CHEMISTRY
FINAL EXAM REVIEW GUIDE 2012
Scott
DISCLAIMER: THIS IS A SAMPLE OF QUESTIONS/PROBLEMS FROM THIS YEAR. PLEASE BE SURE
TO REVIEW YOUR NOTES, LABS, COJ’S, OBJECTIVE SHEETS, ETC. DO NOT RELY ONLY ON THIS
REVIEW GUIDE.
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1) Which of the following observations is quantitative?
a. The liquid turns blue litmus paper red.
b. The liquid boils at 100C.
c. The liquid tastes bitter.
d. The liquid is cloudy.
____
2) The symbols for units of length in order from largest to smallest are
a. m, cm, mm, km.
c. km, mm, cm, m.
b. mm, m, cm, km.
d. km, m, cm, mm.
____
3) A volume of 1 milliliter is equivalent to
a. 1 cubic centimeter.
b. 1 gram.
c. 1 liter.
d. 10–1 cubic decimeters.
____
4) Which pair of quantities determines the density of a material?
a. mass and weight
c. volume and concentration
b. volume and weight
d. volume and mass
____
5) A sample of gold has a mass of 96.5 g and a volume of 5.00 cm3. The density of gold is
a. 0.0518 g/cm3.
c. 101.5 g/cm3.
3
b. 19.3 g/cm .
d. 483 g/cm3.
____
6) The number of grams equal to 0.5 kg is
a. 0.0005.
b. 0.005.
c. 500.
d. 5000.
____
7) If 1 inch equals 2.54 cm, how many centimeters equal 1 yard?
a. 0.0706 cm
c. 30.5 cm
b. 14.2 cm
d. 91.4 cm
____
8) A chemist obtained the values 5.2246 g, 5.2353 g, and 5.2501 g for the mass of a sample. Without knowing
the true mass of the sample, it can be said that these values have
a. good precision.
b. good accuracy.
c. poor precision.
d. poor accuracy.
____
9) The measurement 0.035550 g rounded off to two significant figures would be
a. 0.03 g.
c. 0.036 g.
b. 0.35 g.
d. 3.5  102 g.
____ 10) When 64.4 is divided by 2.00, the correct number of significant figures in the result is
a. 1.
b. 3.
c. 4.
d. 6.
____ 11) According to the law of conservation of mass, when sodium, hydrogen, and oxygen react to form a
compound, the mass of the compound is ____ the sum of the masses of the individual elements.
a. equal to
c. less than
b. greater than
d. either greater than or less than
____ 12) Which concept in Dalton's atomic theory has been modified?
a. All matter is composed of atoms.
b. Atoms of different elements have different properties and masses.
c. Atoms can combine in chemical reactions.
d. Atoms cannot be divided.
____ 13) Rutherford's experiments led him to conclude that atoms contain massive central regions that have
a. a positive charge.
c. no charge.
b. a negative charge.
d. both protons and electrons.
____ 14) Experiments with cathode rays led to the discovery of the
a. proton.
c. neutron.
b. nucleus.
d. electron.
____ 15) Whose series of experiments identified the nucleus of the atom?
a. Rutherford
c. Chadwick
b. Dalton
d. Bohr
____ 16) Because most particles fired at gold foil passed straight through, Rutherford concluded that
a. atoms were mostly empty space.
c. electrons formed the nucleus.
b. atoms contained no charged particles.
d. atoms were indivisible.
____ 17) An atom is electrically neutral because
a. neutrons balance the protons and electrons.
b. nuclear forces stabilize the charges.
c. the numbers of protons and electrons are equal.
d. the numbers of protons and neutrons are equal.
____ 18) The atomic number of oxygen, 8, indicates that there are eight
a. protons in the nucleus of an oxygen atom.
b. oxygen nuclides.
c. neutrons outside the oxygen atom's nucleus.
d. energy levels in the oxygen atom's nucleus.
____ 19) Chlorine has atomic number 17 and mass number 35. It has
a. 17 protons, 17 electrons, and 18 neutrons.
b. 35 protons, 35 electrons, and 17 neutrons.
c. 17 protons, 17 electrons, and 52 neutrons.
d. 18 protons, 18 electrons, and 17 neutrons.
