1.
An unused flashbulb contains magnesium and oxygen.
After use, the contents are changed to magnesium
oxide but the total mass does not change. This
observation can best be explained by the
(a) Law of Constant Composition.
(b) Law of Multiple Proportions.
(c) Avogadro's Law.
(d) Law of Conservation of Mass.
Of the following name/symbol combinations of
elements, which one is WRONG?
(a) uranium/U
(b) sulfur/S
(c) nitrogen/N
(d) potassium/K
(e) iron/I
2.
11.
Of the following symbol/name combinations of
elements, which one is WRONG?
(a) C/carbon
(b) B/barium
(c) F/fluorine
(d) N/nitrogen
(e) U/uranium
Which answer includes all the following that are
chemical changes and not physical changes?
I. freezing of water
II. rusting of iron
III. dropping a piece of iron into hydrochloric acid (H2
is produced)
IV. burning a piece of wood
V. emission of light by a kerosene oil lamp
(a) III and IV
(b) II and V
(c) I, II, III, IV, and V
(d) II, III, and V
(e) II, III, IV, and V
3.
The chemical symbol for manganese is
(a) Mn
(b) Mo
(c) Ma
(d) Ga
(e) Mg
12.
4.
Which response lists all of the following properties of
sulfur that are physical properties and not other
properties?
I. It reacts with hydrogen when heated.
II. It is a yellow solid at room temperature.
III. It is soluble in carbon disulfide.
IV. Its density is 2.97 g/cubic centimeter
V. It melts at 112�C.
(a) II, III, IV, and V
(b) II, IV, and V
(c) I
(d) II, III, and IV
(e) III, IV, and V
The number 0.005436 has how many significant
figures?
(a) 7
(b) 3
(c) 4
(d) 5
(e) 6
5.
The number 10.00 has how many significant figures?
(a) 1
(b) 2
(c) 3
(d) 4
(e) 5
6.
What is the volume of a 2.50 gram block of metal
whose density is 6.72 grams per cubic centimeter?
(a) 16.8 cubic centimeters
(b) 2.69 cubic centimeters
(c) 0.0595 cubic centimeters
(d) 0.372 cubic centimeters
(e) 1.60 cubic centimeters
Answers to Chapter 1
1. (e) 2. (b) 3. (a) 4. (c) 5. (d) 6. (d) 7. (d) 8. (e) 9. (c) 10. (d)
11. (e) 12. (a)
The formula weight of the compound, Al2(SO4)3
18H2O is:
(a) 394.4 g
(b) 666.4 g
(c) 110,900 g
(d) 466.8 g
(e) 561.2 g
7.
A cube of 1.2 inches on the side has a mass of 36
grams. What is the density in g/cm3?
(a) 21
(b) 2.2
(c) 30.
(d) 1.3
(e) 14
2.
The weight of a millimole of (NH4)2HPO4 is:
(a) 132 g
(b) 114 g
(c) 1.14 x 10-3 g
(d) 0.132 g
(e) 6.02 x 1020 g
8.
Nitric acid is a very important industrial chemical:
1.612 x 1010 pounds of it were produced in 1992. If
the density of nitric acid is 12.53 pounds/gallon, what
volume would be occupied by this quantity? (1 gallon
= 3.7854 liters)
(a) 7.646 x 1011 liters
(b) 8.388 x 109 liters
(c) 1.287 x 109 liters
(d) 5.336 x 1010 liters
(e) 4.870 x 109 liters
3.
How many moles of alanine, C3H7NO2, are there in
159 g of alanine?
(a) 1.42 x 104
(b) 1.78
(c) 0.992
(d) 0.560
(e) 3.31
9.
Identify the INCORRECT statement.
(a) Helium in a balloon: an element
(b) Paint: a mixture
(c) Tap water: a compound
(d) Mercury in a barometer; an element
10.
4.
How many atoms are in one mole of CH3OH?
(a) 6
(b) 6.0 x 1023
(c) 12.0 x 1023
(d) 3.6 x 1024
(e) 3
5.
The mass in grams of 2.6 x 1022 chlorine atoms is:
(a) 4.4
(b) 11
(c) 0.76
(d) 1.5
(e) 3.2
14.
6.
How many aluminum atoms are there in 3.50 grams
of Al2O3?
(a) 4.13 x 1022
(b) 4.90 x 1022
(c) 2.07 x 1022
(d) 1.68 x 1022
(e) 2.45 x 1022
An oxide of lead contains 90.65% Pb, by weight. The
empirical formula is:
(a) Pb
(b) PbO
(c) Pb3O4
(d) Pb2O3
(e) PbO2
15.
7.
A 0.500 g sample of a compound containing only
antimony and oxygen was found to contain 0.418 g of
antimony and 0.082 g of oxygen. What is the simplest
formula for the compound?
(a) SbO
(b) SbO2
(c) Sb3O4
(d) Sb2O5
(e) Sb2O3
Which one of the samples contains the most atoms?
(a) 1 mol of CO2(g)
(b) 1 mol of UF6(g)
(c) 1 mol of CH3COCH3(l)
(d) 1 mol of He(g)
(e) all contain the same number of atoms
8.
Which one of the samples contains the most
molecules?
(a) 1 mol of CO2(g)
(b) 1 mol of UF6(g)
(c) 1 mol of CH3COCH3(l)
(d) 1 mol of He(g)
(e) all contain the same number of molecules
16.
9.
Which one of the samples has the largest mass?
(a) 1 mol of CO2(g)
(b) 1 mol of UF6(g)
(c) 1 mol of CH3COCH3(l)
(d) 1 mol of He(g)
(e) all have the same mass
17.
11.
Guanidin, HNC(NH2)2, is a fertilizer. To three
significant figures, what is the percent by mass of
nitrogen in the fertilizer?
(a) 45.2%
(b) 49.4%
(c) 54.8%
(d) 65.1%
(e) 71.1%
12.
Calculate the percent, by weight, of carbon in 154 g of
C4H8O3?
(a) 46%
(b) 31%
(c) 72%
(d) 27%
(e) 55%
13.
Analysis of a sample of a covalent compound showed
that it contained 14.4% hydrogen and 85.6% carbon
A compound contains, by mass, 40.0% carbon, 6.71%
hydrogen, and 53.3% oxygen. A 0.320 mole sample of
this compound weighs 28.8 g. The molecular formula
of this compound is:
(a) C2H4O2
(b) C3H6O3
(c) C2H4O
(d) CH2O
(e) C4H7O2
What mass of cerussite, PbCO3, would contain 35.0
grams of lead?
(a) 27.1 g
(b) 45.1 g
(c) 42.4 g
(d) 35.6 g
(e) 51.7 g
10.
Which of the following statements is(are) FALSE?
1. The percent by mass of each element in a
compound depends on the amount of the compound.
2. The mass of each element in a compound depends
on the amount of the compound.
3. The percent by mass of each element in a
compound depends on the amount of element present
in the compound.
(a) 2 and 3
(b) 1 only
(c) 1 and 2
(d) 1, 2 and 3
(e) another combination
by mass. What is the empirical formula for the
compound?
(a) CH
(b) CH2
(c) CH3
(d) C2H3
(e) none of these
Answers to Chapter 2
1. (b) 2. (d) 3. (b) 4. (d) 5. (d) 6. (a) 7. (c) 8. (e) 9. (b) 10. (b)
11. (e) 12. (a) 13. (b) 14. (c) 15. (e) 16. (b) 17. (b)
1.
Balance the following equation with the smallest
whole number coefficients. Choose the answer that is
the sum of the coefficients in the balanced equation.
Do not forget coefficients of "one."
PtCl4 + XeF2
PtF6 + ClF + Xe
(a) 16
(b) 22
(c) 24
(d) 26
(e) 32
2.
Balance the following equation with the smallest
whole number coefficients. Choose the answer that is
the sum of the coefficients in the balanced equation.
Do not forget coefficients of "one."
Cr2(SO4)3 + RbOH
Cr(OH)3 + Rb2SO4
(a) 10
(b) 12
(c) 13
(d) 14
(e) 15
Calculate the mass of hydrogen formed when 25
grams of aluminum reacts with excess hydrochloric
acid.
2Al + 6HCl
Al2Cl6 + 3H2
Balance the following equation using minimum
integral coefficients:
NH3 + O2
NO2 + H2O
(a) 0.41 g
(b) 1.2 g
(c) 1.8 g
(d) 2.8 g
(e) 0.92 g
3.
The stoichiometric coefficient for oxygen gas O2 is:
(a) 1
(b) 4
(c) 3
(d) 7
(e) 5
9.
4.
When iron pyrite (FeS2) is heated in air, the process
known as "roasting" forms sulfur dioxide and iron(III)
oxide. When the equation for this process is
completed and balanced, using the smallest whole
number coefficients, what is the coefficient for "O2"?
___ FeS2 + ___ O2
___ SO2 + ___ Fe2O3
10.
(a) 2
(b) 4
(c) 7
(d) 8
(e) 11
A commercially valuable paint and adhesive stripper,
dimethyl sulfoxide (DMSO), (CH3)2SO, can be
prepared by the reaction of oxygen with dimethyl
sulfide, (CH3)2S, using a ratio of one mole oxygen to
two moles of the sulfide:
O2 + 2(CH3)2S
2(CH3)2SO
5.
How many moles of KBrO3 are required to prepare
0.0700 moles of Br2 according to the reaction:
KBrO3 + 5KBr + 6HNO3
6KNO3 + 3Br2 + 3H2O
If this process is 83% efficient, how many grams of
DMSO could be produced from 65 g of dimethyl
sulfide and excess O2?
(a) 68 g
(b) 75 g
(c) 83 g
(d) 51 g
(e) 47 g
(a) 0.210
(b) 0.0732
(c) 0.0704
(d) 0.220
(e) 0.0233
6.
11.
Which of the following statements is FALSE for the
chemical equation given below in which nitrogen gas
reacts with hydrogen gas to form ammonia gas
assuming the reaction goes to completion?
N2 + 3H2
2NH3
(a) The reaction of one mole of H2 will produce 2/3
moles of NH3.
(b) One mole of N2 will produce two moles of NH3.
(c) One molecule of nitrogen requires three molecules
of hydrogen for complete reaction.
(d) The reaction of 14 g of nitrogen produces 17 g of
ammonia.
(e) The reaction of three moles of hydrogen gas will
produce 17 g of ammonia.
The formation of ethyl alcohol (C2H5OH) by the
fermentation of glucose (C6H12O6) may be
represented by:
C6H12O6
2C2H5OH + 2CO2
If a particular glucose fermentation process is 87.0%
efficient, how many grams of glucose would be
required for the production of 51.0 g of ethyl alcohol
(C2H5OH)?
(a) 68.3 g
(b) 75.1 g
(c) 115 g
(d) 229 g
(e) 167 g
7.
Calcium carbide, CaC2, is an important preliminary
chemical for industries producing synthetic fabrics
and plastics. CaC2 may be produced by heating
calcium oxide with coke:
CaO + 3C
CaC2 + CO
What is the amount of CaC2 which can be produced
from the reaction of excess calcium oxide and 10.2 g
of carbon? (Assume 100% efficiency of reaction for
purposes of this problem.)
(a) 18.1 g
(b) 28.4 g
(c) 20.8 g
(d) 19.8 g
(e) 27.2 g
8.
When 12 g of methanol (CH3OH) was treated with
excess oxidizing agent (MnO4-), 14 g of formic acid
(HCOOH) was obtained. Using the following chemical
equation, calculate the percent yield. (The reaction is
much more complex than this; please ignore the fact
that the charges do not balance.)
3CH3OH + 4MnO43HCOOH + 4MnO2
(a) 100%
(b) 92%
(c) 82%
(d) 70%
(e) 55%
12.
The limiting reagent in a chemical reaction is one
that:
(a) has the largest molar mass (formula weight).
(b) has the smallest molar mass (formula weight).
(c) has the smallest coefficient.
(d) is consumed completely.
(e) is in excess.
13.
If 5.0 g of each reactant were used for the the
following process, the limiting reactant would be:
2KMnO4 +5Hg2Cl2 + 16HCl
10HgCl2 + 2MnCl2 +
2KCl + 8H2O
(a) KMnO4
(b) HCl
(c) H2O
(d) Hg2Cl2
(e) HgCl2
(a) 13 mL
(b) 22 mL
(c) 39 mL
(d) 73 mL
(e) none of these
14.
What mass of ZnCl2 can be prepared from the reaction
of 3.27 grams of zinc with 3.30 grams of HCl?
Zn +2HCl
ZnCl2 + H2
22.
How many grams of Ag2CO3 are required to react with
28.5 mL of 1.00 M NaOH solution?
Ag2CO3 +2NaOH
Ag2O + Na2CO3 + H2O
(a) 7.87 g
(b) 3.93 g
(c) 15.7 g
(d) 10.8 g
(e) 8.16 g
(a) 6.89 g
(b) 6.82 g
(c) 6.46 g
(d) 6.17 g
(e) 6.02 g
15.
How many grams of NH3 can be prepared from 77.3
grams of N2 and 14.2 grams of H2? (Hint: Write and
balance the equation first.)
(a) 93.9 g
(b) 79.7 g
(c) 47.0 g
(d) 120.0 g
(e) 13.3 g
23.
How many milliliters of 0.200 M NH4OH are needed to
react with 12.0 mL of 0.550 M FeCl3?
FeCl3 + 3NH4OH
Fe(OH)3 + 3NH4Cl
(a) 99.0 mL
(b) 33.0 mL
(c) 8.25 mL
(d) 68.8 mL
(e) 132 mL
16.
Silicon carbide, an abrasive, is made by the reaction of
silicon dioxide with graphite.
SiO2 +3C
SiC + 2CO
24.
When 250. mL of a 0.15 M solution of ammonium
sulfide (NH4)2S is poured into 120. mL of a 0.053 M
solution of cadmium sulfate CdSO4, how many grams
of a yellow precipitate of cadmium sulfide CdS are
formed? The other product is (NH4)2SO4. (Hint: Write
out and balance the equation. Is this a limiting reagent
problem? )
(a) 5.4 g
(b) 0.92 g
(c) 2.6 g
(d) 1.9 g
(e) 530 g
If 100 g of SiO2 and 100 g of C are reacted as far as
possible, which one of the following statements will
be correct?
(a) 111 g of SiO2 will be left over.
(b) 44 g of SiO2 will be left over.
(c) 82 g of C will be left over.
(d) 40 g of C will be left over.
(e) Both reactants will be consumed completely, with
none of either left over.
17.
Calculate the mass of 6.00% NiSO4 solution that
contains 40.0 g of NiSO4?
(a) 667 g
(b) 540 g
(c) 743 g
(d) 329 g
(e) none of these
18.
How many grams of water are contained in 75.0
grams of a 6.10% aqueous solution of K3PO4?
(a) 75.0 g
(b) 73.2 g
(c) 70.4 g
(d) 68.1 g
(e) 62.8 g
Answers to Chapter 3
1. (a) 2. (b) 3. (d) 4. (e) 5. (e) 6. (e) 7. (a) 8. (d) 9. (c) 10. (a)
11. (c) 12. (d) 13. (d) 14. (d) 15. (b) 16. (d) 17. (a) 18. (c) 19.
(d) 20. (b) 21. (b) 22. (b) 23. (a) 24. (b)
1.
What alkaline earth metal is located in period 3?
(a) Li
(b) Na
(c) Ca
(d) Mg
(e) Sr
19.
The mass (in grams) of FeSO4 7H2O required for
preparation of 125 mL of 0.90 M solution is:
(a) 16 g
(b) 25 g
(c) 13 g
(d) 31 g
(e) 43 g
2.
Which of the following is classified as a metal?
(a) Ge
(b) As
(c) F
(d) V
(e) Ar
20.
What is the molarity of phosphoric acid in a solution
labeled 20.0% phosphoric acid (H3PO4) by weight
with a density = 1.12 g/mL?
(a) 0.98 M
(b) 2.3 M
(c) 2.7 M
(d) 3.0 M
(e) 3.6 M
21.
How many mL of 17 M NH3 must be diluted to 500.0
mL to make a 0.75 M solution?
3.
Which of the following is a weak acid?
(a) H2SO4
(b) HClO3
(c) HF
(d) HCl
(e) HNO3
4.
Which one of the following is likely to be the most
soluble base?
(a) Ca(OH)2
(b) Cu(OH)2
(c) Ga(OH)3
(d) Zn(OH)2
(e) Zr(OH)3
(e) In a neutralization reaction, an acid reacts with
base to produce a salt and H2O.
12.
What is the net ionic equation for the acid-base
reaction that occurs when acetic acid and potassium
hydroxide solutions are mixed?
(a) H+(aq) + OH-(aq)
H2O(l)
(b) H+(aq) + KOH(s)
K+(aq) + H2O(l)
(c) CH3COOH(aq) + KOH(s)
KCH3COO(aq) +
H2O(l)
(d) CH3COO-(aq) + H+(aq) + K+(aq) + OH-(aq)
K+(aq) + CH3COO-(aq) + H2O(l)
(e) CH3COOH(aq) + OH-(aq)
CH3COO-(aq) +
H2O(l)
5.
Which one of the following statements is TRUE?
(a) One mole of any acid will ionize completely in
aqueous solution to produce one mole of H+ ions.
(b) Solutions of weak acids always have lower
concentrations of H+ than solutions of strong acids.
(c) There are several common acids that are
insoluble.
(d) All of the IA and IIA metal hydroxides are soluble.
(e) All weak acids are insoluble.
6.
13.
Consider the following reaction:
NH3(g) + H2O(l)
NH4+(aq) + OH-(aq)
Which one of the following statements is false?
(a) The double arrows indicate that ammonia, NH3, is
only very slightly soluble in water.
(b) The reaction is reversible.
(c) When ammonia is added to water, NH4+ and OHions are produced in a 1:1 ratio.
(d) When solutions of NH4Cl and NaOH are mixed,
some ammonia is produced.
(e) Ammonia is considered to be a weak base.
7.
Which one of the following salts is insoluble?
(a) NH4Cl
(b) Ca(NO3)2
(c) BaCO3
(d) Na2S
(e) Zn(CH3COO)2
14.
Which of the following statements is FALSE given the
following net ionic equation?
H3PO4(aq) + 3OH-(aq)
PO43-(aq) + 3H2O(l)
(a) If all the water evaporated away, the salt
remaining could possibly be Na3PO4.
(b) The acid, H3PO4, is a weak electrolyte.
(c) The base involved must be a strong soluble base.
(d) This is classified as a neutralization reaction.
(e)This could be the net ionic equation for H3PO4
reacting with Al(OH)3.
8.
What salt is formed in the following acid/base
reaction?
HClO3 + Ba(OH)2
(a) BaCl2
(b) ClOBa
(c) H2O
(d) BaClO3
(e) Ba(ClO3)2
15.
Which of the following statements is FALSE given the
following net ionic equation?
2H+(aq) + Cu(OH)2(s)
Cu2+(aq) + 2H2O(l)
(a) If all the water evaporated away, the salt
remaining could possibly be CuS.
(b) The acid involved must be a strong electrolyte.
(c) The base, Cu(OH)2, is an insoluble base.
(d) This could be the net ionic equation for HNO3
reacting with Cu(OH)2.
(e) This is classified as a neutralization reaction.
9.
The precipitate formed when barium chloride is
treated with sulfuric acid is _______ .
(a) BaS2O4
(b) BaSO3
(c) BaSO2
(d) BaSO4
(e) BaS
16.
Determine the oxidation number of carbon in K2CO3.
(a) 0
(b) +2
(c) +4
(d) -2
(e) some other value
10.
The spectator ion(s) in the following reaction is/are:
Na2CO3(aq) + Ba(NO3)2(aq)
BaCO3(s) +
2NaNO3(aq)
(a) Na+ and Ba2+
(b) Ba2+ and CO32(c) CO32- and NO3(d) Na+ only
(e) Na+ and NO3-
17.
Which assignment of oxidation number is
INCORRECT for the blinking element?
(a) K2Cr2O7; +6
(b) NH3; +3
(c) H2PO2-; +1
(d) SeO32-; +4
(e) Cu(NO3)2; +2
11.
Which one of the following statements is FALSE?
(a) For the reaction of a strong acid with a strong
soluble base, the net ionic equation is always
H+ + OHH2O
(b) "Spectator ions" appear in the total ionic equation
for a reaction, but not in the net ionic equation.
