Score _____ / 100 pts.
Name _________________________________
Class _________ Date ____________________
Standard 1 Practice Exam
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1. Scientists know the universe is expanding because of the
a. red shift of distant galaxies.
c. blue shift of distant galaxies.
b. red shift of the brightest galaxies.
d. blue shift of the brightest galaxies.
____
2. According to the big bang theory, the entire universe began as a
a. supernova explosion.
c. dense, hot, supermassive point.
b. cool, dark expansive nebula.
d. scattered band of space dust.
____
3. Which of the following is supportive evidence of the big bang theory?
a. pulsars
c. galactic clusters
b. cosmic microwave background radiation
d. the different shapes of galaxies
____
4. In the diagram below, the spectral lines of hydrogen gas from three galaxies, A, B, and C, are compared to the
spectral lines of hydrogen gas observed in a laboratory.
What is the best inference that can be made concerning the movement of galaxies A, B, and C?
a. Galaxy A is moving away from Earth, but
c. Galaxies A, B, and C are all moving toward
galaxies B and C are moving toward Earth.
Earth.
b. Galaxy B is moving away from Earth, but
d. Galaxies A, B, and C are all moving away
galaxies A and C are moving toward Earth.
from Earth.
____
5. Most of the radiant energy released by the sun results from the process of
a. nuclear fission.
c. combustion.
b. nuclear fusion.
d. electrical generation.
____
6. Elements more massive than ___________ are only made in the last stage of the life of a very large star when it
collapses and explodes in a supernova.
a. cobalt
c. helium
b. mercury
d. iron
____
7. What are the two most abundant elements in a main sequence star?
a. carbon and hydrogen
c. helium and carbon
b. hydrogen and helium
d. carbon and heavy metals
____
____
8. The first subatomic particle discovered was the _____.
a. proton
c.
b. neutron
d.
electron
nucleus
9. The only subatomic particle that does not carry an electric charge is the _____.
a. proton
c. electron
b. neutron
d. nucleus
____
10. One isotope of carbon has 6 protons and 6 neutrons. The number of protons and neutrons of a second isotope of
carbon would be _____.
a. 7 and 6
c. 7 and 7
b. 6 and 7
d. 6 and 6
____
11. Which of the following statements is not a main point of Dalton's atomic theory?
a. All matter is made up of atoms.
b. Atoms are made up of smaller particles.
c. Atoms are indestructible.
d. All atoms of one element are exactly alike, but they are different from atoms of other elements.
____
12. J.J. Thomson used a cathode ray to discover the _____.
a. atom
c. proton
b. electron
d. neutron
____
13. If a scientist studies a beam of particles, and those particles are attracted to a negatively charged plate, the particles
are most likely _____.
a. atoms
c. protons
b. electrons
d. neutrons
____
14. Iodine-131 and iodine-127 are examples of _____.
a. nuclei
c.
b. isomers
d.
isotopes
neutrons
____
15. The experimentation of _____ led to the theory that the atom is a sphere of mostly empty space, with a positively
charged nucleus with electrons around it.
a. Bohr
c. Rutherford
b. Nagaoka
d. Thomson
____
16. The _____ is where the electron is most likely to be found.
a. energy level
c. electron cloud
b. electron orbit
d. orbit
____
17. Which one is comprised of the other three?
a. proton
b. atom
c.
d.
electron
neutron
18. Which one(s) have a mass of 1 amu?
a. electron
b. proton
c.
d.
proton and neutron
electron and proton
19. Each row in the periodic table ends with a _____.
a. metal
b. nonmetal
c.
d.
metalloid
noble gas
____
____
____
20. In going from left to right in any given row in the periodic table, the size of atoms generally _____.
a. increases
c. stays the same
b. decreases
d. changes randomly
____
21. Compared to the neutral atom from which it is derived, a negative ion is _____.
a. always larger
b. always smaller
c. larger in some cases and smaller in others
d. the same size
____
22. The valence configuration shared by carbon, silicon, and germanium is _____.
a. 1s22s22p2
c. s2p2
2
6
b. 2s 2p
d. s2p4
____
23. Alkaline earth metals lose _____ electrons to achieve the electron configuration of the noble gas in the preceding
period.
