REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
UNIT 1 INTRODUCTORY1. 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. Which involves a chemical change?(A) powdering sugar(B) condensing steam(C) magnetizing an iron bar(D) separating cream from milk(E) exposing photographic film to light
3. Which process is a chemical change?(A) the melting of ice(B) the burning of a candle(C) the magnetizing of steel(D) the liquefaction of oxygen
4. The graph was obtained by plotting the volume of a material vs. the mass of that same material.
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. Which is a unit for expressing volume?(A) mm(B) g(C) cm3
(D) g·cm
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
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
8. Which unit represents l ´ 10–3 mol?(A) decimole(B) kilomole (C) millimole(D) 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. The metric prefix for 10–6 is(A) mega–(B) kilo–(C) micro–(D) milli–
UNIT 2 MOLES12. Which expression represents the number of atoms in
1.0 ´ 10–3 g of lead?
(A)
(B)
(C)
(D)
(E)
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
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
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
(C) 48.0 g(D) 4.00 g
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) raising the temperature and lowering the pressure.(B) raising the pressure and lowering the temperature.(C) lowering both the temperature and pressure.(D) raising both the temperature and pressure.(E) lowering the temperature and keeping the presure
unchanged.
18. If the temperature and pressure are the same, one gram of hydrogen has about the same number of atoms as
Atomic Molar Masses
H 1.0 g·mol–1
O 16.0 g·mol–1
(A) 1 g of oxygen. .(B) 2 g of oxygen. .(C) 8 g of oxygen.(D) 16 g of oxygen(E) 32 g of oxygen
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) atoms at standard conditions.(B) molecules at all conditions.(C) molecules only at standard conditions.(D) atoms if conditions in both containers are the same.(E) molecules if conditions in both containers are the same.
21. Which is STP?
(A) 0 °C and 76 mmHg(B) 0 K and 76 mmHg(C) 0 K and 760 mmHg (D) 100 °C and 76 cmHg(E) 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.
CH4 O 2 SO2
Which gas exerts the greatest pressure? Assume ideal behavior.
(A) CH4
(B) O2(C) SO2(D) 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) 370(B) 444(C) 760(D) 1230(E) 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) 6.5(B) 13(C) 15(D) 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) 0.250(B) 2.00(C) 1.25(D) 4.00
27. For a given amount of dry gas at constant temperature, when the pressure is doubled the volume is
(A) halved.(B) unchanged.(C) doubled.(D) increased, but not doubled.