(c) HF, HCl, and HNO3 are all examples of strong acids.
(d) Titration is a process which can be used to
determine the concentration of a solution.
What is the net ionic equation for the acid-base
reaction that occurs when nitric acid is added to
copper(II) hydroxide?
(a) H+(aq) + OH-(aq)
H2O(l)
(b) 2H+(aq) + Cu(OH)2(s)
Cu2+(aq) + 2H2O(l)
(c) 2HNO3(aq) +Cu(OH)2(s)
Cu(NO3)2(s) +
2H2O(l)
(d) 2H+(aq) + 2NO3-(aq) + Cu2+(aq) + 2OH-(aq)
Cu(NO3)2(s) + 2H2O(l)
(e) 2H+(aq) + 2NO3-(aq) + Cu2+(aq) + 2OH-(aq)
Cu2+(aq) + 2NO3-(aq) + 2H2O(l)
18.
Consider the following reaction:
4NH3 + 5O2
4NO + 6H2O
The element being oxidized and the oxidizing agent
are:
(a) N and NH3
(b) N and O2
(c) O and NH3
(d) O and O2
(e) H and NH3
(d) the same mass numbers.
(e) the same masses.
5.
19.
Given that the Activity Series is: Na>Mg>Cu>Ag>Au,
which one of the following answers represents the
ions that would not be displaced from aqueous
solution (reduced) by metallic magnesium?
(a) Na+
(b) Cu2+
(c) Cu2+ and Au+
(d) Cu2+, Ag+ and Au+
(e) Na+, Cu2+, Ag+ and Au+
What is the atomic weight of a hypothetical element
consisting of two isotopes, one with mass = 64.23
amu (26.0%), and one with mass = 65.32 amu?
(a) 65.3 amu
(b) 64.4 amu
(c) 64.9 amu
(d) 65.0 amu
(e) 64.8 amu
6.
Naturally occurring rubidium consists of just two
isotopes. One of the isotopes consists of atoms having
a mass of 84.912 amu; the other of 86.901 amu. What
is the percent natural abundance of the heavier
isotope?
(a) 15%
(b) 28%
(c) 37%
(d) 72%
(e) 85%
20.
Which name/formula combination is WRONG?
(a) phosphorous acid / H3PO4
(b) nitrogen oxide / NO
(c) acetate ion / CH3COO(d) sodium chromate / Na2CrO4
(e) calcium hypobromite / Ca(BrO)2
21.
Which name/formula combination is WRONG?
(a) chlorous acid / HClO2
(b) dinitrogen tetroxide / N2O4
(c) ammonium nitrate / NH4NO3
(d) copper(II) periodate / CuIO4
(e) potassium permanganate / KMnO4
Answers to Chapter 4
7.
What is the frequency of light having a wavelength of
4.50 x 10-6 cm?
(a) 2.84 x 10-12 s-1
(b) 2.10 x 104 s-1
(c) 4.29 x 1014 s-1
(d) 1.06 x 1022 s-1
(e) 6.67 x 1015 s-1
8.
The emission spectrum of gold shows a line of
wavelength 2.676 x 10-7 m. How much energy is
emitted as the excited electron falls to the lower
energy level?
(a) 7.43 x 10-19 J
(b) 5.30 x 10-20 J
(c) 6.05 x 10-19 J
(d) 3.60 x 10-20 J
(e) 5.16 x 10-20 J
1. (d) 2. (d) 3. (c) 4. (a) 5. (b) 6. (a) 7. (c) 8. (e) 9. (d) 10. (e)
11. (c) 12. (e) 13. (b) 14. (e) 15. (a) 16. (c) 17. (b) 18. (b) 19.
(a) 20. (a) 21. (d)
1.
Which of the following has a positive charge?
(a) proton
(b) neutron
(c) anion
(d) electron
(e) atom
9.
Which of the responses contains all the statements
that are consistent with the Bohr theory of the atom
(and no others)?
(1) An electron can remain in a particular orbit as
long as it continually absorbs radiation of a definite
frequency.
(2) The lowest energy orbits are those closest to the
nucleus.
(3) An electron can jump from the K shell (n = 1 major
energy level) to the M shell (n = 3 major energy level)
by emitting radiation of a definite frequency.
(a) 1,2,3
(b) 2 only
(c) 3 only
(d) 1,2
(e) 2,3
2.
Rutherford carried out experiments in which a beam
of alpha particles was directed at a thin piece of metal
foil. From these experiments he concluded that:
(a) electrons are massive particles.
(b) the positively charged parts of atoms are moving
about with a velocity approaching the speed of light.
(c) the positively charged parts of atoms are
extremely small and extremely heavy particles.
(d) the diameter of an electron is approximately equal
to that of the nucleus.
(e) electrons travel in circular orbits around the
nucleus.
3.
4.
10.
The Heisenberg Principle states that _____________.
(a) no two electrons in the same atom can have the
same set of four quantum numbers.
(b) two atoms of the same element must have the
same number of protons.
(c) it is impossible to determine accurately both the
position and momentum of an electron
simultaneously.
(d) electrons of atoms in their ground states enter
energetically equivalent sets of orbitals singly before
they pair up in any orbital of the set.
(e) charged atoms (ions) must generate a magnetic
field when they are in motion.
Consider the species 72Zn, 75As and 74Ge. These
species have:
(a) the same number of electrons.
(b) the same number of protons.
(c) the same number of neutrons.
(d) the same number of protons and neutrons.
(e) the same mass number.
The neutral atoms of all of the isotopes of the same
element have
(a) different numbers of protons.
(b) equal numbers of neutrons.
(c) the same number of electrons.
11.
Which statement about the four quantum numbers
which describe electrons in atoms is incorrect?
(a) n = principal quantum number, n = 1, 2, 3, ......
(b) l = subsidiary (or azimuthal) quantum number, l =
1, 2, 3, ... , (n+1)
(c) ml = magnetic quantum number, ml = (-l), .... , 0, .... ,
(+l)
(d) ms = spin quantum number, ms = +1/2 or -1/2.
(e) The magnetic quantum number is related to the
orientation of atomic orbitals in space.
(a) 12
(b) 18
(c) 24
(d) 9
(e) 6
20.
A neutral atom of an element has 2 electrons in the
first energy level, 8 in the second energy level and 8 in
the third energy level. This information does not
necessarily tell us:
(a) the atomic number of the element.
(b) anything about the element's chemical properties.
(c) the total number of electrons in s orbitals.
(d) the total number of electrons in p orbitals.
(e) the number of neutrons in the nucleus of an atom
of the element.
12.
Which atomic orbital is spherical in shape? (Note: you
should know and be able to recognize the shapes of
the s orbital, px, py, and pz orbitals, and dxy, dyz, dxz,
dx2-y2 and dz2 orbitals.)
(a) 2s
(b) 3p
(c) 3d
(d) 4f
(e) they are all spherical
13.
The maximum number of electrons that can be
accommodated in a sublevel for which l = 3 is:
(a) 2
(b) 10
(c) 6
(d) 14
(e) 8
Answers to Chapter 5
1. (a) 2. (c) 3. (c) 4. (c) 5. (d) 6. (b) 7. (e) 8. (a) 9. (b) 10. (c)
11. (b) 12. (a) 13. (d) 14. (e) 15. (e) 16. (e) 17. (a) 18. (c) 19.
(b) 20. (e)
1.
The atom having the valence-shell configuration 4s2
4p5 would be in:
(a) Group VIA and Period 5
(b) Group IVB and Period 4
(c) Group VIB and Period 7
(d) Group VIIA and Period 4
(e) Group VIIB and Period 4
14.
15.
The ground state electron configuration for arsenic is:
(a) [Ar] 4s2 4p13
(b) [Kr] 4s2 4p1
(c) 1s2 2s2 2p6 3s2 3p6 3d12 4s2 4p1
(d) 1s2 2s2 2p6 3s2 3p6 4s2 3d8 4p5
(e) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3
Which of the following electron configurations is
correct for nickel?
(a) [Ar] 4s1 3d8
(b) [Kr] 4s1 4d8
(c) [Kr] 4s1 3d8
(d) [Kr] 4s2 3d8
(e) [Ar] 4s2 3d8
16.
The outer electronic configuration ns2np4
corresponds to which one of the following elements in
its ground state?
(a) As
(b) Ca
(c) Cr
(d) Br
(e) S
2.
Select the term best describing the series of elements:
Mn, Fe, Co, Ni, Cu.
(a) d-transition metals
(b) representative elements
(c) metalloids
(d) alkaline earth metals
(e) halogens
3.
Which element has the largest atomic radius?
(a) Li
(b) Na
(c) Rb
(d) F
(e) I
4.
Which of the following terms accurately describes the
energy associated with the process:
Li(g)
Li+(g) + e(a) electron affinity
(b) binding energy
(c) ionization energy
(d) electronegativity
(e) none of these
17.
In the ground state of a cobalt atom there are _____
unpaired electrons and the atom is _____.
(a) 3, paramagnetic
(b) 5, paramagnetic
(c) 2, diamagnetic
(d) 0, diamagnetic
(e) 2, paramagnetic
18.
Which one of the following sets of quantum numbers
could be those of the distinguishing (last) electron of
Mo?
(a) n = 4, l = 0, ml = 0, ms = +1/2
(b) n = 5, l = 1, ml = 9, ms = -1/2
(c) n = 4, l = 2, ml = -1, ms = +1/2
(d) n = 5, l = 2, ml = +2, ms = -1/2
(e) n = 3, l = 2, ml = 0, ms = +1/2
19.
How many p electrons are there in an atom of
rubidium?
5.
The species that contains 24 protons, 26 neutrons and
22 electrons would be represented by the symbol:
(a) 50V3+
(b) 26Cr2+
(c) 50Cr2+
(d) 50Mn2+
(e) none of these
6.
Which element has the lowest first ionization energy?
(a) He
(b) Ne
(c) Ar
(d) Kr
(e) Xe
(a) in s orbitals only.
(b) located in the outermost occupied major energy
level.
(c) located closest to the nucleus.
(d) located in d orbitals.
(e) located in the innermost occupied shell.
7.
Which element has the highest first ionization
energy? (Note: this is an exception to the general
trend - see notes- you may be responsible for this.)
(a) Be
(b) B
(c) C
(d) N
(e) O
2.
8.
Which of these isoelectronic species has the smallest
radius?
(a) Br(b) Sr2+
(c) Rb+
(d) Se2(e) They are all the same size because they have the
same number of electrons.
9.
Which of the following elements has the greatest
attraction for electrons in a covalent bond?
(a) Ge
(b) As
(c) Se
(d) Br
(e) Bi
3.
With regard to the species 16O2-, 19F- and 20Ne, which
of the following statements is correct?
(a) All three species contain 10 electrons.
(b)The sum of the neutrons in all three species is 27.
(c) The sum of the protons in all three species is 28.
(d) Both 19F- and 20Ne contain 20 neutrons.
(e) none of the above
4.
Which of the following does not have a noble gas
electron configuration? (or Which of the following is
not isoelectronic with a noble gas?)
(a) S2(b) Ba+
(c) Al3+
(d) Sb3(e) Sc3+
10.
Which statement is wrong?
(a) The atomic weight of carbon is about 12.
(b) The most stable ion of lithium is Li+.
(c) A phosphorus atom is larger than an antimony
atom.
(d) The radius of a sodium atom is larger than that of
a sodium cation.
(e) Oxygen has a less negative electron affinity than
fluorine.
5.
Which one of the formulas for ionic compounds below
is incorrect?
(a) SrCl2
(b) Cs2S
(c) AlCl3
(d) Al3P2
(e) CaSe
11.
All of the following properties of the alkaline earth
metals increase going down the group except
(a) atomic radius
(b) first ionization energy
(c) ionic radius
(d) atomic mass
(e) atomic volume
12.
Which of the following pairs of elements and valence
electrons is incorrect?
(a) Al - 3
(b) Br - 7
(c) S - 4
(d) Sr - 2
(e) Tl - 3
6.
Which is classified as nonpolar covalent?
(a) the H-I bond in HI
(b) the H-S bond in H2S
(c) the P-Cl bond in PCl3
(d) the N-Cl bond in NCl3
(e) the N-H bond in NH3
7.
Which of the following is an ionic hydride?
(a) PH3
(b) H2S
(c) HI
(d) KH
(e) CH4
Which one of the compounds below is most likely to
be ionic?
(a) GaAs
(b) ScCl3
(c) NO2
(d) CCl4
(e) ClO2
13.
Which of the following is the most basic oxide?
(a) N2O3
(b) N2O5
(c) P4O6
(d) P4O10
(e) Bi2O5
8.
The correct electron-dot formulation for hydrogen
cyanide shows:
(a) 2 double bonds and two lone pairs of electrons on
the N atom.
(b) 1 C-H bond, 1 C=N bond, 1 lone pair of electrons
on the C atom and 1 lone pair of electrons on the N
atom.
(c) 1 C-H bond, 1 C-N bond, 2 lone pairs of electrons
on the C atom and 3 lone pairs of electrons on the N
atom.
(d) 1 triple bond between C and N, 1 N-H bond and 2
lone pairs of electrons on the C atom.
(e) 1 triple bond between C and N, 1 C-H bond and 1
lone pair of electrons on the N atom.
Answers to Chapter 6
1. (d) 2. (a) 3. (c) 4. (c) 5. (c) 6. (e) 7. (d) 8. (b) 9. (d) 10. (c)
11. (b) 12. (d) 13. (e)
9.
1.
The valence electrons of representative elements are
The correct dot formulation for nitrogen trichloride
has:
10.
(a) 3 N-Cl bonds and 10 lone pairs of electrons.
(b) 3 N=Cl bonds and 6 lone pairs of electrons.
(c) 1 N-Cl bond, 2 N=Cl bonds and 7 lone pairs of
electrons.
(d) 2 N-Cl bonds, 1 N=Cl bond and 8 lone pairs of
electrons.
(e) 3 N-Cl bonds and 9 lone pairs of electrons.
What is the total number of electrons in the correct
Lewis dot formula of the sulfite ion?
(a) 8
(b) 24
(c) 26
(d) 30
(e) 32
(d) NH3 - tetrahedral
(e) PF3 - pyramidal
2.
Choose the species that is incorrectly matched with
the electronic geometry about the central atom.
(a) NO2- - trigonal planar
(b) ClO4- - tetrahedral
(c) SO32- - pyramidal
(d) ClO3- - tetrahedral
(e) BrO4- - tetrahedral
3.
Which of the following pairs of molecules and their
molecular geometries is WRONG?
(a) NF3 - trigonal planar
(b) H2O - bent
(c) BF3 - trigonal planar
(d) AsF5 - trigonal bipyramidal
(e) SeF6 - octahedral
11.
In the Lewis structure for the OF2 molecule, the
number of lone pairs of electrons around the central
oxygen atom is
(a) 0
(b) 1
(c) 2
(d) 3
(e) 4
4.
Which molecule has a linear arrangement of all
component atoms?
(a) CH4
(b) H2O
(c) CO2
(d) NH3
(e) BF3
12.
The electronic structure of the SO2 molecule is best
represented as a resonance hybrid of ____ equivalent
structures.
(a) 2
(b) 3
(c) 4
(d) 5
(e) This molecule does not exhibit resonance.
5.
Which polyatomic ion is incorrectly matched with its
ionic geometry?
(a) SiCl62- - trigonal bipyramidal
(b) PH4+ - tetrahedral
(c) ClO2- - angular
(d) NH4+ - tetrahedral
(e) SO42- - tetrahedral
13.
Consider the bicarbonate ion (also called the
hydrogen carbonate ion). After drawing the correct
Lewis dot structure(s), you would see:
(a) two double bonds around the central carbon atom.
(b) three single bonds around the central carbon
atom.
(c) four single bonds around the central carbon atom.
(d) two equivalent resonance forms.
(e) three equivalent resonance forms.
6.
7.
14.
Draw one of the resonance structures of SO3. The
formal charge of S is
(a) +2
(b) +1
(c) 0
(d) -1
(e) -2
Which of the following species is planar?
(a) NH3
(b) H3O+
(c) SO32(d) PF3
(e) NO3-
8.
What kind of hybrid orbitals are utilized by the
carbon atom in CF4 molecules?
(a) sp
(b) sp2
(c) sp3
(d) sp3d
(e) sp3d2
A neutral molecule having the general formula AB3
has two unshared pair of electrons on A. What is the
hybridization of A?
(a) sp
(b) sp2
(c) sp3
(d) sp3d
(e) sp3d2
15.
Which one of the following violates the octet rule?
(a) PCl3
(b) CBr4
(c) NF3
(d) OF2
(e) AsF5
9.
What hybridization is predicted for sulfur in the HSO3ion?
(a) sp
(b) sp2
(c) sp3
(d) sp3d
(e) sp3d2
Answers to Chapter 7
1. (b) 2. (c) 3. (a) 4. (b) 5. (d) 6. (d) 7. (b) 8. (e) 9. (a) 10. (c)
11. (c) 12. (a) 13. (d) 14. (a) 15. (e)
1.
Choose the molecule that is incorrectly matched with
the electronic geometry about the central atom.
(a) CF4 - tetrahedral
(b) BeBr2 - linear
(c) H2O - tetrahedral
10.
Which of the following four molecules are polar: PH3
OF2 HF SO3?
(a) all except SO3
(b) only HF
(c) only HF and OF2
(d) none of these
(e) all of these
1.
Which statement is false? A sigma molecular orbital
(a) may result from overlap of p atomic orbitals
perpendicular to the molecular axis (side-on).
(b) may result from overlap of p atomic orbitals along
the molecular axis (head-on).
(c) may result from overlap of two s atomic orbitals.
(d) may result from overlap of one s and one p atomic
orbitals.
(e) may be either bonding or antibonding.
11.
Which molecule is nonpolar?
(a) H2Se
(b) BeH2
(c) PF3
(d) CHCl3
(e) SO2
12.
The F-S-F bond angles in SF6 are ______.
(a) 109o28'
(b) 120o only
(c) 90o and 120o
(d) 45o and 90o
(e) 90o and 180o
2.
Carbon monoxide has ten bonding electrons and four
antibonding electrons. Therefore it has a bond order
of
(a) 3
(b) 7
(c) 1
(d) 5/2
(e) 2
13.
Which response contains all of the characteristics
listed that should apply to phosphorus trichloride,
PCl3, and no other characteristics?
(1) trigonal planar
(2) one unshared pair of electrons on P
(3) sp2 hybridized at P
(4) polar molecule
(5) polar bonds
(a) 1, 4, 5
(b) 2, 3, 4
(c) 1, 2, 4
(d) 2, 4, 5
(e) another combination
3.
Which of the following is the correct electron
configuration for C2?
(a) 1s2 2s2 2py2 *1s2 *2s2 *2py2
(b) 1s2 *1s2 2s2 *2s2 2py2 *2pz1 2p1
(c) 1s2 *1s2 2s2 *2s2 2py2 2pz2
(d) 1s2 *1s2 2s2 *2s2 2py1 2pz1
(e) 1s2 *1s2 2s2 *2s2 2py1 *2py1 2pz1 *2pz1
4.
14.
A (pi) bond is the result of the
(a) overlap of two s orbitals.
(b) overlap of an s and a p orbital.
(c) overlap of two p orbitals along their axes.
(d) sidewise overlap of two parallel p orbitals.
(e) sidewise overlap of two s orbitals.
5.
Draw the molecular orbital diagram for the molecular
ion, N2+. The number of electrons in the 2p molecular
orbital is:
(a) 0
(b) 1
(c) 2
(d) 3
(e) 4
15.
A triple bond contains ___ sigma bond(s) and ___ pi
bond(s).
(a) 0, 3
(b) 3, 0
(c) 2, 1
(d) 1, 2
(e) 3, 2
6.
16.
The perchloric acid molecule contains:
(a) 13 lone pairs, 1 bond, and 4 bonds.
(b) 9 lone pairs, no bonds, and 6 bonds.
(c) 8 lone pairs, 2 bonds, and 7 bonds.
(d) 2 lone pairs, 3 bonds, and 4 bonds.
(e) 11 lone pairs, no bonds, and 5 bonds.
What is the bond order in O2+?
(a) 3.5
(b) 2.0
(c) 1.5
(d) 2.5
(e) 0
7.
