REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
UNIT 1 INTRODUCTORY
1. An example of a chemical change is
(A) freezing of water.
(B) burning a match.
(C) boiling carbon tetrachloride.
(D) dissolving alcohol in water.
(E) stretching a rubber band.
2.
(A)
(B)
(C)
(D)
(E)
Which involves a chemical change?
powdering sugar
condensing steam
magnetizing an iron bar
separating cream from milk
exposing photographic film to light
3.
(A)
(B)
(C)
(D)
Which process is a chemical change?
the melting of ice
the burning of a candle
the magnetizing of steel
the liquefaction of oxygen
4.
The graph was obtained by plotting the volume of a
material vs. the mass of that same material.
5
Which unit represents l  10–3 mol?
decimole
kilomole
millimole
micromole
9.
The volume of one milliliter most nearly equals
(A)
454 g
(B)
1000 L
(C)
1 mg
(D)
1 in3
(E)
1 cm3
10. 10.0 mL of a pure liquid substance has a mass of 25.0
g. What is the mass of 3.00 L of the substance?
(A) 83 g
(B) 120 g
(C) 1,200 g
(D) 7,500 g
(E) 25,000 g
11.
(A)
(B)
(C)
(D)
The metric prefix for 10–6 is
mega–
kilo–
micro–
milli–
UNIT 2 MOLES
12. Which expression represents the number of atoms in
1.0  10–3 g of lead?
3
1
1
3
5 7
mass (g)
9
What is the density of the material?
(A) 1.5 g·cm
(B) 2.0 g·cm
(C) 0.67 g·cm
(D) 0.50 g·cm
5.
(A)
(B)
(C)
(D)
8.
(A)
(B)
(C)
(D)
Which is a unit for expressing volume?
mm
g
cm3
g·cm
(A)
207
(22.4)  (1 .0  10 -3 )
-3
(B) (22.4)  (1 .0  10 )
207
(C)
207
(6.02  10 )  (1 .0  10 3 )
23
23
-3
(D) (6.02  10 )  (1.0  10 )
207
23
3
(E) (6.02  10 )  (1.0  10 )
207
6.
The number 149,000,000 is usually written in
scientific notation as
(A) 0.149  109
(B) 149 x 106
(C) 1.49 x 108
(D) 1490 x 105
13. How many atoms are in one mole of hydrogen
sulfide, H2S?
(A) 34  6.02  1023
(B) 3  6.02  1023
(C) 3
(D) 34
7. Which measurement is the most uncertain?
(A) 1.00 ± 0.01 cm
(B) 2.00 ± 0.05 L
(C) 10 ± 1 g
(D) 200 ± 1 mL
14. A substance whose density is 4.00 g·cm–3 occupies a
volume of 12.0 cm3. What is its mass?
(A) 0.333 g
(B) 8.00 g
(C) 48.0 g
(D) 4.00 g
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
21. Which is STP?
15. How many moles of oxygen atoms are present in one
mole of beryllium sulfate tetrahydrate, BeSO4·4H2O?
(A) eight
(B) five
(C) four
(D) two
16. Which statement best accounts for the fact that gases
can be greatly compressed?
(A) Molecules occupy space.
(B) The collisions of molecules are elastic.
(C) Molecules of gases are in constant motion.
(D) The molecules of a given gas are identical.
(E) Molecules of gases are relatively far from each other.
17. Gases may be most easily liquefied by
(A)
(B)
(C)
(D)
(E)
raising the temperature and lowering the pressure.
raising the pressure and lowering the temperature.
lowering both the temperature and pressure.
raising both the temperature and pressure.
lowering the temperature and keeping the presure
unchanged.
(A)
(B)
(C)
(D)
(E)
0 °C and 76 mmHg
0 K and 76 mmHg
0 K and 760 mmHg
100 °C and 76 cmHg
273 K and 760 mmHg
22. A student collects one liter samples of O2, CO2, and
CH4 at laboratory conditions. What quantity is the
same for all three samples?
(A) number of atoms divided by the number of molecules
in each sample
(B) number of molecules in each sample
(C) number of atoms in each sample
(D) mass of each sample
23. Each of three identical containers holds a mole of
gas, all at the same temperature.
CH
O
SO
18. If the temperature and pressure are the same, one
gram of hydrogen has about the same number of
atoms as
Atomic Molar Masses
(A)
(B)
(C)
(D)
(E)
H
O
1 g of oxygen.
2 g of oxygen.
8 g of oxygen.
16 g of oxygen
32 g of oxygen
1.0 g·mol–1
16.0 g·mol–1
.
.
19. One liter of oxygen at STP contains approximately
the same number of molecules as
(A) 2 L of He at STP.
(B) 1/3 L of O3 at STP.
(C) l L of CO2 at STP
(D) 1/5 L of CH4 at STP.
(E) (D) (E) 250 mL of NH3 at STP.
20. According to the Avogadro Principle, one liter of
gaseous hydrogen and one liter of gaseous ammonia
contain the same number of
(A)
(B)
(C)
(D)
atoms at standard conditions.
molecules at all conditions.
molecules only at standard conditions.
atoms if conditions in both containers are the
same.
(E) molecules if conditions in both containers are the
same.
Which gas exerts the greatest pressure? Assume ideal
behavior.
(A)
(B)
(C)
(D)
CH4
O2
SO2
They all exert the the same pressure.
24. A weather balloon contains 12 L of hydrogen at 740
mmHg pressure. At what pressure in mmHg will the
volume become 20 L (temperature constant)?
(A)
(B)
(C)
(D)
(E)
370
444
760
1230
1480
25. A gas occupies a volume of 2.0 cubic feet at 13 atm.
How many cubic feet does this gas occupy at 1.0 atm,
temperature constant?
(A)
(B)
(C)
(D)
6.5
13
15
26
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
26. A sample of gas at 1.00 atm of pressure occupies a
volume of 500 L. If the volume is decreased to 125 L
and the temperature is held constant, what is the new
pressure in atmospheres?
(A)
(B)
(C)
(D)
Atomic Molar Masses
C
H
Mg
O
0.250
2.00
1.25
4.00
27. For a given amount of dry gas at constant
temperature, when the pressure is doubled the
volume is
(A)
(B)
(C)
(D)
31. The molar mass of magnesium acetate,
Mg(C2H3O2)2, in g·mol–1 is
halved.
unchanged.
doubled.
increased, but not doubled.
(A)
(B)
(C)
(D)
(E)
15
16
83
142
166
32. How many mole(s) of calcium carbonate, CaCO3, is
represented by 50 g of the compound?
28. Approximately how many molecules are in 11 g of
carbon dioxide, CO2, gas?
Atomic Molar Masses
Ca
C
O
Atomic Molar Masses
C
O
(A)
(B)
(C)
(D)
12.0 g·mol–1
16.0 g·mol–1
1.5  1023
3.0  1023
6.0  1023
2.4  1023
29. What is the mass of one mole of calcium nitrate,
Ca(NO3)2?