____ 20) Phosphorus-33 contains
a. 33 protons.
b. 18 neutrons.
c. 33 neutrons.
d. 18 protons.
____ 21) The atomic number of neon is 10. The atomic number of calcium is 20. Compared with a mole of neon, a
mole of calcium contains
a. twice as many atoms.
c. an equal number of atoms.
b. half as many atoms.
d. 20 times as many atoms.
____ 22) The mass of a sample containing 3.5 mol of silicon atoms is approximately
a. 28 g.
c. 72 g.
b. 35 g.
d. 98 g.
____ 23) Refer to the figure below. To which group do fluorine and chlorine belong?
a. alkaline-earth metals
b. transition elements
c. halogens
d. actinides
____ 24) Across a period in the periodic table, atomic radii
a. gradually decrease.
b. gradually decrease, then sharply increase.
c. gradually increase.
d. gradually increase, then sharply decrease.
____ 25) Which is the best reason that the atomic radius generally increases with atomic number in each group of
elements?
a. The nuclear charge increases.
b. The number of neutrons increases.
c. The number of occupied energy levels increases.
d. A new octet forms.
____ 26) The number of valence electrons in Group 17 elements is
a. 7.
c. 17.
b. 8.
d. equal to the period number.
____ 27) Across a period, ionization energies of d-block elements generally
a. increase.
c. remain constant.
b. decrease.
d. drop to zero.
____ 28) Among the d-block elements, as atomic radii decrease, electronegativity values
a. remain constant.
c. decrease.
b. increase.
d. drop to zero.
____ 29) The electrons involved in the formation of a chemical bond are called
a. dipoles.
c. Lewis electrons.
b. s electrons.
d. valence electrons.
____ 30) If two covalently bonded atoms are identical, the bond is
a. nonpolar covalent.
c. dipole covalent.
b. polar covalent.
d. coordinate covalent.
____ 31) The B—F bond in BF3 is
a. polar covalent.
b. ionic.
c. nonpolar covalent.
d. metallic.
____ 32) In a molecule of fluorine, the two shared electrons give each fluorine atom how many electron(s) in the outer
energy level?
a. 1
c. 8
b. 2
d. 32
____ 33) The substance whose Lewis structure shows three covalent bonds is
a. H2O.
c. NH3.
b. CH2Cl2.
d. CCl4.
____ 34) What is the correct Lewis structure for hydrogen chloride, HCl?
a. A
b. B
c. C
d. D
____ 35) The Lewis structure for the ammonium ion, NH4, has
a. nonpolar covalent bond.
c. polar covalent bond.
b. ionic bond.
d. metallic bond.
____ 36) According to VSEPR theory, an AB2 molecule is
a. trigonal-planar.
c. linear.
b. tetrahedral.
d. octahedral.
____ 37) According to VSEPR theory, the structure of the ammonia molecule, NH3, is
a. trigonal-planar.
b. bent.
c. trigonal-pyramidal.
d. tetrahedral.
____ 38) Use VSEPR theory to predict the shape of the hydrogen chloride molecule, HCl.
a. tetrahedral
c. bent
b. linear
d. trigonal-planar
____ 39) Use VSEPR theory to predict the shape of the carbon tetraiodide molecule, CI4.
a. tetrahedral
c. bent
b. linear
d. trigonal-planar
____ 40) The reason the boiling point of water (H2O) is higher than the boiling point of hydrogen sulfide (H2S) is
partially explained by
a. London forces.
c. ionic bonding.
b. covalent bonding.
d. hydrogen bonding.
____ 41) The following molecules contain polar bonds. The only polar molecule is
a. CCl4.
c. NH3.
b. CO2.
d. CH4.