28. Approximately how many molecules are in 11 g of carbon dioxide, CO2, gas?
Atomic Molar Masses
C 12.0 g·mol–1
O 16.0 g·mol–1
(A) 1.5 ´ 1023
(B) 3.0 ´ 1023
(C) 6.0 ´ 1023
(D) 2.4 ´ 1023
29. What is the mass of one mole of calcium nitrate, Ca(NO3)2?
Atomic Molar Masses
Ca 40. g·mol–1
N 14. g·mol–1
O 16. g·mol–1
(A) 82 g(B) 102 g(C) 164 g(D) 204 g
30. The number of moles of water in 1,000 g of water is
Atomic Molar Masses
H 1.0 g·mol–1
O 16.0 g·mol–1
(A) 18.0(B) 55.5(C) 180.0
(D) 1000.0(E) 18,000.0
31. The molar mass of magnesium acetate, Mg(C2H3O2)2, in g·mol–1 is
Atomic Molar Masses
C 12. g·mol–1
H 1. g·mol–1
Mg 24. g·mol–1
O 16. g·mol–1
(A) 15(B) 16(C) 83(D) 142(E) 166
32. How many mole(s) of calcium carbonate, CaCO3, is represented by 50 g of the compound?
Atomic Molar Masses
Ca 40.1 g·mol–1
C 12.0 g·mol–1
O 16.0 g·mol–1
(A) 1.0(B) 2.0(C) 0.20(D) 4.0(E) 0.50
33. The molar mass of aluminum sulfate, Al2(SO4)3, is
Atomic Molar Masses
Al 27 g·mol–1
O 16 g·mol–1
S 32 g·mol–1
(A) 150 g·mol–1
(B) 170 g·mol–1
(C) 278 g·mol–1
(D) 342 g·mol–1
(E) 450 g·mol–1
34. The mass of one mole of ammonium carbonate, (NH4)2CO3, is approximately
Atomic Molar Masses
C 12.0 g·mol–1
H 1.0 g·mol–1
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
N 14.0 g·mol–1
O 16.0 g·mol–1
(A) 43.0 g(B) 72.0 g(C) 78.0 g(D) 96.0 g
35. The number of molecules present in 22.0 g of carbon dioxide at STP is
Atomic Molar Masses
C 12.0 g·mol–1
O 16.0 g·mol–1
(A) 2.01 ´ 1012
(B) 6.03 ´ 1012
(C) 2.06 ´ 1022
(D) 3.01 ´ 1023
(E) 6.02 ´ 1023
36. What mass of nitrogen dioxide, NO2, has the same number of molecules as 18.0 g of water, H2O?
Atomic Molar Masses
H 1.0 g·mol–1
N 14.0 g·mol–1
O 16.0 g·mol–1
(A) 6.02 g(B) 18.0 g(C) 23.0 g(D) 46.0 g
37. Calculate the mass of 12.0 ´ 1023 molecules of chlorine gas, Cl2.
Atomic Molar Mass
Cl 35.5 g·mol–1
(A) 35.5 g(B) 71.0 g(C) 142 g(D) 284 g
UNIT 3 NOMENCLATURE38. The correct formula for iron(III) sulfate is
(A) FeSO4 (D) Fe2(SO4)3(B) Fe(SO4)2 (E) Fe3(SO4)2
(C) Fe2SO4
39. The one correct formula among these is
(A) Na2OH (D) Zn(NO3)3(B) Cu(SO4)2 (E) BaNO3
(C) ZnCl2
40. Which formula is incorrect?
(A) BaHCO3 (B) Ca(OH)2 (C) Al2O3 (D) K2SO4
(E) ZnCO3
41. Which formula is incorrect?
(A) Al2(SO4)3 (D) NH4HSO4(B) BaHCO3 (E) LiH
(C) Ca(OH)2
42. What is the formula for aluminum sulfate?
(A) AlSO4(B) Al2SO4(C) Al3SO4(D) Al3(SO4)2
(E) 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?
(A) SrS (B) Sr2S (C) SrS2 (D) SrS3
45. What is the formula for copper(II) hydroxide?
(A) CuOH (B) Cu(OH)2 (C) Cu2OH (D) CuOH2
46. Which is the formula for ammonium nitrate?
(A) NH3N (B) NH4N (C) NH4NO2 (D) NH4NO3
47. What is the formula for sodium carbonate?
(A) NaHCO3 (C) So2CO3
(B) NaCO3 (D) Na2CO3
48. What is the formula for chromium(III) sulfate?
(A) Cr2(SO4)3 (C) Cr2(SO3)3
(B) Cr3(SO4)2 (D) Cr3SO4
49. Which formula is followed by its correct name?
(A) FeCl3, iron(III) chloride(B) FeS, iron(II) sulfite(C) Mg3N2, magnesium nitrite
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
(D) KNO2, potassium nitrate(E) HClO, hydrochloric oxide
50. The compound not properly named is
(A) Fe2O3, iron(III) oxide.(B) Pb3O4, lead(III) tetraoxide.(C) CuCl2, copper(II) chloride.(D) Pb3(PO4)2, lead(III) phosphate.(E) P2S5, diphosphorus pentasulfide.
51. What is the name of the compound having the formula CaH2?
(A) calcium amide (C) calcium hydrate(B) calcium hydride (D) calcium hydroxide
52. What is the name of the compound Fe2(SO4)3?
(A) iron(II) sulfate (C) iron(II) trisulfate(B) iron(III) sulfate (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 (C) Dy(BrO3)3
(B) Dy3BrO3 (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 (D) YbCl3
56. In which pair of anions do both names end in ‘–ate’?
(A) Cl–, ClO3– (C) NO2
–, NO3–
(B) ClO3–, NO3
– (D) HS–, HSO4–
57. Barium perrhenate has this formula: Ba(ReO4)2. The perrhenate ion is
(A) ReO4– (B) ReO4
2– (C) ReO43– (D) ReO4
4–
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) 4(B) 10(C) 11(D) 16(E) 3661. The total number of atoms represented by
5Al(C2H3O2)3 is
(A) 22(B) 60(C) 71(D) 84(E) 110
62. The number of atoms of oxygen indicated by the formula Ca3(PO4)2 is
(A) 12(B) 8(C) 7(D) 4(E) 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 X2Z 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 (C) NaCl, CH4, Br2(B) H3O+, Cl–, I3