17.
Draw a complete line-bond or electron-dot formula
for acetic acid and then decide which statement is
incorrect.
(a) One carbon is described by sp2 hybridization.
(b) The molecule contains only one bond.
(c) The molecule contains four lone pairs of valence
electrons.
(d) One carbon is described by sp3 hybridization.
(e) Both oxygens are described by sp3 hybridization.
Answers to Chapter 8
1. (e) 2. (c) 3. (a) 4. (c) 5. (a) 6. (e) 7. (c) 8. (d) 9. (c) 10. (a) 11.
(b) 12. (e) 13. (d) 14. (d) 15. (d) 16. (e) 17. (e)
What is the correct electron configuration for the
molecular ion, B2+?
(a) 1s2 *1s2 2s2 *2s2 2p2
(b) 1s2 *1s2 2s2 *2s2 2py2
(c) 1s2 *1s2 2s2 *2s2 2py2 2pz1
(d) 1s2 *1s2 2s2 *2s2 2p1 2py1
(e) none of the above.
Draw the molecular orbital diagram for B2. The
number of unpaired electrons in the B2 molecule is
_____.
(a) zero
(b) 1
(c) 2
(d) 3
(e) 4
8.
Which one of the following statements is false?
(a) Valence bond theory and molecular orbital theory
can be described as two different views of the same
thing.
(b) When one considers the molecular orbitals
resulting from the overlap of any two specific atomic
orbitals, the bonding orbitals are always lower in
energy than the antibonding orbitals.
(c) Molecular orbitals are generally described as
being more delocalized than hybridized atomic
orbitals.
(d) One of the shortcomings of molecular orbital
theory is its inability to account for a triple bond in
the nitrogen molecule, N2.
(e) One of the shortcomings of valence bond theory is
its inability to account for the paramagnetism of the
oxygen molecule, O2.
9.
Antibonding molecular orbitals are produced by
(a) constructive interaction of atomic orbitals.
(b) destructive interaction of atomic orbitals.
(c) the overlap of the atomic orbitals of two negative
ions
(d) all of these
(e) none of these
Which is the strongest acid?
(a) HClO4
(b) HClO3
(c) HClO2
(d) HClO
(e) HF
6.
Which of these species is probably the weakest acid?
(a) HCl
(b) H3PO4
(c) H2PO4(d) HPO42(e) HNO3
7.
10.
Consider the neutralization reactions between the
following acid-base pairs in dilute aqueous solutions:
(1) CH3COOH + NaOH
(2) HNO3 + Mg(OH)2
(3) H3PO4 + Ba(OH)2
(4) HCl + KOH
(5) H2CO3 + LiOH
Which statement regarding stable heteronuclear
diatomic molecules is false?
(a) All have bond orders greater than zero.
(b) The antibonding molecular orbitals have more of
the character of the more electropositive element
than of the more electronegative element.
(c) Their molecular orbital diagrams are more
symmetrical than those of homonuclear diatomic
molecules.
(d) The bonding molecular orbitals have more of the
character of the more electronegative element than of
the less electronegative element.
(e) The greater is the difference in energy between
two overlapping atomic orbitals, the more polar the
resulting bond is, due to electrons occupying the
resulting bonding molecular orbital.
For which of the reactions is the net ionic equation:
H+ + OHH2O ?
(a) 1, 3
(b) 1, 4, 5
(c) 2, 3
(d) 4
(e) 1
8.
Answers to Chapter 9
(a) 2. (a) 3. (c) 4. (e) 5. (b) 6. (d) 7. (c) 8. (d) 9. (b) 10. (c)
1.
Arrhenius defined an acid as:
(a) a species that can donate a proton.
(b) a species that can accept a proton.
(c) a source of OH- ions in water.
(d) a sourse of H+ ions in water.
(e) a species that can accept a pair of electrons.
9.
Which one of the following is an amphoteric metal
hydroxide?
(a) KOH
(b) Ba(OH)2
(c) Pb(OH)2
(d) LiOH
(e) Mg(OH)2
2.
In the Bronsted-Lowry system, a base is defined as:
(a) a proton donor.
(b) a hydroxide donor.
(c) an electron-pair acceptor.
(d) a water-former.
(e) a proton acceptor.
10.
According to the Lewis theory, a base _____ .
(a) is a proton acceptor.
(b) is a proton donor.
(c) makes available a share in a pair of electrons.
(d) produces OH- ions in aqueous solution.
(e) accepts a share in a pair of electrons.
3.
In the equation: HF + H2O
H3O+ + F(a) H2O is a base and HF is its conjugate acid.
(b) H2O is an acid and HF is the conjugate base.
(c) HF is an acid and F- is its conjugate base.
(d) HF is a base and H3O+ is its conjugate acid.
(e) HF is a base and F- is its conjugate acid.
4.
For the system shown here: HOBr + OHH2O +
OBrBronsted would classify the base species as:
(a) OH- and HOBr
(b) H2O and OH(c) OBr- and OH(d) OBr- and HOBr
(e) H2O and HOBr
5.
Which one of the following represents the net ionic
equation for the reaction of nitric acid with aluminum
hydroxide?
(a) 3H+ + Al(OH)3
Al3+ + 3H2O
(b) 3HNO3 + Al(OH)3
Al(NO3)3 + 3H2O
(c) HNO3 + OHNO3- + H2O
(d) H+ + OHH2O
(e) 3NO3- + Al3+
Al(NO3)3
Answers to Chapter 10
1. (d) 2. (e) 3. (c) 4. (c) 5. (a) 6. (d) 7. (d) 8. (a) 9. (c) 10. (c)
1.
How many grams of Ca(OH)2 are contained in 1500
mL of 0.0250 M Ca(OH)2 solution?
(a) 3.17 g
(b) 2.78 g
(c) 1.85 g
(d) 2.34 g
(e) 4.25 g
2.
(b) 0.270 g
(c) 1.32 g
(d) 0.660 g
(e) 0.859 g
10.
What volume of 12.6 M HCl must be added to enough
water to prepare 5.00 liters of 3.00 M HCl?
(a) 1.19 L
(b) 21.0 L
(c) 0.840 L
(d) 7.56 L
(e) 2.14 L
What is the oxidation number for carbon in CaC2O4?
(a) 0
(b) +2
(c) +3
(d) +4
(e) +6
11.
3.
4.
Balance the molecular equation for the following
redox reaction. What is the sum of the coefficients?
Don't forget coefficients of one. Use the smallest
whole number coefficients possible.
H2SO4(aq) + HI(aq)
I2(s) + SO2(g)
(a) 7
(b) 9
(c) 11
(d) 13
(e) 5
What is the molarity of the salt produced in the
reaction of 200 mL of 0.100 M HCl with 100 mL of
0.500 M KOH?
(a) 0.0325 M
(b) 0.0472 M
(c) 0.0667 M
(d) 0.0864 M
(e) 0.0935 M
What volume of 0.50 M KOH would be required to
neutralize completely 500 mL of 0.25 M H3PO4
solution?
(a) 2.5 x 102 mL
(b) 1.4 x 103 mL
(c) 83 mL
(d) 7.5 x 102 mL
(e) 5.2 x 102 mL
12.
For the reaction between permanganate ion and
sufite ion in basic solution, the unbalanced equation
is:
MnO4- + SO32MnO2 + SO42When this equation is balanced using the smallest
whole number coefficients possible, the number of
OH- ions is
(a) two on the right.
(b) two on the left.
(c) three on the right.
(d) four on the right.
(e) four on the left.
5.
A 0.6745 gram sample of KHP reacts with 41.75 mL of
KOH solution for complete neutralization. What is the
molarity of the KOH solution? (Molecular weight of
KHP = 204 g/mol. KHP has one acidic hydrogen.)
(a) 0.158 M
(b) 0.099 M
(c) 0.139 M
(d) 0.079 M
(e) 0.061 M
13.
Balance the following redox equation in acidic
solution with the smallest whole number coefficients
possible. What is the sum of all the coefficients? (Do
not forget coefficients of one.)
Cu + SO42Cu2+ + SO2 (in acidic solution)
(a) 9
(b) 10
(c) 11
(d) 12
(e) 13
6.
How many equivalents of phosphoric acid are
contained in 300 mL of 4.00 M phosphoric acid?
(Assume the acid is to be completely neutralized by a
base.)
(a) 0.600 eq
(b) 1.20 eq
(c) 2.40 eq
(d) 3.60 eq
(e) 4.80 eq
14.
When the following equation is balanced with the
smallest possible set of integers, what is the sum of all
the coefficients? (Do not forget coefficients of one.)
Cr2O72- + H2S
Cr3+ + S (in acidic solution)
(a) 13
(b) 24
(c) 19
(d) 7
(e) 29
7.
Calculate the normality of a solution that contains 4.5
g of (COOH)2 in 3000 mL of solution? (Assume the
(COOH)2 is to be completely neutralized in an acidbase reaction.)
(a) 0.033 N
(b) 0.045 N
(c) 0.066 N
(d) 0.090 N
(e) 0.12 N
15.
When the following equation is balanced with the
smallest possible set of integers, what is the sum of all
the coefficients? (Do not forget coefficients of one.)
MnO4- + Se2MnO2 + Se (in basic solution)
(a) 20
(b) 22
(c) 24
(d) 26
(e) 28
8.
What volume of 0.100 N HNO3 is required to
neutralize 50.0 mL of a 0.150 N solution of Ba(OH)2?
(a) 50.0 mL
(b) 75.0 mL
(c) 100. mL
(d) 125 mL
(e) 150. mL
9.
How many grams of NaOH would be required to
neutralize all the acid in 75.0 mL of 0.0900 N H2SO4?
(a) 0.540 g
16.
Consider the following unbalanced equation in acidic
solution:
NaClO3 + H2O + I2
HIO3 + NaCl
A 25.0 mL sample of 0.0833 M NaClO3 reacted with
30.0 mL of an aqueous solution of I2. How many
grams of I2 were contained in the I2 solution?
(a) 0.264 g
(b) 0.397 g
(c) 0.236 g
(d) 0.159 g
(e) 0.317 g
Calculate the normality of a NaClO solution if 35.00
mL of the solution is required to react with 0.615 g of
Zn according to the following unbalanced equation:
Zn + ClOZn(OH)2 + Cl- (in basic solution)
(a) 0.537 N
(b) 0.275 N
(c) 0.108 N
(d) 0.366 N
(e) 0.791 N
17.
Consider the following unbalanced net ionic equation:
NO2- + MnO4NO3- + Mn2+ (in acidic solution)
What is the molarity of a sodium nitrite, NaNO2,
solution if 30.0 mL of it just reacts with 0.238 grams
of KMnO4?
(a) 0.410 M
(b) 0.126 M
(c) 0.0502 M
(d) 0.251 M
(e) 0.0316 M
24.
A solution of nitrous acid was standardized in a
reaction where HNO2
NO3- and its concentration
was determined to be 0.100 N nitrous acid. What
volume of this 0.100 N nitrous acid solution would be
required to oxidation of 0.200 g of CoCl2 to CoCl3
according to the following net ionic equation?
Co2+ + HNO2
Co3+ + NO (in acidic solution)
(a) 33.9 mL
(b) 15.4 mL
(c) 7.70 mL
(d) 67.8 mL
(e) 30.8 mL
18.
What is the equivalent weight (in grams) of copper(II)
nitrate for use in a reaction involving the conversion
of copper(II) to copper metal?
(a) 46.9 g/eq
(b) 93.8 g/eq
(c) 187.6 g/eq
(d) 375.2 g/eq
(e) 562.8 g/eq
25.
What is the sum of all coefficients when the following
net ionic equation is balanced using the smallest
whole number coefficients possible? Do not forget
coefficients of one.
MnO4- + Mn2+
MnO2 (in basic solution)
(a) 19
(b) 16
(c) 13
(d) 11
(e) 7
19.
What is the normality of a K2Cr2O7 solution prepared
by dissolving 5.00 g of K2Cr2O7 in 200 mL of solution,
which will be used in the following unbalanced
reaction?
Cr2O72- + SO32Cr3+ + SO42- (in acidic solution)
(a) 0.733 N
(b) 0.226 N
(c) 0.510 N
(d) 0.441 N
(e) 0.810 N
20.
What mass of KMnO4 must be dissolved to prepare
1.25 L of 0.110 N KMnO4 solution? It is used in the
reaction in which MnO4- ions oxidize Fe2+ into Fe3+
ions and are reduced to Mn2+ ions under acidic
conditions?
(a) 4.34 g
(b) 23.8 g
(c) 115 g
(d) 19.1 g
(e) 70.6 g
Answers to Chapter 11
1. (b) 2. (a) 3. (c) 4. (d) 5. (d) 6. (d) 7. (a) 8. (b) 9. (b) 10. (c)
11. (a) 12. (a) 13. (b) 14. (b) 15. (b) 16. (e) 17. (b) 18. (b) 19.
(c) 20. (a) 21. (e) 22. (e) 23. (a) 24. (e) 25. (b)
1.
Which statement is false?
(a) The density of a gas is constant as long as its
temperature remains constant.
(b) Gases can be expanded without limit.
(c) Gases diffuse into each other and mix almost
immediately when put into the same container.
(d) The molecular weight of a gaseous compound is a
non-variable quantity.
(e) Pressure must be exerted on a sample of a gas in
order to confine it.
21.
A 0.250 M solution of Na2C2O4 is to be used in a
reaction in which the C2O42- will be oxidized to CO2.
What is the normality of this Na2C2O4 solution?
(a) 0.250 N
(b) 1.00 N
(c) 0.125 N
(d) 0.0625 N
(e) 0.500 N
2.
22.
23.
What volume of a 0.150 N KI solution is required to
react in basic solution with 34.1 mL of a 0.216 N
solution of KMnO4? The products in the reaction
include MnO2 and IO3-.
(a) 25.4 mL
(b) 37.9 mL
(c) 12.6 mL
(d) 98.2 mL
(e) 49.1 mL
3.
A sample of oxygen occupies 47.2 liters under a
pressure of 1240 torr at 25oC. What volume would it
occupy at 25oC if the pressure were decreased to 730
torr?
(a) 27.8 L
(b) 29.3 L
(c) 32.3 L
(d) 47.8 L
(e) 80.2 L
A sample of nitrogen occupies 5.50 liters under a
pressure of 900 torr at 25oC. At what temperature will
it occupy 10.0 liters at the same pressure?
(a) 32oC
(b) -109oC
(c) 154oC
(d) 269oC
(e) 370oC
(b) C2H6
(c) C2H5
(d) C4H10
(e) C4H12
4.
Under conditions of fixed temperature and amount of
gas, Boyle's law requires that
I. P1V1 = P2V2
II. PV = constant
III. P1/P2 = V2/V1
(a) I only
(b) II only
(c) III only
(d) I, II, and III
(e) another combination
5.
12.
A mixture of 90.0 grams of CH4 and 10.0 grams of
argon has a pressure of 250 torr under conditions of
constant temperature and volume. The partial
pressure of CH4 in torr is:
(a) 143
(b) 100
(c) 10.7
(d) 239
(e) 26.6
13.
The volume of a sample of nitrogen is 6.00 liters at
35oC and 740 torr. What volume will it occupy at STP?
(a) 6.59 L
(b) 5.46 L
(c) 6.95 L
(d) 5.67 L
(e) 5.18 L
What pressure (in atm) would be exerted by a
mixture of 1.4 g of nitrogen gas and 4.8 g of oxygen
gas in a 200 mL container at 57oC?
(a) 4.7
(b) 34
(c) 47
(d) 27
(e) 0.030
6.
The density of chlorine gas at STP, in grams per liter,
is approximately:
(a) 6.2
(b) 3.2
(c) 3.9
(d) 4.5
(e) 1.3
14.
A sample of hydrogen gas collected by displacement
of water occupied 30.0 mL at 24oC on a day when the
barometric pressure was 736 torr. What volume
would the hydrogen occupy if it were dry and at STP?
The vapor pressure of water at 24.0oC is 22.4 torr.
(a) 32.4 mL
(b) 21.6 mL
(c) 36.8 mL
(d) 25.9 mL
(e) 27.6 mL
7.
What pressure (in atm) would be exerted by 76 g of
fluorine gas in a 1.50 liter vessel at -37oC?
(a) 26 atm
(b) 4.1 atm
(c) 19,600 atm
(d) 84
(e) 8.2 atm
15.
8.
What is the density of ammonia gas at 2.00 atm
pressure and a temperature of 25.0oC?
(a) 0.720 g/L
(b) 0.980 g/L
(c) 1.39 g/L
(d) 16.6 g/L
(e) 0.695 g/L
9.
A container with volume 71.9 mL contains water
vapor at a pressure of 10.4 atm and a temperature of
465oC. How many grams of the gas are in the
container?
(a) 0.421 g
(b) 0.183 g
(c) 0.129 g
(d) 0.363 g
(e) 0.222 g
16.
A mixture of 0.50 mol H2(g) and 0.50 mol N2(g) is
introduced into a 15.0 liter container having a pinhole
leak at 30oC. After a period of time, which of the
following is true?
(a) The partial pressure of H2 exceeds that of N2 in the
container.
(b) The partial pressure of N2 exceeds that of H2 in the
container.
(c) The partial pressures of the two gases remain
equal.
(d) The partial pressures of both gases increase above
their initial values.
(e) The partial pressure of H2 in the container
increases above the initial value.
10.
What is the molecular weight of a pure gaseous
compound having a density of 4.95 g/L at -35 oC and
1020 torr?
(a) 24
(b) 11
(c) 72
(d) 120
(e) 44
11.
A 0.580 g sample of a compound containing only
carbon and hydrogen contains 0.480 g of carbon and
0.100 g of hydrogen. At STP, 33.6 mL of the gas has a
mass of 0.087 g. What is the molecular (true) formula
for the compound?
(a) CH3
Which one of the following statements is not
consistent with the kinetic-molecular theory of
gases?
(a) Individual gas molecules are relatively far apart.
(b) The actual volume of the gas molecules
themselves is very small compared to the volume
occupied by the gas at ordinary temperatures and
pressures.
(c) The average kinetic energies of different gases are
different at the same temperature.
(d) There is no net gain or loss of the total kinetic
(translational) energy in collisions between gas
molecules.
(e) The theory explains most of the observed
behavior of gases at ordinary temperatures and
pressures.
17.
If helium effuses through a porous barrier at a rate of
4.0 moles per minute, at what rate (in moles per
minute) would oxygen gas diffuse?
(a) 0.20
(b) 0.50
(c) 2.0
(d) 8.0
(e) 1.41
18.
A real gas most closely approaches the behavior of an
ideal gas under conditions of:
(a) high P and low T
(b) low P and high T
(c) low P and T
(d) high P and T
(e) STP
25.
For a gas, which pair of variables are inversely
proportional to each other (if all other conditions
remain constant)?
(a) P, T
(b) P, V
(c) V, T
(d) n, V
(e) n, P
19.
Which one of the following statements about the
following reaction is false?
CH4(g) + 2O2(g)
CO2(g) + 2H2O(g)
(a) Every methane molecule that reacts produces two
water molecules.
(b) If 32.0 g of oxygen reacts with excess methane, the
maximum amount of carbon dioxide produced will be
22.0 g.
(c) If 11.2 liters of methane react with an excess of
oxygen, the volume of carbon dioxide produced at
STP is (44/16)(11.2) liters.
(d) If 16.0 g of methane react with 64.0 g of oxygen,
the combined masses of the products will be 80.0 g.
(e) If 22.4 liters of methane at STP react with 64.0 g of
oxygen, 22.4 liters of carbon dioxide at STP can be
produced.
Answers to Chapter 12
1. (a) 2. (e) 3. (d) 4. (d) 5. (e) 6. (b) 7. (a) 8. (c) 9. (e) 10. (c)
11. (d) 12. (d) 13. (d) 14. (d) 15. (c) 16. (b) 17. (e) 18. (b) 19.
(c) 20. (d) 21. (e) 22. (c) 23. (d) 24. (a) 25. (b)
1.
What type of intermolecular forces are due to the
attraction between temporary dipoles and their
induced temporary dipoles?
(a) metallic bond
(b) London dispersion
(c) hydrogen bond
(d) ionic bond
(e) covalent bond
20.