a. one
c. six
b. two
d. seven
____
24. The most unreactive group of elements is the _____.
a. halogens
c.
b. noble gases
d.
alkali metals
transition elements
____
25. Because transition metals have similar atomic radii, transition metals have _____ chemical properties.
a. similar
c. definitely different
b. no
d. identical
____
26. Which category of elements have the property of being malleable and ductile?
a. gases
c. metalloids
b. metals
d. nonmetals
____
27. Which correctly describes elements in the same group?
a. They have the same number of valence electrons.
b. They have electrons in the same outermost energy level.
c. They have the same atomic radius.
d. They must be in the same state of matter.
____
28. Which metalloid is in the fourth period and the same group as Carbon?
a. Silicon
c. Tin
b. Germanium
d. Boron
____
29. Which is the most important characteristic in detemining an element’s chemical properties?
a. the number of protons and neutrons in its nucleus
b. which period it is found in
c. the number of valence electrons it contains
d. its outermost energy level
____
30. Which of the Group 13 elements is not a metal?
a. boron
b. aluminum
c.
d.
gallium
indium
Standard 2 Practice Exam
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1. If a wave has a high frequency, it also has _____.
a. high wavelength and high energy
b. high wavelength and low energy
c. low wavelength and high energy
low wavelength and low energy
d.
____
2. Electron 1 falls from energy level four to energy level two. Electron 2 falls from energy level three to energy level
two. Which electron is more likely to emit red light?
a. 1
b. 2
c. Neither electron could emit red light.
d. Both electrons emit red light.
____
3. Which of the following is the best evidence for the existence of sublevels?
a. large gaps in a spectrum
b. only four lines in a spectrum
c. all colors of light in a spectrum
d. closely spaced lines in a spectrum
____
4. Which of the following orbitals is closest to the nucleus?
a. 1s
c. 4s
b. 2p
d. 3d
____
5. An element is most likely to have properties similar to those of _____.
a. another element in the same period
c. a noble gas
b. another element in the same group
d. a transition element
____
6. In which group of elements would you find an element with two valence electrons?
a. Group 1, the alkali metals
c. Group 17, the halogens
b. Group 2, the alkaline earth metals
d. Group 18, the noble gases
____
7. Which diagram shows a wave with the highest frequency?
a.
b.
____
A
B
c.
d.
C
D
8. Which type of wave has a frequency of approximately 10 8 hertz?
a.
b.
AM radio
ultraviolet light
c.
d.
X-rays
TV, FM Radio, Cell phone
____
9. A photon is emitted from an atom with an energy of 4.25 x 10 -19 J. What is the wavelength of the photon?
a. 467 nm
c. 1.28 x 10-10 m
-4
b. 2.73 x 10 m
d. 6.42 x 1014 m
____
10. Light is released when an electron moves from higher energy levels to a lower energy level. The resulting spectrum
is a(n) _____ spectrum.
a. absorption
c. excitation
b. emission
d. lower energy
____
11. What is the highest occupied sublevel in the structure of an atom of arsenic?
a. 3s
c. 3d
b. 3p
d. 4p
____
12. The ratio of protons to neutrons in stable isotopes of the lighter elements tends to be approximately _____.
a. 1:1
c. 2:1
b. 1:2
d. unpredictable
____
13. The most difficult radiation to block out is _____.
a. alpha particles
b. beta particles
____
14.
a.
b.
and
are examples of _____.
allotropes
isotopes
____
15. When
a. alpha
b. beta
becomes
____
16. A(n) _____ is a high energy electron.
a. beta particle
b. helium nucleus
c.
d.
gamma rays
visible light rays
c.
d.
particles of radiation
tracers
, what type of decay has taken place?
c. gamma
d. positron
c.
d.
alpha particle
positron
____
17. Which is the only type of radiation that might penetrate the walls of a house?
a. alpha
c. gamma
b. beta
d. All will penetrate.