(A)
(B)
(C)
(D)
(E)
(A)
(B)
(C)
(D)
40. g·mol–1
14. g·mol–1
16. g·mol–1
82 g
102 g
164 g
204 g
30. The number of moles of water in 1,000 g of water is
40.1 g·mol–1
12.0 g·mol–1
16.0 g·mol–1
1.0
2.0
0.20
4.0
0.50
33. The molar mass of aluminum sulfate, Al2(SO4)3, is
Atomic Molar Masses
Al
O
S
Atomic Molar Masses
Ca
N
O
12. g·mol–1
1. g·mol–1
24. g·mol–1
16. g·mol–1
(A)
(B)
(C)
(D)
(E)
27 g·mol–1
16 g·mol–1
32 g·mol–1
150 g·mol–1
170 g·mol–1
278 g·mol–1
342 g·mol–1
450 g·mol–1
34. The mass of one mole of ammonium carbonate,
(NH4)2CO3, is approximately
Atomic Molar Masses
H
O
(A)
(B)
(C)
(D)
(E)
18.0
55.5
180.0
1000.0
18,000.0
Atomic Molar Masses
1.0 g·mol–1
16.0 g·mol–1
C
H
N
O
(A)
(B)
(C)
(D)
43.0 g
72.0 g
78.0 g
96.0 g
12.0 g·mol–1
1.0 g·mol–1
14.0 g·mol–1
16.0 g·mol–1
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
41. Which formula is incorrect?
35. The number of molecules present in 22.0 g of carbon
dioxide at STP is
Atomic Molar Masses
C
O
(A)
(B)
(C)
(D)
(E)
12.0 g·mol–1
16.0 g·mol–1
(A)
(B)
(C)
(D)
(E)
2.01  1012
6.03  1012
2.06  1022
3.01  1023
6.02  1023
AlSO4
Al2SO4
Al3SO4
Al3(SO4)2
Al2(SO4)3
43. What is the formula for chromium(III) oxide?
(A) CrO
(B) Cr2O
(C) Cr3O
(D) Cr2O3
44. What is the formula for strontium sulfide?
Atomic Molar Masses
H
N
O
(D) NH4HSO4
(E) LiH
42. What is the formula for aluminum sulfate?
36. What mass of nitrogen dioxide, NO2, has the same
number of molecules as 18.0 g of water, H2O?
(A)
(B)
(C)
(D)
(A) Al2(SO4)3
(B) BaHCO3
(C) Ca(OH)2
1.0 g·mol–1
14.0 g·mol–1
16.0 g·mol–1
(A) SrS
(B) Sr2S
(C) SrS2
(D) SrS3
45. What is the formula for copper(II) hydroxide?
(A) CuOH
6.02 g
18.0 g
23.0 g
46.0 g
(B) Cu(OH)2 (C) Cu2OH
(D) CuOH2
46. Which is the formula for ammonium nitrate?
37. Calculate the mass of 12.0 
chlorine gas, Cl2.
1023
molecules of
(A) NH3N (B) NH4N
(D) NH4NO3
(C) NH4NO2
47. What is the formula for sodium carbonate?
Atomic Molar Mass
Cl
(A) NaHCO3
(B) NaCO3
35.5 g·mol–1
(A) 35.5 g
(B) 71.0 g
(C) 142 g
(D) 284 g
UNIT 3 NOMENCLATURE
38. The correct formula for iron(III) sulfate is
(A) FeSO4
(B) Fe(SO4)2
(C) Fe2SO4
(D) Fe2(SO4)3
(E) Fe3(SO4)2
39. The one correct formula among these is
(A) Na2OH
(B) Cu(SO4)2
(C) ZnCl2
(D) Zn(NO3)3
(E) BaNO3
40. Which formula is incorrect?
(A) BaHCO3 (B) Ca(OH)2 (C) Al2O3
(E)
ZnCO3
(D) K2SO4
(C) So2CO3
(D) Na2CO3
48. What is the formula for chromium(III) sulfate?
(A) Cr2(SO4)3
(B) Cr3(SO4)2
(C) Cr2(SO3)3
(D) Cr3SO4
49. Which formula is followed by its correct name?
(A)
(B)
(C)
(D)
(E)
FeCl3, iron(III) chloride
FeS, iron(II) sulfite
Mg3N2, magnesium nitrite
KNO2, potassium nitrate
HClO, hydrochloric oxide
50. The compound not properly named is
(A)
(B)
(C)
(D)
(E)
Fe2O3, iron(III) oxide.
Pb3O4, lead(III) tetraoxide.
CuCl2, copper(II) chloride.
Pb3(PO4)2, lead(III) phosphate.
P2S5, diphosphorus pentasulfide.
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
51. What is the name of the compound having the
formula CaH2?
61. The total number of atoms represented by
5Al(C2H3O2)3 is
(A) calcium amide
(B) calcium hydride
(A)
(B)
(C)
(D)
(E)
(C) calcium hydrate
(D) calcium hydroxide
52. What is the name of the compound Fe2(SO4)3?
(A) iron(II) sulfate
(B) iron(III) sulfate
(C) iron(II) trisulfate
(D) iron(II) sulfate(III)
53. What is the correct name for Fe(NO3)2?
(A) iron(II) nitrate
(C) iron(III) nitrate
(B) iron(II) nitrite
(D) iron(III) nitrite
54. The formula for hydrogen bromate is HBrO3, and the
formula for dysprosium oxide is Dy2O3. What is the
formula for dysprosium bromate?
(A) Dy2BrO3
(B) Dy3BrO3
(C) Dy(BrO3)3
(D) Dy2(BrO3)3
55. The formula for ytterbium sulfate is Yb2(SO4)3.
What is the formula for ytterbium chloride?
(A) YbCl2
(B) Yb2Cl3
(C) Yb2Cl2
(C) NO2–, NO3–
(D) HS–, HSO4–
57. Barium perrhenate has this formula: Ba(ReO4)2. The
perrhenate ion is
(A) ReO4–
(B) ReO42–
(C) ReO43– (D) ReO44–
58. What is the total number of oxygen atoms
represented by the formula KAl(SO4)2·12H2O?
(A) 9
(B) 16
(C) 20
(D) 48
(E) 96
59. Which is the number of atoms of hydrogen in one
molecule of glycerine, C3H5(OH)3?
(A) 14
(B) 8
(C) 6
(D) 5
60. The total number of atoms represented by the
formula K3Fe(CN)6 is
(A)
(B)
(C)
(D)
(E)
4
10
11
16
36
62. The number of atoms of oxygen indicated by the
formula Ca3(PO4)2 is
(A)
(B)
(C)
(D)
(E)
12
8
7
4
3
63. How many atoms are in one molecule of acetone,
CH3COCH3 ?
(A) 1
(B) 6
(C) 3
(D) 10
64. Using only these formulas,
XY2
(D) YbCl3
56. In which pair of anions do both names end in ‘–ate’?
(A) Cl–, ClO3–
(B) ClO3–, NO3–
22
60
71
84
110
X2 Z
QZ
what formula would you expect for a compound of
elements Q and Y?
(A) QY
(B) QY2
(C) Q2Y
(D) QY4
65. Which set consists only of compounds?
(A) Na, Ca, He
(B) H3O+, Cl–, I3–
(C) NaCl, CH4, Br2
(D) H2S, CuCl2, KI
66. Which substance contains only one kind of atom?
(A) water
(C) aluminum
(B) ethanol
(D) carbon dioxide
UNIT4 BALANCE EQUATIONS (REG, IONIC, NETIONIC STOICHIOMETRY AND LIMITNG
REACTANTS
67. Which property is always conserved during a
chemical reaction?
