogen
_____17. Cyclohexylbenzene has the empirical formula C3H4. Its molar mass is 160.3 g/mol.
Its molecular formula is
A. C3H4 B. C6H8
C. C9H12
D. C12H16
______ 18. The correct name of NO3 is
A. nitrate
B. nitrite
C. nitrogen trioxide
D. mononitrogen trioxide
E. nitrogen oxid
_____19. 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.
_____20. All of the following have covalent bonds EXCEPT
A. HCl B. CCl4
C. H2O
D. CsF
E. CO2
_____21. Which of the following is (are) the WEAKEST attractive force(s)?
A. Dipole-dipole forces
D. Polar covalent bonding
B. Coordinate covalent bonding
E. Ionic bonding
C. Covalent bonding
_____22. Which of the following pairs of compounds can be used to illustrate the Law of
Multiple Proportions?
A. NO and NO2
D. NH3 and NH4Cl
B. CH4 and CO2
E. H2O and HCl
C. ZnO2 and ZnCl2
Statement I
A nonpolar molecule can have polar
bonds.
33.
BECAUSE
BECAUSE
T
F
The bond in an O2 molecule is
nonpolar.
34.
T
F
Water is a polar molecule.
35.
T
F
BECAUSE
BECAUSE
Statement II
Polar bonds can be symmetrically
arranged in a molecule so that there
are no net poles.
T
F
The oxygen atoms in an O2 molecule
share a double bond.
T
F
Water has polar bonds.
T
F
Chemical Reactions & Stoichiometry
_____ 1. 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%
B. 50.0%
C. 73.5%
D. 85.0%
_____ 2. 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)
_____3. Which of the following diagrams represents a single displacement (replacement)
reaction?
A.
B.
C.
D.
_____4. 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.
_____ 5. What is the sum of the coefficients of the following equation when it is balanced?
C6H12O6 + O2  CO2 + H2O
A) 20
B) 38
C) 21
D) 19
E) 18
_____ 6. What mass of CF4 is formed by the reaction of 8.00 g of methane with an excess of
fluorine?
CH4(g) + 4F2(g)  CF4(g)+ 4HF(g)
A) 19 g
B) 22 g
C) 38 g
D) 44 g
E) 88 g
_____ 7. How many moles of water are formed by a mixture of 100 grams of H2 and 100
grams of O2? (Assume the reaction goes to completion.)
(A) 100/32+100/64
(B) 100 + 2(100/32)
(C) 2(100/32)
(D) 100(100/32)
(E) 200(100/32)
_____ 8. All of the following involve a chemical change EXCEPT
(A) the formation of HCl from H2 and Cl2
(B) the color change when NO is exposed to air
(C) the formation of steam from burning H2 and O2
(D) the solidification of vegetable oil at low temperatures
(E) the odor of NH3 when NH4Cl is rubbed together with Ca(OH)2 powder
_____ 9. When hydrocarbons burn, the products include carbon dioxide and
(A) oxygen
(D) hydroxide
(B) hydrogen
(E) hydrogen peroxide
(C) water
_____ 10. A balanced equation is shown below.
C6H12O6(l)  2C2H5OH(l) + 2CO2(g)
Which of the following statements correctly compares the mass of the reactant with the
mass of the products in this equation?
A. The mass of the reactant is half the mass of the products.
B. The mass of the reactant is twice the mass of the products.
C. The mass of the reactant is one-fourth the mass of the products.
D. The mass of the reactant is the same as the mass of the products.
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. 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
_____ 6. 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
B. beta, alpha, gamma
C. beta, gamma, alpha
D. gamma, beta, alpha
_____ 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. 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.
_____ 10. Consider the following reaction:
60
60
27 Co  Z X + e
What is the value of Z in the beta decay reaction below:
(A) 25
(B) 26
(C) 27
(D) 28
(E) 29
_____ 11. Element 102
20  is formed as a result of 3  and 2  decays. Which of the following is
the parent element?
A. 90
B. 114
C. 114
D. 128  90
16 
28 
12 
24 
_____ 12. A physicist starts out with 320 grams of a radioactive element Z and after 20
minutes he has only 20 grams left. What is the half-life of element Z?