What total gas volume (in liters) at
and 880
torr would result from the decomposition of 33 g of
potassium bicarbonate according to the equation:
2KHCO3(s)
K2CO3(s) + CO2(g) + H2O(g)
(a) 56 L
(b) 37 L
(c) 10 L
(d) 19 L
(e) 12 L
520oC
2.
What type of interparticle forces holds liquid N2
together?
(a) ionic bonding
(b) London forces
(c) hydrogen bonding
(d) dipole-dipole interaction
(e) covalent bonding
21.
Calculate the weight of KClO3 that would be required
to produce 29.5 L of oxygen measured at 127oC and
760 torr.
2KClO3(s)
2KCl(s) + 3O2(g)
(a) 7.82 g
(b) 12.2 g
(c) 14.6 g
(d) 24.4 g
(e) 73.5 g
3.
22.
Which of the following statements is false?
(a) The properties of N2(g) will deviate more from
ideality at -100oC than at 100oC.
(b) Van der Waal's equation corrects for the nonideality of real gases.
(c) Molecules of CH4(g) at high pressures and low
temperatures have no attractive forces between each
other.
(d) Molecules of an ideal gas are assumed to have no
significant volume.
(e) Real gases do not always obey the ideal gas laws.
23.
The ideal gas law predicts that the molar volume
(volume of one mole) of gas equals:
(a) gRT/PV
(b) (MW)P/RT
(c) 1/2ms-2
(d) RT/P
(e) 22.4 L at any temperature and pressure
24.
Three 1.0 liter flasks are filled with H2, O2 and Ne,
respectively, at STP.
Which of the following statements is true?
(a) Each flask has the same number of gas molecules.
(b) The velocity of the gas molecules is the same in
each flask.
(c) The density of each gas is the same.
(d) There are twice as many O2 and H2 molecules as
Ne atoms.
(e) None of the above is true.
4.
Which response includes only those compounds that
can exhibit hydrogen bonding?
CH4, AsH3, CH3NH2, H2Te, HF
(a) AsH3, H2Te
(b) AsH3, CH3NH2
(c) CH4, AsH3, H2Te
(d) CH3NH2, HF
(e) HF, H2Te
Which of the following boils at the highest
temperature?
(a) CH4
(b) C2H6
(c) C3H8
(d) C4H10
(e) C5H12
5.
Which probably has the lowest boiling point at 1.00
atm?
(a) HF
(b) HCl
(c) HBr
(d) HI
(e) H2SO4
6.
The normal boiling point of a liquid is
(a) the temperature at which the vapor pressure
equals 760 torr.
(b) the temperature above which the substance
cannot exist as a liquid regardless of the pressure.
(c) the temperature at which the gas molecules have
more kinetic energy than the molecules in the liquid.
(d) the only temperature at which there can be
equilibrium between liquid and gas.
(e) the temperature at which the liquid will usually
boil.
7.
Which of the following changes would increase the
vapor pressure of a liquid?
1. an increase in temperature
2. an increase in the intermolecular forces in the
liquid
3. an increase in the size of the open vessel containing
the liquid
(a) 1 and 2 only
(b) 1 and 3 only
(c) 1 only
(d) 2 only
(e) 3 only
(a) At the temperature and pressure at point 4, Y(g)
will spontaneously convert to Y(l).
(b) At 0oC and 1200 torr, Y exists as a solid.
(c) At the pressure and temperature of point 1, Y(s)
will spontaneously convert to Y(g) and no Y(l) is
possible.
(d) At the pressure and temperature at point 3, Y(s)
Y(g).
(e) At the temperature and pressure at point 2, Y(l)
8.
For water (m.p. 0oC, b.p. 100oC)
Heat of fusion = 333 J/g @ 0oC
Heat of vaporization = 2260 J/g @ 100oC
Specific Heat (solid) = 2.09 J/goC
Specific Heat (liquid) = 4.18 J/goC
Specific Heat (gas) = 2.03 J/goC
Calculate the amount of heat (in kJ) that must be
absorbed to convert 108 g of ice at 0oC to water at
70oC.
(a) 77
(b) 68
(c) 64
(d) 57
(e) 50
12.
Where on a phase diagram can you locate conditions
under which only one phase exists?
(a) at an intersection of two lines
(b) at the normal boiling point
(c) at an intersection of three lines
(d) in an area bounded by lines
(e) at the triple point
13.
In any cubic lattic, an atom lying at the corner of a
unit cell is shared equally by how many unit cells?
(a) one
(b) two
(c) eight
(d) four
(e) sixteen
9.
For mercury (m.p. -39oC, b.p. 357oC)
Heat of fusion = 11.6 J/g @ -39oC
Heat of vaporization = 292 J/g @ 357oC
Specific Heat (solid) = 0.141 J/goC
Specific Heat (liquid) = 0.138 J/goC
Specific Heat (gas) = 0.104 J/goC
Calculate the amount of heat that must be released to
convert 20.0 g of mercury vapor at 387 oC to liquid
mercury at 307oC (in kJ).
(a) 61.9
(b) 6.56
(c) 6.04
(d) 5.69
(e) 5.10
14.
Which statement is false?
(a) Molecular solids generally have lower melting
points than covalent solids.
(b) Metallic solids exhibit a wide range of melting
points because metallic bonds cover a wide range of
bond strength.
(c) The metallic solid can be viewed as positive ions
closely packed in a sea of valence electrons.
(d) Most molecular solids melt at lower temperatures
than metallic solids.
(e) The interactions among the molecules in
molecular solids are generally stronger than those
among the particles that define either covalent or
ionic crystal lattices.
10.
Which of the following phase changes is(are)
endothermic?
1. melting
3. sublimation 5. deposition
2. vaporization 4. condensation 6. freezing
(a) 1, 2, and 3
(b) 4, 5, and 6
(c) 1 and 2 only
(d) 4 and 6 only
(e) some other combination
15.
Which one of the following classifications is
incorrect?
(a) H2O(s), molecular solid
(b) C4H10(s), molecular solid
(c) KF(s), ionic solid
(d) SiC(s), covalent solid
(e) S(s), metallic solid
11.
According to the phase diagram given for Compound
Y, what description is correct?
Y(g)
16.
Which of the following compounds would be expected
to have the highest melting point?
(a) BaF2
(b) BaCl2
(c) BaBr2
(d) BaI2
(e) H2O
17.
Which observation(s) reflect(s) colligative
properties?
(I) A 0.5 m NaBr solution has a higher vapor pressure
than a 0.5 m BaCl2 solution.
(II) A 0.5 m NaOH solution freezes at a lower
temperature than pure water.
(III) Pure water freezes at a higher temperature than
pure methanol.
(a) only I
(b) only II
(c) only III
(d) I and II
(e) I and III
Which one of the following substances can be melted
without breaking chemical bonds?
(a) sodium sulfate
(b) zinc chloride
(c) sulfur dioxide
(d) silicon dioxide
(e) diamond
Answers to Chapter 13
1. (b) 2. (b) 3. (d) 4. (e) 5. (b) 6. (a) 7. (c) 8. (b) 9. (c) 10. (a)
11. (a) 12. (d) 13. (c) 14. (e) 15. (e) 16. (a) 17. (c)
7.
The vapor pressure of pure water at 85oC is 434 torr.
What is the vapor pressure at 85oC of a solution
prepared from 100 mL of water (density 1.00 g/mL)
and 150 g of diglyme, C6H14O3, a nonvolatile
substance?
(a) 361 torr
(b) 390 torr
(c) 425 torr
(d) 388 torr
(e) 317 torr
1.
Consider the three statements below. Which
statement(s) is(are) true?
1. Hydration is a special case of solvation in which the
solvent is water.
2. The oxygen end of water molecules is attracted
toward Ca2+ ions.
3. The hydrogen end of water molecules is attracted
toward Cl- ions.
(a) 1 only
(b) 2 only
(c) 3 only
(d) 1 and 2 only
(e) 1, 2, and 3
2.
Consider the following pairs of liquids. Which pairs
are miscible?
1. benzene, C6H6, and hexane, C6H12
2. water and methanol, CH3OH
3. water and hexane
(a) 1, 2 only
(b) 2 only
(c) 1 only
(d) 1, 2, 3
(e) 2, 3 only
3.
Calculate the molality of a solution that contains 51.2
g of naphthalene, C10H8, in 500 mL of carbon
tetrachloride. The density of CCl4 is 1.60 g/mL.
(a) 0.250 m
(b) 0.500 m
(c) 0.750 m
(d) 0.840 m
(e) 1.69 m
4.
What is the molality of a solution labeled "8.6%
glucose (C6H12O6) by weight?" (Note: If the question
does not give the solvent, assume it is water.)
(a) 0.26 m
(b) 0.34 m
(c) 0.44 m
(d) 0.52 m
(e) 0.67 m
5.
Calculate the mole fraction of C2H5OH in a solution
that contains 46 grams of ethanol, C2H5OH, and 64
grams of methanol, CH3OH.
(a) 1/3
(b) 0.42
(c) 1/2
(d) 2/3
(e) none of these
8.
The vapor pressure of a solution containing a
nonvolatile solute is directly proportional to the
(a) molality of the solvent.
(b) osmotic pressure of the solute.
(c) molarity of the solvent.
(d) mole fraction of solvent.
(e) mole fraction of solute.
9.
If 4.27 grams of sucrose, C12H22O11, are dissolved in
15.2 grams of water, what will be the boiling point of
the resulting solution? (Kb for water = 0.512 oC/m)
(Note: If the Kf and Kb are not given on the exam, you
can find them on the back of the exam envelope.)
(a) 101.64 oC
(b) 100.42 oC
(c) 99.626 oC
(d) 100.73 oC
(e) 101.42 oC
10.
What are the ideal van't Hoff factors for the following
compounds:
Ba(OH)2, C6H12O6, K3PO4, HNO3 ?
(a) 1, 1, 1, 1
(b) 2, 1, 2, 2
(c) 3, 1, 4, 2
(d) 6, 3, 5, 5
(e) none of the above
11.
Calculate the approximate initial boiling point (in oC)
of a solution of 285 g of magnesium chloride in 2.0 kg
of water. (Assume complete dissociation of the salt.)
(a) 103.1 oC
(b) 101.6 oC
(c) 102.3 oC
(d) 100.8 oC
(e) 104.8 oC
12.
6.
13.
A solution made by dissolving 9.81 g of a nonvolatile
nonelectrolyte in 90.0 g of water boiled at 100.37 oC
at 760 mm Hg. What is the approximate molecular
weight of the substance? (For water, Kb = 0.51 oC/m)
(a) 240 g/mol
(b) 150 g/mol
(c) 79 g/mol
(d) 61 g/mol
(e) 34 g/mol
What is the freezing point of an aqueous 1.00 m NaCl
solution? (Kf = 1.86 oC/m) (Assume complete
dissociation of the salt.)
(a) -1.86 oC
(b) +1.86 oC
(c) -3.72 oC
(d) -0.93 oC
(e) 0.0 oC
14.
A system suffers an increase in internal energy of 80 J
and at the same time has 50 J of work done on it.
What is the heat change of the system?
(a) +130 J
(b) +30 J
(c) -130 J
(d) -30 J
(e) 0 J
4.
A 17.3 mg sample of an organic compound (a nonelectrolyte) was ground up with 420 mg of camphor
to form a homogeneous mixture melting at 170.0 oC.
What is the apparent formula weight of the organic
compound? (Kf of camphor = 37.7 oC/m, m.p. of
camphor = 178.4 oC) (Note: This is a freezing point
depression problem - note the Kf of camphor camphor is the solvent.)
(a) 353 g/mol
(b) 285 g/mol
(c) 231 g/mol
(d) 185 g/mol
(e) 166 g/mol
15.
A 5.000 g sample of methanol, CH3OH, was combusted
in the presence of excess oxygen in a bomb
calorimeter conaining 4000 g of water. The
temperature of the water increased from 24.000 oC to
29.765 oC. The heat capacity of the calorimeter was
2657 J/oC. The specific heat of water is 4.184 J/goC.
Calculate
E for the reaction in kJ/mol.
(a) -314 kJ/mol
(b) -789 kJ/mol
(c) -716 kJ/mol
(d) -121 kJ/mol
(e) -69.5 kJ/mol
5.
A coffee cup calorimeter having a heat capacity of 451
J/oC was used to measure the heat evolved when
0.0300 mol of NaOH(s) was added to 1000 mL of
0.0300 M HNO3 initially at 23.000 oC. The
temperature of the water rose to 23.639 oC. Calculate
H (in kJ/mol NaNO3) for this reaction. Assume the
specific heat of the final solution is 4.18 J/goC; the
density of each solution is 1.00 g/mL; and the
addition of solid does not appreciably affect the
volume of the solution.
HNO3(aq) + NaOH(s)
NaNO3(aq) + H2O(l)
(a) -63.7 kJ/mol
(b) -151 kJ/mol
(c) -2.55 kJ/mol
(d) -81.4 kJ/mol
(e) -98.6 kJ/mol
Calculate the osmotic pressure associated with 50.0 g
of an enzyme of molecular weight 98,000 g/mol
dissolved in water to give 2600 mL of solution at 30.0
oC.
(a) 0.484 torr
(b) 1.68 torr
(c) 1.96 torr
(d) 2.48 torr
(e) 3.71 torr
16.
A 250 mL solution containing 21.4 g of a polymer in
toluene had an osmotic pressure of 0.055 atm at 27
oC. What is the apparent formula weight of the
polymer?
(a) 15,000 g/mol
(b) 18,000 g/mol
(c) 26,000 g/mol
(d) 32,000 g/mol
(e) 38,000 g/mol
6.
The
kJ.
Ho for the following reaction at 298 K is -36.4
1/2 H2(g) + 1/2 Br2(l)
HBr(g)
Calculate
Eo at 298 K. The universal gas constant,
R, is 8.314 J/mol K.
(a) -35.2 kJ
(b) +35.2 kJ
(c) -36.4 kJ
(d) -37.6 kJ
(e) +37.6 kJ
Answers to Chapter 14
1. (e) 2. (a) 3. (b) 4. (d) 5. (a) 6. (d) 7. (a) 8. (d) 9. (b) 10. (c)
11. (c) 12. (b) 13. (c) 14. (d) 15. (e) 16. (e)
7.
Calculate the amount of work done for the conversion
of 1.00 mole of Ni to Ni(CO)4 in the reaction below, at
75oC. Assume that the gases are ideal. The value of R
1.
Which one of the following thermodynamic quantities
is not a state function?
(a) Gibbs free energy
(b) enthalpy
(c) entropy
(d) internal energy
(e) work
2.
At a constant temperature, an ideal gas is compressed
from 6.0 liters to 4.0 liters by a constant external
pressure of 5.0 atm. How much work is done on the
gas?
(a) w = +10 liter atm
(b) w = -10 liter atm
(c) w = +30 liter atm
(d) w = -30 liter atm
(e) The answer cannot be calculated.
3.
is 8.31 J/mol K.
Ni(s) + 4 CO (g)
(a) 1.80 x 103 J
(b) 8.68 x 103 J
(c) -1.80 x 103 J
(d) -8.68 x 103 J
(e) -494 J
Ni(CO)4(g)
8.
9.
All of the following have a standard heat of formation
value of zero at 25oC and 1.0 atm except:
(a) N2(g)
(b) Fe(s)
(c) Ne(g)
(d) H(g)
(e) Hg(l)
10.
For which of the following reactions would the
Ho
for the reaction be labeled
Hfo?
(a) Al(s) + 3/2 H2(g) + 3/2 O2(g)
Al(OH)3(s)
(b) PCl3(g) + 1/2 O2(g)
POCl3(g)
(c) 1/2 N2O(g) + 1/4 O2(g)
NO(g)
(d) CaO(s) + SO2(g)
CaSO3(s)
(e) The
Ho for all these reactions would be labeled
o
Hf .
N2(g) + 3H2(g)
So298 (J/mol K) 191.5
(a) -198.7 J/K
(b) 76.32 J/K
(c) 303.2 J/K
(d) -129.7 J/K
(e) 384.7 J/K
The entropy will usually increase when
I. a molecule is broken into two or more smaller
molecules.
II. a reaction occurs that results in an increase in the
number of moles of gas.
III. a solid changes to a liquid.
IV. a liquid changes to a gas.
(a) I only
(b) II only
(c) III only
(d) IV only
(e) I, II, III, and IV
Ho for the reaction:
Na2O(s) + SO3(g)
Na2SO4(g)
given the following information:
Ho
NaOH(s) + 1/2 H2(g)
-146 kJ
(2) Na2SO4(s) + H2O(l)
SO3(g)
2NaOH(s) +
+418
kJ
(3) 2Na2O(s) + 2H2(g)
4Na(s) + 2H2O(l)
+259
kJ
16.
Calculate
Go for the reaction given the following
information:
2SO2(g) + O2(g)
2SO3(g)
Gfo for SO2(g) = -300.4 kJ/mol
Gfo for SO3(g) = -370.4 kJ/mol
(a) -70.0 kJ
(b) +70.0 kJ
(c) -670.8 kJ
(d) -140.0 kJ
(e) +140.0 kJ
(a) +255 kJ
(b) -435 kJ
(c) -581 kJ
(d) +531 kJ
(e) -452 kJ
11.
Calculate
Horxn for the following reaction at 25.0 oC:
Fe3O4(s) + CO(g)
3FeO(s) + CO2(g)
Hfo
(kJ/mol)
(a) -263 kJ
(b) 54 kJ
(c) 19 kJ
(d) -50 kJ
(e) 109 kJ
-1118
110.5
-272
-393.5
17.
For the following reaction at 25oC,
Ho = +115 kJ
and
So = +125 J/K. Calculate
Go for the reaction
at 25o.
SBr4(g)
S(g) + 2Br2(l)
(a) +152 kJ
(b) -56.7 kJ
(c) +77.8 kJ
(d) +37.1 kJ
(e) -86.2 kJ
12.
13.
Calculate the standard heat of formation,
Hfo, for
FeS2(s), given the following information:
2FeS2(s) + 5O2(g)
2FeO(s) + 4SO2(g)
Horxn = -1370 kJ
Hfo for SO2(g) = -297 kJ/mol
Hfo for FeO(s) = -268 kJ/mol
(a) -177 kJ
(b) -1550 kJ
(c) -774 kJ
(d) -686 kJ
(e) +808 kJ
Estimate the heat of reaction at 298 K for the reaction
shown, given the average bond energies below.
Br2(g) + 3F2(g)
2BrF3(g)
Bond Bond Energy
Br-Br
192 kJ
F-F
158 kJ
Br-F
197 kJ
(a) -516 kJ
(b) -410 kJ
(c) -611 kJ
(d) -665 kJ
(e) -720 kJ
14.
What is the standard entropy change of the reaction
below at 298 K with each compound at the standard
pressure?
192.3
15.
Calculate
(1) Na(s) + H2O(l)
130.6
2NH3(g
)
18.
The heat of vaporization of freon, CCl2F2, is 17.2
kJ/mol at 25oC. What is the change of entropy for one
mole of liquid freon when it vaporizes at 25oC? (Hint:
The vaporization process is at equilibrium and what
is true for
G at equilibrium?)
(a) 57.7 J/K
(b) 0.688 J/K
(c) 5.13 x 103 kJ/K
(d) 3.16 J/K
(e) 239 J/K
19.
Estimate the boiling point of Br2(l) (
S = 93.0 J/K).
Br2(l)
Br2(g)
(a) 85oC
(b) 373oC
(c) 177oC
(d) 59oC
(e) 44oC
H = 30.9 kJ;
20.
For the reaction, A + B
C,
Ho = +30 kJ;
So =
+50 J/K.
Therefore the reaction is:
(a) spontaneous at all temperatures.
(b) nonspontaneous at all temperatures.
(c) spontaneous at temperatures less than 600 K.
(d) spontaneous at temperatures greater than 600 K.
(e) spontaneous only at 25oC.
(a) the order with respect to A is 1 and the order
overall is 1.
(b) the order with respect to A is 2 and the order
overall is 2.
(c) the order with respect to A is 2 and the order
overall is 3.
(d) the order with respect to B is 2 and the order
overall is 2.
(e) the order with respect to B is 2 and the order
overall is 3.
21.
How much heat is absorbed in the complete reaction
of 3.00 grams of SiO2 with excess carbon in the
reaction below?
Ho for the reaction is +624.7 kJ.