– (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, NET-IONIC STOICHIOMETRY AND LIMITNG REACTANTS
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
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 (B) 2NO2 (C) 3NO2 (D) 4NO2
(E) 2NO
69. When the equation
? Sb + ? Cl2 ® ? SbCl3
is correctly balanced, the sum of the coefficients is
(A) 2 (B) 3 (C) 6 (D) 7
(E) 9
70. Which expression is correctly balanced?
(A) Na2O2 + 2H2O ® 2NaOH + O2(B) 2Na2O2 + 2H2O ® 4NaOH + 2O2(C) 4Na2O2 + 3H2O ® 4NaOH + 2O2(D) 2Na2O2 + 2H2O ® 4NaOH + O2(E) 3Na2O2 + 2H2O ® 6NaOH + O2
71. Which set of coefficients balances this equation?
? CH4(g) ? C3H8(g) + ? H2(g)
(A) 3, 1, 1 (B) 3, 2, 1 (C) 3, 1, 2 (D) 6, 2, 2
(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 (D) 1, 8, 3, 3(B) 3, 8, 6, 6 (E) 1, 4, 3, 3(C) 1, 4, 3, 2
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 (D) 3, l, 8, 7, 4, 2(B) 2, 4, 3, 8, 3, 3 (E) 3, 2, 8, 3, 2, 4(C) 3, 2, 8, 7, 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 (C) 1, 8, 5, 12(B) 2, 8, 10, 6 (D) 1, 11, 5, 12
75. Which set of coefficients correctly balances the equation?
? Al(s) + ? H+(aq) ® ? Al3+(aq) + ? H2(g)
(A) 1, 2, 1, 2 (C) 3, 2, 3, 2(B) 2, 6, 2, 3 (D) 2, 3, 2, 3
76. Which equation represents the complete combustion of acetylene in an excess of air?
(A) C2H2 + 2O2 ® 2CO2 + H2(B) C2H2 + O2 ® 2CO + H2(C) C2H2 + O ® 2C + H2O(D) C2H2 + O2 ® 2C + H2O2(E) 2C2H2 + 5O2 ® 4CO2 + 2H2O
77. Dysprosium oxide, Dy2O3, reacts with hydrochloric acid to produce only water and a salt. The salt is
(A) Dy2Cl3 (B) DyCl2 (C) DyCl3 (D) DyCl6
78. Which equation represents the dissolving of sodium sulfate, Na2SO4, in water?
(A) Na2SO4(s) ® Na2+(aq) + SO42–(aq)
(B) Na2SO4(s) ® 2Na+(aq) + SO42–(aq)
(C) Na2SO4(s) ® Na22+(aq) + S2–(aq) + 4O2–(aq)
(D) 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) Na+(aq) + NO3–(aq) ® Na(s) + 1/2N2(g) + 3/2O2(g)
(B) Ag+(aq) + Cl–(aq) ® Ag(s) + 1/2Cl2(g)(C) Ag+(aq) + Cl–(aq) Ag+(aq) + Cl–(aq)(D) Ag+(aq) + Cl–(aq) ® AgCl(s)
80. Which equation represents the dissolving (dissociation) of aluminum sulfate, Al2(SO4)3, in water?
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
(A) Al2(SO4)3(s) ® 2Al3+(aq) + 3S6+(aq) + 4O2–(aq)
(B) Al2(SO4)3(s) ® 2Al3+(aq) + 3SO42–(aq)
(C) Al2(SO4)3(s) ® 2Al2+(aq) + 3SO43–(aq)
(D) Al2(SO4)3(s) ® Al3+(aq) + SO43–(aq)
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) Ag+(aq) + Cl–(aq) AgCl(s)(B) K+(aq) + Cl–(aq) KCl(s)(C) K+(aq) + NO3
–(aq) KNO3(s)(D) 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) 2Ag+(aq) + CrO42–(aq) ® Ag2CrO4(s)
(B) Ag+(aq) + CrO42–(aq) ® Ag2CrO4(s)
(C) Ag0(aq) + CrO42–(aq) ® Ag2CrO4(s)
(D) Ag2CrO4(s) ® 2Ag+(aq) + CrO42–(aq)
84. Which two ions do not participate in the reaction between solutions of silver nitrate, AgNO3, and potassium chloride, KCl?
(A) K+ and Ag+ (C) K+ and Cl–(B) K+ and NO3
– (D) Ag+ and Cl–
85. In the equation:
BaCl2(aq) + Na2SO4(aq) ® BaSO4(s) +2NaCl (aq)
What is the net ionic equation for this reaction?