A. 2 minutes
C. 4 minutes
E. 10 minutes
B. 3 minutes
D. 5 minutes
_____ 13. What is the daughter element produced by technetium-99 (atomic number-43;
mass number-99) after gamma decay?
99
A. 98
C. 44
E. 99
Ru
43Tc
43Tc
B.
99
42
D.
Mo
95
41
Nb
Questions 14-15 refer to the following equation:
14
14
6C  7 N + X
_____ 14. What is X?
A. 24 He
C. 10 e
B.
0
1
E. 126C
D. 11 H
e
_____ 15. This nuclear reaction is an example of:
A.  decay
C. + decay
B.  decay
D. fusion
16.
17.
The “bullet” usually used to initiate
the fusion of 235U is a neutron
T
F
The radioactive decay of
shown as
U
238
92
T
18.
U can be
238
92
Th He
234
90
4
2
BECAUSE
F
Nuclear fusion in the sun converts
hydrogen to helium with a release of
energy
T
BECAUSE
F
BECAUSE
E.  decay
Capture of the neutron by the 235U
nucleus causes an unstable condition
that leads to its disintegration.
T
F
The radioactive decay of 238
92 U is
accompanied by the release of a beta
particle.
T
F
Some mass is converted to energy in
solar fusion.
T
F
States of Matter
____ 1. The table below contains data from one trial in an experiment
designed to determine the molar mass of a sample of an unidentified compound
X in the gaseous state.
Trial 1
Mass of gas (g)
6.42
Gas volume (L)
4.48
Density (g/L)
1.43
0.0
Temperature (°C)
Pressure (atm)
1.0
Based on the data gathered in this first trial, what is the molar mass of the
compound?
A. 19.4 g/mol
C. 32.1 g/mol
B. 28.8 g/mol
D. 144 g/mol
_____ 2. 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
_____ 3. 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.
_____ 4. 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
_____ 5. 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.
_____ 6. Which of the following correctly describes molecules of two different
gases if they are at the same temperature and pressure?
A. They must have the same mass.
B. They must have the same velocity.
C. They must have the same average kinetic energy.
D. They must have the same average potential energy.
_____ 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 the molecules increases.
D. The number of collisions between the molecules increases.
_____ 8. The illustration shows a hot air balloon. The pilot can change the
altitude of the balloon by changing the temperature of the gas inside the balloon.
When the gas is heated, the balloon rises.
Which of the following best explains this phenomenon?
A. Heating the gas reduces its pressure.
B. Heating the gas decreases its density.
C. Heating the gas decreases its molecular motion.
D. Heating the gas reduces the frequency of the gas molecules’ collisions.
_____ 9. Boyle’s Law can be used for which of the following?
A. Predicting the expected volume of two party balloons
B. Predicting the relative pressures inside a hot air balloon
C. Predicting the change in volume of an inflatable toy from summer
to winter
D. Predicting the height of a mercury barometer column in a lowpressure system
E. Predicting the change in volume of a party balloon inside a bell
jar as a vacuum is being drawn
_____ 10. If the pressure of a gas sample is doubled at constant temperature, the
volume will be
A. 4 times the original
B. 2 times the original
C. 1/2 of the original
D. 1/4 of the original
E. 1/8 of the original
_____ 11. All of the following statements underlie the kinetic molecular theory of
gases EXCEPT:
A. Gas molecules have no intermolecular forces.
B. Gas particles are in random motion.
C. The collisions between gas particles are elastic.
D. Gas particles have negligible volume.
E. The average kinetic energy directly depends on the volume of the gas.
_____ 12. Which law states that the total pressure of a gaseous mix is equal to the sum of the
partial pressures?
A. Boyle’s Law
B. Charles’s Law
C. Graham’s Law
D. Ideal Gas Law
E. Dalton’s Law
_____ 13. Which law states that volume is directly proportional to temperature?
A. Boyle’s Law
B. Charles’s Law
C. Gay-Lussac’s Law
D. Graham’s Law
E. Dalton’s Law
_____ 14. The gas in a large cylinder is at a pressure of 3,040 torr. Assuming
constant temperature and ideal gas behavior, what volume of this gas could you
compress into a 100 L box at 8 atm?