SiO2(s) + 3C(s)
SiC(s) + 2CO(g)
(a) 366 kJ
(b) 1.13 x 105 kJ
(c) 5.06 kJ
(d) 1.33 x 104 kJ
(e) 31.2 kJ
4.
Given the following data for this reaction:
NH4+(aq) + NO2-(aq)
N2(g) + 2H2O(l)
EXPT [NH4+] [NO2-]
RATE
22.
The standard heat of combustion of ethanol, C2H5OH,
is 1372 kJ/mol ethanol. How much heat (in kJ) would
be liberated by completely burning a 20.0 g sample?
(a) 686 kJ
(b) 519 kJ
(c) 715 kJ
(d) 597 kJ
(e) 469 kJ
24.
Which statement is false?
(a) The thermodynamic quantity most easily
measured in a "coffee cup" calorimeter is
H.
(b) No work is done in a reaction occurring in a bomb
calorimeter.
(c)
H is sometimes exactly equal to
E.
(d)
H is often nearly equal to
E.
(e)
H is equal to
E for the reaction:
2H2(g) + O2(g)
2H2O(g)
6.
3.
For a reaction 2A + B
Rate = k[A]2[B]
2C, with the rate equation:
0.015 M 0.020 M 0.030 M/s
What are the units of k for the rate law: Rate =
k[A][B]2, when the concentration unit is mol/L?
(a) s-1
(b) s
(c) L mol-1 s-1
(d) L2 mol-2 s-1
(e) L2 s2 mol-2
Given: A + 3B
2C + D
This reaction is first order with respect to reactant A
and second order with respect to reactant B. If the
concentration of A is doubled and the concentration
of B is halved, the rate of the reaction would _____ by a
factor of _____.
(a) increase, 2
(b) decrease, 2
(c) increase, 4
(d) decrease, 4
(e) not change
7.
The decomposition of carbon disulfide, CS2, to carbon
monosulfide, CS, and sulfur is first order with k = 2.8 x
10-7 s-1 at 1000oC.
CS2
CS + S
What is the half-life of this reaction at 1000oC?
(a) 5.0 x 107 s
(b) 4.7 x 10-6 s
(c) 3.8 x 105 s
(d) 6.1 x 104 s
(e) 2.5 x 106 s
8.
The decomposition of dimethylether at 504 oC is first
order with a half-life of 1570 seconds. What fraction
of an initial amount of dimethylether remains after
4710 seconds?
(a) 1/3
(b) 1/6
(c) 1/8
(d) 1/16
(e) 1/32
2.
The speed of a chemical reaction
(a) is constant no matter what the temperature is.
(b) is independent of the amount of contact surface of
a solid involved.
(c) between gases should in all cases be extremely
rapid because the average kinetic energy of the
molecules is great.
(d) between ions in aqueous solution is extremely
rapid because there are no bonds that need to be
broken.
(e) varies inversely with the absolute temperature.
2
5.
1.
The combustion of ethane (C2H6) is represented by
the equation:
2C2H6(g) + 7O2(g)
4CO2(g) + 6H2O(l)
In this reaction:
(a) the rate of consumption of ethane is seven times
faster than the rate of consumption of oxygen.
(b) the rate of formation of CO2 equals the rate of
formation of water.
(c) water is formed at a rate equal to two-thirds the
rate of formation of CO2.
(d) the rate of consumption of oxygen equals the rate
of consumption of water.
(e) CO2 is formed twice as fast as ethane is consumed.
0.010 M 0.020 M 0.020 M/s
3
0.010 M 0.010 M 0.005 M/s
The rate law for the reaction is:
(a) Rate = k[NH4+][NO2-]
(b) Rate = k[NH4+]2[NO2-]2
(c) Rate = k[NH4+]2[NO2-]
(d) Rate = k[NH4+][NO2-]2
(e) none of the above
23.
Which statement is incorrect?
(a) At constant pressure,
H=
E+P V
(b) The thermodynamic symbol for entropy is S.
(c) Gibbs free energy is a state function.
(d) For an endothermic process,
H is negative.
(e) If the work done by the system is greater than the
heat absorbed by the system,
E is negative.
1
9.
The half-life for a first-order reaction is 32 s. What
was the original concentration if, after 2.0 minutes,
the reactant concentration is 0.062 M?
(a) 0.84 M
(b) 0.069 M
(c) 0.091 M
(d) 0.075 M
(e) 0.13 M
15.
At 300 K, the following reaction is found to obey the
rate law: Rate = k[NOCl]2:
2NOCl
2NO + Cl2
Consider the three postulated mechanisms given
below. Then choose the response that lists all those
that are possibly correct and no others.
Mechanism
slow
NOCl
NO + Cl
1
10.
Given that a reaction absorbs energy and has an
activation energy of 50 kJ/mol, which of the following
statements are correct? (Hint: Draw the potential
energy diagram.)
(1) The reverse reaction has an activation energy
equal to 50 kJ/mol.
(2) The reverse reaction has an activation energy less
than 50 kJ/mol.
(3) The reverse reaction has an activation energy
greater than 50 kJ/mol.
(4) The change in internal energy is less than zero.
(5) The change in internal energy is greater than zero.
(a) (1) and (4)
(b) (2) and (4)
(c) (3) and (4)
(d) (2) and (5)
(e) (3) and (5)
11.
12.
Cl + NOCl
Step 1
A+B
Step 2
AB + B
AB
Mechanism
2
Mechanism
3
(2) A2 + A
A2
A3
16.
NOCl2
NO + Cl2
fast
NOCl
NO + Cl
NO +
fast,
equilibrium
slow
(a) 2, 3
(b) 3
(c) 1
(d) 2
(e) 1, 2
17.
Suppose the activation energy of a certain reaction is
250 kJ/mol. If the rate constant at T1 = 300 K is k1 and
the rate constant at T2 = 320 K is k2, then the reaction
is __ times faster at 320 K than at 300 K. (Hint: Solve
for k2/k1.)
(a) 3 x 10-29
(b) 0.067
(c) 15.0
(d) 525
(e) 3 x 10-28
18.
What is the activation energy (in kJ) of a reaction
whose rate constant increases by a factor of 100 upon
increasing the temperature from 300 K to 360 K?
(a) 27
(b) 35
(c) 42
(d) 53
(e) 69
C+
fast, equilibrium
(3) A3 + B
A + C + D fast
According to the mechanism, the rate law will be:
(a) Rate = k[A]2
(b) Rate = k[A][B]
(c) Rate = k[A]2[B]
(d) Rate = k[A]
(e) Rate = k[A]3
slow
A correct reaction mechanism for a given reaction
usually is:
(a) the same as its balanced chemical equation.
(b) obvious if its heat of reaction is known.
(c) obvious if its reaction order is known.
(d) sometimes difficult to prove.
(e) obvious if its activation energy is known.
AB2 fast
slow
NOCl2 +
Overall: 2NOCl
2NO + Cl2
14.
(1) A + A
2NOCl
NO
NOCl + Cl
Cl2
slow
A possible mechanism for the reaction, 2A + B
D, is:
fast
Overall: 2NOCl
2NO + Cl2
AB2 occurs by the
Overall A + 2B
AB2
The rate law expression must be Rate = _________.
(a) k[A]
(b) k[B]
(c) k[A][B]
(d) k[B]2
(e) k[A][B]2
2NO
Overall: 2NOCl
2NO + Cl2
13.
Suppose the reaction: A + 2B
following mechanism:
fast
NOCl2 + NO
+ Cl2
If reaction A has an activation energy of 250 kJ and
reaction B has an activation energy of 100 kJ, which of
the following statements must be correct?
(a) If reaction A is exothermic and reaction B is
endothermic then reaction A is favored kinetically.
(b) At the same temperature the rate of reaction B is
greater than the rate of reaction A.
(c) The energy of reaction A must be greater than the
energy of reaction B.
(d) The energy of reaction B must be greater than the
energy of reaction A.
(e) The rate of reaction A at 25 oC equals the rate of
reaction B at 100 oC.
If the activation energy in the forward direction of an
elementary step is 52 kJ and the activation energy in
the reverse direction is 74 kJ, what is the energy of
reaction
E for this step?
(a) 22 kJ
(b) -22 kJ
(c) 52 kJ
(d) -52 kJ
(e) 126 kJ
NOCl2
19.
Most reactions are more rapid at high temperatures
than at low temperatures. This is consistent with:
(I) an increase in the activation energy with
increasing temperature.
(II) an increase in the rate constant with increasing
temperatures.
(III) an increase in the percentate of "high energy"
collisions with increasing temperature.
(a) only I
(b) only II
(c) only III
(d) only I and II
(e) only II and III
(d) The rate law is: rate = k[H2S][O2].
(e) The rate law cannot be determined from the
information given.
20.
Which items correctly complete the following
statment?
A catalyst can act in a chemical reaction to:
(I) increase the equilibrium constant.
(II) lower the activation energy.
(III) decrease
E for the reaction.
(IV) provide a new path for the reaction.
(a) only I & II
(b) only II & III
(c) only III & IV
(d) only I & III
(e) only II & IV
Answers to Chapter 16
1. (e) 2. (d) 3. (c) 4. (d) 5. (d) 6. (b) 7. (e) 8. (c) 9. (a) 10. (d)
11. (b) 12. (b) 13. (c) 14. (e) 15. (d) 16. (d) 17. (d) 18. (e) 19.
(e) 20. (e) 21. (a) 22. (e) 23. (b) 24. (b) 25. (e)
1.
When the system A + B
C + D is at equilibrium,
(a) the sum of the concentrations of A and B must
equal the sum of the concentrations of C and D.
(b) the forward reaction has stopped.
(c) both the forward and the reverse reactions have
stopped.
(d) the reverse reaction has stopped.
(e) neither the forward nor the reverse reaction has
stopped.
21.
A catalyst:
(a) actually participates in the reaction.
(b) changes the equilibrium concentration of the
products.
(c) does not affect a reaction energy path.
(d) always decreases the rate for a reaction.
(e) always increases the activation energy for a
reaction.
22.
23.
Which statement is false?
(a) If a reaction is thermodynamically spontaneous it
may occur rapidly.
(b) If a reaction is thermodynamically spontaneous it
may occur slowly.
(c) Activation energy is a kinetic quantity rather than
a thermodynamic quantity.
(d) If a reaction is thermodynamically
nonspontaneous, it will not occur spontaneously.
(e) If a reaction is thermodynamically spontaneous, it
must have a low activation energy.
Which of the following statements are true?
(1) Reactions with more negative values of
Go are
spontaneous and proceed at a higher rate than those
with less negative values of
Go .
(2) The activation energy, Ea, is usually about the
same as
E for a reaction.
(3) The activation energy for a reaction does not
change significantly as temperature changes.
(4) Reactions usually occur at faster rates at higher
temperatures.
(a) 1, 2, 4
(b) 3, 4
(c) 1, 2, 3
(d) 2, 3, 4
(e) 1, 2, 3, 4
2.
2SO3(g)
2SO2(g) + O2(g)
The conventional equilibrium constant expression
(Kc) for the system as described by the above
equation is:
(a) [SO2]2/[SO3]
(b) [SO2]2[O2]/[SO3]2
(c) [SO3]2/[SO3]2[O2]
(d) [SO2][O2]
(e) none of these
3.
Consider the following reversible reaction. In a 3.00
liter container, the following amounts are found in
equilibrium at 400 oC: 0.0420 mole N2, 0.516 mole H2
and 0.0357 mole NH3. Evaluate Kc.
N2(g) + 3H2(g)
4.
If the equilibrium constant for the reaction
A + 2B
C + 5/2 D
has a value of 4.0, what is the value of the equilibrium
constant for the reaction
2C + 5D
at the same temperature?
(a) 0.25
(b) 0.063
(c) 2.0
(d) 8.0
(e) 16
24.
When the concentration of reactant molecules is
increased, the rate of reaction increases. The best
explanation is: As the reactant concentration
increases,
(a) the average kinetic energy of molecules increases.
(b) the frequency of molecular collisions increases.
(c) the rate constant increases.
(d) the activation energy increases.
(e) the order of reaction increases.
25.
For the reaction, 2H2S(g) + O2(g)
2S(s) + 2H2O(l),
which one of the following statements is absolutely
true?
(a) The reaction is first order with respect to H2S and
second order with respect to O2.
(b) The reaction is fourth order overall.
(c) The rate law is: rate = k[H2S]2[O2].
2NH3(g)
(a) 0.202
(b) 1.99
(c) 16.0
(d) 4.94
(e) 0.503
2A + 4B
5.
At 445oC, Kc for the following reaction is 0.020.
2HI(g)
H2(g) + I2(g)
A mixture of H2, I2, and HI in a vessel at 445oC has the
following concentrations: [HI] = 2.0 M, [H2] = 0.50 M
and [I2] = 0.10 M. Which one of the following
statements concerning the reaction quotient, Qc, is
TRUE for the above system?
(a) Qc = Kc; the system is at equilibrium.
(b) Qc is less than Kc; more H2 and I2 will be produced.
(c) Qc is less than Kc; more HI will be produced.
(d) Qc is greater than Kc; more H2 and I2 will be
produced.
(e) Qc is greater than Kc; more HI will be produced.
2H2O(g)
2H2(g) + O2(g)
Given that the forward reaction (the conversion of
"left-hand" species to "right-hand" species) is
endothermic, which of the following changes will
decrease the equilibrium amount of H2O?
(a) adding more oxygen
(b) adding a solid phase calalyst
(c) decreasing the volume of the container (the total
pressure increases)
(d) increasing the temperature at constant pressure
(e) adding He gas
6.
Nitrosyl chloride, NOCl, dissociates on heating as
shown below. When a 1.50 gram sample of pure NOCl
is heated at 350oC in a volume of 1.00 liter, the
percent dissociation is found to be 57.2%. Calculate Kc
for the reaction as written.
NOCl(g)
(a) 0.876
(b) 9.26
(c) 0.107
(d) 1.75 x 10-4
(e) 0.0421
NO(g) + 1/2 Cl2(g)
12.
The conventional equilibrium constant expression
(Kc) for the system below is:
2ICl(s)
(a) [I2][Cl2]/[ICl]2
(b) [I2][Cl2]/2[ICl]
(c) [Cl2]
(d) ([I2] + [Cl2])/2[ICl]
(e) [Cl2]/[ICl]2
7.
A quantity of HI was sealed in a tube, heated to 425oC
and held at this temperature until equilibrium was
reached. The concentration of HI in the tube at
equilibrium was found to be 0.0706 mol/L. Calculate
the equilibirum concentration of H2 (and I2). For the
gas-phase reaction,
H2 + I2
2HI
13.
Consider the equilibrium system:
Kc = 54.6 at 425oC
2ICl(s)
I2(s) + Cl2(g)
Which of the following changes will increase the total
amount of of Cl2 that can be produced?
(a) removing some of the I2(s)
(b) adding more ICl(s)
(c) removing the Cl2 as it is formed
(d) decreasing the volume of the container
(e) all of the above
10-3
(a) 9.55 x
M
(b) 1.17 x 10-3 M
(c) 1.85 x 10-4 M
(d) 4.78 x 10-3 M
(e) 2.34 x 10-3 M
8.
Consider the reaction:
N2(g) + O2(g)
2NO(g)
14.
Kc = 0.10 at 2000oC
At equilibrium, a 1.0 liter container was found to
contain 0.20 moles of A, 0.20 moles of B, 0.40 moles of
C and 0.40 mole of D. If 0.10 moles of A and 0.10
moles of B are added to this system, what will be the
new equilibrium concentration of A?
Starting with initial concentrations of 0.040 mol/L of
N2 and 0.040 mol/L of O2, calculate the equilibrium
concentration of NO in mol/L
(a) 0.0055 mol/L
(b) 0.0096 mol/L
(c) 0.011 mol/L
(d) 0.080 mol/L
(e) 0.10 mol/L
A(g) + B(g)
(a) 0.37 mol/L
(b) 0.47 mol/L
(c) 0.87 mol/L
(d) 0.23 mol/L
(e) 0.15 mol/L
9.
Kc = 0.040 for the system below at 450oC. If a reaction
is initiated with 0.40 mole of Cl2 and 0.40 mole of PCl3
in a 2.0 liter container, what is the equilibrium
concentration of Cl2 in the same system?
PCl5(g)
Consider the following system in a 1.00 L container:
A(g) + B(g)
2C(g)
The equilibrium concentrations at 200oC were
determined to be:
[A] = 0.200 M
[B] = 3.00 M
The reversible reaction:
16.
Consider the reversible reaction at equilibrium at
392oC:
2A(g) + B(g)
C(g)
The partial pressures are found to be: A: 6.70 atm, B:
10.1 atm, C: 3.60 atm. Evaluate Kp for this reaction.
(a) 7.94 x 10-3
(b) 0.146
(c) 0.0532
(d) 54.5
(e) 121
11.
Consider the gas-phase equilibrium system
represented by the equation:
[C] = 0.500 M
How many moles of A must be added to increase the
concentration of C to 0.700 M at 200oC?
(a) 0.225 mol
(b) 0.305 mol
(c) 0.417 mol
(d) 0.610 mol
(e) 0.700 mol
10.
2SO2(g) + O2(g)
2SO3(g)
has come to equilibrium in a vessel of specific volume
at a given temperature. Before the reaction began, the
concentrations of the reactants were 0.060 mol/L of
SO2 and 0.050 mol/L of O2. After equilibrium is
reached, the concentration of SO3 is 0.040 mol/L.
What is the equilibrium concentration of O2?
(a) 0.010 M
(b) 0.020 M
(c) 0.030 M
(d) 0.040 M
(e) none of these
C(g) + D(g)
15.
PCl3(g) + Cl2(g)
(a) 0.07 M
(b) 0.16 M
(c) 0.11 M
(d) 0.04 M
(e) 0.26 M
I2(s) + Cl2(g)
17.
Kc = 0.040 for the system below at 450oC:
PCl5(g)
PCl3(g) + Cl2(g)
Evaluate Kp for the reaction at 450oC.
(a) 0.40
(b) 0.64
(c) 2.4
(d) 0.052
(e) 6.7 x 10-4
(e) (CH3)3N
4.
In a sample of pure water, only one of the following
statements is always true at all conditions of
temperature and pressure. Which one is always true?
(a) [H3O+] = 1.0 x 10-7 M
(b) [OH-] = 1.0 x 10-7 M
(c) pH = 7.0
(d) pOH = 7.0
(e) [H3O+] = [OH-]
18.
What is the equilibrium constant for a reaction that
has a value of
Go = -41.8 kJ at 100oC?
(a) 1.01
(b) 7.1 x 105
(c) -5.87
(d) 1.4 x 10-6
(e) 13.5
5.
If Kw is 2.9 x 10-15 at 10oC, what is the pH of pure
water at 10oC?
(a) 6.72
(b) 7.00
(c) 7.27
(d) 7.53
(e) none of these
19.
The equilibrium constant at 427oC for the reaction:
N2(g) + 3H2(g)
2NH3(g)
is Kp = 9.4 x 10-5. Calculate the value of
Go for the
reaction at 427o.
(a) -33 kJ
(b) -54 kJ
(c) 54 kJ
(d) 33 kJ
(e) 1.3 J
20.
For a specific reaction, which of the following
statements can be made about K, the equilibrium
constant?
(a) It always remains the same at different reaction
conditions.
(b) It increases if the concentration of one of the
products is increased.
(c) It changes with changes in the temperature.
(d) It increases if the concentration of one of the
reactants is increased.
(e) It may be changed by the addition of a catalyst.
Answers to Chapter 17
6.
The pOH of a solution of NaOH is 11.30. What is the
[H+] for this solution?
(a) 2.0 x 10-3
(b) 2.5 x 10-3
(c) 5.0 x 10-12
(d) 4.0 x 10-12
(e) 6.2 x 10-8
7.
The [H3O+] in a 0.050 M solution of Ba(OH)2 is:
(a) 1.0 x 10-5 M
(b) 5.0 x 10-2 M
(c) 1.0 x 10-13 M
(d) 5.0 x 10-10 M
(e) 2.0 x 10-5 M
8.
What is the approximate pH of a solution labeled 6 x
10-5 M HBr?
(a) 4.2
(b) 4.5
(c) 5.8
(d) 9.8
(e) 8.2
9.
