Captain Fordo
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- Posted: Tue, 24 Nov 2009 16:35:40 +0000
I need help with these problems. ninja
____ 1. Which of the following properties are consistent with liquids?
1. The volume of a liquid is determined by the size of its container.
2. A liquid has a rigid shape and fixed volume.
3. A liquid has a fixed volume that varies little with temperature or pressure changes.
a. 1 only b. 2 only c. 3 only d. 1, 2, and 3
____ 2. Concrete is composed of sand, gravel, and calcium oxide. Concrete is best described as
a. a heterogeneous mixture. b. a homogenous mixture. c. a pure substance. d. a chemical compound.
____ 3. Which of the following statements concerning the kinetic-molecular theory are correct?
1. Gas particles move faster when they are heated.
2. Particles in a liquid are closely spaced, but are not confined to specific positions.
3. Particles in a solid are closely spaced and are confined to specific positions.
a. 1 only b. 2 only c. 3 only d. 1, 2, and 3
____ 4. Nitrogen, N2, is a(n) ____ that is composed of two nitrogen ____.
a. atom, molecules b. compound, molecules c. element, atoms d. atom, elements
____ 5. Which of the following are chemical properties of iodine?
1. Iodine is a purple solid at 25 C.
2. Iodine reacts with sodium metal to form sodium iodide.
3. The density of iodine is 4.93 g/cm3.
a. 1 only b. 2 only c. 3 only d. 1 and 3
____ 6. Which of the following are physical properties of potassium?
1. Potassium reacts with water, producing hydrogen gas and aqueous potassium hydroxide.
2. Potassium conducts electricity.
3. Potassium is malleable at room temperature.
a. 1 only b. 2 only c. 3 only d. 2 and 3
____ 7. You can identify a metal by carefully determining its density. A 23.1 g piece of an unknown metal is 1.23 cm long, 2.11 cm wide, and 1.00 cm thick. What is the identity of the element?
a. nickel, 8.90 g/cm3 b. aluminum, 2.70 g/cm3 c. zirconium, 6.51 g/cm3 d. chromium 7.20 g/cm3
____ 8. The density of liquid mercury is 13.5 g/cm3. What mass of mercury (in kg) is required to fill a hollow cylinder having an inner diameter of 2.00 cm to a height of 25.0 cm? a.1.06 kg b. 4.24 kg c. 0.171 kg d.1.71 10-4 kg
____ 9. Which one of the following lists contains only extensive properties? a. melting point, density and color b. electrical conductivity and mass c. density, boiling point and volume d. volume and mass
____ 10. Thermostats are often set to 22 C. What is this temperature in Kelvin?a. 251 K b. 284 K c. 295 K d. 321 K
____ 11. Helium boils at 4.3 K. What is this temperature in Celsius?
a. 268.9 C b. 277.4 C c. -277.4 C d. -268.9 C
____ 12. The radius of a carbon atom is 7.7 10-11 m. What is the radius in picometers?
a. 7.7 pm b. 77 pm c. 7.7 102 pm d. 7.7 103 pm
____ 13. A typical volumetric flask holds a volume of 0.250 L. What is this volume in cubic centimeters?
a. 0.25 cm3 b. 2.50 cm3 c. 2.50 102 cm3 d. 2.50 103 cm3
____ 14. A student determines the density of a bar of silver by measuring its dimensions (2.00 cm by 1.15 cm by 1.00 cm) and determining its mass (25.3 g). If the true density of silver is 10.5 g/cm3, what is the percent error in the student's measurement? a. 1% b. 3% c. 5% d. 10%
___ 15. You and your lab partner are asked to determine the mass of a bar of silver. You use an analytical balance that measures mass to four decimal places (Method A). Your partner uses a top-loader balance that measures mass to two decimal places (Method B). The results are tabulated below.
Method A (g) Method B (g)
Measurement #1 3.3682 3.41
Measurement #2 3.3684 3.71
Measurement #3 3.3682 3.35
Measurement #4 3.3681 3.92
Average mass 3.3682 3.60
Percent Error 6.495% 0.1%
The actual mass of the silver bar is 3.6022 g. Which statements best describe the results?
a. Method A has good precision and poor accuracy. Method B has poor precision and good accuracy. b. Method A has poor precision and good accuracy. Method B has good precision and poor accuracy. c. Method A has poor precision and poor accuracy. Method B has good precision and good accuracy. d. Method A has good precision and good accuracy. Method B has good precision and poor accuracy.