____
18. How much hydrogen-3 will remain after 60 years if the original sample had a mass of 80.0 g and the half-life of
hydrogen-3 is 12 years?
a. 1.25 g
c. 5.00 g
b. 2.50 g
d. 10.0 g
____
19. When one large nucleus is split into two smaller nuclei, the process is nuclear _____.
a. decay
c. fusion
b. fission
d. tracing
20. Two or more nuclei combine to form one larger nucleus in the process of nuclear _____.
a. decay
c. fusion
b. fission
d. tracing
____
____
21. The product of -ray emission from a radioactive isotope of lead is
a. bismuth
c. mercury
b. lead
d. thallium
____
____
____
22. Alpha particles have a _____ charge.
a. -1
b. 0
c.
d.
+1
+2
23. Beta particles have a _____ charge.
a. -1
b. 0
c.
d.
+1
+2
24. Gamma rays have a _____ charge.
a. -1
b. 0
c.
d.
+1
+2
____
25. An electron in an excited energy level, n = 5, can drop to lower energy levels, n = 3 or n = 1. One of these drops
will emit a yellow photon; the other will emit a violet photon. If the electron drops from n = 5 to n = 3, which color
will the photon emit?
a. It will randomly emit both yellow and violet. c. Violet
b. Yellow
d. It is impossible to determine.
____
26. An element emits three photons: orange, red and violet. Which answer lists these photons in order of increasing
energy?
a. red, violet, orange
c. orange, red violet
b. red, orange, violet
d. violet, red, orange
____
27. Which portion of the electromagnetic spectrum is more energetic than visible light?
a. microwaves
c. radio waves
b. x-rays
d. infrared
____
28. What name is given to the discrete packet of energy an electron must absorb to move to a higher energy level?
a. quantum
c. electromagnetic wave
b. frequency
d. poppinstance
____
29. As observed in flame tests, energy needed for an electron to jump from the ground state to the excited state
a. is the same for every element.
c. is unique to each element.
b. is never consistently the same.
d. is impossible to measure.
____
30. You run a flame test on the elements calcium and copper. The calcium produced a reddish orange flame and the
copper produced a bluish green flame. Which element emitted the higher energy photons?
a. copper
c. you cannot tell from the color, you would need
to measure the temperature of the flame.
b. calcium
d. both are comparable.
Standard 3 Practice Exam
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1. The properties of a compound are _____ the properties of the elements that form it.
a. similar to
c. identical to
b. different from
d. derived from
____
2. A colorless, odorless gas combines with a magnetic, metallic element. What can you predict about the product?
a. It will also be magnetic.
b. A gas and a solid produce a liquid.
c. The compound will be shiny and odorless.
d. It is impossible to predict its specific properties.
____
3. Reactions between atoms involve only _____.
a. valence electrons
b. inner electrons
c.
d.
neutrons
protons
____
4. An atom of magnesium has two valence electrons. It will most likely react with one atom of _____, which contains
_____ valence electrons.
a. carbon, 4
c. oxygen, 6
b. nitrogen, 5
d. chlorine, 7
____
5. Oppositely charged ions attract each other, forming a(n) _____ bond.
a. covalent
c. ionic
b. crystal
d. molecular
____
6. The strong crystal structure of an ionic compound is one reason ionic compounds have _____ melting points.
a. high
c. moderate
b. low
d. variable
____
7. The formula for iron(III) oxide, Fe2Cl3, shows that one unit of the compound contains _____ iron atoms.
a. 2
c. 5
b. 3
d. 6
____
8. Lithium has much less attraction for any valence electrons than does fluorine. Atoms of these two elements would
form _____ bonds.
a. covalent
c. crystal
b. ionic
d. molecular
____
9. Two atoms of bromine react with each other to form a(n) _____ bond.
a. covalent
c. crystal
b. ionic
d. molecular
____
____
10. Electron sharing produces _____.
a. crystals
b. ions
c.
d.