1. (e) 2. (b) 3. (b) 4. (b) 5. (b) 6. (c) 7. (a) 8. (c) 9. (a) 10. (c)
11. (d) 12. (c) 13. (c) 14. (d) 15. (b) 16. (a) 17. (c) 18. (b) 19.
(c) 20. (c)
10.
The pH of a solution is 4.80. What is the concentration
of hydroxide ions in this solution?
(a) 4.2 x 10-9 M
(b) 1.6 x 10-5 M
(c) 3.6 x 10-12 M
(d) 6.3 x 10-10 M
(e) 2.0 x 10-8 M
1.
Which one of the following is a weak acid?
(a) HNO3
(b) HI
(c) HBr
(d) HF
(e) HClO3
11.
A solution in which [H+] = 10-8 M has a pH of ___ and is
___.
(a) 8, acidic
(b) 6, basic
(c) -6, basic
(d) -8, neutral
(e) 8, basic
2.
Which salt is not derived from a strong acid and a
strong soluble base?
(a) MgCl2
(b) Ba(NO3)2
(c) LiClO4
(d) CsBr
(e) NaI
3.
Which one of the following is a strong electrolyte?
(a) H2O
(b) KF
(c) HF
(d) HNO2
What is the pH of 500 mL of solution containing
0.0124 grams of Ca(OH)2?
(a) 11.04
(b) 9.68
(c) 2.96
(d) 3.17
(e) 10.83
12.
The pH of a 0.02 M solution of an unknown weak acid
is 3.7. what is the pKa of this acid?
(a) 5.7
(b) 4.9
(c) 3.2
(d) 2.8
(e) 3.7
13.
What is the approximate pH of a solution labeled
0.050 M HClO?
(a) 5.1
(b) 3.9
(c) 4.4
(d) 2.1
(e) 7.6
23.
What is the pH of 0.060 M NH4Cl?
(a) 5.06
(b) 5.12
(c) 5.18
(d) 5.24
(e) 5.35
14.
What is the pH of a solution labeled 0.30 M (CH3)3N?
(a) 9.5
(b) 10.8
(c) 9.2
(d) 11.7
(e) 12.2
24.
What is the concentration of ammonium chloride in a
solution if its pH is 4.80?
(a) 0.25 M
(b) 0.30 M
(c) 0.45 M
(d) 0.60 M
(e) 0.15 M
15.
Which of the following solutions has the lowest pH at
25oC? (No calculations required.)
(a) 0.2 M sodium hydroxide
(b) 0.2 M hypochlorous acid
(c) 0.2 M ammonia
(d) 0.2 M benzoic acid
(e) pure water
What is the concentration of a sodium acetate
solution if the pH of the solution is 9.19?
(a) 0.30 M
(b) 0.43 M
(c) 2.1 M
(d) 0.068 M
(e) 0.59 M
25.
The pH of 0.15 M trimethylammonium chloride,
(CH3)3NHCl, a salt, is 5.34. What is the percent
hydrolysis?
(a) 0.0031 %
(b) 0.0068 %
(c) 0.0094 %
(d) 0.011 %
(e) 0.022 %
16.
A 0.10 M solution of a weak acid, HX, is 0.059%
ionized. Evaluate Ka for the acid.
(a) 3.8 x 10-9
(b) 6.5 x 10-7
(c) 7.0 x 10-6
(d) 4.2 x 10-6
(e) 3.5 x 10-8
17.
What is the percent ionization of an 1.2 M HF
solution?
(a) 2.4 %
(b) 4.2 %
(c) 0.84 %
(d) 0.082 %
(e) 0.22 %
Answers to Chapter 18
1. (d) 2. (a) 3. (b) 4. (e) 5. (c) 6. (a) 7. (c) 8. (a) 9. (e) 10. (d)
11. (e) 12. (a) 13. (c) 14. (d) 15. (d) 16. (e) 17. (a) 18. (e) 19.
(c) 20. (a) 21. (d) 22. (b) 23. (d) 24. (c) 25. (a)
18.
Which of the following weak acids ionizes to give the
strongest conjugate base?
(a) HClO
(b) CH3COOH
(c) HF
(d) HNO2
(e) HCN
19.
Which of the following is true about a 0.10 M solution
of a weak acid, HX?
(a) [X-] = 0.10 M
(b) pH = 1
(c) [HX] > [H+]
(d) [H+] = 0.10 M
(e) both b and d
1.
Which of the following combinations cannot produce
a buffer solution?
(a) HNO2 and NaNO2
(b) HCN and NaCN
(c) HClO4 and NaClO4
(d) NH3 and (NH4)2SO4
(e) NH3 and NH4Br
2.
What is the pH of a solution composed of 0.20 M NH3
and 0.15 M NH4Cl?
(a) 2.15
(b) 4.62
(c) 8.26
(d) 9.38
(e) 8.89
20.
Calculate the hydrolysis constant for the cyanide ion,
CN-.
(a) 2.5 x 10-5
(b) 1.0 x 10-7
(c) 4.0 x 10-10
(d) 5.6 x 10-10
(e) none of these
3.
Calculate the ratio [CH3COOH]/[NaCH3COO] that gives
a solution with pH = 5.00?
(a) 0.28
(b) 0.36
(c) 0.44
(d) 0.56
(e) 0.63
21.
Calculate the pH of a 0.50 M solution of NaNO2.
(a) 12.18
(b) 5.48
(c) 1.82
(d) 8.52
(e) 7.00
22.
4.
Consider a solution which is 0.10 M in CH3COOH and
0.20 M in NaCH3COO. Which of the following
statements is true?
(a) If a small amount of NaOH is added, the pH
decreases very slightly.
(b) If NaOH is added, the OH- ions react with the
CH3COO- ions.
(c) If a small amount of HCl is added, the pH decreases
very slightly.
(d) If HCl is added, the H+ ions react with CH3COOH
ions.
(e) If more CH3COOH is added, the pH increases.
(d) 2.52
(e) 2.80
10.
Which indicator (identified by a letter) could be used
to titrate aqueous NH3 with HCl solution?
Indicator
Acid Range Color
Color-Change pH
5.
A buffer was prepared by mixing 1.00 mole of
ammonia and 1.00 mole of ammonium chloride to
form an aqueous solution with a total volume of 1.00
liter. To 500 mL of this solution was added 30.0 mL of
1.00 M NaOH. What is the pH of this solution?
(a) 8.96
(b) 9.83
(c) 9.31
(d) 9.11
(e) 9.57
6.
How many grams of NaF would have to be added to
2.00 L of 0.100 M HF to yield a solution with a pH =
4.00?
(a) 300 g
(b) 36 g
(c) 0.84 g
(d) 6.9 g
(e) 60. g
HIn
+
H2O
H3O+
+
9.
Calculate the pH of the solution resulting from the
addition of 20.0 mL of 0.100 M NaOH to 30.0 mL of
0.100 M HNO3.
(a) 1.35
(b) 1.70
(c) 1.95
3.4 - 4.6
(c)
yellow
6.5 - 7.8
(d)
colorless
8.3 - 9.9
(e)
none of these
(a) HCl + NH3
less than 7
(b) HNO3 + Ca(OH)2
equal to 7
(c) HClO4 + NaOH
equal to 7
(d) HClO + NaOH
less than 7
(e) CH3COOH + KOH
greater than
12.
What is the pH at the equivalence point in the
titration of 100.0 mL of 0.20 M ammonia with 0.10 M
hydrochloric acid?
(a) 4.6
(b) 5.2
(c) 7.0
(d) 5.5
(e) 4.9
13.
Calculate the pH of a solution prepared by mixing 300
mL of 0.10 M HF and 200 mL of 0.10 M KOH.
(a) 2.82
(b) 2.96
(c) 3.32
(d) 3.44
(e) 3.53
14.
What is the approximate pH of a solution prepared by
mixing equal volumes of 0.05 M methylamine and
0.20 M hydrochloric acid?
(a) 2.57
(b) 1.12
(c) 1.63
(d) 10.5
(e) 9.8
In-
yellow
red
Which of the responses contains all the true
statements and no others?
(1) The predominant color in its acid range is yellow.
(2) In the middle of the pH range of its color change a
solution containing the indicator will probably be
orange.
(3) At pH = 7.00, a solution containing this indicator
(and no other colored species) will be red. (Hint:
Write the equilibrium constant expression for the
indicator.)
(4) At pH = 7.00, most of the indicator is in the unionized form.
(5) The pH at which the indicator changes color is pH
= 4.
(a) 1, 3, 5
(b) 2, 4
(c) 3, 4, 5
(d) 1, 2, 3, 5
(e) another combination
1.2 - 2.8
blue
Consider the titrations of the pairs of aqueous acids
and bases listed on the left. For which pair is the pH at
the equivalence point stated incorrectly?
Acid-Base Pair
pH at Equiv
8.
Consider an indicator that ionized as shown below for
which its Ka = 1.0 x 10-4
pink
(b)
11.
7.
Calculate the pH that results when the following
solutions are mixed.
(1) 35 mL of 0.20 M formic acid
(2) 55 mL of 0.10 M sodium formate
(3) 110 mL of water
(a) 3.64
(b) 3.11
(c) 4.58
(d) 3.39
(e) 4.20
(a)
15.
Which of the following salts give acidic aqueous
solutions?
(1) KNO3
(2) KCH3COO
(5) (NH4)2SO4
(a) 2, 7, 8
(b) 3, 5
(c) 2, 4, 6
(d) 1, 4, 7, 8
(e) 1, 4, 6
(6) BaCl2
16.
The following titration curve is the kind of curve
expected for the titration of a ____ acid with a ____ base.
(3) NH4N
(7) NaCN
(b) 6.4 x 10-7
(c) 4.1 x 10-8
(d) 3.4 x 10-6
(e) 1.4 x 10-5
5.
The solubility of silver sulfate in water at 100oC is
approximately 1.4 g per 100 mL. What is the solubility
product of this salt at 100oC?
(a) 5.7 x 10-8
(b) 3.5 x 10-7
(c) 8.3 x 10-6
(d) 4.1 x 10-5
(e) 3.6 x 10-4
(a) strong, strong
(b) weak, strong
(c) strong, weak
(d) weak, weak
(e) none of these
6.
17.
Consider the titration of 30.0 mL of 0.20 M nitrous
acid by adding 0.0500 M aqueous ammonia to it. The
pH at the equivalence point is _____. (Note: This is the
titration of a weak acid with a weak base.)
(a) greater than 7
(b) equal to 7
(c) less than 7
(d) cannot be determined without more data (not
including Ka and Kb)
(e) is impossible to predict
7.
What is the molar solubility, s, of Ba3(PO4)2 in terms
of Ksp?
(a) s = Ksp1/2
(b) s = Ksp1/5
(c) s = [Ksp/27]1/5
(d) s = [Ksp/108]1/5
(e) s = [Ksp/4]5
For Cu(OH)2, Ksp = 1.6 x 10-19. What is the molar
solubility of Cu(OH)2?
(a) 3.4 x 10-7 M
(b) 6.4 x 10-7 M
(c) 2.7 x 10-11 M
(d) 5.1 x 10-10 M
(e) 1.7 x 10-10 M
Answers to Chapter 19
8.
Many lead salts are often used as pigments. If PbSO4
were used in an unglazed ceramic bowl, how many
milligrams of lead(II) could dissolve per liter of
water?
(a) 43
(b) 35
(c) 11
(d) 28
(e) 53
1. (c) 2. (d) 3. (d) 4. (c) 5. (c) 6. (e) 7. (a) 8. (d) 9. (b) 10. (b)
11. (d) 12. (b) 13. (d) 14. (b) 15. (b) 16. (a) 17. (c)
1.
The solubility product expression for tin(II)
hydroxide, Sn(OH)2, is
(a) [Sn2+][OH-]
(b) [Sn2+]2[OH-]
(c) [Sn2+][OH-]2
(d) [Sn2+]3[OH-]
(e) [Sn2+][OH-]3
9.
Ag3PO4 would be least soluble at 25oC in
(a) 0.1 M AgNO3
(b) 0.1 M HNO3
(c) pure water
(d) 0.1 M Na3PO4
(e) solubility in (a), (b), (c), or (d) is not different
2.
The solubility product expression for silver(I) sulfide,
using x to represent the molar concentration of
silver(I) and y to represent the molar concentration of
sulfide, is formulated as:
(a) xy
(b) x2y
(c) xy2
(d) x2y2
(e) xy3
3.
Consider the following solubility data for various
chromates at 25oC.
Ksp
Ag2CrO4
BaCrO4
9.0 x 10-12
2.0 x 10-10
PbCrO4
1.8 x 10-14
The chromate that is the most soluble in water at
25oC on a molar basis is:
(a) Ag2CrO4
(b) BaCrO4
(c) PbCrO4
(d) impossible to determine
(e) none of these
4.
The molar solubility of PbBr2 is 2.17 x 10-3 M at a
certain temperature. Calculate Ksp for PbBr2.
(a) 6.2 x 10-6
10.
The molar solubility of PbCl2 in 0.20 M Pb(NO3)2
solution is:
(a) 1.7 x 10-4 M
(b) 9.2 x 10-3 M
(c) 1.7 x 10-5 M
(d) 4.6 x 10-3 M
(e) 8.5 x 10-5 M
11.
When we mix together, from separate sources, the
ions of a slightly soluble ionic salt, the salt will
precipitate if Qsp _____ Ksp, and will continue to
precipitate until Qsp _____ Ksp.
(a) is greater than; equals
(b) is less than; is greater than
(c) is less than; equals
(d) equals; is less than
(e) equals; is greater than
12.
Which of the following pairs of compounds gives a
precipitate when aqueous solutions of them are
mixed? Assume that the concentrations of all
compounds are 1.0 M immediately after mixing.
(a) CuBr2 and K2CO3
(b) HNO3 and NH4I
(c) BaCl2 and KClO4
(d) Na2CO3 and H2SO4
(e) KCl and KNO3
chemical change that occurs at this electrode is called
_______.
(a) anode, oxidation
(b) anode, reduction
(c) cathode, oxidation
(d) cathode, reduction
(e) cannot tell unless we know the species being
oxidized and reduced.
13.
A swimming pool was sufficiently alkaline so that CO2
absorbed from the air produced in the pool a solution
which was 2 x 10-4 M in CO32- M. If the pool water was
originally 4 x 10-3 M in Mg2+, 6 x 10-4 M in Ca2+ and 8 x
10-7 M in Fe2+, then a precipitate should form of:
(a) only MgCO3
(b) only CaCO3
(c) only FeCO3
(d) only CaCO3 and FeCO3
(e) MgCO3, CaCO3 and FeCO3
2.
Which of the following statements is FALSE?
(a) Oxidation and reduction half-reactions occur at
electrodes in electrochemical cells.
(b) All electrochemical reactions involve the transfer
of electrons.
(c) Reduction occurs at the cathode.
(d) Oxidation occurs at the anode.
(e) All voltaic (galvanic) cells involve the use of
electricity to initiate nonspontaneous chemical
reactions.
14.
15.
When equal volumes of the solutions indicated are
mixed, precipitation should occur only for:
(a) 2 x 10-3 M Mg2+ + 2 x 10-3 M OH(b) 2 x 10-1 M Ba2+ + 2 x 10-3 M F(c) 2 x 10-3 M Ca2+ + 2 x 10-2 M OH(d) 2 x 10-3 M Ca2+ + 2 x 10-3 M OH(e) 2 x 10-4 M Pb2+ + 2 x 10-5 M SO42At what pH will Cu(OH)2 start to precipitate from a
solution with [Cu2+] = 0.0015 M?
(a) 9.0
(b) 8.0
(c) 6.0
(d) 9.4
(e) 4.6
3.
4.
16.
What is the pH of a saturated solution of Mg(OH)2?
(a) 3.5
(b) 10.1
(c) 10.9
(d) 10.5
(e) 9.2
17.
Which solid will precipitate first if an aqueous
solution of Na2CrO4 at 25oC is slowly added to an
aqueous solution containing 0.001 M Pb(NO3)2 and
0.100 M Ba(NO3)2 at 25oC?
(a) BaCrO4(s)
(b) NaNO3(s)
(c) PbCrO4(s)
(d) Pb(NO3)2(s)
(e) none of these
Answers to Chapter 20
(2) the
(3) 2 Cl-
(4) Cl2
Cl2 + 2 eO2 + 4 H+ + 4 e-
(8) ele
circuit
(9) oxidation
(a) 2, 6, 8, 9
(b) 1, 5, 7, 9
(c) 2, 5, 7, 9
(d) 1, 6, 8, 10
(e) 2, 6, 8, 10
(10) re
What mass (in grams) of nickel could be electroplated
from a solution of nickel(II) chloride by a current of
0.25 amperes flowing for 10 hours?
(a) 12 g
(b) 5.5 g
(c) 0.046 g
(d) 2.7 g
(e) 6.0 g
6.
Molten AlCl3 is electrolyzed for 5.0 hours with a
current of 0.40 amperes. Metallic aluminum is
produced at one electrode and chlorine gas, Cl2, is
produced at the other. How many liters of Cl2
measured at STP are produced when the electrode
efficiency is only 65%?
(a) 0.55 L
(b) 0.63 L
(c) 0.84 L
(d) 0.98 L
(e) 1.02 L
1.
7.
(6) 2 H
(7) electrons flow from the electrode to the external
circuit
5.
1. (c) 2. (b) 3. (a) 4. (c) 5. (e) 6. (d) 7. (a) 8. (d) 9. (a) 10. (d)
11. (a) 12. (a) 13. (d) 14. (a) 15. (c) 16. (d) 17. (c) 18. (e)
In an electrolytic cell the electrode at which the
electrons enter the solution is called the ______ ; the
During the electrolysis of aqueous KCl solution using
inert electrodes, gaseous hydrogen is evolved at one
electrode and gaseous chlorine at the other electrode.
The solution around the electrode at which hydrogen
gas is evolved becomes basic as the electrolysis
proceeds. Which of the following responses describe
or are applicable to the cathode and the reaction that
occurs at the cathode?
(1) the positive electrode
(5) 2 H2O
18.
A solution is 0.0010 M in both Ag+ and Au+. Some solid
NaCl is added slowly until the solid AgCl just begins to
precipitate. What is the concentration of Au+ ions at
this point? Ksp for AgCl = 1.8 x 10-10 and for AuCl = 2.0
x 10-13.
(a) 2.0 x 10-10 M
(b) 4.5 x 10-7 M
(c) 1.8 x 10-7 M
(d) 3.0 x 10-4 M
(e) 1.1 x 10-6 M
The half-reaction that occurs at the anode during the
electrolysis of molten sodium bromide is:
(a) 2 BrBr2 + 2 e(b) Br2 + 2 e2 Br(c) Na+ + eNa
(d) Na
Na+ + e(e) 2 H2O + 2 e2 OH- + H2
How long (in hours) must a current of 5.0 amperes be
maintained to electroplate 60 g of calcium from
molten CaCl2?
(a) 27 hours
(b) 8.3 hours
(c) 11 hours
(d) 16 hours
(e) 5.9 hours
8.
How long, in hours, would be required for the
electroplating of 78 g of platinum from a solution of
[PtCl6]2-, using an average current of 10 amperes at an
80% electrode efficiency?
(a) 8.4
(b) 5.4
(c) 16.8
(d) 11.2
(e) 12.4
9.
What is the reduction potential for the half-reaction at
25o C:
Al3+ + 3eAl, if [Al3+] = 0.10 M and Eo = -1.66 V ?
(a) -1.84 V
(b) -1.60 V
(c) -1.68 V
(d) -1.66 V
(e) -1.72 V
15.
What is the value of E for the half-cell:
MnO4- (0.010 M) + 8H+ (0.20 M) + 5eM) + 4H2O ?
(a) 1.50 V
(b) 1.86 V
(c) 1.44 V
(d) 1.58 V
(e) 1.52 V
16.
Calculate the potential (in volts) for the voltaic (or
galvanic) cell indicated at 25oC.
Ga / Ga3+ (10-6 M) || Ag+ (10-4 M) / Ag
(a) 1.29 V
(b) 0.97 V
(c) 1.45 V
(d) 1.21 V
(e) 1.37 V
How many faradays are required to reduce 1.00 g of
aluminum(III) to the aluminum metal?
(a) 1.00
(b) 1.50
(c) 3.00
(d) 0.111
(e) 0.250
10.
Which of the following is the strongest oxidizing
agent?