____ 42) When Group 2A elements form ions, they ____.
a. lose two protons
c. lose two electrons
b. gain two protons
d. gain two electrons
____ 43) What is the correct name for the N
a. nitrate ion
b. nitrogen ion
ion?
c. nitride ion
d. nitrite ion
____ 44) Aluminum is a group 3A metal. Which ion does A1 typically form?
a. Al
c. Al
b. Al
d. Al
____ 45) The nonmetals in Groups 6A and 7A ____.
a. lose electrons when they form ions
b. have a numerical charge that is found by subtracting 8 from the group number
c. all have ions with a –1 charge
d. end in -ate
____ 46) Which of the following compounds contains the Mn ion?
a. MnS
c. Mn O
b. MnBr
d. MnO
____ 47) How are chemical formulas of ionic compounds generally written?
a. cation then anion
b. anion then cation
c. Roman numeral first, then anion, then cation
d. subscripts first, then ions
____ 48) Which of the following formulas represents an ionic compound?
a. CS
c. N O
b. BaI
d. PCl
____ 49) Which element, when combined with fluorine, would most likely form an ionic compound?
a. lithium
c. phosphorus
b. carbon
d. chlorine
____ 50) Which of the following shows correctly an ion pair and the ionic compound the two ions form?
a. Sn , N ; Sn N
c. Cr , I ; CrI
b. Cu , O ; Cu O
d. Fe , O ; Fe O
____ 51) Which of the following compounds contains the lead(II) ion?
a. PbO
c. Pb2O
b. PbCl4
d. Pb2S
____ 52) Which of the following shows both the correct formula and correct name of an acid?
a. HClO , chloric acid
c. H PO , phosphoric acid
b. HNO , hydronitrous acid
d. HI, iodic acid
____ 53) What is the formula for phosphoric acid?
a. H PO
b. H PO
c. HPO
d. HPO
____ 54) Select the correct formula for sulfur hexafluoride.
a. S F
c. F S
b. F SO
d. SF
____ 55) What is the correct name for the compound CoCl ?
a. cobalt(I) chlorate
c. cobalt(II) chlorate
b. cobalt(I) chloride
d. cobalt(II) chloride
____ 56) What is the correct formula for barium chlorate?
a. Ba(ClO)
c. Ba(ClO )
b. Ba(ClO )
d. BaCl
____ 57) Which of the following is the correct name for N O ?
a. nitrous oxide
c. nitrogen dioxide
b. dinitrogen pentoxide
d. nitrate oxide
____ 58) To balance a chemical equation, it may be necessary to adjust the
a. coefficients.
c. formulas of the products.
b. subscripts.
d. number of products.
____ 59) A chemical equation is balanced when the
a. coefficients of the reactants equal the coefficients of the products.
b. same number of each kind of atom appears in the reactants and in the products.
c. products and reactants are the same chemicals.
d. subscripts of the reactants equal the subscripts of the products.
____ 60) Which equation is not balanced?
a.
b.
c.
d.
____ 61) Which coefficients correctly balance the formula
a. 1,2,2
c. 2,1,1
b. 1,1,2
d. 2,2,2
____ 62) The units of molar mass are
a. g/mol.
b. mol/g.
?
c. amu/mol.
d. amu/g.
____ 63) In the reaction represented by the equation 2Al2O3  4Al + 3O2, what is the mole ratio of aluminum to
oxygen?
a. 10:6
c. 2:3
b. 3:4
d. 4:3
____ 64) The Haber process for producing ammonia commercially is represented by the equation N2(g) + 3H2(g) 
2NH3(g). To completely convert 9.0 mol hydrogen gas to ammonia gas, how many moles of nitrogen gas are
required?
a. 1.0 mol
c. 3.0 mol
b. 2.0 mol
d. 6.0 mol
____ 65) For the reaction represented by the equation 2Fe + O2  2FeO, how many grams of iron(II) oxide are
produced from 8.00 mol of iron in an excess of oxygen?
a. 71.8 g
c. 712 g
b. 575 g
d. 1310 g
____ 66) For the reaction represented by the equation 2Na + Cl2  2NaCl, how many grams of sodium chloride can be
produced from 500. g each of sodium and chlorine?
a. 112 g
c. 409 g
b. 319 g
d. 824 g
____ 67) For the reaction represented by the equation CH4 + 2O2  2H2O + CO2, calculate the percentage yield of
carbon dioxide if 1000. g of methane react with excess oxygen to produce 2300. g of carbon dioxide.