(A) Cl–(aq) + Na+(aq) ® NaCl (aq)(B) Cl22–(aq) + Na2
2+(aq) ® 2NaCl (aq)(C) Ba2+(aq) + SO4
2–(aq) ® BaSO4(s)
(D) BaCl2(s) + Na2SO4(s) ® Ba2+(aq) + 2Cl–(aq) + 2Na+(aq) + SO4
2–(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) + CrO4
2–(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) Pb2+(aq) + S2–(aq) ® PbS(s)(B) K+(aq) + NO3
–(aq) ® KNO3(aq)
(C) Pb(NO3)2(aq) + K2S(aq) ® PbS(s) + 2KNO3(aq)
(D) 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 (B) 2930 J (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
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?
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(A) 0.400 cm3 (C) 7.00 cm3
(B) 2.50 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
92. In the reaction
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 (D) 4
93. Given the equation
N2 + 3H2 2NH3
Theoretically, the number of moles of ammonia produced from 2 mol of nitrogen is
(A) 1 (B) 2 (C) 3 (D) 4
(E) 5
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) C4H4O4 + NaOH ® NaC4'H3O4 + H2O(B) C4H4O4 + 2NaOH ® Na2C4H2O4 + 2H2O(C) C4H4O4 + 3NaOH ® Na3C4HO4 + 3H2O(D) C4H4O4 + 4NaOH ® Na4C4O4 + 4H2O
95. In neutralizing 0.015 mol of H3PO3, 0.030 mol of NaOH was consumed. Which equation describes this reaction?
(A) H3PO3 + NaOH ® NaPO3 + H2O(B) H3PO3 + NaOH ® NaH2PO3 + H2O(C) H3PO3 + 3NaOH ® NaPO3 + 3H2O(D) 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 (B) H2O (C) HF (D) SiO2
97. How many grams of calcium carbonate, CaCO3, would be needed to produce 44.8 L of carbon dioxide gas, CO2, measured at STP?
Atomic Molar Masses
Ca 40.1 g·mol–1
C 12.0 g·mol–1
O 16.0 g·mol–1
CaCO3 + 2HCl ® CaCl2 + H2O + CO2
(A) 50.0 (B) 100 (C) 111 (D) 200
98. What volume of oxygen, O2, at STP can be prepared by the complete decomposition of 0.100 mol of potassium chlorate, KClO3?
2KClO3 ® 2KCl + 3O2
(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 (D) 10000 mL(B) 1600 mL (E) 20000 mL(C) 2000 mL
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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 1.0 g·mol–1
O 16.0 g·mol–1
(A) 1.0 g of hydrogen (D) 16 g of oxygen(B) 1.0 g of oxygen (E) 24 g of oxygen(C) 8.0 g of oxygen
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 (D) 4.00 L
(E) 13.0 L
103.The mass of potassium chloride formed by the complete decomposition of 490 g of potassium chlorate is
Atomic Molar Masses
Cl 35.5 g·mol–1
K 39.1 g·mol–1
O 16.0 g·mol–1
(A) 96 g (B) 122.5 g (C) 149 g (D) 298 g
(E) 490 g
104.In the reaction
2Al + 3H2SO4 ® 3H2 + Al2(SO4)3
the mass of aluminum that reacts with 1 mol of hydrogen ions is approximately
(A) 3.0 g (B) 9.0 g (C) 13.5 g (D) 27.0 g
(E) 81.0 g
105.What is the maximum mass of tungsten (W) obtained from the use of 18 g of hydrogen according to the equation
WO3 + 3H2 ® W + 3H2O
Atomic Molar Masses
H 1. g·mol–1
W 184. g·mol–1
(A) 1 ´ 184 g (D) 18 ´ 184 g(B) 3 ´ 184 g (E) 184 g + 3 ´ 16 g(C) 9 ´ 184 g
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
NaI 150. g·mol–1
I2 254. g·mol–1
(A) 5.00 g (B) 25.4 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?
3Fe(s) + 4H2O(g) ® Fe3O4(s) + 4H2(g)
Molar MassFe3O4 231. g·mol–1
(A) 154 g (B) 231 g (C) 462 g (D) 693 g
108.What mass of calcium hydroxide, Ca(OH)2, is obtained from 18.7 g of calcium oxide, CaO?
Atomic Molar Masses
Ca 40.1 g·mol–1
H 1.0 g·mol–1
O 16.0 g·mol–1
CaO + H2O ® Ca(OH)2
(A) 18.7 g (B) 24.7 g (C) 56.1 g (D) 74.1 g
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
109.Consider the equation:
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
Al 27.0 g·mol–1
H 1.0 g·mol–1
O 16.0 g·mol–1
(A) 3.86 g (B) 5.19 g (C) 4.20 g (D) 22.5 g
110.What mass of water is produced by complete combustion of 126 g of propene, C3H6?
