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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


(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


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


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


(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


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?


(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?


(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


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


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


(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:


(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


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


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


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?


(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.


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


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?


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.

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