a. 83.88%
c. 92.76%
b. 89.14%
d. 96.78%
____ 68) A catalyst is ____.
a. the product of a combustion reaction
b. not used up in a reaction
c. one of the reactants in single-replacement reactions
d. a solid product of a reaction
____ 69) Chemical equations must be balanced to satisfy ____.
a. the law of definite proportions
c. the law of conservation of mass
b. the law of multiple proportions
d. Avogadro’s principle
____ 70) In every balanced chemical equation, each side of the equation has the same number of ____.
a. atoms of each element
c. moles
b. molecules
d. coefficients
____ 71) When potassium hydroxide and barium chloride react, potassium chloride and barium hydroxide are formed.
The balanced equation for this reaction is ____.
a. KH BaCl
KCl BaH
c. 2KOH BaCl
2KCl Ba(OH)
b. KOH BaCl
KCl BaOH
d. KOH BaCl
KCl
BaOH
____ 72) In order to predict whether or not a single-replacement reaction takes place, you need to consult a chart that
shows the ____.
a. periodic table
b. activity series of metals
c. common polyatomic ions
d. ionic charges of representative elements
____ 73) In a double-replacement reaction, the ____.
a. products are always molecular
b. reactants are two ionic compounds
c. reactants are two elements
d. products are a new element and a new compound
____ 74) Which of the following is a balanced equation representing the decomposition of lead(IV) oxide?
a. PbO
c. Pb O
Pb 2O
2Pb O
b. PbO
d. PbO
Pb O
Pb O
____ 75) The equation 2C H OH 9O
a. combustion reaction
b. single-replacement reaction
6CO
8H O is an example of which type of reaction?
c. double-replacement reaction
d. decomposition reaction
____ 76) A double-replacement reaction takes place when aqueous cobalt(III) chloride reacts with aqueous lithium
hydroxide. One of the products of this reaction is ____.
a. Co(OH)
c. LiCo
b. Co(OH)
d. LiCl
____ 77) If a synthesis reaction takes place between rubidium and bromine, the chemical formula for the product is
____.
a. RuBr
c. RbBr
b. Rb Br
d. RbBr
____ 78) What is the balanced chemical equation for the reaction that takes place between bromine and sodium iodide?
a. Br
c. Br NaI
NaI
NaBr
I
NaBrI
b. Br
d. Br NaI
2NaI
2NaBr I
NaBr I
____ 79) Pressure is the force per unit
a. volume.
b. surface area.
c. length.
d. depth.
____ 80) If force is held constant as surface area decreases, pressure
a. remains constant.
b. decreases.
c. increases.
d. increases or decreases, depending on the volume change.
____ 81) Why does a can collapse when a vacuum pump removes air from the can?
a. The inside and outside forces balance out and crush the can.
b. The unbalanced outside force from atmospheric pressure crushes the can.
c. The atmosphere exerts pressure on the inside of the can and crushes it.
d. The vacuum pump creates a force that crushes the can.
____ 82) Convert the pressure 0.840 atm to mm Hg.
a. 365 mm Hg
b. 437 mm Hg
c. 638 mm Hg
d. 780 mm Hg
____ 83) To correct for the partial pressure of water vapor in a gas collection bottle, the vapor pressure of H2O at the
collecting temperature is generally
a. subtracted from the partial pressure of the collected gas.
b. added to the pressure of the collected gas.
c. subtracted from the atmospheric pressure.
d. added to the atmospheric pressure.
____ 84) A sample of oxygen occupies 560. mL when the pressure is 800.00 mm Hg. At constant temperature, what
volume does the gas occupy when the pressure decreases to 700.0 mm Hg?
a. 80.0 mL
c. 600. mL
b. 490. mL
d. 640. mL
____ 85) Chlorine is produced by the reaction 2HCl(g)  H2(g) + Cl2(g). How many grams of HCl (36.5 g/mol) must
be used to produce 10.0 L of chlorine at STP?
a. 15.8 g
c. 32.6 g
b. 30.2 g
d. 36.5 g
____ 86) Calculate the approximate volume of a 0.600 mol sample of gas at 15.0°C and a pressure of 1.10 atm.
a. 12.9 L
c. 24.6 L
b. 22.4 L
d. 139 L
____ 87) Why is a gas easier to compress than a liquid or a solid?
a. Its volume increases more under pressure than an equal volume of liquid does.
b. Its volume increases more under pressure than an equal volume of solid does.
c. The space between gas particles is much less than the space between liquid or solid
particles.
d. The volume of a gas’s particles is small compared to the overall volume of the gas.