2C3H6 + 9O2 ® 6H2O + 6CO2
Atomic Molar Masses
C 12.0 g·mol–1
H 1.0 g·mol–1
O 16.0 g·mol–1
(A) 18.0 g (B) 54.0 g (C) 126 g (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
Fe 55.8 g·mol–1
S 32.1 g·mol–1
Fe + S ® FeS
(A) 10.0 (B) 15.7 (C) 27.6 (D) 88.0
112.The equation for the complete combustion of propane, C3H8, is
C3H8(g) + 5O2(g) ® 3CO2(g) + 4H2(g)
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?
Atomic Molar Masses
C 12.0 g·mol–1
O 16.0 g·mol–1
(A) 22.0 g (B) 29.3 g (C) 44.0 g (D) 66.0 g(E) 132. g
113.Consider the equation:
CH4(g) + 2O2(g) ® CO2(g) + 2H2O(l)
How many moles of reactant are in excess when 2.0 mol of CH4(g) are ignited in 2.0 mol of O2(g)?
(A) l.0 mol CH4 (C) 0.5 mol CH4(B) 2.0 mol O2 (D) no excess of either
reactant
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
H 1.0 g·mol–1
O 16.0 g·mol–1
(A) 18.0 g (B) 36.0 g (C) 68.0 g (D) 72.0 g
EMPIRICAL FORMULAS
115.Upon analysis a compound is found to contain 22.8% sodium, 21.8% boron, and 55.4% oxygen. Its simplest formula is
Atomic Molar Masses
B 11 g·mol–1
Na 23 g·mol–1
O 16 g·mol–1
(A) Na2B4O7 (D) Na3B4O(B) NaBO (E) Na3BO4
(C) NaB2O5
116.A compound contains 85.71% carbon and 14.29% hydrogen by mass. Its simplest formula is
Atomic Molar Masses
C 12 g·mol–1
H 1 g·mol–1
(A) CH2 (B) CH (C) C2H (D) C2H2
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(E) C2H4
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?
Atomic Molar Masses
C 12.0 g·mol–1
H 1.0 g·mol–1
(A) CH2 (B) CH4 (C) C2H5 (D) C3H7
(E) C3H9
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 19.0 g·mol–1
S 32.0 g·mol–1
(A) SF (B) SF2 (C) SF3 (D) SF6
119.A compound has the empirical formula CH2O and the molecular mass 180 g·mol–1. What is its molecular formula?
(A) CH8O10 (C) C12H4O2
(B) C6H12O6 (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 12.0 g·mol–1
H 1.0 g·mol–1
(A) (CH3)1 (B) (CH3)2 (C) (CH3)3 (D) (CH3)4
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 12.0 g·mol–1
H 1.0 g·mol–1
(A) CH2 (B) C2H4 (C) C2H2 (D) CH4
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?
Atomic Molar Masses
F 19. g·mol–1
S 32. g·mol–1
(A) SF (B) S2F2 (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. (D) lost two electrons.(B) gained one electron. (E) gained two electrons.(C) gained one ion.124.An atom that loses or gains an electron becomes
(A) an ion. (D) a molecule.(B) a radical. (E) an electrolyte.(C) an isotope.
125.Metallic atoms become ions by
(A) losing protons. (C) gaining protons.(B) losing electrons. (D) gaining electrons.
126.How many electrons are in a chromium(III) ion, Cr3+?
(A) 52 (B) 27 (C) 24 (D) 21
127.The number of neutrons in the nucleus of an atom of Be is
(A) 36 (B) 13 (C) 9 (D) 5(E) 4
128.Which symbol represents an atom that contains the largest number of neutrons?
(A) U (B) U (C) Np (D) Pu
(E) Pa
129.An ion has 13 electrons, 12 protons, and 14 neutrons. What is the mass of the ion?
(A) 14 u (B) 25 u (C) 26 u (D) 27 u(E) 39 u
130.The symbol that represents 11 protons, 12 neutrons, and 10 electrons would be:
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(A) Na+ (B) Na (C) Mg2+ (D) Mg
131.The atomic number of an element is determined by the number of
(A) protons in each of its atoms.(B) neutrons in each of its atoms.(C) particles in each of its atoms.(D) protons plus neutrons in each of its atoms.(E) protons plus electrons in each of its atoms.
132.All positive ions differ from their corresponding atoms by having
(A) larger diameters.(B) fewer electrons.(C) a charge of +1.(D) greater atomic masses.(E) stronger metallic properties133.Which group represents particles that contain the
same number of electrons?