(a) Pb2+
(b) I2
(c) Ag+
(d) Pb
(e) Cu2+
11.
As the cell given below operates, the strip of silver
gains mass (only silver) and the concentration of
silver ions in the solution around the silver strip
decreases, while the strip of lead loses mass and the
concentration of lead increases in the solution around
the lead strip. Which of the following represents the
reaction that occurs at the negative electrode in the
above cell?
Pb / Pb(NO3)2 (1.0 M) || AgNO3 (1.0 M) / Ag
(a) Pb2+ + 2 ePb
(b) Pb
Pb2+ + 2 e(c) Ag+ + eAg
+
(d) Ag
Ag + e(e) none of the above
17.
A concentration cell is constructed by placing
identical Cu electrodes in two Cu2+ solutions. If the
concentrations of the two Cu2+ solutions are 1.0 M
and 0.0020 M, calculate the potential of the cell.
(a) 0.020 V
(b) 1.2 V
(c) 0.030 V
(d) 1.0 V
(e) 0.080 V
18.
What is
Go per mole of dichromate ions for the
reduction of dichromate ions, Cr2O72-, to Cr3+ by
bromide ions, Br-, in acidic solution? (Hint: Use the
standard cell potential.)
(a) +26.3 kJ
(b) -145 kJ
(c) +145 kJ
(d) -26.3 kJ
(e) -53.6 kJ
19.
Estimate the equilibrium constant for the system
indicated at 25oC.
12.
13.
14.
Mn2+ (0.020
3 Mg2+ + 2Al
For a voltaic (or galvanic) cell using Ag,Ag+ (1.0 M)
and Zn,Zn2+ (1.0 M) half-cells, which of the following
statements is incorrect?
(a) The zinc electrode is the anode.
(b) Electrons will flow through the external circuit
from the zinc electrode to the silver electrode.
(c) Reduction occurs at the zinc electrode as the cell
operates.
(d) The mass of the zinc electrode will decrease as the
cell operates.
(e) The concentration of Ag+ will decrease as the cell
operates.
20.
Consider the standard voltaic (or galvanic) cell:
Fe,Fe2+ versus Au,Au3+. Which answer identifies the
cathode and gives the Eo for the cell?
(a) Fe, -0.44 V
(b) Au, 1.94 V
(c) Fe, 1.06 V
(d) Au, 1.06 V
(e) Fe, 1.94 V
21.
3Mg + 2Al3+
~1069
(a)
(b) ~1023
(c) ~10-24
(d) ~10-36
(e) ~10-72
In voltaic cells, such as those diagrammed in your
text, the salt bridge _______ .
(a) is not necessary in order for the cell to work
(b) acts as a mechanism to allow mechanical mixing of
the solutions
(c) allows charge balance to be maintained in the cell
(d) is tightly plugged with firm agar gel through
which ions cannot pass
(e) drives free electrons from one half-cell to the
other
Which of the following statements is(are) true for all
voltaic (or galvanic) cells?
(I) Reduction occurs at the cathode.
(II) The anode gains mass during discharge (note: this
means operation of the cell.)
(III) The voltage is less than or equal to zero.
(a) only III
(b) only II
(c) only I
(d) II and III
(e) I, II, and III
(e) Cu
6.
How much magnesium can be obtained from one
pound of seawater if the concentration
of Mg2+ is 0.13 weight percent? Assume 100%
recovery.
(a) 1.7 grams
(b) 1.3 grams
(c) 0.35 grams
(d) 0.59 grams
(e) 2.5 grams
22.
In the standard notation for a voltaic cell, the double
vertical line "||" represents:
(a) a phase boundary
(b) gas electrode
(c) a wire (metal) connection
(d) a salt bridge
(e) a standard hydrogen electrode
7.
For every 100 pounds of iron ore there are 27.8
pounds of magnetite, Fe3O4. What is the weight
percent iron in this ore?
(a) 72.3%
(b) 20.1%
(c) 27.8%
(d) 16.7%
(e) 23.1%
Answers to Chapter 21
1. (d) 2. (e) 3. (a) 4. (e) 5. (d) 6. (a) 7. (d) 8. (b) 9. (d) 10. (c)
11. (b) 12. (c) 13. (b) 14. (c) 15. (c) 16. (d) 17. (e) 18. (b) 19.
(e) 20. (c) 21. (c) 22. (d)
8.
Calculate the amount of coke necessary to produce
800 g of Fe.
(a) 114 g
(b) 1030 g
(c) 258 g
(d) 172 g
(e) 544 g
1.
Some metals are found in the uncombined free state
while other metals are found in the combined state.
What is a deciding factor?
(a) Metals with negative reduction potentials can
occur in the free state while metals with positive
reduction potentials occur in the combined state.
(b) The active metals can occur in the free state while
the less active metals occur in the combined state.
(c) Metals with positive reduction potentials can
occur in the free state while metals with negative
reduction potentials can occur in the combined state.
(d) There is no way we can predict which metals will
be free or combined.
(e) none of the above
2.
9.
10.
What is the charge on the copper ion in the mineral
azurite, Cu3(CO3)2(OH)2?
(a) 2+
(b) 1+
(c) 0
(d) 1(e) 2How many coulombs of electricity are required to
produce one metric ton (1000 kg) of magnesium?
There are 96,500 coulombs in one faraday.
MgCl2
Mg(l) + Cl2
(a) 4.0 x 109 coulombs
(b) 2.0 x 109 coulombs
(c) 7.9 x 109 coulombs
(d) 1.2 x 1012 coulombs
(e) 5.3 x 108 coulombs
Soluble metal compounds tend to be found in the
_____, whereas insoluble metal compounds tend to be
found in the _____ .
(a) oceans; earth's crust
(b) earth's crust; oceans
(c) salt beds; oceans
(d) oceans; salt beds
(e) rivers; oceans
3.
In the process known as 'roasting,' a(n) _____ is
chemically converted to a(n) _____.
(a) sulfide; oxide
(b) carbonate; oxide
(c) hydroxide; oxide
(d) oxide; sulfate
(e) phosphate; phosphide
A reaction sequence for the reduction of one of the
iron ores is as follows:
2 C(coke) + O2
2 CO
Fe2O3 + 3 CO
2 Fe + 3 CO2
Answers to Chapter 22
1. (c) 2. (a) 3. (a) 4. (e) 5. (c) 6. (d) 7. (b) 8. (c) 9. (a) 10. (c)
4.
5.
Which metal can be found as the free element?
(a) Na
(b) Mn
(c) Fe
(d) Cr
(e) Pt
The Hall-Heroult process is used in the production of:
(a) Mg
(b) Fe
(c) Al
(d) Au
1.
Which of the following is NOT true for the Group 1A
elements?
(a) Most of them are soft, silvery corrosive metals.
(b) Their atomic radii increases with increasing
molecular weight.
(c) They are named the alkaline earth metals.
(d) They are excellent conductors of heat and
electricity.
(e) They exhibit a +1 oxidation state in compounds.
2.
3.
Which element group is the most reactive of all the
metallic elements?
(a) alkali metals
(b) alkaline earth metals
(c) coinage metals
(d) transition metals
(e) Group 2B metals
In a surprisingly large number of their properties
beryllium resembles aluminum, and boron resembles
silicon. Such a relationship is called:
(a) amphoterism
(b) an allotropic relationship
(c) a diagonal relationship
(d) the periodic law
(e) an isoelectronic series
4.
Which of the following properties of the alkaline earth
metals decreases with increasing atomic weight?
(a) ionic radii
(b) ionization energy
(c) atomic radii
(d) activity
(e) atomic number
The nitrate of which of the following cations would
exhibit paramagnetism to the GREATEST extent?
(a) Co3+
(b) Cr3+
(c) Fe3+
(d) Mn3+
(e) V3+
Answers to Chapter 23
1. (c) 2. (a) 3. (c) 4. (b) 5. (b) 6. (c) 7. (e) 8. (d) 9. (b) 10. (d)
11. (c)
1.
Some element groups of the periodic table are more
likely to contain elements that are gases than other
groups. Which of the following groups contains the
greatest number of gaseous elements?
(a) IA
(b) IIA
(c) IVA
(d) VIA
(e) VIII (or 0)
5.
Of the following oxides, the most basic is:
(a) MgO.
(b) Na2O.
(c) P2O3.
(d) BeO.
(e) SO2.
2.
Which of the following is NOT true for the halogens?
(a) They are nonmetals.
(b) They show the -1 oxidation number in most of
their compounds.
(c) The electronic configuration of their outermost
electrons is ns2 np6.
(d) Their compounds with metals are generally ionic
in nature.
(e) Elemental halogens exist as diatomic molecules.
6.
A 300 g sample of CaCO3 was heated until 10.0 L of
CO2 was collected at 50.0oC and 742 torr. What
percentage of the CaCO3 had decomposed?
(a) 6.84%
(b) 9.10%
(c) 12.3%
(d) 15.8%
(e) 20.6%
3.
Which of the following substances is the strongest
reducing agent?
(a) Cl2
(b) Cl(c) Br2
(d) Br(e) I2
7.
What mass of lithium nitride could be formed from
104 g of lithium and excess nitrogen gas?
(a) 35 g
(b) 60 g
(c) 105 g
(d) 140 g
(e) 174 g
8.
The most abundant metal in the earth's crust is:
(a) Cu
(b) Fe
(c) Na
(d) Al
(e) Ca
4.
Chlorine gas is prepared commercially by:
(a) electrolysis of carbon tetrachloride.
(b) oxidation of chloride ion with F2(g).
(c) electrolysis of NaCl(aq).
(d) oxidation of chloride ion with Br2(aq).
(e) electrolysis of AlCl3(aq).
5.
9.
Which element has the electron configuration [Ar] 3d7
4s2?
(a) Fe
(b) Co
(c) Cr
(d) Ti
(e) Zn
6.
Of the oxyacids listed below, which one possesses the
greatest acid strength in water?
(a) HClO4
(b) H2CO3
(c) H3BO3
(d) HClO
(e) HBrO
10.
What is the electron configuration of Mn3+ ion?
(a) [Ar] 4s2 3d10
(b) [Ar] 4s2 3d2
(c) [Ar] 3d5
(d) [Ar] 3d4
(e) [Ar] 4d1 3d3
11.
Which one of the following does not correctly
describe one or all of the hydrogen halides, HX?
(a) Their aqueous solutions are acidic.
(b) HF has the lowest of the H-X bond energies.
(c) HI is the largest.
(d) HCl has the lowest boiling point.
(e) HF exhibits hydrogen bonding.
7.
Draw the correct Lewis formula for chlorous acid. The
structure contains ___ single bonds, ___ double bonds
and ___ lone pairs of electrons.
(a) 2, 1, 5
(b) 3, 0, 7
(c) 1, 2, 4
(d) 2, 1, 5
(e) none of these
(b) polar molecule
(c) extremely soluble in water
(d) forms basic aqueous solutions
(e) none of these
16.
8.
Which of the following has a pyramidal structure
(molecular geometry)?
(a) CBr4
(b) PF3
(c) BF3
(d) OF2
(e) BrCl
Which compound gives photochemical smog a
brownish color?
(a) NO
(b) HNO2
(c) NO2
(d) N2O4
(e) N2O3
17.
9.
Which statement about the Group VIA elements is
false?
(a) All have an outer electronic configuration of ns2
np4.
(b) The electronegativity of Group VIA elements
decreases as one goes down the group.
(c) Most are found in sulfide deposits.
(d) Oxygen has the highest boiling point and melting
point.
(e) Polonium has the smallest first ionization energy.
10.
Which statement about the Group VIA hydrides is
false?
(a) H2S, H2Se and H2Te are all gases at room
temperature and atmospheric pressure.
(b) All are colorless.
(c) All except H2O are toxic.
(d) H2Po has the lowest boiling point.
(e) All are covalent compounds.
What is the major mineral present in phosphate rock?
(a) Ca3(PO4)2
(b) Na2HPO4
(c) Ca10(PO4)6F2
(d) NaH2PO4
(e) Ca10(PO4)6(OH)2
Answers to Chapter 24
1. (e) 2. (c) 3. (d) 4. (c) 5. (b) 6. (a) 7. (b) 8. (b) 9. (d) 10. (d)
11. (e) 12. (c) 13. (b) 14. (b) 15. (e) 16. (c) 17. (a)
1.
The ______ sphere is enclosed in brackets in formulas
for complex species, and it includes the central metal
ion plus the coordinated groups.
(a) ligand
(b) donor
(c) oxidation
(d) coordination
(e) chelating
11.
Which acid listed on the right cannot be obtained by
adding water to the substance on the left?
(a) H2S2O7 - sulfuric acid
(b) SeO2 - selenous acid
(c) SO3 - sulfuric acid
(d) SO2 - sulfurous acid
(e) TeO2 - tellurous acid
2.
Which of the following statements about sulfuric acid
is false?
(a) It is a strong acid.
(b) One mole of sulfuric acid reacts completely with
two moles of potassium hydroxide.
(c) The sulfur atom is sp2 hybridized.
(d) It is often present in acid rain.
(e) During the dilution of sulfuric acid, the correct
method is to add sulfuric acid to water.
3.
In coordination chemistry, the donor atom of a
ligand is
(a) a Lewis acid.
(b) the counter ion
(c) the central metal atom.
(d) the atom in the ligand that shares an electron pair
with the metal.
(e) the atom in the ligand that accepts a share in an
electron pair from the metal.
12.
Consider the coordination compound, Na2[Pt(CN)4].
The Lewis acid is
(a) [Pt(CN)4]2(b) Na+
(c) Pt
(d) Pt2+
(e) CN-
13.
What maximum mass of sulfuric acid can be produced
from the sulfur contained in 100 kilograms of iron
pyrite that is 75.0% FeS2?
(a) 84.4 kg
(b) 123 kg
(c) 136 kg
(d) 144 kg
(e) 168 kg
4.
14.
In which one of the following is the oxidation state of
nitrogen given incorrectly?
(a) N2O3, +3
(b) N2H4, +2
(c) HNO3, +5
(d) NaNO2, +3
(e) H2N2O2, +1
15.
Which of the following does not correctly describe
ammonia?
(a) pyramidal molecule
5.
Consider the coordination compound, K2[Cu(CN)4]. A
coordinate covalent bond exists between
(a) K+ and CN(b) Cu2+ and CN(c) K+ and [Cu(CN)4]2(d) C and N in CN(e) K+ and Cu2+
Given the list of ligands and their corresponding
names, choose the pair that disagree.
LIGAND NAME
(a) OH-
hydroxo
(b) CN-
cyanide
(c) Cl(d) H2O
chloro
aqua
(e) NH3
ammine
14.
(Valence Bond Theory) The coordination complex,
[Cu(OH2)6]2+ has one unpaired electron. Which of the
following statements are true?
(1) The complex is octahedral.
(2) The complex is an outer orbital complex.
(3) The complex is d2sp3 hybridized.
(4) The complex is diamagnetic.
(5) The coordination number is 6.
(a) 1, 4
(b) 1, 2, 5
(c) 2, 3, 5
(d) 2, 3
(e) 4, 5
6.
Select the correct IUPAC name for: [FeF4(OH2)2](a) diaquatetrafluoroiron(III) ion
(b) diaquatetrafluoroferrate(III) ion
(c) diaquatetrafluoroiron(I) ion
(d) diaquatetrafluoroferrate(I) ion
(e) none of these
7.
8.
9.
10.
11.
12.
13.
Select the correct IUPAC name for: [Co(NH3)6]2+
(a) hexammoniacobaltate(II) ion
(b) hexaamminecobaltate(II) ion
(c) hexammoniacobalt(II) ion
(d) hexaamminecobalt(II) ion
(e) hexammoniacobalt ion
15.
(Crystal Field Theory) Which one of the following
statements is FALSE?
(a) In an octahedral crystal field, the d electrons on a
metal ion occupy the eg set of orbitals before they
Which name-formula combination is NOT correct?
occupy the t2g set of orbitals.
(b) Diamagnetic metal ions cannot have an odd
FORMULA
NAME
number of electrons.
(a) [Co(NH3)4(OH2)I]SO4
tetraammineaquaiodocobalt(III) sulfate
(c) Low spin complexes can be paramagnetic.
(b) K[Cr(NH3)2Cl4]
potassium diamminetetrachlorochromate(III)
(d) In high spin octahedral complexes,
oct is less
(c) [Mn(CN)5]2pentacyanomanganate(II) ion
than the electron pairing energy, and is relatively very
small.
(d) [Ni(CO)4]
tetracarbonylnickel(0)
(e) Ca[PtCl4]
calcium tetrachloroplatinate(II) (e) Low spin complexes contain strong field ligands.
16.
(Crystal Field Theory) When the valence d orbitals of
What is the oxidation number of the central metal
the central metal ion are split in energy in an
atom in the coordination compound
octahedral ligand field, which orbitals are raised least
[Pt(NH3)3Cl]Cl?
in energy?
(a) -1
(a) dxy and dx2-y2
(b) 0
(b) dxy, dxz and dyz
(c) +1
(c) dxz and dyz
(d) +2
(d) dxz, dyz and dz2
(e) +3
(e) dx2-y2 and dz2
17.
(Valance Bond Theory) Magnetic measurements
(Crystal Field Theory) How many unpaired electrons
indicate that [Co(OH2)6]2+ has 3 unpaired electrons.
are there in a strong field iron(II) octahedral
Therefore, the hybridization of the metal's orbitals in
complex?
[Co(OH2)6]2+ is:
(a) 0
3
(a) sp
(b) 1
(b) sp2d
(c) 2
(c) dsp2
(d) 4
3
2
(d) sp d
(e) 6
(e) d2sp3
18.
(Crystal Field Theory) Consider the complex ion
Which one of the following complexes can exhibit
[Mn(OH2)6]2+ with 5 unpaired electrons. Which
geometrical isomerism?
response includes all the following statements that
(a) [Pt(NH3)2Cl2] (square planar)
are true, and no false statements?
I. It is diamagnetic.
(b) [Zn(NH3)2Cl2] (tetrahedral)
II. It is a low spin complex.
(c) [Cu(NH3)4]2+ (square planar)
III. The metal ion is a d5 ion.
(d) [Co(NH3)5Cl]2+ (octahedral)
IV. The ligands are weak field ligands.
(e) [Cu(CN)2](linear)
V. It is octahedral.
(a) I, II
(b) III, IV, V
A molecule that cannot be superimposed on its mirror
(c) I, IV
image is said to exhibit which of the following?
(d) II, V
(a) geometrical isomerism
(e) III, IV
(b) optical isomerism
19.
(c) linkage isomerism
(Crystal Field Theory) Consider the violet-colored
(d) reactive isomerism
compound, [Cr(OH2)6]Cl3 and the yellow compound,
(e) coordination isomerism
[Cr(NH3)6]Cl3. Which of the following statements is
false?
In which one of the following species does the
transition metal ion have d3 electronic configuration?
(a) [Cr(NH3)6]3+
(b) [Co(OH2)6]2+
(c) [CoF6]3(d) [Fe(CN)6]3(e) [Ni(OH2)6]2+
(e) none of these
3.
The mass defect for an isotope was found to be 0.410
amu/atom. Calculate the binding energy in kJ/mol of
atoms. (1 J = 1 kg m2/s2)
(a) 3.69 x 1010 kJ/mol
(b) 1.23 x 1020 kJ/mol
(c) 3.69 x 1013 kJ/mol
(d) 1.23 x 103 kJ/mol
(e) 1.23 x 1023 kJ/mol
4.
Calculate the binding energy per nucleon (in units of
MeV) for 9Be, for which the atomic mass is 9.01219
amu. Particle masses in amu are: proton = 1.007277;
neutron = 1.008665; electron = 0.0005486.
Conversion factor for E = mc2 is 931 MeV/amu.
(a) 6.46 MeV
(b) 6.33 MeV
(c) 6.23 MeV
(d) 11.39 MeV
(e) 56.93 MeV
(a) Both chromium metal ions are paramagnetic with
3 unpaired electrons.
(b)
oct for [Cr(NH3)6]3+ is calculated directly from
the energy of yellow light.
(c)
oct for [Cr(OH2)6]3+ is less than
oct for
[Cr(NH3)6]3+.
(d) A solution of [Cr(OH2)6]Cl3 transmits light with an
approximate wavelength range of 4000 - 4200
angstroms.