____ 88) Why does the pressure inside a container of gas increase if more gas is added to the container?
a. There is an increase in the number of collisions between particles and the walls of the
container.
b. There is an increase in the temperature of the gas.
c. There is a decrease in the volume of the gas.
d. There is an increase in the force of the collisions between the particles and the walls of the
container.
____ 89) If a balloon is heated, what happens to the pressure of the air inside the balloon if the volume remains
constant?
a. It increases.
c. It decreases.
b. It stays the same.
d. The change cannot be predicted.
____ 90) If a sealed syringe is heated, in which direction will the syringe plunger move?
a. out
c. The plunger will not move.
b. in
d. The direction cannot be predicted.
____ 91) If the atmospheric pressure on Mt. Everest is one-third the atmospheric pressure at sea level, the partial
pressure of oxygen on Everest is ____.
a. one-sixth its pressure at sea level
c. one-half its pressure at sea level
b. one-third its pressure at sea level
d. equal to its pressure at sea level
____ 92) Which of the following best describes temperature?
a. energy as heat absorbed or released in a chemical or physical change
b. a measure of the average kinetic energy of the particles in a sample of matter
c. energy in the form of heat
d. energy of change
____ 93) How is a Celsius temperature reading converted to a Kelvin temperature reading?
a. by adding 273.15
c. by dividing by 273.15
b. by subtracting 273.15
d. by multiplying by 273.15
____ 94) What is the energy required to raise the temperature of 1 g of a substance by 1°C or 1 K?
a. specific heat
c. heat capacity
b. heat energy
d. enthalpy of formation
____ 95) How much energy does a copper sample absorb as energy in the form of heat if its specific heat is 0.384
J/(g·°C), its mass is 8.00 g, and it is heated from 10.0°C to 40.0°C?
a. 0.0016 J/(g·°C)
c. 92.2 J
b. 0.0016 J
d. 92.2 J/(g·°C)
____ 96) Find the specific heat of a material if a 6.0 g sample absorbs 50. J when it is heated from 30°C to 50°C.
a. 0.60 J
c. 0.42 J
b. 0.60 J/(g·°C)
d. 0.42 J/(g·°C)
____ 97)
is always positive for a
a. spontaneous reaction.
b. nonspontaneous reaction.
____ 98) Ice melting is an example of a(n)
a. exothermic reaction.
b. negative entropic reactions.
c. exothermic reaction.
d. endothermic reaction.
c. endothermic reaction.
d. catalysed reaction.
____ 99) A piece of metal is heated, then submerged in cool water. Which statement below describes what happens?
a. The temperature of the metal will increase.
b. The temperature of the water will increase.
c. The temperature of the water will decrease.
d. The temperature of the water will increase and the temperature of the metal will decrease.
Short Answer
100) What is the law of conservation of mass?
101) Describe the nucleus of an atom.
102) What is the atomic number of an atom?
103) What is the mass number of an atom?
104) The element chromium has four naturally occurring isotopes. Use the relative abundance of each to calculate
the average atomic mass of chromium.
105)
106)
107)
108)
109)
110)
111)
112)
113)
114)
115)
116)
117)
118)
119)
120)
Cr = 4.34%, Cr = 83.79%, Cr = 9.50%, Cr = 2.37%.
In terms of the periodic law, explain which two of these elements are most similar: sodium (element 11),
phosphorus (element 15), and sulfur (element 16).
What can you predict about the properties of xenon and helium, both in Group 18 in the periodic table? Why?
Differentiate/ compare/ contrast an ionic compound and a molecular compound.
Why must a chemical equation be balanced to solve stoichiometry problems?
Give at least three reasons why the actual yield of a chemical reaction could be less than the theoretical yield.