molecules
liquids
11. A covalent compound is most likely formed from _____.
a. two metals
c. two metalloids
b. two nonmetals
d. a metal and a nonmetal
____
12. Nitrogen atoms each have five valence electrons. How many pairs of electrons must be shared in a molecule of N 2?
a. 1
c. 4
b. 3
d. 6
____
13. A tug-of-war in which neither side is able to move the other side could be used to model a(n) _____.
a. covalent bond
c. ionic bond
b. crystal
d. transfer of electrons
____
14. A compound has the formula X3Y. For every 15 X atoms present in this compound, how many Y atoms are there?
a. 3
c. 45
b. 5
d. 15
____
15. Which is a physical property of ionic compounds in their solid state?
a. good conductor of electricity
b. weak attractive forces between ions
c. low boiling point
d. high melting point
____
16. Covalent compounds display which of these properties?
a. They are hard, brittle solids
b. They have high melting and boiling points
c. They display luster.
d. Their intermolecular forces are relatively weak.
____
17. The reaction of hydrogen with oxygen produces __________.
a. salt
c. carbon dioxide
b. water
d. sodium
____
18. To achieve a noble gas configuration, oxygen requires __________ electron(s), while hydrogen receives
__________ electron(s).
a. 1,2
c. 2,1
b. 1,1
d. 2,2
____
19. Which is not a property of an ionic compound?
a. high melting point
b. hard
c.
d.
brittle
smooth
____
20. Which of the following is the correct chemical formula for a formula unit of aluminum bromide?
a. AlBr3
c. Al3Br9
b. Al2Br6
d. Al4Br12
____
21. Based on its position in the periodic table, the most likely charge of an iodide ion is _____.
a. 1+
c. 2+
b. 1d. 7-
____
22. Which of the following formulas is incorrect?
a. Al2(SO4)3
b. AlOH3
c.
d.
Ca(OH)2
(NH4)2S
23. The correct name for Fe2S3 is _____.
a. iron(III) sulfide
b. iron sulfide
c.
d.
iron(II) sulfide
iron(I) sulfide
____
____
24. In order to separate two liquids from each other by distillation, they must _____.
a. evaporate at the same temperature
c. both be molecular substances
b. evaporate at different temperatures
d. both be inorganic compounds
____
25. Which is the correct formula for the ionic compound that results from these two atoms?
a.
b.
X2Y5
X5Y2
c.
d.
X2Y3
X3Y2
____
26. Which is the charge that results when oxygen becomes an ion?
a. +2
c. +3
b. -3
d. -2
____
27. Which is the correct formula for the compound formed between beryllium and nitrogen?
a. BeN
c. Be3N2
b. Be3N
d. Be2N3
____
28. Which is the correct name for the compound N2O3?
a. dinitro trioxide
c.
b. nitrogen (II) oxide (III)
d.
____
29. Methane consists of a single carbon atom bonded to four hydrogen atoms. What is the shape of the methane
molecule?
a.
b.
____
trinitrogen dioxide
dinitrogen trioxide
triangular
square
c.
d.
tetrahedral
triangular pyramidal
30. Methane and water have very different properties because methane is __________ and water is __________.
a.
b.
ionic, covalent
covalent, ionic
c.
d.
polar, non-polar
non-polar, polar
Standard 4 Practice Exam
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1. In this equation, which are the reactant(s)?
A
B
C
a.
b.
____
A only
C only
c.
d.
A+B
B only
c.
d.
synthesis
decomposition
2. Which type of chemical reaction is shown?
a.
b.
double replacement
combustion
____
3. In a chemical equation, (aq) after one of the substances means that it is which of these?
a. a solid formed from two ionic substances
b. released as gas
c. dissolved in water
d. produced from nothing
____
4. How many total atoms are in 3Na2SO4?
a. 21
b. 10
c.
d.