(A) mass
(B) volume
(C) pressure (D) solubility
68. The equation
Cu + 4HNO3  Cu(NO3)2 + 2H2O + ?
would be completed and balanced by using
(A) NO2
(E) 2NO
(B) 2NO2
(C) 3NO2
(D) 4NO2
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
69. When the equation
75. Which set of coefficients correctly balances the
equation?
? Sb + ? Cl2  ? SbCl3
? Al(s) + ? H+(aq)  ? Al3+(aq) + ? H2(g)
is correctly balanced, the sum of the coefficients is
(A) 2
(B) 3
(C) 6
(D) 7
(E) 9
(A) 1, 2, 1, 2
(B) 2, 6, 2, 3
(C) 3, 2, 3, 2
(D) 2, 3, 2, 3
76. Which equation represents the complete combustion
of acetylene in an excess of air?
70. Which expression is correctly balanced?
(A)
(B)
(C)
(D)
(E)
(A)
(B)
(C)
(D)
(E)
Na2O2 + 2H2O  2NaOH + O2
2Na2O2 + 2H2O  4NaOH + 2O2
4Na2O2 + 3H2O  4NaOH + 2O2
2Na2O2 + 2H2O  4NaOH + O2
3Na2O2 + 2H2O  6NaOH + O2
71. Which set of coefficients balances this equation?
? CH4(g)
(A) 3, 1, 1


? C3H8(g) + ? H2(g)
(B) 3, 2, 1
(C) 3, 1, 2
C2H2 + 2O2  2CO2 + H2
C2H2 + O2  2CO + H2
C2H2 + O  2C + H2O
C2H2 + O2  2C + H2O2
2C2H2 + 5O2  4CO2 + 2H2O
77. Dysprosium oxide, Dy2O3, reacts with hydrochloric
acid to produce only water and a salt. The salt is
(A) Dy2Cl3
(D) 6, 2, 2
(B) DyCl2
(C) DyCl3
(D) DyCl6
78. Which equation represents the dissolving of sodium
sulfate, Na2SO4, in water?
(E) 6, 2, 6
72. Consider the unbalanced expression:
? CH3CH2CHO(l) + ? O2(g)  ? CO2(g) + ? H2O(g)
Which set of coefficients balances the equation?
(A) 2, 8, 3, 6
(B) 3, 8, 6, 6
(C) 1, 4, 3, 2
(D) 1, 8, 3, 3
(E) 1, 4, 3, 3
73. Consider the unbalanced expression:
? Cu(s) + ? NO3–(aq) + ? H+(aq) 
? Cu2+(aq) + ? NO(g) + ? H2O(l)
Which set of coefficients correctly balances the equation?
(A) 4, 5, 3, 8, 2, 3
(B) 2, 4, 3, 8, 3, 3
(C) 3, 2, 8, 7, 2, 4
(D) 3, l, 8, 7, 4, 2
(E) 3, 2, 8, 3, 2, 4
74. The expression for pentane, C5H12, burning in
oxygen is
? C5H12(g) + ? O2(g)  ? CO2(g) + ? H2O(g)
What set of coefficients balances the equation?
(A) 1, 8, 5, 6
(B) 2, 8, 10, 6
(C) 1, 8, 5, 12
(D) 1, 11, 5, 12
(A)
(B)
(C)
(D)
Na2SO4(s)  Na2+(aq) + SO42–(aq)
Na2SO4(s)  2Na+(aq) + SO42–(aq)
Na2SO4(s)  Na22+(aq) + S2–(aq) + 4O2–(aq)
Na2SO4(s)  2Na2+(aq) + S2–(aq) + O2–(aq)
79. What is the net ionic equation for the reaction
between solutions of sodium chloride, NaCl, and
silver nitrate, AgNO3?
(A)
(B)
(C)
(D)
Na+(aq) + NO3–(aq)  Na(s) + 1/2N2(g) + 3/2O2(g)
Ag+(aq) + Cl–(aq)  Ag(s) + 1/2Cl2(g)
+
–
Ag+(aq) + Cl–(aq) 
 Ag (aq) + Cl (aq)
Ag+(aq) + Cl–(aq)  AgCl(s)
80. Which equation represents the dissolving
(dissociation) of aluminum sulfate, Al2(SO4)3, in
water?
(A)
(B)
(C)
(D)
Al2(SO4)3(s)  2Al3+(aq) + 3S6+(aq) + 4O2–(aq)
Al2(SO4)3(s)  2Al3+(aq) + 3SO42–(aq)
Al2(SO4)3(s)  2Al2+(aq) + 3SO43–(aq)
Al2(SO4)3(s)  Al3+(aq) + SO43–(aq)
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
81. The overall equation for the reaction between KCl
and AgNO3 is
K+(aq) + Cl–(aq) + Ag+(aq) + NO3–(aq)
K+(aq) + NO3–(aq) + AgCl(s)


What is the net ionic equation?
(A)
(B)
(C)
(D)
Ag+(aq) + Cl–(aq) 
 AgCl(s)
K+(aq) + Cl–(aq) 
 KCl(s)
K+(aq) + NO3–(aq) 
 KNO3(s)
K+(aq) + Cl–(aq) + Ag+(aq) + NO3–(aq) 

Ag+(aq) + K+(aq) + Cl–(aq) + NO3–(aq)
82. Complete the equation for the reaction between
solutions of lead nitrate, Pb(NO3)2, and ammonium
sulfide, (NH4)2S.
Pb2+(aq) + 2NO3–(aq) + 2NH4+(aq) + S2–(aq) 
(A) 2NH4NO3(s) + Pb2+(aq) + S2–(aq)
(B) Pb(NO3)2(s) + 2 NH4+(aq) + S2–(aq)
(C) (NH4)2S(s) + Pb2+(aq) + 2NO3–(aq)
(D) PbS(s) + 2NH4+(aq) + 2NO3–(aq)
83. Which is the balanced net ionic equation for the
formation of the precipitate silver chromate,
Ag2CrO4?
(A)
(B)
(C)
(D)
2Ag+(aq) + CrO42–(aq)  Ag2CrO4(s)
Ag+(aq) + CrO42–(aq)  Ag2CrO4(s)
Ag0(aq) + CrO42–(aq)  Ag2CrO4(s)
Ag2CrO4(s)  2Ag+(aq) + CrO42–(aq)
86. Which is the net ionic equation for the reaction of
lead(II) nitrate and sodium chromate?
(A) Pb2+(aq) + CrO42–(aq)  PbCrO4(s)
(B) Pb(NO3)2(aq) + Na2CrO4(aq) 
PbCrO4(s) + 2NaNO3(aq)
(C) 2Na+(aq) + CrO42–(aq)  Na2CrO4(aq)
(D) Pb2+(aq) + NO3–(aq) + Na+(aq) + CrO42–(aq) 
PbCrO4(s) + Na+(aq) + NO3–(aq)
87. What is the net ionic equation for the reaction
between lead(II) nitrate and potassium sulfide?