____ 16. Which of the following statements are correct?
1. Electrons and protons have identical masses but opposite charges.
2. Most of an atom's mass is concentrated in a small, positively charged, nucleus.
3. Atoms have equal numbers of protons and neutrons.
a. 1 only b. 2 only c. 3 only d. 1 and 2
____ 17. Which of the following statements are correct?
1. Atomic number equals number of protons plus neutrons.
2. Mass number equals the number of neutrons.
3. An atomic mass unit equals 1/12 the mass of a carbon-12 atom.
a. 1 only b. 2 only c. 3 only d. 2 and 3
____ 18. Which of the following atoms has the greatest number of protons? a. 12C b. 15O c. 14C d. 15N
____ 19. Which of the following atoms contains the greatest number of neutrons? a. b. c. d.
____ 20. Silver has an average atomic mass of 107.9 u and is known to have only two naturally occurring isotopes. If 51.84% of Ag exists as Ag-107 (106.9051 u), what is the identity and the atomic mass of the other isotope?
a. Ag-110; 110.1 u b. Ag-110; 109.9 u c. Ag-108; 107.9 u d. Ag-109; 109.0 u
____ 21. An element consists of two isotopes. The abundance of one isotope is 60.40% and its atomic mass is 68.9257 u. The atomic mass of the second isotope is 70.9249 u. What is the average atomic mass of the element?
a. 69.72 u b. 69.93 u c. 70.13 u d. 139.9 u
____ 22. You have 0.500 g of the following elements: He, Ne, Ar, and Kr. Which sample contains the largest number of atoms? a. He b. Ne c. Ar d. Kr
____ 23. What is the mass of 0.442 mol Fe? a. 7.91 10-3 g b. 24.7 g c. 4.05 10-2 g d. 126 g
____ 24. Calculate the moles of Na in a 4.15 mg sample.
a. 1.81 10-4 mol b. 9.54 10-2 mol c. 1.05 10-2 mol d. 5.54 103 mol
____ 25. The molar mass of cesium is 132.9 g/mol. What is the mass of a single Cs atom?
a. 2.207 10-22 g b. 1.249 10-26 g c. 2.763 10-23 g d. 4.531 1021 g
____ 26. What is the mass of 5.8 1018 atoms of Ne?a. 9.6 10-6 g b. 1.9 10-4 g c. 4.8 10-7 g d. 5.1 103 g
____ 27. The density of silver is 10.5 g/cm3. What is the volume of a piece of Ag that contains 2.8 1022 atoms?
a. 0.48 cm3 b. 53 cm3 c. 4.8 10-2 cm3 d. 2.1 cm3
____ 28. What alkaline earth metal is located in the fourth period? a. K b. Ca c. Ga d. Ge
____ 29. Which chalcogen is located in the second period? a. P b. N c. O d. S
____ 30. Which grouping of elements is composed entirely of nonmetals? a. iodine, indium, and xenon b. aluminum, silicon, and phosphorus c. sulfur, neon, and bromine d. gallium, argon, and oxygen
____ 31. Which grouping of elements is composed entirely of metalloids?
a. B, As, and Sb b. Si, P, and Ge c. As, Ge, and Pb d. In, Sn, and Ge
____ 32. The following lists of elements are all found in the human body. Which three elements are found in the highest concentrations? a. sodium, oxygen, and magnesium b. oxygen, carbon, and hydrogen c. selenium, oxygen, and potassium d. carbon, oxygen, and iron
____ 33. The formula for acetic acid, CH3CO2H, is an example of a(n)
a. condensed formula. b. mathematical formula. c. structural formula. d. molecular formula.
____ 34. C2H6O is the formula for two possible molecules, ethanol and dimethyl ether. This type of formula is known as a(n) a. condensed formula. b. mathematical formula. c. structural formula. d. molecular formula.