(A) F, Ne, Na (D) O2–, S2–, Se2–
(B) Mg, Al, Si (E) Ca2+, Fe2+, Zn2+
(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?
AY BX CYDXCYBX
ppt
ppt
S S S S
(A) DY (B) BY (C) AX (D) CX
135.An odorless, colorless, tasteless gas is suspected to be oxygen. Which result would support this hypothesis?
(A) The gas would extinguish a flame.(B) The gas would turn limewater milky.(C) The gas would burn in air producing only water.(D) A glowing splint would burst into flame in the gas.
136.The chemical properties of atoms depend principally upon
(A) their atomic masses.(B) the masses of the nuclei involved.(C) the number of neutrons in their nuclei.(D) 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) the same number of electrons in the outermost shell.(B) the same number of electrons.(C) the same number of protons.(D) similar nuclear structures.(E) a common origin
138.The best explanation of the extreme activity of fluorine as compared to other halogens is that the fluorine atom
(A) has the smallest atomic radius.(B) has the smallest nuclear charge.(C) has seven valence electrons.(D) is the strongest reducing agent.(E) needs one electron to complete its outermost shell.139.In the modern periodic table the elements are
arranged in the order of increasing
(A) atomic masses. (C) atomic numbers.(B) atomic radii. (D) atomic volumes.
140. In which set are the three elements in the same family?
(A) B, C, N (C) Hg, Ga, Sr(B) N, O, F (D) Zn, Cd, Hg
141.Which scientist is given credit for developing the periodic table?
(A) Rutherford (C) Dalton(B) Mendeleev (D) Planck
142.lf XO2 is the correct formula for an oxide, the formula for the chloride of X is
(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
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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 (C) transition metals(B) alkali metals (D) alkaline–earth metals
148.Which element is the most electronegative?(A) Be (B) Mg (C) Ca (D) Sr(E) Ba149.Since sodium and potassium are both members of
Group 1A in the periodic table, a sodium and a potassium atom have the same
(A) atomic mass.(B) number of protons in their nuclei.(C) atomic number and the same nuclear charge.(D) 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
(A) Na (B) Be (C) O (D) Cl
(E) Ar
151.In which part of the periodic table are the most electronegative elements found?
(A) upper left (C) upper right
(B) lower left (D) lower right
152.Consider a plot of a property of the alkaline earth metals.
Which property is plotted on this graph?
(A) first ionization energy(B) atomic radius(C) atomic mass(D) number of valence electrons
153.As the atomic numbers of the elements in a family increase, the
(A) atomic radii decrease.(B) atomic masses decrease.(C) ionization energies decrease.(D) elements become less metallic.(E) number of electrons in the outermost energy level
increases.
154.Which of these atoms has the smallest radius?
(A) K (B) Cl (C) Br (D) Cs
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155.Which characteristic of fluorine causes it to be the most active member of the halogen family, Group 7A?
(A) It forms diatomic molecules.(B) It has the smallest atomic radius.(C) It has no naturally occuring isotopes.(D) 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. (C) isotopes.(B) neutrons. (D) electrons.
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
(A) (D) ´ 100
(B) ´ 100 (E) ´ 100
(C) ´ 100
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% (C) 0.400%(B) 0.104% (D) 1.04%
159.The number of protons in the atom whose atomic mass is 89 and atomic number is 39, is
(A) 39 (B) 50 (C) 51 (D) 89
(E) 128
160.The particles present in the orbitals of an atom are
(A) mesons. (D) positrons.(B) protons. (E) electrons.(C) neutrons.
161.A neutral atom whose outermost electron shell contains eight electrons
(A) is very active.(B) has a combining number of one(C) is classified as a metal.(D) is chemically inert.(E) is more active than hydrogen.
162.When the halogens form ions, the result is
(A) colored ions.
(B) positive ions.(C) diatomic molecules.(D) covalent compounds.(E) a completed outer shell of electrons.
163.The correct electronic configuration for the sodium atom, Na, is
(A) 1s22s22p6
(B) 1s22s22p63s1
(C) 1s22s22p43s23p1
(D) 1s22s22p82d103s1(E) 1s22s22p62d103s23p1
164.Which element has the electron configuration 1s22s22p63s23p6 4s13d5?
(A) zinc (D) chromium(B) copper (E) potassium(C) nickel
165.The electron arrangement that represents the most active metallic element in this list is
(A) 2)7 (B) 2)8)1 (C) 2)8)2 (D) 2)8)3
(E) 2)8)6
166.What is the electronic configuration of an aluminum atom, Al?