(A)
(B)
(C)
(D)
Pb2+(aq) + S2–(aq)  PbS(s)
K+(aq) + NO3–(aq)  KNO3(aq)
Pb(NO3)2(aq) + K2S(aq)  PbS(s) + 2KNO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2K+(aq) + S2–(aq) 
PbS(s) + 2K+(aq) + 2NO3–(aq)
STOICHIOMETRY W/LIMITING REACTANTS
88. 50.0 g of water is heated from 22.0 °C to 36.0 °C.
How much heat is absorbed?
Specific Heat Capacity for Water
4.18 J·°C–1·g–1
(A) 1510 J
(C) K+ and Cl–
(D) Ag+ and Cl–
(C) 4520 J
(D) 4600 J
(E) 7520 J
89. How much heat is required to raise the temperature of
25.0 g of iron from 10.0 °C to 40.0 °C?
Specific Heat Capacity of Iron
84. Which two ions do not participate in the reaction
between solutions of silver nitrate, AgNO3, and
potassium chloride, KCl?
(A) K+ and Ag+
(B) K+ and NO3–
(B) 2930 J
0.444 J·g–1·°C–1
(A) 750 J
(B) 444 J
(C) 333 J
(D) 313 J
90. What volume is occupied by 2.00 g of a substance
having a density of 5.00 g·cm–3?
85. In the equation:
BaCl2(aq) + Na2SO4(aq)  BaSO4(s) +2NaCl (aq)
What is the net ionic equation for this reaction?
(A)
(B)
(C)
(D)
Cl–(aq) + Na+(aq)  NaCl (aq)
Cl22–(aq) + Na22+(aq)  2NaCl (aq)
Ba2+(aq) + SO42–(aq)  BaSO4(s)
BaCl2(s) + Na2SO4(s) 
Ba2+(aq) + 2Cl–(aq) + 2Na+(aq) + SO42–(aq)
(A) 0.400 cm3
(B) 2.50 cm3
(C) 7.00 cm3
(D) 10.0 cm3
91. If 50 mL of a 200 mL sample of 0.10 M sodium
chloride solution is spilled, what is the concentration
of the remaining solution?
(A) 0.20 M
(B) 0.10 M
(C) 0.075 M (D) 0.025 M
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
92. In the reaction
97. How many grams of calcium carbonate, CaCO3,
would be needed to produce 44.8 L of carbon dioxide
gas, CO2, measured at STP?
4Al + 3O2  2Al2O3
how many moles of aluminum oxide, Al2O3, are
produced from one mole of aluminum, Al?
(A) 0.5
(B) 2
(C) 3
Atomic Molar Masses
(D) 4
93. Given the equation
N2 + 3H2


CaCO3 + 2HCl  CaCl2 + H2O + CO2
2NH3
(A) 50.0
Theoretically, the number of moles of ammonia produced
from 2 mol of nitrogen is
(A) 1
(B) 2
(C) 3
(D) 4
C4H4O4 + NaOH  NaC4'H3O4 + H2O
C4H4O4 + 2NaOH  Na2C4H2O4 + 2H2O
C4H4O4 + 3NaOH  Na3C4HO4 + 3H2O
C4H4O4 + 4NaOH  Na4C4O4 + 4H2O
95. In neutralizing 0.015 mol of H3PO3, 0.030 mol of
NaOH was consumed. Which equation describes this
reaction?
H3PO3 + NaOH  NaPO3 + H2O
H3PO3 + NaOH  NaH2PO3 + H2O
H3PO3 + 3NaOH  NaPO3 + 3H2O
H3PO3 + 2NaOH  Na2HPO3 + 2H2O
96. Silica, SiO2, reacts with hydrofluoric acid, HF,
according to this equation
SiO2 + 4HF  2H2O + SiF4
Which reagent is completely consumed when 2 mol of
SiO2 is added to 6 mol of HF?
(A) SiF4
(C) 111
(D) 200
2KClO3  2KCl + 3O2
94. In an experiment, 0.0041 mol of maleic acid,
C4H4O4, reacted with 0.0082 mol of sodium
hydroxide, NaOH. Which equation describes the
reaction?
(A)
(B)
(C)
(D)
(B) 100
98. What volume of oxygen, O2, at STP can be prepared
by the complete decomposition of 0.100 mol of
potassium chlorate, KClO3?
(E) 5
(A)
(B)
(C)
(D)
40.1 g·mol–1
12.0 g·mol–1
16.0 g·mol–1
Ca
C
O
(B) H2O
(C) HF
(D) SiO2
(A) 1.49 L
(B) 3.36 L
(C) 4.80 L
(D) 6.72 L
99. The equilibrium equation for the Haber process at
500 °C is
N2 + 3H2


2NH3 + heat
When one liter of nitrogen combines with three liters of
hydrogen the maximum volume of ammonia produced is
(A) 1 L
(B) 2 L
(C) 3 L
(D) 4 L
(E) 6 L
100. The volume of pure oxygen needed to burn
completely 800 mL of acetylene (C2H2) gas is
(A) 800 mL
(B) 1600 mL
(C) 2000 mL
(D) 10000 mL
(E) 20000 mL
101. A mixture of 2.0 g of hydrogen and 32 g of oxygen is
exploded and produces water. What mass of gas
remains uncombined?
Atomic Molar Masses
H
O
(A) 1.0 g of hydrogen
(B) 1.0 g of oxygen
(C) 8.0 g of oxygen
1.0 g·mol–1
16.0 g·mol–1
(D) 16 g of oxygen
(E) 24 g of oxygen
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
102. The equation for the complete combustion of butane
gas, C4H10, is
2C4H10 + 13O2  8CO2 + 10H2O
How many liters of carbon dioxide is produced when a
mixture of 1.00 L of butane gas and 13.0 L of oxygen is
burned? (measured under the same conditions)
(A) 1.00 L
(B) l. 63 L
(C) 8.00 L
106. In the reaction represented by the equation
COCl2 + 2NaI  2NaCl + CO + I2
what is the maximum mass of iodine that can be liberated
from 60.0 g of sodium iodide?
Molar Masses
150. g·mol–1
254. g·mol–1
NaI
I2
(D) 4.00 L
(E) 13.0 L
(A) 5.00 g
103. The mass of potassium chloride formed by the
complete decomposition of 490 g of potassium
chlorate is
(A) 96 g
35.5 g·mol–1
39.1 g·mol–1
16.0 g·mol–1
(B) 122.5 g
(C) 149 g
(C) 50.8 g
(D) 127 g
(E) 254 g
107. What mass of iron oxide, Fe3O4, is produced from
2.00 mol of iron, Fe?
Atomic Molar Masses
Cl
K
O
(B) 25.4 g
3Fe(s) + 4H2O(g)  Fe3O4(s) + 4H2(g)
Molar Mass
Fe3O4
(D) 298 g
231. g·mol–1
(E) 490 g
(A) 154 g
104. In the reaction
108. What mass of calcium hydroxide, Ca(OH)2, is
obtained from 18.7 g of calcium oxide, CaO?