____ 35. Which of the following statements are correct?
1. Metals generally lose electrons to become cations.
2. Nonmetals generally gain electrons to become anions.
3. Group 2A metals form ions with a 2+ charge.
a. 1 only b. 2 only c. 3 only d. 1,2, and 3
____ 36. What charge is most commonly observed for a calcium ion? a. 2- b. 1- c. 1+ d. 2+
____ 37. What charge is most commonly observed for an oxide ion? a. 3- b. 2- c. 1- d. 1+
____ 38. What charge is most commonly observed for a nickel ion? a. 4+ b. 2+ c. 1+ d. 2-
____ 39. What is the formula for magnesium oxide? a. MgO b. Mg2O c. MgO2 d. Mg2O3
____ 40. What is the formula for aluminum chloride? a. AlCl b. AlCl2 c. AlCl3 d. Al2Cl3
____ 41. What is the formula for calcium fluoride? a. CaF b. CaF2 c. Ca2F d. Ca2F3
____ 42. What is the formula for lead(IV) sulfate? a. PbS b. Pb(SO3)2 c. PbSO4 d. Pb(SO4)2
____ 43. What is the name of Mg(OH)2?
a. magnesium oxide b. magnesium dihydroxide c. magnesium hydroxide d. magnesium(II) hydroxide
____ 44. What is the name of KHCO3? a. hydrogen carbonate potassium ion b. potassium hydrogen carbon trioxide c. potassium bicarbonate d. potassium hydrocarbonate
____ 45. What is the name of (NH4)2Cr2O7? a. ammonium dichromium heptaoxide b. dichromate diammine c. diammonia dichromate d. ammonium dichromate
____ 46. What is the name of Cu2S? a. copper sulfide b. copper(I) sulfide c. copper(II) sulfide d. dicopper sulfide
____ 47. What is the name of PCl5? a. phosphorus chloride b. phosphorus pentachlorine c. phosphorus pentachloride d. phosphorus(V) chloride
____ 48. What is the molecular formula for silicon dioxide? a. SiO b. Si2O c. SiO2 d. (SiO)2
____ 49. What is the common name for NH3?
a. ammonia b. nitrogen trihydride c. trihydrogen nitride d. ammonium
____ 50. Which of the following statements are correct.
1. The attractive forces between ions of opposite charges increase with increasing separation.
2. The greater the charges on ions, the greater the attractive forces between ions.
3. Ions combine to form small molecules, such as NaCl.
a. 1 only b. 2 only c. 3 only d. 1 and 2
____ 51. Calculate the moles present in 10.5 g CaCO3. a. 0.105 mol b. 0.410 mol c. 9.53 mol d. 1.05 103 mol
____ 52. Determine the mass of 0.350 mol Na3PO4. a. 2.14 10-3 g b. 0.0174 g c. 57.4 g d. 163.9 g
____ 53. How many oxygen atoms are in 0.20 g CO2? a. 2.4 1023 oxygen atoms b. 2.7 1021 oxygen atoms c. 5.5 1021 oxygen atoms d. 1.2 1023 oxygen atoms