(A) ls22s22p63d3
(B) 1s22s22p63s23p1
(C) ls22s22p62d13s2
(D) 1s22s22p62d103s23p5
(E) 1s22s22p63s23p63d74s2
167.Which atom contains a partially filled 3p orbital?
(A) iron (D) calcium(B) argon (E) aluminum(C) boron
168.Which element has the electron configuration 1s22s22p63s2?
(A) aluminum (C) magnesium(B) calcium (D) sodium
169.Which electron configuration represents an atom in an excited state?
(A) 1s22s22p6
(B) 1s22s22p63s2
(C) 1s22s22p63s23p64s23d1
(D) 1s22s22p63s23p64s24p1
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170.When two electrons occupy the same orbital, they must have
(A) opposite spins.(B) mutual attraction.(C) four identical quantum numbers.(D) different magnetic quantum numbers.(E) different principal quantum numbers.
171.What is the maximum number of electrons allowed in an orbital?
(A) 1 (B) 2 (C) 3 (D) 6
(E) 10
172.What neutral atom has the electron configuration 1s22s22p63s23p64s1?
(A) Na (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 (C) 1s22s22p63s23p64s1(B) ls22s22p63s23p6 (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?
(A) A (B) B (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) return to lower energy levels.(B) orbit the nucleus in a single energy level.(C) remain in their normal energy levels and move faster.(D) remain in their normal energy levels and move
slower.
178.Which electron configuration represents a transition element?
(A) 1s22s22p63s2
(B) 1s22s22p63s23p6
(C) 1s22s22p63s23p64s1(D) 1s22s22p63s23p63d3 4s2
BONDING
179.In which pair do both compounds exhibit ionic bonding?
(A) SO2, HCl (D) KCl, CO2(B) KNO3, CH4 (E) NaCl, H2O(C) NaF, KBr
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.
181.Which bond has the least ionic character?
(A) P—Cl (B) H—Cl (C) Br—Cl (D) S—Cl
(E) Cl—Cl
182.Which type of bonding predominates in solid potassium chloride, KCl?
(A) ionic (C) hydrogen(B) metallic (D) covalent (molecular)
183.Which pair of elements react to form a compound that has the greatest ionic character?
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(A) xenon and fluorine (C) cesium and chlorine(B) carbon and oxygen (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
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 (C) metallic bonding(B) covalent bonding (D) van der Waals forces
187.The type of bond formed when two atoms share a pair of electrons is called
(A) ionic. (D) bivalent.(B) double. (E) electrovalent.(C) covalent
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. (C) metallic.(B) network. (D) covalent (molecular).
189.Which pair of atoms forms a covalent bond?
(A) Li and Br (C) K and Br(B) Na and Br (D) H and Br
190.When a chlorine molecule, Cl2, is formed, the orbital overlap may be represented by the designation
(A) p – p (B) s – p (C) s – s (D) s – d
(E) p – d
POLARITY OF MOLECULES
191.Which represents a polar molecule?
(A) F2 (B) O2 (C) CH4 (D) CO2
(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) low molar mass of water.(B) low electrical conductivity of water.(C) covalent bonds in the water molecule.(D) stability of the bonds in the water molecules.(E) hydrogen bonds between the water molecules
194.The graph below shows the boiling points of four hydrogen compounds.
What type of bonding explains the large difference between the boiling points of H2O and the other hydrogen compounds?
(A) ionic bonding (C) hydrogen bonding(B) covalent 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) ethane has dipolar molecules.(B) water is more dense than liquid ethane.(C) water has a higher boiling point than ethane.(D) water molecules are lighter than ethane molecules.(E) energy is needed to break the hydrogen bonding
between water molecules.
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
196.The higher boiling point of HF compared with HCl, HBr, and HI is caused by
(A) covalently bonded molecules.(B) the size of the molecules.(C) the shape of the molecules.(D) hydrogen bonding between molecules.(E) 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. (C) isotopes.(B) polymers. (D) allotropes.
198.A molecule is said to be polar if it
(A) has a north and south pole.(B) has a symmetrical electron distribution.(C) exhibits a polar spin under certain conditions.(D) may exhibit a positive or negative charge.(E) exhibits a partial positive charge at one end and a
partial negative charge at the other.
199.Which represents a polar molecule?
(A) H–Cl (D) H–H
(B) O=C=O (E)
C ClCl
ClCl
(C) NºN
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. (D) tetrahedral.(B) cubical. (E) planar triangular.(C) octahedral.
204.Which molecule is nonpolar?