24
18
____
5. Which are the product(s) of this reaction? Mg + N2 
a. NMg
c. Mg2N3
b. MgN
d. Mg3N2
____
6. What are the correct coefficients when this equation is balanced? K + Br 2  KBr
a. 1, 1, 1
c. 2, 1, 2
b. 1, 2, 1
d. 2, 1, 1
____
7. What are the correct coefficients when this chemical equation is balanced?
P4 + O2  P2O5
a. 4, 2, 7
c. 2, 5, 4
b. 1, 1, 1
d. 1, 5, 2
____
8. Which element has a molar mass of 30.974 g/mol?
a. Potassium
c.
b. Phosphorus
d.
Gallium
Palladium
9. Which is the mass of 8 moles of sodium chloride?
a. 7.3 grams
b. 468 grams
0.137 g/mole
468 moles
____
____
c.
d.
10. How many moles of atoms are in one mole of AlPO4?
a. 122.1 moles
c.
b. 7.35 x 1025 moles
d.
6 moles
6.022 x 1023 moles
____
____
11. How many molecules are in 3.6 grams of NaCl?
a. 0.06
b. 1.0 X 1021
c.
d.
1.3 x 1026
3.7 x 1022
12. How many grams are in 1.946 moles of NaCl?
a. 113.8 g
b. 30.1 g
c.
d.
0.033 g
44.7 g
____
13. Which change in state requires that energy be added to the substance?
a. freezing
c. condensing
b. deposition
d. vaporizing
____
14. In the Kelvin temperature scale, 0.00 indicates
a. absolute zero
b. the freezing point of water
c.
d.
the freezing point of alcohol
the freezing point of sea water
____
15. The burning of gasoline in an automobile engine is an example of a(n) _____.
a. photosynthesis reaction
c. exothermic reaction
b. endothermic reaction
d. reversible reaction
____
16. The most common form of energy encountered in chemical reactions is _____.
a. electrical energy
c. light energy
b. nuclear energy
d. heat
____
17. In a chemical change, energy can be _____.
a. created, but not destroyed
b. destroyed, but not created
c.
d.
either created or destroyed
neither created nor destroyed
____
18. The two terms below that are identical in meaning are _____.
a. calorie and Calorie
c. Calorie and joule
b. calorie and joule
d. kilocalorie and Calorie
____
19. In a(n) _____ reaction, the products are at a higher energy level than are the reactants.
a. activation
c. endothermic
b. catalytic
d. exothermic
____
20. Even in an exothermic reaction, _____ is needed to get the reaction started.
a. activation energy
c. an endothermic reaction
b. a catalyst
d. an inhibitor
____
21. The energy involved in endothermic and exothermic reactions is _____.
a. chemical
c. light
b. heat
d. electrical
____
22. If the heat of reaction is negative, the reaction is _____.
a. endothermic
c. negative
b. exothermic
d. positive
____
23. If the energy graphs of a reaction, catalyzed and uncatalyzed, are examined, the peak representing activation energy
is _____ for the catalyzed reaction.
a. equal
c. lower
b. higher
d. unchanged
____
24. When methane burns, the main products are
a. oxygen, water, and heat
b. carbon dioxide, water, and heat
c.
d.
oxygen and water
carbon dioxide and water
Sample
1
2
3
4
Gas Samples
Volume of Nitrogen (L)
2.00
1.01
1.98
1.98
Volume of Oxygen (L)
3.02
1.99
4.99
1.02
____
25. Samples of gases containing nitrogen and oxygen were decomposed by electricity and the volume of each gas
produced was measured. Based on these results, which is the correct chemical formula for Sample 2?
a. N2O3
c. NO2
b. NO
d. N2O
____
26. Which is NOT a chemical change?
a. a piece of wood is burned
b. an egg is cooked
c.
d.
a bar of iron rusts
ingredients for a cake are stirred together
____
27. A 25.0-gram sample of magnesium oxide contains 10.8 grams of magnesium. What is the percent of oxygen by
mass in this compound?
a. 89.2 %
c. 43.2 %
b. 14.2 %
d. 56.8 %
____
28. According to the Law of Conservation of Mass, in any chemical change
a. the mass of reactants is greater than the mass of products
b. the mass of reactants is less than the mass of products
c. the mass of reactants is equal to the mass of products
d. the masses of products and reactants have no predictable relationship
____
29. Which is a chemical property of copper?
a. has a boiling point of 2567° C
b. reacts in air to form a green layer
c. has the symbol Cu
d. is a solid at room temperature
____
30. Which of the following substances is an element?
a. Baking soda
b. Iron
c.
d.