____ 54. What is the mass percent of oxygen in acetic acid, CH3CO2H? a. 26.64% b. 46.71% c. 53.29% d. 73.36%
____ 55. Toluene is composed of 91.25% C and 8.75% H. Determine the empirical formula for toluene.
a. CH b. CH3 c. C4H5 d. C7H8
____ 56. Benzene, an organic solvent, has the empirical formula CH. If the molar mass of benzene is 78.11 g/mol, what is the molecular formula of benzene? a. C4H30 b. C5H18 c. C6H6 d. C7H8
____ 57. You want to determine the value of x in hydrated iron(II) sulfate, FeSO4 • x H2O. In the laboratory you weigh out 1.983 g of the hydrated salt. After thorough heating, 1.084 g of the anhydrous salt remains. What is the value of x? a. 2 b. 3 c. 6 d. 7
____ 58. Balance the following chemical equation: Na2S(s) + HCl(aq) H2S(g) + NaCl(aq)
a. Na2S(s) + HCl(aq) H2S(g) + NaCl(aq) b. Na2S(s) + 2 HCl(aq) H2S(g) + 2 NaCl(aq) c. NaS(s) + HCl(aq) HS(g) + NaCl(aq) d. Na2S(s) + H2Cl(aq) H2S(g) + Na2Cl(aq)
____ 59. Balance the following chemical equation: CH3OH() + O2(g) CO2(g) + H2O(g)
a. CH3OH() + O2(g) CO2(g) + H2O(g) b. CH3OH() + 2 O2(g) CO2(g) + 2 H2O(g) c. 2 CH3OH() + 3 O2(g) 2 CO2(g) + 4 H2O(g) d. CH3OH() + O3(g) CO2(g) + 2 H2O(g)
____ 60. Potassium metal reacts with water to form aqueous potassium hydroxide and hydrogen gas. Write a balanced chemical equation for this reaction. a. K(s) + H2O() KOH(aq) + H2(g) b. 2 K(s) + 2 H2O() 2 KOH(aq) + H2(g) c. 2 K(s) + H2O() 2 KOH(aq) + H2(g) d. 4 K(s) + 2 H2O() 4 KOH(aq) + H2(g)
____ 61. What amount of carbon dioxide (in moles) is produced from the reaction of 2.24 moles of ethanol with excess oxygen? C2H5OH(aq) + 3 O2(g) 2 CO2(g) + 3 H2O(g) a. 1.12 mol b. 2.24 mol c. 4.48 mol d. 6.72 mol
____ 62. What amount of bromine (in moles) reacts with 4.0 mol of aluminum to produce AlBr3?
2 Al(s) + 3 Br2(g) 2 AlBr3(s) a. 2.0 mol b. 4.0 mol c. 6.0 mol d. 12 mol
____ 63. What amount of sodium (in moles) reacts with 1.5 moles of fluorine gas? The unbalanced chemical equation is given below. Na(s) + F2(g) NaF(s) a. 0.75 mol b. 1.5 mol c. 3.0 mol d. 6.0 mol
____ 64. What mass of magnesium iodide is produced from the reaction of 0.555 g Mg with excess iodine?
Mg(s) + I2(s) MgI2(s) a. 0.0228 g b. 0.0485 g c. 0.555 g d. 6.35 g
____ 65. What mass of sodium will react with 1.23 g of chlorine gas to produce sodium chloride?
2 Na(s) + Cl2(g) 2 NaCl(s) a. 0.199 g b. 0.399 g c. 0.798 g d. 2.46 g
____ 66. The unbalanced equation for the process of reducing iron ore to the metal is given below. What mass of iron metal is produced from the reduction of 10.0 g Fe2O3? Fe2O3(s) + CO(g) Fe(s) + CO2(g)
a. 3.50 g b. 6.99 g c. 24.6 g d. 57.2 g
____ 67. What amount of ammonia (in moles) is produced by the reaction of 4.00 mol H2 with 3.00 mol N2?
3 H2(g) + N2(g) 2 NH3(g) a. 2.67 mol b. 2.00 mol c. 6.00 mol d. 7.00 mol
____ 68. What mass of H2S is produced by the reaction of 1.55 g Na2S with 2.11 g HCl?
Na2S(s) + 2 HCl(g) 2 NaCl(s) + H2S(g)
a. 0.677 g b. 0.986 g c. 1.97 g d. 3.66 g
____ 69. The combustion of methane produces carbon dioxide and water. If 5.00 g CH4 reacts with 25.0 g O2, what mass of water is produced? a. 2.82 g b. 5.62 g c. 11.2 g d. 14.1 g
____ 70. The decomposition of 3.32 g CaCO3 yields 1.24 g CaO. What is the percent yield of this reaction?
CaCO3(s) CaO(s) + CO2(g) a. 1.86% b. 37.3% c. 62.7% d. 66.7%
____ 71. The reaction of 23.1 g NH3 and 18.3 g O2 produces 4.10 g NO. What is the percent yield of this reaction?
4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g) a. 9.90% b. 10.1% c. 22.4% d. 29.9%
____ 72. A mixture of CaCO3 and CaO has a mass of 2.500 g. After heating, the mass of the sample is reduced to 2.016 g. What is the mass percent of CaCO3 in the mixture? CaCO3(s) CaO(s) + CO2(g)
a. 19.4% b. 44.0% c. 54.6% d. 80.6%
____ 73. Cyclooctene is a hydrocarbon containing only C and H atoms. When burned in oxygen, 1.000 g of cyclooctene produces 3.195 g CO2 and 1.144 g H2O. What is the empirical formula of cyclooctene?