(A) H2O (B) HF (C) NF3 (D) CF4
205.The arrangement of atoms in a water molecule, H2O, is best described as
(A) ring. (B) bent. (C) linear. (D) spherical.
206.What is the shape of the ammonia, NH3, molecule?
(A) bent (C) planar(B) linear (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) NH3209. Which term best describes the shape of the ammonia,
NH3, molecule?
(A) linear (C) tetrahedral(B) pyramidal (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
(A) finding the mass of the electron.(B) showing the existence of the neutron.(C) showing that the electron carries a negative charge.(D) showing that the atom has a concentrated central
charge
211.The symbol Zn indicates this isotope contains
(A) 30 protons and 35 neutrons.(B) 35 protons and 30 neutrons.(C) 35 protons and 35 neutrons.(D) 65 protons and 30 neutrons.(E) 95 protons and 30 electrons.
212.Isotopes differ in
(A) atomic number. (D) number of neutrons.(B) nuclear charge. (E) number of electrons.(C) number of protons
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 (B) 40.8 u (C) 41.0 u (D) 41.5 u
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
214.Copper has an atomic molar mass of 63.5 g·mol–1. Why is the atomic molar mass not a whole number?
(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.
215.The difference between the atomic number of an atom and its mass number gives the number of
(A) protons. (D) orhitals.(B) neutrons. (E) electrons.(C) energy levels.
216.Two kinds of emission from radioactive substances that are considered to be particles of matter are
(A) alpha and beta emission.(B) alpha and gamma emission.
(C) beta and gamma emission.(D) gamma emission and X–radiation.(E) alpha emission and X–radiation.
217.What type of reaction is illustrated by this equation?
H + H ® He + energy
(A) a chemical reaction (C) a fission reaction(B) radioactive decay (D) a fusion reaction
218.A radioactive element having atomic number 82 and atomic mass 214 loses a beta particle, . The resulting element has
Atomic No. Atomic Mass
(A) 80 210 u(B) 81 213 u(C) 81 214 u(D) 82 213 u(E) 83 214 u
219.If the radioactive atom U emits an alpha particle, the atom remaining is represented by
(A) U (B) Th (C) U (D) Th
(E) Pa
220.Which particle completes the equation?
O + n ® C + ?
(A) beta (B) alpha (C) proton (D) neutron
(E) deuteron
221.Which nuclide is produced when a radioactive carbon–14 atom emits an electron?
C ® ? + e
(A) C (B) 147 N (C) C (D) B
222.Given the nuclear reaction
Th ® Pa + XWhat is X?
(A) A proton, p (C) A positron, e(B) A neutron, n (D) A beta particle, e
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?
(A) l g (B) 1/2 g (C) 1/3 g (D) 1/8 g
(E) 1/16 g
224.Strontium–90 has a half–life of 28 years. What fraction of a sample remains as strontium–90 after 84 years?
(A) 1/28 (B) 1/8 (C) 1/4 (D) 1/3
LAB TECHNIQUES/PROCEDURES
225.A barometer is used to measure the
(A) pressure of the air at 0 °C only.(B) mass of a column of mercury.(C) temperature of the air at standard pressure.(D) density of mercury.(E) pressure of the air.
226.Which apparatus delivers 50.00 mL of liquid most accurately?
(A) 50–mL buret(B) 50–mL beaker(C) 50–mL test tube(D) 50–mL graduated cylinder
227.Most student thermometers have an uncertainty of 0.2 Celsius degrees. Which is the proper reading of the thermometer shown in the illustration?
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
17161514
(A) l6. °C (B) 16.4 °C(C) 16.40 °C (D)16.45 °C
228.A narrow–necked, glass–stoppered bottle contains sulfuric acid. When the acid is being poured, the stopper should be
(A) placed on the lab table.(B) put into the reaction vessel.(C) held in the palm of the hand.(D) held inverted between the index and middle fingers.
229.Which device is commonly used to measure liquid volumes most precisely?
(A) graduated cylinder (C) balance(B) graduated beaker (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?
5 10152025
30
2010
30
20
10 25 mL
A B C D
(A) A (B) B (C) C (D) D
232.In the laboratory, never dip a stirring rod into a reagent bottle because
(A) the bottle may tip.(B) the rod might break.(C) the rod may puncture the bottle.(D) the contents of the bottle may become contaminated.(E) the amount of liquid remaining on the rod is too
small to be used.
233.The purpose of filtration is to
(A) form precipitates.(B) remove water from solutions.(C) separate dissolved ions from the solvent.(D) separate insoluble substances from a solution.
REVIEW SHEET 1ST SEMESTER HONORS CHEMISTRY
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