Salt
Sugar
Standard 5 Practice Exam
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1. Which of the following statements about a catalyst is true?
a. A catalyst can initiate a reaction.
b. A catalyst can accelerate a reaction.
c. A catalyst can be consumed during a reaction.
d. A catalyst can be changed during a reaction.
____
2. Which of the following factors does NOT affect the rate of the reaction?
a. the amount of the reactants
c. the size of the container used
b. the physical state of the reactants
d. temperature
____
3. A/An _____ lowers the activation energy required for a reaction to take place.
a. catalyst
c. reactant
b. inhibitor
d. product
____
4. A collision requires _____ to be effective.
a. only enough energy
b. favorable orientation
c. enough energy and favorable orientation
d. a reaction mechanism
____
5. If a collision between molecules is very gentle, the molecules are
a. more likely to be oriented favorably.
b. less likely to be oriented favorably.
c. likely to react.
d. likely to rebound without reacting.
____
6. Which of the following factors does not affect the rate of reaction?
a. The physical state of the reactants.
b. The amount of the reactants.
c. The size of the container used.
d. The temperature at which the reaction is carried out.
____
7. Hydrogen iodide (HI) decomposes as
Average rate
–
. The average reaction rate is expressed as:
.
The negative sign used in the rate expression indicates that:
a. There are repulsive forces between the reactants.
b. The concentration of HI decreases with time.
c. The concentration of the reactants is less than that of the product.
d. The reaction rate is decreasing with time.
____
8. A _____ reaction is a chemical reaction that can occur in both the forward and reverse directions.
a. complete
c. reversible
b. forward
d. incomplete
____
9. A decrease in the concentration of reactants causes the rate of the _____ reaction to slow.
a. complete
c. reverse
b. forward
d. incomplete
____
10. In an exothermic reaction, equilibrium shifts _____ when temperature is lowered.
a. toward products
c. to the center
b. toward reactants
d. not at all
____
11. What change can result in a shift in equilibrium during a reaction?
a. change in concentration
c. change in volume and pressure
b. change in temperature
d. all of the above
____
12. Which of the following factors will NOT change the concentration of ammonia (NH 3) in the reaction?
?
a. Decrease in the amount of N2.
c. Decrease in pressure.
b. Increase in the amount of catalyst.
d. Decrease in temperature.
____
13. Which of the following are necessary for successful collisions to occur?
1. Favorable collision geometry
2. Sufficient Kinetic Energy
3. Large H
a. 1 only
c. 2 and 3 only
b. 1 and 2 only
d. 1, 2 and 3
____
14. Collision theory states that
a. All collisions lead to chemical reactions.
b. Most collisions lead to chemical reactions.
c.
d.
Very few reactions involve particle collisions.
Effective collisions lead to chemical reactions.
____
15. Milk is refrigerated in order to slow the rate of decomposition by bacterial action. This decrease in reaction rate is due to
a. a decrease in surface area.
c. a decrease in the fraction of collisions occurring
with sufficient energy.
b. a decrease in the enthalpy change of the reaction. d. the introduction of an alternate pathway to
reaction requiring greater activation energy.
____
16. In general, chemical reactions requiring large activation energy proceed
a. at fast rates.
c. only when entropy increases.
b. at slow rates.
d. only at low concentrations.
____
17. When a lit match is touched to the wick of a candle, the candle begins to burn. When the match is removed, the candle
continues to burn. In this reaction, the match
a. behaves as the catalyst for the reaction.
c. Is part of the rate-determining step.
b. supplies activation energy.
d. lowers the activation energy barrier.