a. CH2 b. C2H5 c. C3H7 d. C4H7
____ 74. Benzoic acid contains C, H, and O atoms. When 1.500 g benzoic acid is burned in oxygen, 3.784 g CO2 and 0.6639 g H2O are produced. What is the empirical formula of benzoic acid?
a. C3H5O2 b. C4H6O c. C5H6O2 d. C7H6O2
____ 75. Which of the following compounds is a nonelectrolyte?
a. calcium chloride, CaCl2 b. acetic acid, CH3CO2H c. potassium iodide, KI d. ethanol, CH3CH2OH
____ 76. Which of the following statements concerning electrolytes are correct?
1. All ionic compounds are strong electrolytes.
2. Weak acids, such as acetic acid, are weak electrolytes.
3. Molecular species are weak electrolytes, provided they dissolve in water.
a. 1 only b. 2 only c. 3 only d. 1, 2, and 3
____ 77. Which compound is insoluble in water? a. BaSO4 b. K2HPO4 c. NaOH d. Ni(NO3)2
____ 78. Which compound is soluble in water? a. AgI b. Cu3(PO4)2 c. Zn(OH)2 d. (NH4)2CO3
____ 79. What ions are formed when NaHCO3 dissolves in water?
a. Na+, H+, and CO3- b. NaH+, and CO32- c. Na+ and HCO3- d. Na-, H+, and CO32-
____ 80. Which of the following acids is a weak acid? a. HF b. HCl c. HNO3 d. H2SO4
____ 81. What ions are produced from the dissolution of sodium hydroxide in water?
a. Na+ and H+ b. Na- and H+ c. Na+ and OH- d. Na+ and O2-
____ 82. Metal oxides are often referred to as basic oxides. Write a balanced chemical equation for the reaction of CaO and water. a. CaO(s) + H2O() CaO2(s) + H2(g) b. CaO(s) + H2O() Ca(OH)2(aq) c. CaO(s) + H2O() CaOH(aq) + OH-(aq) d. CaO(s) + H2O() H2CO2(aq)
____ 83. Nonmetal oxides are often referred to as acidic oxides. Write a balanced chemical equation for the reaction of SO3 and water. a. SO3(g) + H2O() H2SO4(aq) b. SO3(g) + H2O() SO2(g) + 2 OH-(aq) c. SO3(g) + H2O() SO2(g) + H2O2(aq) d. SO3(g) + H2O() H2SO3(aq) + O2-(aq)
____ 84. Write a balanced net ionic equation for the reaction of NaOH with FeCl3.
a. Na+(aq) + Cl-(aq) NaCl(s) b. Fe3+(aq) + 3 Na+(aq) Fe(Na)3(s) c. Fe3+(aq) + 3 OH-(aq) Fe(OH)3(s) d. Cl-(aq) + OH-(aq) HOCl(s)
____ 85. Write a balanced equation for the reaction of aqueous solutions of silver nitrate with calcium bromide.
a. 2 AgNO3(aq) + CaBr2(aq) CaAg2(s) + 2 BrNO3(aq) b. Ag(NO3)2(aq) + CaBr2(aq) AgBr2(aq) + Ca(NO3)2(s) c. AgNO3(aq) + CaBr2(aq) AgBr2(s) + CaNO3(aq) d. 2 AgNO3(aq) + CaBr2(aq) 2 AgBr(s) + Ca(NO3)2(aq)
____ 86. Write a balanced equation for the reaction of acetic acid with potassium hydroxide.
a. CH3CO2H(aq) + KOH(aq) KCH3CO2(aq) + H2O() b. CH3CO2H(aq) + KOH(aq) KCH3O(aq) + OH-(aq) + CO2(g) c. CH3CO2H(aq) + KOH(aq) KCH3(aq) + H2O() + CO2(g) d. CH3CO2H(aq) + KOH(aq) KCH3CO3H(aq)