2Al + 3H2SO4  3H2 + Al2(SO4)3
(B) 9.0 g
(C) 13.5 g
Ca
H
O
(D) 27.0 g
(E) 81.0 g
(D) 693 g
40.1 g·mol–1
1.0 g·mol–1
16.0 g·mol–1
CaO + H2O  Ca(OH)2
105. What is the maximum mass of tungsten (W) obtained
from the use of 18 g of hydrogen according to the
equation
(A) 18.7 g
Atomic Molar Masses
H
W
1. g·mol–1
184. g·mol–1
(B) 24.7 g
(C) 56.1 g
(D) 74.1 g
109. Consider the equation:
WO3 + 3H2  W + 3H2O
(A) 1  184 g
(B) 3  184 g
(C) 9  184 g
(C) 462 g
Atomic Molar Masses
the mass of aluminum that reacts with 1 mol of hydrogen
ions is approximately
(A) 3.0 g
(B) 231 g
2Al(OH)3  Al2O3 + 3H2O
When 15.0 g of aluminum hydroxide, Al(OH) 3 is
decomposed, how many grams of water will be formed?
Atomic Molar Masses
(D) 18  184 g
(E) 184 g + 3  16 g
Al
H
O
(A) 3.86 g
(B) 5.19 g
27.0 g·mol–1
1.0 g·mol–1
16.0 g·mol–1
(C) 4.20 g
(D) 22.5 g
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
110. What mass of water is produced by complete
combustion of 126 g of propene, C3H6?
2C3H6 + 9O2  6H2O + 6CO2
114. How many grams of water, H2O, can be prepared
when 2.00 mol of hydrogen, H2, and 2.00 mol of
oxygen, O2, are mixed and reacted in this process?
2H2 + O2  2H2O
Atomic Molar Masses
12.0 g·mol–1
1.0 g·mol–1
16.0 g·mol–1
C
H
O
(A) 18.0 g
(B) 54.0 g
(C) 126 g
Atomic Molar Masses
H
O
(D) 162 g
111. If 10.0 g of iron, Fe, and 10.0 g of sulfur, S, are
heated together, how many grams of iron(II) sulfide,
FeS, could be formed?
Atomic Molar Masses
(A) 18.0 g
(B) 36.0 g
(B) 15.7
(C) 27.6
Atomic Molar Masses
B
Na
O
C3H8(g) + 5O2(g)  3CO2(g) + 4H2(g)
(A) Na2B4O7 
(B) NaBO
(C) NaB2O5
Atomic Molar Masses
C
H
Atomic Molar Masses
(B) 29.3 g
12.0 g·mol–1
16.0 g·mol–1
(C) 44.0 g
(A) CH2
(E) C2H4
(D) 66.0 g
113. Consider the equation:
(C) C2H
(D) C2H2
Atomic Molar Masses
C
H
How many moles of reactant are in excess when 2.0 mol
of CH4(g) are ignited in 2.0 mol of O2(g)?
(C) 0.5 mol CH4
(D) no excess of either
reactant
(B) CH
12 g·mol–1
1 g·mol–1
117. Decomposition of 12 g of a compound containing
only carbon and hydrogen yields 9 g of carbon and 3
g of hydrogen. What is the simplest formula of the
compound?
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
(A) l.0 mol CH4
(B) 2.0 mol O2
(D) Na3B4O
(E) Na3BO4
116. A compound contains 85.71% carbon and 14.29%
hydrogen by mass. Its simplest formula is
What is the maximum mass of carbon dioxide produced
when a mixture of 0.500 mol of propane and 3.00 mol of
oxygen is ignited?
(A) 22.0 g
(E) 132. g
11 g·mol–1
23 g·mol–1
16 g·mol–1
(D) 88.0
112. The equation for the complete combustion of
propane, C3H8, is
C
O
(D) 72.0 g
115. Upon analysis a compound is found to contain 22.8%
sodium, 21.8% boron, and 55.4% oxygen. Its
simplest formula is
Fe + S  FeS
(A) 10.0
(C) 68.0 g
EMPIRICAL FORMULAS
55.8 g·mol–1
32.1 g·mol–1
Fe
S
1.0 g·mol–1
16.0 g·mol–1
(A) CH2
(E) C3H9
(B) CH4
12.0 g·mol–1
1.0 g·mol–1
(C) C2H5
(D) C3H7
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
118. A sample of a compound contains 3.21 g of sulfur, S,
and 11.4 g of fluorine, F. Find the empirical formula
of the compound.
Atomic Molar Masses
F
S
(A) SF
(B) SF2
19.0 g·mol–1
32.0 g·mol–1
(C) SF3
(D) SF6
(C) C12H4O2
(D) C12H24O12
120. A substance has an empirical (simplest) formula of
CH3 and a molar mass of 30 g·mol–1. The molecular
(true) formula is
Atomic Molar Masses
C
H
(A) an ion.
(B) a radical.
(C) an isotope.
(A) losing protons.
(B) losing electrons.
(B) (CH3)2
52 3+
24Cr ?
(A) 52
(B) 27
(C) 24
(D) 21
127. The number of neutrons in the nucleus of an atom of
9
4Be
is
(A) 36
(E) 4
(B) 13
(C) 9
(D) 5
128. Which symbol represents an atom that contains the
largest number of neutrons?
g·mol–1
12.0
1.0 g·mol–1
(C) (CH3)3
(C) gaining protons.
(D) gaining electrons.
126. How many electrons are in a chromium(III) ion,
235
(A) 92U
(A) (CH3)1
(D) a molecule.
(E) an electrolyte.
125. Metallic atoms become ions by
119. A compound has the empirical formula CH2O and
the molecular mass 180 g·mol–1. What is its
molecular formula?
(A) CH8O10
(B) C6H12O6
124. An atom that loses or gains an electron becomes
239
(B) 92U
239
(C) 93Np
239
(D) 94Pu
(D) (CH3)4
231
121. A compound whose empirical formula is CH2 has a
molar mass of 28 g·mol–1. What is the molecular
formula?
Atomic Molar Masses
C
H
(A) CH2
(B) C2H4
129. An ion has 13 electrons, 12 protons, and 14 neutrons.
What is the mass of the ion?
(A) 14 u
(E) 39 u
12.0 g·mol–1
1.0 g·mol–1
(C) C2H2
(E) 91Pa
(D) CH4
F
S
(A) SF
(B) S2F2
19. g·mol–1
32. g·mol–1
(C) S3F3
(D) SF4
UNIT 7 atomic theory 9-10 PERIOD TABLE/TRENDS
123. A calcium ion is a calcium atom that has
(A) lost one electron.
(B) gained one electron.
(C) gained one ion.
(A) 11Na+
(D) lost two electrons.
(E) gained two electrons.
(D) 27 u
23
(B) 11Na
23
23
(C) 12Mg2+ (D) 12Mg
131. The atomic number of an element is determined by
the number of
(A)
(B)
(C)
(D)
(E)
Atomic Molar Masses
(C) 26 u
130. The symbol that represents 11 protons, 12 neutrons,
and 10 electrons would be:
23
122. A gaseous compound contains a ratio of one atom of
sulfur to one atom of fluorine. A mole of this gas has
a mass of approximately 102 g. What is the
molecular formula?
(B) 25 u
protons in each of its atoms.
neutrons in each of its atoms.
particles in each of its atoms.
protons plus neutrons in each of its atoms.
protons plus electrons in each of its atoms.