A. 20 L
B. 200 L
C. 5,000 L
D. 50,000 L
E. 500,000 L
_____ 15. Gas A is at 30°C and gas B is at 20°C. Both gases are at 1
atmosphere. What is the ratio of the volume of 1 mole of gas A to 1 mole
of gas B?
A. 1:1
D. 303:293
B. 2:3
E. 606:293
C. 3:2
_____ 16. A 200 mL flask contains oxygen at 200 mm Hg, and a 300 mL flask
contains neon at 100 mm Hg. The two flasks are connected so that each gas fills the
combined volume of 500 mL. Assuming no change in temperature, what is the
partial pressure of neon after the mixing is complete?
A. 60 mm Hg
D. 150 mm Hg
B. 80 mm Hg
E. 200 mm Hg
C. 100 mm Hg
_____ 17. Which of the following characteristics of gases allows for scuba-diving
tanks to contain large amounts of oxygen?
A. Effusion
B. Diffusion
C. Fluidity
D. Compressiblity
E. Expansion
_____ 18. Which of the following laws explains why a helium balloon tied to a
mailbox would sink lower to the ground once the sun goes down and the
temperature cools?
A. Boyle’s
B. Charles’
C. Gay-Lussac’s
D. Avagadro’s
_____ 19. If one mole of a gas originally at STP is placed in a container where
the pressure is doubled and the temperature in K is tripled, what is the
new volume in L?
A. 2.2
B. 5.6
C. 7.5
D. 11.2
E. 33.6
_____ 20. For a sample of an ideal gas of fixed weight at a fixed temperature,
I. the volume varies directly with the pressure exerted on it
II. the volume varies inversely with the pressure exerted on it
III. the pressure varies directly with the density of the gas
A. I only
B. II only
C. III only
D. I and II
E. II and III
Quest
Statement I
ion #
21.
The ideal gas law does
not hold under low
temperatures and
high pressure
22.
T
F
When an ideal gas is
cooled, its volume will
increase
T
23.
BECAUSE
F
Statement II
interactions between particles
cannot be neglected under
these conditions.
T
BECAUSE
F
In the kinetic theory
of gases, collisions
between gas particles
and the walls of the
container are
considered elastic
T
BECAUSE
BECAUSE
F
temperature and volume are
directly proportional.
T
F
gas molecules are considered
pointlike, volumeless particles
with no intermolecular forces
and in constant, random
motion.
T
F
______1. Which of the following will be electrically conductive?
I. Salt dissolved in water
II. Pure water
III. Pure solid salt
A. I only
B. III only
C. I and II only
D. I and III only
E. I, II, and III
______2. 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
______3. The table below gives information about four aqueous solutions of sodium nitrate
(NaNO3).
Beaker
1
2
3
4
Concentration of NaNO3 (%)
20
20
2
2
Temperature (°C)
0
40
80
100
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
______4. Which of the following solutes and solvents would be expected to form stable
solutions?
I
II
Solute
Ethanol (CH3CH2OH)
Table Salt
Solvent
Water
Water
III
IV
Bromine
Oil
Carbon tetrachloride
Water
A.
B.
C.
D.
E.
I only
I and III only
III only
I, II, and III only
I, II, and IV only
______5. What mass of sodium carbonate, Na2CO3, (molar mass= 106 amu), is needed to
make 120 mL of a 1.5 M solution?
A. 295 g
B. 9.5 g
C. 19 g
D. 589 g
E. 19,000 g
_______6. If you mix 3 liters of 0.5 M NaCl with 9 liters of 0.2777 M NaCl, what will the
concentration of the final solution be, assuming that volumes are additive?
A. 0.33M
B. 0.39M
C. 0.5777M
D. 0.5777m
E. None of the above
______7. 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 higher freezing points than
pure substances.
D. The distilled water froze because pure substances have higher freezing
points than solutions.