Balance the following equation. Identify the reaction type.
Mg H PO
Mg (PO )
H
Balance the following equation. Identify the reaction type.
C H
O
CO H O
Balance the following equation. Identify the reaction type.
Au O
Au O
Balance the following equation. Identify the reaction type.
Na PO
ZnSO
Na SO
Zn (PO )
Complete and balance the following equation. Identify the reaction type.
Al Cl
Complete and balance the following equation. Identify the reaction type.
Fe (SO )
Ba(OH)
What are standard temperature and pressure? Why is a standard necessary?
A gas occupies a volume of 140 mL at 35.0 C and 97 kPa. What is the volume of the gas at STP?
How many moles of N are in a flask with a volume of 250 mL at a pressure of 300.0 kPa and a temperature
of 300.0 K?
How can a calorimeter measure energy?
Explain the difference between heat and temperature.
121) What mass of KCl (solubility =
at 20 C) can dissolve in 3.30
10 g of water?
122) What is the molarity of a solution containing 9.0 moles of solute in 2500 mL of solution?
123) If 1.0 mL of 6.0M HCl is added to 499 mL of water to give exactly a 500-mL solution, what is the molarity of
the dilute solution?
124) If the volume of solute is 6.0 mL and the volume of solution is 300.0 mL, what is the solute's percent by
volume?
125) Calculate the molality of a solution prepared by dissolving 175 g of KNO in 750 g of water.
126) If you supply 36 kJ of heat, how many moles of ice at 0 C can be melted, heated to its boiling point, and
completely boiled away? ( H
127)
128)
129)
130)
131)
132)
133)
= 40.5 kJ/mol; H
= 6.0 kJ/mol; specific heatwater = 0.0753
Use the periodic table to determine the number of electrons in a neutral atom of lithium.
How many protons are present in an atom of Be-9?
Write the charge of a chloride ion.
How many atoms are present in 80.0 mol of zirconium?
How many moles of iron are equivalent to 1.11  1025 atoms?
Determine the mass in grams of 5.00 mol of oxygen.
Draw a ball-and-stick model of a water molecule. Include the polarities of the bonds.What type of
intermolecular force is present in water?
)
134) What mass of PCl3 forms in the reaction of 75.0 g P4 with 275 g Cl2 ?
135) The reaction of 100. g of salicylic acid, C7H6O3, with excess acetic anhydride produces 50.0 g of aspirin,
C9H8O4, according to the equation below. What is the percentage yield for this reaction?
136) Explain nuclear charge and how it affects the general trend in radii of atoms of elements going from left to
right across a period in the periodic table.
137) What determines whether one metal will replace another metal from a compound in a single-replacement
reaction?
138) Predict the precipitate that forms when aqueous solutions of silver nitrate and potassium chloride react to
form products in a double-replacement reaction. Include a discussion of how to write the complete chemical
equation describing this reaction.
139) Explain Boyle’s pressure-volume relationship in terms of the kinetic-molecular theory.
140) When a mixture of sulfur and metallic silver is heated, silver sulfide is produced. What mass of silver sulfide
is produced from a mixture of 3.0 g Ag and 3.0 g S ?
16Ag(s) + S (s)
8Ag S(s)
141) A 500-g sample of Al (SO ) is reacted with 450 g of Ca(OH) . A total of 596 g of CaSO is produced.
What is the limiting reagent in this reaction, and how many moles of excess reagent are unreacted?
Al (SO ) (aq) + 3Ca(OH) (aq)
2Al(OH) (s) + 3CaSO (s)
142) Discuss the different factors that can affect the solubility of a substance. Include specific examples in your
discussion.
143) What is a diatomic molecule? Identify the elements that exist as diatomic molecules in nature.
144) Explain why the conversion factor
cannot be used for the reaction represented by the equation
145) Determine the maximum number of moles of product that can be produced from 7.0 mol Al and 8.0 mol Cl2
according to the equation 2Al + 3Cl2
2 AlCl3.
Describe in words the method used. Then show the calculation(s).
146) Explain the difference between a limiting reactant and an excess reactant.
147) How can H be treated like other values in a stoichiometry problem even though it is not a physical product?