132. All positive ions differ from their corresponding
atoms by having
(A)
(B)
(C)
(D)
(E)
larger diameters.
fewer electrons.
a charge of +1.
greater atomic masses.
stronger metallic properties
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
133. Which group represents particles that contain the
same number of electrons?
(D) O2–, S2–, Se2–
(E) Ca2+, Fe2+, Zn2+
(A) F, Ne, Na
(B) Mg, Al, Si
(C) Cl–, Ar, K+
134. Note the chart of interactions of equal volumes of
various 0.100 M aqueous solutions. (Symbols of
elements or ions have been replaced by capital letters,
and soluble products are indicated by “S”) What is
the formula of the precipitate?
139. In the modern periodic table the elements are
arranged in the order of increasing
(A) atomic masses.
(B) atomic radii.
(C) atomic numbers.
(D) atomic volumes.
140. In which set are the three elements in the same
family?
(A) B, C, N
(B) N, O, F
(C) Hg, Ga, Sr
(D) Zn, Cd, Hg
141. Which scientist is given credit for developing the
periodic table?
AY BX CY
D X pp t
S
CY
S
S
B X pp t
(A) DY
(B) BY
(A) Rutherford
(B) Mendeleev
S
142. lf XO2 is the correct formula for an oxide, the
formula for the chloride of X is
(C) AX
(D) CX
135. An odorless, colorless, tasteless gas is suspected to be
oxygen. Which result would support this hypothesis?
(A)
(B)
(C)
(D)
The gas would extinguish a flame.
The gas would turn limewater milky.
The gas would burn in air producing only water.
A glowing splint would burst into flame in the gas.
136. The chemical properties of atoms depend principally
upon
(A)
(B)
(C)
(D)
their atomic masses.
the masses of the nuclei involved.
the number of neutrons in their nuclei.
the ratio in which the atoms combine with other
atoms.
(E) the number of electrons in their outermost shells.
137. The similar chemical behavior of the elements in a
given family in the periodic table is best accounted
for by the fact that atoms of these elements have
(A)
(B)
(C)
(D)
(E)
the same number of electrons in the outermost shell.
the same number of electrons.
the same number of protons.
similar nuclear structures.
a common origin
138. The best explanation of the extreme activity of
fluorine as compared to other halogens is that the
fluorine atom
(A)
(B)
(C)
(D)
(E)
(C) Dalton
(D) Planck
has the smallest atomic radius.
has the smallest nuclear charge.
has seven valence electrons.
is the strongest reducing agent.
needs one electron to complete its outermost shell.
(A) XCl2
(B) XCl4
(C) XCl
(D) X2Cl3
(E) XCl3
143. M represents a metallic element, the oxide of which
has the formula M2O. The formula of the chloride of
M is
(A) MCl
(B) MCl2
(C) MCl3
(D) MCl4
(E) M2Cl
144. What is the most probable formula for a compound of
silicon, Si, and hydrogen, H?
(A) SiH
(B) SiH2
(C) SiH6
(D) SiH4
145. A hypothetical element, Z, forms a chloride with the
formula ZCl5. What is the most probable formula for
its oxide?
(A) ZO2
(B) ZO5
(C) Z2O5
(D) Z5O2
146. Based on the position of the elements in the periodic
chart, the most likely formula for strontium nitride is
(A) Sr2N5
(B) Sr5N2
(C) Sr2N3
(D) Sr3N2
147. Which family of elements always forms ions with an
oxidation number of +2 in compounds?
(A) halogens
(B) alkali metals
(C) transition metals
(D) alkaline–earth metals
148. Which element is the most electronegative?
(A) Be
(B) Mg
(C) Ca
(D) Sr
(E) Ba
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
149. Since sodium and potassium are both members of
Group 1A in the periodic table, a sodium and a
potassium atom have the same
152. Consider a plot of a property of the alkaline earth
metals.
(A)
(B)
(C)
(D)
atomic mass.
number of protons in their nuclei.
atomic number and the same nuclear charge.
characteristic of losing one electron per atom to form
an ion.
(E) total number of electrons around the nucleus.
150. The element requiring the least amount of energy to
remove one electron from an atom is
H
Li
Na
Be
Mg
(A) Na
B
Al
(B) Be
C
Si
N
P
O
S
(C) O
F
Cl
(D) Cl
(E) Ar
151. In which part of the periodic table are the most
electronegative elements found?
He
Ne
Ar
4 12 20
38
Be Mg Ca
Sr
Atomic number
56
Ba
Which property is plotted on this graph?
(A)
(B)
(C)
(D)
first ionization energy
atomic radius
atomic mass
number of valence electrons
153. As the atomic numbers of the elements in a family
increase, the
(A)
(C)
(B)
(D )
(A)
(B)
(C)
(D)
(E)
atomic radii decrease.
atomic masses decrease.
ionization energies decrease.
elements become less metallic.
number of electrons in the outermost energy level
increases.
154. Which of these atoms has the smallest radius?
(A) K
(A) upper left
(B) lower left
(C) upper right
(D) lower right
(B) Cl
(C) Br
(D) Cs
155. Which characteristic of fluorine causes it to be the
most active member of the halogen family, Group
7A?
(A)
(B)
(C)
(D)
It forms diatomic molecules.
It has the smallest atomic radius.
It has no naturally occuring isotopes.
It has seven electrons in its outer shell.
UNIT 11-13 ATOMIC STRUCTURE/DIAGRAMMING
ELECTRONS
156. The chemical activity of an atom is most closely
related to the number and arrangement of its
(A) protons.
(B) neutrons.
(C) isotopes.
(D) electrons.
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
157. The molar mass of a compound is 75 g·mol–1. A
student reported an experimental value of 78 g·mol–1.
The percent error is
164. Which element has the electron configuration
1s22s22p63s23p6 4s13d5?
(A)
78 – 75
78
(D)
78 – 75
 100
75
(A) zinc
(B) copper
(C) nickel
(B)
75
 100
78
(E)
78 – 75
 100
78
165. The electron arrangement that represents the most
active metallic element in this list is
(C)
78
 100
75
(A) 2)7
(B) 2)8)1
(C) 2)8)2
(D) 2)8)3
(E) 2)8)6
158. A student reads a balance as 38.81 g. The correct
reading is 38.41 g. What is the percent error?
(A) 0.0104%
(B) 0.104%
(C) 0.400%
(D) 1.04%
(B) 50
(C) 51
166. What is the electronic configuration of an aluminum
27
atom, 13Al?
159. The number of protons in the atom whose atomic
mass is 89 and atomic number is 39, is
(A) 39
(D) chromium
(E) potassium
(D) 89
(A)
(B)
(C)
(D)
(E)
ls22s22p63d3
1s22s22p63s23p1
ls22s22p62d13s2
1s22s22p62d103s23p5
1s22s22p63s23p63d74s2
(E) 128
167. Which atom contains a partially filled 3p orbital?
160. The particles present in the orbitals of an atom are
(A) iron
(B) argon
(C) boron
(A) mesons.
(B) protons.
(C) neutrons.
(D) positrons.
(E) electrons.
161. A neutral atom whose outermost electron shell
contains eight electrons
(A)
(B)
(C)
(D)
(E)
is very active.
has a combining number of one
is classified as a metal.
is chemically inert.
is more active than hydrogen.