______8. Which of the following solutions would probably have the highest boiling point?
A. 0.100 m KOH
B. 0.100 m Na2SO4
C. 0.100 m C6H12O6
D. 0.200 m CaCl2
E. 0.200 m CH3CH2OH
______9. 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.
______10. 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 solute in 1.5 L solvent
_____11. 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
E.
_____12. 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.
______13. 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. The kinetic energy of the gas particles increases.
B. The kinetic energy of the gas particles decreases.
C. The velocity with which the gas particles hit the container wall increases.
D. The frequency with which the gas particles hit the container wall increases.
______14. Which of the following statements explains what happens to the gas inside a
balloon as the external pressure on the balloon decreases and the temperature stays
constant?
A. The volume increases
B. The volume decreases
C. The molecular speed increases
D. The molecular speed decreases
______ 15. 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
Quest
Statement I
BECAUSE
Statement II
ion #
17.
A nonelectrolyte will
BECAUSE
Dissociation is the process of an
dissociate in a polar
ionic compound separating into
solvent.
its component ions when
dissolved.
18.
19.
20.
T
F
A saturated solution
of glucose in boiling
water crystallizes as it
cools
T
F
Salt dissolved in
water depresses the
freezing point
T
F
The rate at which
sugar dissolves in
water increases with
stirring
T
F
BECAUSE
T
F
the solubility increases as the
temperature decreases.
BECAUSE
T
F
the solute particles interfere
with ice crystal formation.
BECAUSE
T
F
stirring exposes the surface of a
solute crystal to a less
concentrated layer of solution.
T
F
Acids & Bases
_____ 1. Which is not an acid-base conjugate pair?
a) HS- and S2b) H3O+ and OHc) HNO2 and NO2d) CH3NH4+ and CH3NH3
e) C6H5COOH and C6H5COO_____ 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
b) 3
c) 7
d) 10
_____ 3. 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)
_____ 4. 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)
b)
c)
d)
NaOH is a neutralizer.
NaOH is a proton acceptor.
NaOH is a hydroxide ion donor.
NaOH is an electron pair provider.
_____ 5. The equation below represents the reaction of hydrogen iodide with water.
HI + H2O --> H3O+ + IWhich reactant in this equation acts as a Bronsted base?
a) HI
b) H2O
c) H3O+
d) I_____ 6. 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
_____ 7. An aqueous solution contains hydronium ion, H3O+, at a concentration of
2.4 x 10-6 M. What is the hydroxide ion, OH-, concentration in this solution?
a) 2.4 x 108 M
b) 2.4 x 10-6 M
c) 1.0 x 10-7 M
d) 4.2 x 10-9 M
_____ 8. What is the pH of a solution of 0.034 M HCl(aq)?
a) -1.47
b) 0.47
c) 1.47
d) 2.54
_____ 9. Which of the following is a proton acceptor?
a) Bronsted-Lowry Acid
d) Bronsted-Lowry Base
b) Arrhenius Acid
e) Lewis Base
c) Lewis Acid
_____ 10. In the reaction of hydrobromic acid (HBr) and ammonia (NH3), ammonia acts as a
Bronsted base. Which of the following ions is formed?
a) N+
b) NH2+
c) NH2d)NH4+
_____ 11. According to 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.
_____ 12. 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 the
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.
_____ 13. The table below shows pH values of some foods.
pH Values of Some Important Foods
Vegetables pH
Citrus
pH
Dairy/Egg pH
Grapefru
Asparagus
5.6
3.2
Butter
6.2
it
Beans
5.5
Lemons 2.3
Cheese
5.6
Peas
6.1
Limes
1.9
Eggs (fresh) 7.8
Spinach
5.4
Oranges 3.5
Milk
6.5
Starches
Bread
(white)
Corn
Crackers
Potatoes
pH
5.5
6.2
7.5
5.8
A patient has chronic indigestion due to an overproduction of stomach acid. Which foods
should the patient avoid until the condition is resolved?
a) vegetables
c) dairy/egg
b) citrus
d) starches
_____ 14. 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