148) Sulfur in gasoline can produce sulfuric acid, H2SO4, according to the two-step process shown below. For each
125 g of sulfur in gasoline, how many moles of H2SO4 will be produced?
149) What mass in grams of sodium hydroxide is produced if 20.0 g of sodium metal react with excess water
according to the chemical equation 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)?
150) In the reaction represented by the equation 2NH3 + CO2
CO(NH2)2 + H20, 30.7 g of CO(NH2)2 forms per
1.00 mol of CO2 that reacts when NH3 is in excess. What is the percentage yield?
151) Draw a Lewis structure for carbon disulfide, CS2.
CHEMISTRY
Answer Section
FINAL EXAM REVIEW GUIDE 2012
Scott
MULTIPLE CHOICE
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ANS: B
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D
A
D
B
C
D
C
C
B
A
D
A
D
A
A
C
A
A
B
C
D
C
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D
A
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SHORT ANSWER
100) The statement that mass cannot be created or destroyed in ordinary chemical reactions.
101) An atom’s very small central region, which is made up of protons and neutrons.
102) The number of protons in the nucleus of an atom.
103) The sum of the number of protons and neutrons in an atom.
104) 50 amu 0.0434 = 2.17 amu
52 amu 0.8379 = 43.57 amu
53 amu 0.0950 = 5.04 amu
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54 amu 0.0237 = 1.28 amu
_________________________
= 52.06 amu
52 amu
105) Their locations in the periodic table indicate that phosphorus and sulfur are nonmetals and sodium is a
metal. Nonmetals are a group with characteristic properties, so phosphorus and sulfur are the most similar
elements of the three.
106) In the periodic table, elements in the same column or group have similar properties. Because helium and
xenon are in the same group, they have similar properties.
107) Atoms in a molecular compound share electrons to achieve stability. Atoms in an ionic compound gain or
lose electrons to form ions, which combine so that the number of positive and negative charges is equal.
108) Only a balanced equation reveals the correct mole ratios of the reacting substances.
109) The actual yield could be less than a theoretical yield for these reasons: reactants may form byproducts in competing side reactions, reactants may contain impurities, and reactions may not go to
completion.
110)
3Mg 2H PO
Mg (PO )
3H
111)
112)
113)
114)
115)
C H
3O
3CO 3H O Combustion
2Au O
4Au 3O
2Na PO
3ZnSO
3Na SO
Zn (PO )
2Al 3Cl
2AlCl Synthesis
Fe (SO )
3Ba(OH)
2Fe(OH)
3BaSO Double Replacement.
116)
Standard temperature is 0C, and standard pressure is 1 atm. Scientists have agreed upon
standard conditions for temperature and pressure to compare volumes of gases.
117)
T = 35.0 C + 273 = 308 K
T = 0.0 C + 273 = 273 K
V =P
V
V = 97 kPa
118)
n=P
140 mL
250 mL
=
= 120 mL
= 0.25 L
= 0.030 mol
119)
When energy is released or absorbed inside the chamber of a calorimeter, the
temperature of the water surrounding the chamber changes. The change in energy as heat can be calculated by
multiplying the mass of the water by the temperature change and by the specific heat of water.
120)
Energy as heat is energy that is transferred between objects. Temperature is a measure of
the average kinetic energy of all the particles of a substance.
121)
122)
123)
124)
125)
% = 2.0%
molar mass KNO :
K: 1 39.1 g = 39.1 g
N: 1 14.0 g = 14.0 g
O: 3 16.0 g = 48.0 g
molar mass = 101.1 g
126) Total heat = heat to melt ice + heat to warm water to 100 C + heat to evaporate water
Total heat = (moles ice
H ) + (moles water C T) + (moles water
H )
36 kJ = (moles of H O
6.0 kJ/mol) + (moles of H O
0.0753
kJ/mol)
36 kJ = moles H O (6.0 kJ/mol + 0.0753
100 C + 40.5 kJ/mol)
36 kJ = moles H O (54.0 kJ/mol)
moles H O =
moles H O = 0.67 mol
127)
128)
129)
130)
3
4
–1
4.82 1025 atoms Zr
Solution:
131)
18.4 mol Fe
Solution:
132)
80.0 g O
Solution:
133)
-
+
134)
Assuming that P4 is the limiting reagent:
Assuming that Cl2 is the limiting reagent:
100 C) + (moles of H O
40.5
Since the smaller amount of product is formed from P4, it is the limiting reagent. The mass of product formed
is:
135)
136)
Nuclear charge is the attraction an atomic nucleus has on the electrons surrounding it. As you move
from left to right across a period, the atomic number increases, and therefore the number of protons in the
nucleus increases. The more protons within a nucleus, the greater is the nuclear charge. A greater nuclear
charge pulls the electrons closer to the nucleus, decreasing the atomic radius.