(e) The two complexes absorb their complementary
colors.
5.
Which isotope below has the highest nuclear binding
energy per gram? No calculation is necessary.
(a) 4He
(b) 16O
(c) 32S
(d) 55Mn
(e) 238U
6.
Which of the following statements is incorrect?
(a) Mass defect is the amount of matter that would be
converted into energy if a nucleus were formed from
initially separated protons and neutrons.
(b) Nuclear binding energy is the energy released in
the formation of an atom from subatomic particles.
(c) Nuclei with highest binding energies are the most
stable nuclei.
(d) Einstein postulated the Theory of Relativity in
which he stated that matter and energy are
equivalent.
(e) Mass number is the sum of all protons and
electrons in an atom.
20.
(Crystal Field Theory) Strong field ligands such as CN:
(a) usually produce high spin complexes and small
crystal field splittings.
(b) usually produce low spin complexes and small
crystal field splittings.
(c) usually produce low spin complexes and high
crystal field splittings.
(d) usually produce high spin complexes and high
crystal field splittings.
(e) cannot form low spin complexes.
7.
A positron has a mass number of _____, a charge of
_____, and a mass equal to that of a(an) _____.
(a) 0, 1+, proton
(b) 1, 2+, proton
(c) 0, 1+, electron
(d) 1, 2+, electron
(e) 0, 0, proton
Answers to Chapter 25
1. (d) 2. (d) 3. (d) 4. (b) 5. (b) 6. (b) 7. (d) 8. (c) 9. (d) 10. (d)
11. (a) 12. (b) 13. (a) 14. (b) 15. (a) 16. (b) 17. (a) 18. (b) 19.
(b) 20. (c)
8.
1.
The "magic numbers" for atoms are
(a) numbers of electrons that confer atomic stability.
(b) numbers of protons and/or neutrons that confer
nuclear stability.
(c) n/p ratios that confer nuclear stability.
(d) atomic masses that confer nuclear stability.
(e) atomic masses that indicate fissile isotopes.
2.
The actual mass of a 37Cl atom is 36.966 amu.
Calculate the mass defect (amu/atom) for a 37Cl atom.
(a) 0.623 amu
(b) 0.388 amu
(c) 0.263 amu
(d) 0.341 amu
9.
Emission of which one of the following leaves both
atomic number and mass number unchanged?
(a) positron
(b) neutron
(c) alpha particle
(d) gamma radiation
(e) beta particle
Which type of radiation is the least penetrating?
(a) alpha
(b) beta
(c) gamma
(d) x-ray
(e) neutron
10.
A radioisotope of argon, 35Ar, lies below the "band of
stability: (n/p ratio too low). One would predict that it
decays via _____.
(a) neutron emission
(b) beta emission
(c) positron emission
(d) alpha emission
(e) fission
(d) 3820 years
(e) 9080 years
18.
11.
Which of the following describes what occurs in the
fission process?
(a) A heavy nucleus is fragmented into lighter ones.
(b) A neutron is split into a neutron and proton.
(c) Two light nuclei are combined into a heavier one.
(d) A proton is split into three quarks.
(e) A particle and anti-particle appear in an area of
high energy density.
A Geiger-Muller tube is a _____ .
(a) gas ionization detector
(b) cloud chamber
(c) fluorescence detector
(d) spectrophotometer
(e) photographic detector
12.
The half life of 231Pa is 3.25 x 104 years. How much of
an initial 10.40 microgram sample remains after 3.25
x 105 years?
(a) 0.0102 micrograms
(b) 0.240 micrograms
(c) 2.18 micrograms
(d) 0.0240 micrograms
(e) 1.04 micrograms
19.
Which of the following statements about nuclear
fission is always correct?
(a) Very little energy is released in fission processes.
(b) Nuclear fission is an energetically favorable
process for heavy atoms.
(c) Due to its instability, 56Fe readily undergoes
fission.
(d) In fission reactions, a neutron is split into a proton
and an electron.
(e) All nuclear fission reactions are spontaneous.
13.
Consider the case of a radioactive element X which
decays by electron (beta) emission with a half-life of 4
days to a stable nuclide of element Z. Which of the
following statements is CORRECT?
(a) After 8 days the sample will consist of one-fourth
element Z and three-fourths element X.
(b) Element Z will weigh exactly the same as element
X when decay is complete (weighed to an infinite
number of significant figures).
(c) 2.0 g of element X is required to produce 1.5 g of
element Z after 8 days (to 2 significant figures).
(d) If element X as an atomic number equal to n, then
element X has an atomic number equal to n-1.
(e) None of the above.
14.
Carbon-11 is a radioactive isotope of carbon. Its halflife is 20 minutes. What fraction of the initial number
of C-11 atoms in a sample will have decayed away
after 80 minutes?
(a) 1/16
(b) 1/8
(c) 1/4
(d) 7/8
(e) 15/16
20.
Which one of the following would be most likely to
undergo thermonuclear fusion?
(a) 2H
(b) 4He
(c) 56Fe
(d) 141Ba
(e) 235U
21.
Which one of the following statements about nuclear
reactions is false?
(a) Particles within the nucleus are involved.
(b) No new elements can be produced.
(c) Rate of reaction is independent of the presence of
a catalyst.
(d) Rate of reaction is independent of temperature.
(e) They are often accompanied by the release of
enormous amounts of energy.
22.
Complete and balance the following equation. The
missing term is _____ .
239Pu + alpha particle
_____ + neutron
(a) 2 115Ag
(b) 2 106Rh
(c) 235U
(d) 233Pa
(e) 242Cm
15.
How old is a bottle of wine if the tritium (3H) content
(called activity) is 25% that of a new wine? The halflife of tritium is 12.5 years.
(a) 1/4 yr
(b) 3.1 yr
(c) 25 yr
(d) 37.5 yr
(e) 50 yr
23.
When 59Cu undergoes positron emission, what is the
immediate nuclear product?
(a) 59Ni
(b) 58Ni
(c) 58Cu
(d) 59Zn
(e) 58Zn
16.
A Geiger counter registered 1000 counts/second from
a sample that contained a radioactive isotope of
polonium. After 5.0 minutes, the counter registered
281 counts/second. What is the half-life of this
isotope in seconds?
(a) 87
(b) 110
(c) 164
(d) 264
(e) 2.18
24.
As a result of the process of electron capture ("Kcapture") by 211At, the new isotope formed is:
(a) 210At
(b) 212At
(c) 211Po
(d) 211Rn
(e) 207Bi
17.
The 14C activity of some ancient Peruvian corn was
found to be 10 disintegrations per minute per gram of
C. If present-day plant life shows 15 dpm/g, how old
is the Peruvian corn? The half-life of 14C is 5730 years.
(a) 1455 years
(b) 1910 years
(c) 3350 years
25.
When 235U is bombarded with one neutron, fission
occurs and the products are three neutrons, 94Kr, and
_____ .
(a) 139Ba
(b) 141Ba
(c) 139Ce
(d) 139Xe
(e) 142I
Select the correct IUPAC name for:
Answers to Chapter 26
1. (b) 2. (d) 3. (a) 4. (a) 5. (d) 6. (e) 7. (c) 8. (d) 9. (a) 10. (c)
11. (a) 12. (a) 13. (c) 14. (e) 15. (c) 16. (c) 17. (c) 18. (a) 19.
(b) 20. (a) 21. (b) 22. (e) 23. (a) 24. (c) 25. (a)
(a) 1,4-dimethylcyclopentane
(b) 1,3-dimethylcyclopentane
(c) 2,5-dimethylcyclopentane
(d) 2,3-dimethylcyclopentane
(e) 2,4-dimethylcyclopentane
7.
1.
What makes carbon such a unique element?
(a) Elemental carbon comes in two forms, diamond
and graphite.
(b) Carbon forms four bonds, although the ground
state configuration would predict the formation of
fewer bonds.
(c) Carbon forms covalent bonds rather than ionic
bonds.
(d) To a greater extent than any other element,
carbon can bond to itself to form straight chains,
branched chains and rings.
(e) Carbon has two stable isotopes, carbon-12 and
carbon-13.
8.
The correct name for the compound given below is:
2.
The hybridization of carbon atoms in alkanes is
(a) sp
(b) sp2
(c) sp3
(d) sp3d
(e) sp3d2
The general formula for noncyclic alkenes is:
(a) CnH2n+2
(b) CnH2n
(c) CnH2n-2
(d) CnHn+2
(e) CnHn
(a) 2-methyl-1-butene
(b) 2-ethyl-1-propene
(c) 2-ethyl-1-pentane
(d) 3-methyl-2-butene
(e) pentene
9.
Select the best name for:
3.
A molecule with the formula C3H8 is a(n):
(a) hexane
(b) propane
(c) decane
(d) butane
(e) ethane
4.
Select the correct IUPAC name for:
(a) 4-ethyl-cis-3-octene
(b) 4-ethyl-trans-3-octene
(c) 4-butyl-cis-3-hexene
(d) 5-ethyl-trans-5-octene
(e) 5-ethyl-cis-5-octene
10.
Name the following compound:
5.
(a) 5-methyl-5-ethyloctane
(b) 5-methyl-5-propylheptane
(c) 4-ethyl-4-methyloctane
(d) 3-methyl-3-propyloctane
(e) 3-methyl-3-propylheptane
Select the correct IUPAC name for:
(a) 6-ethyl-4-methylcyclohexene
(b) 6-ethyl-3-methylcyclohexene
(c) 3-ethyl-5-methylcyclohexene
(d) 6-ethyl-4-methylcyclohex-1-ene
(e) 6,4-dialkylcyclohexene
(a) 1,1,3-trimethylpentane
(b) 1-ethyl-1,3-dimethylbutane
(c) 2,4-dimethylhexane
(d) 3,5-dimethylhexane
(e) 3,5,5-trimethylpentane
6.
11.
What is the IUPAC name of the following compound?
(a) CH3CH2OH
(b) CH3OH
(c) CH3CH(OH)CH3
(d) (CH3)C3OH
(e) none of these
18.
Select the IUPAC name for:
(CH3)2CHCH(OH)CH2C(CH3)3.
(a) 2,5,5-trimethyl-3-hexanol
(b) 1,1,4,4-pentamethylbutanol
(c) 1,1-dimethylisopentanol
(d) 2,5-dimethyl-4-hexanol
(e) none of these
(a) 2,6-diethyl-3-nonyne
(b) 2,5-diethyl-3-nonyne
(c) 3,7-dimethyl-5-nonyne
(d) 3,7-dimethyl-4-nonyne
(e) 2,6-diethyl-3-heptyne
12.
The following chemical structure represents a
molecule of what molecular formula?
(a) C8H10
(b) C6H6
(c) C6H8
(d) C8H12
(e) C8H6
19.
Which is NOT a physical property of alcohols or
phenols?
(a) Phenols are generally only slightly soluble in
water.
(b) The solubilities of normal primary alcohols in
water decrease with increasing molecular weight.
(c) The hydroxyl group of an alcohol is nonpolar.
(d) Due to hydrogen bonding, boiling points of
alcohols are much higher than those of corresponding
alkanes.
(e) Boiling points of normal primary alcohols increase
with increasing molecular weight.
20.
Give the IUPAC name of this compound: CH3OCH2CH3.
(a) dimethyl ether
(b) methoxyethane
(c) methylethyloxide
(d) propyl ether
(e) none of the above
13.
How many actual double bonds does the benzene
ring possess?
(a) None, carbon-carbon bonds in benzene are
delocalized around the ring
(b) 1 double bond
(c) 2 double bonds
(d) 3 double bonds
(e) 4 double bonds
21.
The compound below is classified as a _____ .
14.
15.
Para-xylene is the same as:
(a) 1,2-dimethylbenzene
(b) 1,3-diethylbenzene
(c) 1,3-dimethylbenzene
(d) 1,4-diethylbenzene
(e) 1,4-dimethylbenzene
Which of the following formulas represents an
alkene?
(a) CH3CH2CH3
(b) CH3CH3
(c) CH3CH2CHCH2
(d) CH3CH2Cl
(e) CHCH
16.
(a) primary amine
(b) secondary amine
(c) tertiary amine
(d) primary amide
(e) secondary amide
22.
The systematic name for the compound in Problem 21
is _____ .
(a) pentyl amine
(b) methyl-n-propyl amine
(c) diethyl amine
(d) 2-aminopentane
(e) isobutylamine
23.
Select the IUPAC name for the compound below.
What is the name of the following compound?
(a) 1,3-dibromophenol
(b) 2,5-dibromophenol
(c) 2,6-dibromophenol
(d) m-dibromophenol
(e) o-dibromophenol
17.
Which one of the following is a secondary alcohol?
(a) 2,4-dimethylpentanoic acid
(b) 1,1,3-trimethylbutanoic acid
(c) 1-hydroxy-2,4-dimethylpentanone
(d) 2-carboxyisohexane
(e) none of these
24.
Select the best name for:
(c) 20. (b) 21. (b) 22. (d) 23. (a) 24. (a) 25. (b) 26. (b) 27. (d)
28. (a)
1.
Two isomeric forms of a saturated hydrocarbon
(a) have the same structure.
(b) have different compositions of elements.
(c) have the same molecular formula.
(d) have a different content of the isotopes of
hydrogen.
(e) react vigorously with one another.
(a) m-chlorobenzoic acid
(b) o-chlorobenzaldehyde
(c) p-chlorobenzoate
(d) m-chlorosalicylic acid
(e) none of these
2.
25.
Which of the following hydrocarbons does not have
isomers?
(a) C7H16
(b) C6H14
(c) C5H10
(d) C4H8
(e) C3H8
The best classification for the following compound is:
_____ .
3.
(a) aldehyde
(b) ester
(c) ketone
(d) carboxylic acid
(e) alcohol
The name of the alkane isomer of cis-3-hexene is:
(a) 2-methylpentane
(b) 3-methylpentane
(c) n-hexane
(d) 2,3-dimethylbutane
(e) cyclohexane
26.
The compound given below is called _____ .
4.
How many aromatic isomers of dibromobenzene
exist?
(a) 2
(b) 3
(c) 4
(d) 6
(e) 8
(a) butyl acetate
(b) ethyl pentanoate
(c) propyl pentanoate
(d) ethyl butanoate
(e) butyl ethanoate
5.
Which one of the following compounds is an isomer of
CH3CH2CH2CH2OH?
(a) CH3CH2CH2OH
(b) CH3CH(OH)CH3
(c) CH3CH2CH2CHO (Note: This is one way to write an
aldehyde.)
(d) CH3CH2CH2CH3
(e) none of the above
27.
The compound illustrated below is called _____ .
6.
(a) acetamide
(b) formyl acetamide
(c) dimethyl acetate
(d) N,N-dimethylformamide
(e) dimethylamine
Which of the following compounds is a functional
group isomer of C2H5OH, ethanol (ethyl alcohol)?
(a) ethanal, CH3CHO
(b) acetic acid, CH3COOH
(c) diethyl ether, (C2H5)2O
(d) dimethyl ether, (CH3)2O
(e) propanol, C3H7OH
28.
The functional group given below is characteristic of
organic _____ .
7.
CH3CH=CH2
(a) ketones
(b) acids
(c) aldehydes
(d) esters
(e) alcohols
For which of the compounds below are cis-trans
isomers possible?
CH3CH=CHCH2CH3
CH3CH=CHCH3
(1)
(2)
(a) only 2
(b) both 1 and 2
(c) both 2 and 3
(d) all three
(e) only 3
8.
Answers to Chapter 27
1. (d) 2. (c) 3. (b) 4. (c) 5. (c) 6. (b) 7. (b) 8. (a) 9. (a) 10. (c)
11. (d) 12. (a) 13. (a) 14. (e) 15. (c) 16. (c) 17. (c) 18. (a) 19.
Which of the following does NOT exhibit geometric
isomerism? (Hint: draw them!)
(a) 4-octene
(b) 2-pentene
(c) 3-hexene
(d) 2-hexene
(3)
(e) 1-hexene
(d) CH3CH2COOH
(e) CH3CH2OH
9.
Which of the following compounds displays optical
isomerism?
(a) CH2(OH)-CH2(OH)
(b) CH3-CHCl-COOH
(c) CH2=CHCl
(d) CHCl=CHCl
(e) CH3-O-C2H5
16.
How many isomeric alkanes of the molecular formula
C5H12 are there?
(a) 1
(b) 2
(c) 3
(d) 4
(e) 5
17.
How many alcohols are structural isomers with the
formula: C5H11OH?
(a) 5
(b) 6
(c) 7
(d) 8
(e) 9
18.
What is the relationship between the structures
shown?
19.
What is the expected product formed from the
reaction between 2-butene and Cl2?
(a) 1-chlorobutane
(b) 2-chlorobutane
(c) 2,3-dichlorobutane
(d) 2,2-dichlorobutane
(e) 3,3-dichlorobutane
10.
The reaction of ethyne with which of the following
gives CH2Br-CHBrCl?
(a) HCl, then HBr
(b) HCl, then Br2
(c) Cl2, then HBr
(d) Cl2, then Br2
(e) H2, then Br2
11.
Dehydration of an alcohol leads to the formation of an
_____ .
(a) alkene
(b) alkane
(c) alkyne
(d) alkyl halide
(e) aldehyde
12.
A reaction in which a carboxylic acid reacts with a
base to form a salt and water is called _____ .
(a) ionization
(b) esterification
(c) hydrolysis
(d) saponification
(e) neutralization
20.
(a) structural isomers
(b) geometric isomers
(c) conformational structures
(d) identical structures
(e) optical isomers
How many moles of sodium hydroxide will react with
one mole of:
13.
Which of the following statements concerning
conformations is (are) TRUE?
(1) Ethane has an infinite number of conformations.
(2) The eclipsed conformation of a molecule is slightly
more stable and energetically favored than the
staggered conformation.
(3) A conformation is one specific geometry of a
molecule.
(a) 1 only
(b) 2 only
(c) 1 and 2
(d) 2 and 3
(e) 1 and 3
(a) 5
(b) 4
(c) 3
(d) 2
(e) 1
21.
Ethanol can be oxidized stepwise. What is the first
stable intermediate product when ethanol is oxidized
with a mild oxidation agent?
(a) CH3COOH
(b) CO2
(c) CH3CHO
(d) CH3CH2OH
(e) CH3OCH3
14.
Which of the following statements is FALSE regarding
the reaction between Cl2 and C2H6?
(a) It is a substitution reaction.
(b) The reaction will give a single product of C2H5Cl.
(c) The reaction mechanism involves free radicals.
(d) The reaction can be initiated with either sunlight
or heat.
(e) The first step in the mechanism is the cleavage of
the Cl-Cl bond to give chlorine atoms.
22.
Which of the following alcohols forms a ketone when
oxidized?
(a) 1-propanol
(b) methanol
(c) 2-methyl-2-propanol
(d) 2-propanol
(e) all of the above
15.
Which of the following will undergo an addition
reaction with chlorine?
(a) CH3CH2CH2CH3
(b) CH3CH2CH=CHCH3
(c) C6H6
23.
What is the sum of the coefficients in the balanced
equation for the complete combustion of 2-
24.
methylbutane? Use smallest whole number
coefficients. Do not forget coefficients of 1.
(a) 10
(b) 13
(c) 17
(d) 20
(e) 23
The organic starting materials for the preparation of
an ester could be _______ .
(a) an acid and an alcohol
(b) a ketone and an alcohol
(c) an alkane and a ketone
(d) only an acid
(e) an amine and an acid
25.
Hydrolysis (saponification) of a fat would yield ______ .
(a) water and an alkene
(b) ethanol and propanoic acid
(c) glycerol and soap
(d) ethanol and a soap
(e) a triester of glycerol with fatty acids
26.
The segment -CH2CH2CH2CH2CH2CH2- represents the
polymer named _______ .
(a) polybutylene
(b) polyhexene
(c) polypropylene
(d) polystyrene
(e) polyethylene
Answers to Chapter 28
1. (c) 2. (e) 3. (e) 4. (b) 5. (e) 6. (d) 7. (c) 8. (e) 9. (b) 10. (c)
11. (d) 12. (c) 13. (e) 14. (b) 15. (b) 16. (c) 17. (b) 18. (a) 19.
(e) 20. (e) 21. (c) 22. (d) 23. (d) 24. (a) 25. (c) 26. (e)