162. When the halogens form ions, the result is
(A)
(B)
(C)
(D)
(E)
colored ions.
positive ions.
diatomic molecules.
covalent compounds.
a completed outer shell of electrons.
163. The correct electronic configuration for the sodium
23
atom, 11Na, is
(A)
(B)
(C)
(D)
(E)
1s22s22p6
1s22s22p63s1
1s22s22p43s23p1
1s22s22p82d103s1
1s22s22p62d103s23p1
(D) calcium
(E) aluminum
168. Which element has the electron configuration
1s22s22p63s2?
(A) aluminum
(B) calcium
(C) magnesium
(D) sodium
169. Which electron configuration represents an atom in
an excited state?
(A)
(B)
(C)
(D)
1s22s22p6
1s22s22p63s2
1s22s22p63s23p64s23d1
1s22s22p63s23p64s24p1
170. When two electrons occupy the same orbital, they
must have
(A)
(B)
(C)
(D)
(E)
opposite spins.
mutual attraction.
four identical quantum numbers.
different magnetic quantum numbers.
different principal quantum numbers.
171. What is the maximum number of electrons allowed in
an orbital?
(A) 1
(E) 10
(B) 2
(C) 3
(D) 6
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
172. What neutral atom has the electron configuration
1s22s22p63s23p64s1?
178. Which electron configuration represents a transition
element?
(A) Na
(A)
(B)
(C)
(D)
(B) K
(C) Ca
(D) Ba
173. Which sublevel becomes filled when a chloride ion,
Cl–, is formed?
(A) 2p
(B) 3p
(C) 4p
(D) 3s
174. Which electron configuration represents a noble gas?
(A) ls22s22p63s23p5
(B) ls22s22p63s23p6
(C) 1s22s22p63s23p64s1
(D) ls22s22p63s23p64s2
175. When an electron shifts from one energy level to a
higher level in the same atom, energy is absorbed.
Which of the electron transitions represented below
absorbs (that is, requires) the most energy?
BONDING
179. In which pair do both compounds exhibit ionic
bonding?
(A) SO2, HCl
(B) KNO3, CH4
(C) NaF, KBr
(D) KCl, CO2
(E) NaCl, H2O
180. A chemical bond is considered to be predominantly
ionic if
(A) atoms of the same element combine.
(B) the reaction forming the bond is endothermic.
(C) atoms of an active metal combine with the atoms of
an active nonmetal.
(D) the bond is between atoms of elements which are of
the same family.
(E) atoms of one metal combine with atoms of another
metal.
4
3
A
1s22s22p63s2
1s22s22p63s23p6
1s22s22p63s23p64s1
1s22s22p63s23p63d3 4s2
D
2
181. Which bond has the least ionic character?
B
C
(A) P—Cl
(B) H—Cl
(C) Br—Cl
(D) S—Cl
(E) Cl—Cl
1
(A) A
(B) B
182. Which type of bonding predominates in solid
potassium chloride, KCl?
(C) C
(D) D
176. A single burst of light is released from an atom.
Which statement explains what happens in the atom?
(A) An electron is changed from a particle to a wave.
(B) An electron moved from a higher to a lower energy
level.
(C) An electron pulled a proton out of the nucleus.
(D) An electron pulled a neutron out of the nucleus.
177. Neon atoms produce characteristic spectral lines
when their electrons
(A)
(B)
(C)
(D)
return to lower energy levels.
orbit the nucleus in a single energy level.
remain in their normal energy levels and move faster.
remain in their normal energy levels and move
slower.
(A) ionic
(B) metallic
(C) hydrogen
(D) covalent (molecular)
183. Which pair of elements react to form a compound
that has the greatest ionic character?
(A) xenon and fluorine
(B) carbon and oxygen
(C) cesium and chlorine
(D) iron and sulfur
184. Which compound contains both ionic and covalent
bonds?
(A) CO2
(B) KNO3
(C) NaCl
(D) CCl2F2
185. The electronegativity of francium is 0.7 and that of
fluorine is 4.0. The difference in electronegativity
suggests that the predominant bonding between Fr
and F is
(A) ionic.
(B) metallic. (C) covalent.
(D) very weak. (E) coordinate covalent
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
186. A solid has no electrical conductivity at room
temperature. It is heated to 600 °C, melts, and then
has electrical conductivity. The solid has which type
of bonding?
(A) ionic bonding
(B) covalent bonding
194. The graph below shows the boiling points of four
hydrogen compounds.
120
100
80
60
40
20
0
-20
-40
-60
-80
(C) metallic bonding
(D) van der Waals forces
187. The type of bond formed when two atoms share a
pair of electrons is called
(A) ionic.
(B) double.
(C) covalent
(D) bivalent.
(E) electrovalent.
188. A pure substance melts at 113 °C and does not
conduct electricity in either the solid or liquid state.
The bonding in this substance is primarily
(A) ionic.
(B) network.
(A) Li and Br
(B) Na and Br
(C) K and Br
(D) H and Br
190. When a chlorine molecule, Cl2, is formed, the orbital
overlap may be represented by the designation
(B) s – p
(C) s – s
(D) s – d
(E) p – d
POLARITY OF MOLECULES
191. Which represents a polar molecule?
(A) F2
H Te
HS
0
(C) metallic.
(D) covalent (molecular).
189. Which pair of atoms forms a covalent bond?
(A) p – p
HO
(B) O2
(C) CH4
(D) CO2
20
H Se
40 60 80 100 120 140
Molar Mass, g mol
What type of bonding explains the large difference
between the boiling points of H2O and the other hydrogen
compounds?
(A) ionic bonding
(B) covalent bonding
(C) hydrogen bonding
(D) van der Waals
attractions
195. An explanation of the heat of vaporization of water
being much higher than the heat of vaporization of
ethane (C2H6) is that
(A)
(B)
(C)
(D)
(E)
ethane has dipolar molecules.
water is more dense than liquid ethane.
water has a higher boiling point than ethane.
water molecules are lighter than ethane molecules.
energy is needed to break the hydrogen bonding
between water molecules.
(E) HCl
192. Which molecule is essentially nonpolar?
(A) CH4
(B) HCl
(C) HBr
(D) H2O
(E) NH3
193. The compounds H2S, H2Se, and H2Te boil below 0
°C at standard pressure. Water (H2O) boils at 100 °C.
This abnormally high boiling point of water is a
consequence of the
(A)
(B)
(C)
(D)
(E)
low molar mass of water.
low electrical conductivity of water.
covalent bonds in the water molecule.
stability of the bonds in the water molecules.
hydrogen bonds between the water molecules
196. The higher boiling point of HF compared with HCl,
HBr, and HI is caused by
(A)
(B)
(C)
(D)
(E)
covalently bonded molecules.
the size of the molecules.
the shape of the molecules.
hydrogen bonding between molecules.
weak van der Waals forces between HF molecules.
MOLECULAR SHAPES
197. Compounds that have the same molecular formula
but different structural formulas are known as
(A) isomers.
(B) polymers.
(C) isotopes.
(D) allotropes.