137) Whether one metal will replace another is determined by the relative reactivity of the two metals. The
activity series of metals lists metals in order of decreasing reactivity. A reactive metal will replace any metal
found below it in the activity series.
138) Because the reaction is a double-replacement type, cations are exchanged among compounds during the
reaction. The first step is to write the equation in skeleton form:
AgNO + KCl
AgCl + KNO
Inspection of this equation shows that the insoluble precipitate silver chloride forms in an aqueous solution of
potassium nitrate. The relative amounts of elements are the same on either side of the equation, so the
complete equation is:
AgNO (aq) + KCl(aq)
AgCl(s) + KNO (aq) (balanced)
139) The pressure of a gas is caused by collisions of the gas particles with the container walls. If the volume
of a container is decreased, but the same quantity of gas is present at the same temperature, there will be more
molecules per unit volume. This means that there will be more gas particle-wall collisions for a given unit of
wall surface. In other words, there will be greater pressure (force/area on a surface).
Conversely, if the volume of a container is increased, but the same quantity of gas is present at the same
temperature, there will be fewer molecules per unit volume. This means that there will be fewer gas particlewall collisions for a given unit of wall surface. In other words, there will be decreased pressure (force/area on
a surface).
140) The limiting reagent is silver.
3.0 g Ag 1 mol Ag/108 g Ag = 0.03 mol Ag
3.0 g S
1 mol S /256 g S = 0.01 mol S
0.03 mol Ag 8 mol Ag S/16 mol Ag 248 g Ag S/1 mol Ag S = 3.72 g Ag S
3.72 g of silver sulfide is produced.
141) 500 g Al (SO )
1 mol Al (SO ) /342 g Al (SO )
3 mol Ca(OH) /1 mol Al (SO )
Ca(OH) /1 mol Ca(OH) = 325 g Ca(OH)
450 g – 325 g = 125 g
There are 125 g excess Ca(OH) . Al (SO ) is the limiting reagent.
125 g Ca(OH)
1 mol Ca(OH) /74 g Ca(OH) = 1.69 mol Ca(OH)
There are 1.69 mol Ca(OH)2 remaining of unreacted excess reagent.
74 g
142) The factors are temperature, pressure, and the nature of the solute and solvent. Specific examples
include the following. Potassium nitrate is more soluble in water at high temperature than at low temperature.
Gases are less soluble at high temperatures than at low temperatures. The solubility of a particular gas
increases as the partial pressure of that gas above the solution increases. Sodium nitrate is much more soluble
in water than is barium sulfate, regardless of temperature, because the attractive forces between the ions in
BaSO are stronger than the attractive force of the solvent molecules.
143) HONClBrIF
144) This conversion factor uses coefficients to compare mass directly. Ratios of moles must be used to
solve stoichiometry problems.
145) Find the limiting reagent first. Assume that the other reactant is in excess when you calculate the moles
of product formed from the first reactant. The reactant that gives the smaller amount of product is limiting. It
gives the maximum amount of product for the reaction. Set up mole ratios and make them equal to each other.
This means that chlorine is limiting.
146) In a reaction that goes to completion, a limiting reactant is used up, and an excess reactant is not used
up.
147) H is the energy that is released in a reaction, and it is proportional to the reactants. These proportions
can be used as conversion factors in stoichiometry.
148) 3.90 mol H2SO4
149) 34.8 g NaOH
150) First find the theoretical yield:
151)