____
18. Consider the following factors related to reaction rate:
1. Reactant particles collide
2. Sufficient kinetic energy is present
3. A favorable geometric orientation occurs in a collision.
4. A catalyst is present.
Which combination of the above factors is required for product formation?
a. 1 only
c. 1, 2 and 3 only
b. 2 and 3 only
d. 1, 2, 3 and 4 are all required.
____
19. Which of the following would change the amount of required activation energy for a heterogeneous reaction?
a. adding a catalyst.
c. changing the reactant concentration.
b. changing the surface area.
d. changing the average kinetic energy.
____
20. ______________________ indicates that forward and reverse reactions are occuring at the same rate.
a. equilibrium
c. reversible
b. synthesis
d. electrochemistry
____
21. An equilibrium constant with a value less than one indicates
a. that more reactants are present at equilibrium.
c. that a reaction will move at a very slow rate.
b. that more products are present at equilibrium.
d. that a reaction has a limiting reactant present.
____
22. In this reaction, which of the following factors will NOT change the concentration of CO 2 in the reaction?
a.
b.
a decrease in the volume of H2O
an increase in the amount of catalyst
c.
d.
a decrease in pressure
a decrease in temperature
____
23. In an exothermic reaction, equilibrium shifts _____ when temperature is raised.
a. to the left
c. to the center
b. to the right
d. none
____
24. In an endothermic reaction, equilibrium shifts _____ when temperature is raised.
a. to the left
c. to the center
b. to the right
d. none
____
25. The equilibrium constant for the reaction
carbon dioxide gas, if
a.
b.
0.16 mol/L
0.0014 mol/L
and
c.
d.
at 700.0 K is 0.44. What is the concentration of
?
0.075 mol/L
0.75 mol/L
____
26. According to Le Chatelier’s principle, stress applied to an equlibrium will cause a shift that alleviates the stress. For which
of the following reversible reactions will the equilibrium shift toward the products when volume is increased?
a.
b.
c.
d.
____
27. In the reaction
heat is evolved. What happens when chlorine (Cl2) is added to the equilibrium
mixture at constant volume?
a. The temperature of the system increases.
b. The temperature of the system decreases.
c. More chlorine is produced.
d. The temperature remains unaffected.
____
28. Which experimental disturbance will produce more
a.
b.
decrease in pressure
addition of more ammonia
(ammonia) in the provided equilibrium?
c.
d.
decrease in volume
addition of a catalyst
____
29. Which statement most accurately describes the state of a reaction after it has reached chemical equilibrium?
a. At chemical equilibrium, equal amounts of
c. The forward and reverse reactions are occurring
products and reactants are present.
at equal rates.
b. The forward and reverse reactions are producing d. At equilibrium, the reaction is continuing in
equal concentrations
either the forward or reverse direction.
____
30. Which of the following would NOT increase the rate of a reaction?
a. an increase in the concentration of the reactants c. and increase in the solvent volume
b. the removal of a catalyst
d. an increase in temperature
Standard 6 Practice Exam
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
____
1. More solute can be dissolved in a _____ solution:
a. saturated
b. supersaturated
c.
d.
2. A substance that dissolves in a solvent is said to be:
a. insoluble
c.
b. immiscible
d.
suspended
unsaturated
miscible
soluble
____
3. A _____ contains the maximum amount of dissolved solute for a given amount of solvent.
a. saturated
c. suspended
b. supersaturated
d. unsaturated
____
4. A _____ solution contains more dissolved solute than a saturated solution at the same temperature.
a. saturated
c. suspended
b. supersaturated
d. unsaturated
____
5. Suppose 8 moles of solute is dissolved in 2 liters of solution. What is the molarity of the solution?
a. 2M
c. 8M
b. 4M
d. 16M
____
6. Which of the following methods for measuring concentration would apply to calculations of changes in colligative
properties?
a. molarity
c. mole fraction
b. molality
d. solubility
____
7. Which of the following approaches would result in an increase in the rate at which salt dissolves in water?
a. reducing the amount of water
c. grinding the salt to a powder
b. increasing the amount of salt
d. lowering the temperature
____
8. If you were to sketch a solution of salt water at the particle level, which of the following descriptions would best
describe what you would draw?