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
198. A molecule is said to be polar if it
(A)
(B)
(C)
(D)
(E)
209. Which term best describes the shape of the ammonia,
NH3, molecule?
has a north and south pole.
has a symmetrical electron distribution.
exhibits a polar spin under certain conditions.
may exhibit a positive or negative charge.
exhibits a partial positive charge at one end and a
partial negative charge at the other.
(D) H–H
Cl
(B) O=C=O
(E)
Cl
(C) tetrahedral
(D) trigonal planar
UNIT 8 RADIOACTIVITY/ATOMIC STRUCTURE
210. Rutherford’s alpha–particle bombardment of gold foil
helped develop our current model of the atom by
199. Which represents a polar molecule?
(A) H–Cl
(A) linear
(B) pyramidal
C Cl
(A)
(B)
(C)
(D)
Cl
finding the mass of the electron.
showing the existence of the neutron.
showing that the electron carries a negative charge.
showing that the atom has a concentrated central
charge
(C) NN
65
200. Which formula represents a nonpolar molecule?
(A) HCl
(B) CF4
(C) NH3
(D) H2S
201. Which is an example of a nonpolar molecule that
contains polar covalent bonds?
(A) CCl4
(B) N2
(C) H2O
(D) NH3
202. Which molecule is nonpolar?
(A) H2O
(B) HF
(C) NF3
(D) CF4
203. The shape of a chloroform molecule, CHCl3, is
(A) linear.
(B) cubical.
(C) octahedral.
(D) tetrahedral.
(E) planar triangular.
(B) HF
(C) NF3
(A)
(B)
(C)
(D)
(E)
30 protons and 35 neutrons.
35 protons and 30 neutrons.
35 protons and 35 neutrons.
65 protons and 30 neutrons.
95 protons and 30 electrons.
212. Isotopes differ in
(A) atomic number.
(B) nuclear charge.
(C) number of protons
(D) CF4
205. The arrangement of atoms in a water molecule, H2O,
is best described as
(D) number of neutrons.
(E) number of electrons.
213. A hypothetical element X has three isotopes: 40X,
41X, and 42X. Their abundances are 72.0%, 9.00%,
and 19.0% respectively. What is the atomic mass of
X?
(A) 40.5 u
204. Which molecule is nonpolar?
(A) H2O
211. The symbol 30Zn indicates this isotope contains
(B) 40.8 u
(C) 41.0 u
(D) 41.5 u
214. Copper has an atomic molar mass of 63.5 g·mol–1.
Why is the atomic molar mass not a whole number?
206. What is the shape of the ammonia, NH3, molecule?
(A) All copper atoms have identical chemical properties.
(B) The fractional number results from the fact that
protons and neutrons have different masses.
(C) There are at least two naturally occurring isotopes of
copper.
(D) Every copper atom has an atomic mass of 63.5 u.
(A) bent
(B) linear
215. The difference between the atomic number of an
atom and its mass number gives the number of
(A) ring.
(B) bent.
(C) linear.
(D) spherical.
(C) planar
(D) pyramidal
207. The shape of the CH4 molecule is most similar to the
shape of a molecule of
(A) H2O
(B) N2H4
(C) SiH4
(D) C2H4
208. Which molecule has all of its atoms in one plane?
(A) H2SO4
(B) CH4
(C) BF3
(D) NH3
(A) protons.
(B) neutrons.
(C) energy levels.
(D) orhitals.
(E) electrons.
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
216. Two kinds of emission from radioactive substances
that are considered to be particles of matter are
222. Given the nuclear reaction
234
234
90Th  91Pa
(A) alpha and beta emission.
(B) alpha and gamma emission.
+X
What is X?
1
0
(C) beta and gamma emission.
(D) gamma emission and X–radiation.
(E) alpha emission and X–radiation.
(A) A proton, 1p
217. What type of reaction is illustrated by this equation?
223. The half–life of radium is 1600 years. If a given
sample contains one gram of radium, how much
radium remains after 4800 years?
1
1H
3
1
4
(A) l g
(C) a fission reaction
(D) a fusion reaction
218. A radioactive element having atomic number 82 and
0
atomic mass 214 loses a beta particle, –1. The
resulting element has
Atomic No.
Atomic Mass
80
81
81
82
83
210 u
213 u
214 u
213 u
214 u
(A)
(B)
(C)
(D)
(E)
(E)
234
91Pa
(B)
236
92Th
234
92U
(C)
(D)
(A) beta
1
234
90Th
(D) 1/8 g
224. Strontium–90 has a half–life of 28 years. What
fraction of a sample remains as strontium–90 after 84
years?
(A)
(B)
(C)
(D)
(E)
(B) 1/8
(C) 1/4
(D) 1/3
pressure of the air at 0 °C only.
mass of a column of mercury.
temperature of the air at standard pressure.
density of mercury.
pressure of the air.
226. Which apparatus delivers 50.00 mL of liquid most
accurately?
(A)
(B)
(C)
(D)
50–mL buret
50–mL beaker
50–mL test tube
50–mL graduated cylinder
13
+ 0n  6C + ?
(B) alpha
(C) 1/3 g
225. A barometer is used to measure the
220. Which particle completes the equation?
16
8O
(B) 1/2 g
LAB TECHNIQUES/PROCEDURES
238
238
92U
(D) A beta particle, –1e
(E) 1/16 g
(A) 1/28
219. If the radioactive atom 92U emits an alpha particle,
the atom remaining is represented by
(A)
0
(B) A neutron, 0n
+ 1H  2He + energy
(A) a chemical reaction
(B) radioactive decay
(C) A positron, +1e
(C) proton
(D) neutron
227. Most student thermometers have an uncertainty of
0.2 Celsius degrees. Which is the proper reading of
the thermometer shown in the illustration?
(E) deuteron
221. Which nuclide is produced when a radioactive
carbon–14 atom emits an electron?
14
6C
14
(A) 6C
(B)
14
7
17
16
15
14
0
 ? + –1e
N
13
(C) 6C
13
(D) 5B
(A) l6. °C
(B) 16.4 °C(C)
16.40 °C (D) 16.45 °C
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
228. A narrow–necked, glass–stoppered bottle contains
sulfuric acid. When the acid is being poured, the
stopper should be
(A)
(B)
(C)
(D)
placed on the lab table.
put into the reaction vessel.
held in the palm of the hand.
held inverted between the index and middle fingers.
229. Which device is commonly used to measure liquid
volumes most precisely?
(A) graduated cylinder
(B) graduated beaker
(C) balance
(D) buret
230. This drawing shows the surface of water in a 10 mL
graduated cylinder. How much water is in the
cylinder?
8
7
6
(A) 6.20 mL (B) 6.25 mL (C) 6.40 mL (D) 7.80 mL
231. Which device should be used to measure 22.5 mL of
an aqueous solution?
A
(A) A
B
(B) B
C
(C) C
D
(D) D
232. In the laboratory, never dip a stirring rod into a
reagent bottle because
(A)
(B)
(C)
(D)
(E)
the bottle may tip.
the rod might break.
the rod may puncture the bottle.
the contents of the bottle may become contaminated.
the amount of liquid remaining on the rod is too
small to be used.
233. The purpose of filtration is to
(A)
(B)
(C)
(D)
form precipitates.
remove water from solutions.
separate dissolved ions from the solvent.
separate insoluble substances from a solution.