a. water molecules dividing to absorb salt
c. water molecules surrounding individual
molecules
sodium and chloride ions
b. salt crystals changing into water molecules
d. salt ions surrounding water molecules
____
9. Which of the following is a colligative property?
a. hardness
b. osmotic pressure
c.
d.
intermolecular attractions
density
____
10. Which of the following properties would increase with the dissolution of more solute in a solution?
a. boiling point
c. vapor pressure
b. freezing point
d. none
____
11. To calculate the amount number of moles of solute in a solution with its given molarity you should
a. multiply volume by molarity
c. subtract volume from molarity
b. divide molarity by volume
d. use the molar mass of the solute to convert to
grams.
____
____
____
12. Mole fraction is a measure of
a. total moles of solute total moles of solvent
b. total moles of solute or solvent to total moles
of solution
c.
d.
13. Parts per million is a method used to measure
a. very small amounts of solute in a solution
c.
b. changes in colligative properties of a substance d.
moles of solute to kilograms of solvent
moles of solute to liters of solution
saturated solutions
supersaturated solutions
14. Calculate the hydrogen ion concentration of an aqueous solution, given the concentration of hydroxide ions is 1 x
10-5 M and the ion constant for water is 1 x 10-14.
a.
c.
b.
d.
____
15. When acids react with metals, they produce _____ gas.
a. hydrogen
c. sulfur
b. nitrogen
d. oxygen
____
16. A basic solution contains more _____ ions than hydrogen.
a. oxygen
c. hydroxide
b. nitrogen
d. sulfide
____
17. A _____ is produced when a base accepts a hydrogen ion from an acid.
a. conjugate acid
c. acid
b. conjugate base
d. base
____
18. The _____ of a weak acid is strong.
a. conjugate acid
b. conjugate base
c.
d.
acid
base
____
19. What is the pH of blood, given the hydrogen ion concentration is 4.0 x 10 -8 M?
a. 7.0
c. 7.4
b. 7.2
d. 7.6
____
20. An acid that can donate only one hydrogen ion is called a _____ acid.
a. monoprotic
c. triprotic
b. diprotic
d. polyprotic
____
21. In the Bronsted-Lowry model of acids and bases, an _____ is a hydrogen donor and a _____ is a hydrogen acceptor.
a. acid, base
c. conjugate acid, conjugate base
b. base, acid
d. conjugate base, conjugate acid
____
22. Which model states that an acid is a substance that contains hydrogen and ionizes to produce hydrogen ions?
a. Arrhenius
c. Lewis
b. Bronsted-Lowry
d. Hydrogen
____
23. What is the ph of 0.45 M of H2SO4?
a. 0.0045
b. 0.045
____
c.
d.
0.45
4.50
24. What is the hydrogen ion concentration of 0.050 M H3PO4?
a. 0.015 M
c. 1.50 M
b. 0.15 M
d. 15.0 M
____
25. A solution that contains equal concentrations of hydrogen and hydroxide ions is _____.
a. an acid
c. neutral
b. a base
d. ionized
____
26. The pH scale allows scientists to define the acidity of a substance on a scale of _____.
a. 0-7
c. 1-7
b. 0-14
d. 1-14
____
27. Identify the acid and conjugate base pair in the following equation:
a. HF & H20
c. HF & Fb. HF & H30
d. H20 & H3O+
____
28. Which of these decreases as the pH of a solution increases?
a. basicity of a solution
c. value of
b.
number of hydrogen ions
d.
number of hydroxide ions
____
29. Strong acids or bases make the best electrolytes because they
a. do not ionize in solution
c. ionize completely in solution
b. react in a equilibrating manner
d. have extremely small ionization constants
____
30. The neutralization of a strong acid by a strong base always involves the products
a. water and a salt
c. water and an ion
b. an anion and a salt
d. a weak acid and a strong base