____ 87. Write a net ionic equation for the reaction of hydrobromic acid with sodium hydroxide.
a. Na+(aq) and Br-(aq) NaBr(aq) b. HBr(aq) + NaOH(aq) H2O() + NaBr(aq) c. H+(aq) + OH-(aq) H2O() d. Na+(aq) + Br-(aq) NaBr(s)
____ 88. Write a balanced chemical equation for the reaction of zinc(II) carbonate with nitric acid.
a. ZnCO3(s) + 2 HNO3(aq) Zn(NO3)2(aq) + H2O() + CO2(g) b. ZnCO3(s) + HNO3(aq) ZnO(s) + CO2(g) + HNO3(aq) c. ZnCO3(s) + 2 HNO3(aq) Zn(OH)2(s) + N2O5(s) + CO2(g) d. ZnCO3(s) + 2 HNO3(aq) Zn(NO3)2(aq) + H2CO3(aq)
____ 89. The chemical equation below, 2 Na(s) + 2 H2O() 2 NaOH(aq) + H2(g)
is an example of a(n) ____ reaction.
a. precipitation b. oxidation-reduction c. acid-base d. both oxidation-reduction and acid-base
____ 90. The chemical equation below, AlCl3(aq) + 3 NaOH(aq) Al(OH)3(s) + 3 NaCl(aq)
is an example of a(n) ____ reaction.
a. precipitation b. oxidation-reduction c. strong acid-strong base d. gas-forming
____ 91. Determine the oxidation number of sulfur in SO42-. a. +2 b. +4 c. +6 d. +8
____ 92. Determine the oxidation number of each atom in CH2Cl2.
a. C = 0, H = +1, Cl = -1 b. C = +2, H = 0, Cl = -1 c. C = 0, H = 0, Cl = 0 d. C = -2, H = +1, Cl = 0
____ 93. Which of the following lists contains only common oxidizing agents?
a. K, Ca, and F2 b. Li, Na, and K c. F2, MnO4-, and HNO3 d. F-, Cl-, and Br-
____ 94. If 8.33 g Ca(NO3)2 is dissolved in enough water to make 0.250 L of solution, what is the molar concentration of Ca(NO3)2? a. 0.0127 M b. 0.203 M c. 3.33 M d. 4.92 M
____ 95. What volume of 0.250 M KOH contains 25.0 g KOH? a. 0.111 L b. 0.781 L c. 0.625 L d. 1.78 L
____ 96. If 45.0 mL of 0.244 M NiCl2 is diluted to 275 mL with pure water, what is the molar concentration of NiCl2 in the diluted solution? a. 0.0399 M b. 0.105 M c. 0.671 M d. 1.49 M
____ 97. If 0.40 M of acetic acid solution has a pH of 2.57, what is the H+ concentration of the solution?
a. 2.7 10-3 M b. 0.039 M c. 0.40 M d. 3.7 102 M
____ 98. What is the pH of a 0.34 M HNO3 solution? a. -0.47 b. 0.34 c. 0.47 d. 2.13
____ 99. If you combine 35.0 mL of 0.100 M AgNO3 with 45.0 mL of 0.0800 M NaBr, what mass of AgBr is produced? a. 0.537 g b. 0.657 g c. 0.676 g d. 0.818 g
____ 100. A 25.00 mL sample of sulfuric acid, H2SO4, requires 42.13 mL of 0.1533 M NaOH for titration to the equivalence point. What is the concentration of the sulfuric acid?
H2SO4(aq) + 2 NaOH Na2SO4(aq) + 2 H2O()
a. 0.04558 M b. 0.1292 M c. 0.2583 M d. 0.5167 M
____ 1. Which of the following properties are consistent with liquids?
1. The volume of a liquid is determined by the size of its container.
2. A liquid has a rigid shape and fixed volume.
3. A liquid has a fixed volume that varies little with temperature or pressure changes.
a. 1 only b. 2 only c. 3 only d. 1, 2, and 3
____ 2. Concrete is composed of sand, gravel, and calcium oxide. Concrete is best described as
a. a heterogeneous mixture. b. a homogenous mixture. c. a pure substance. d. a chemical compound.
____ 3. Which of the following statements concerning the kinetic-molecular theory are correct?
1. Gas particles move faster when they are heated.
2. Particles in a liquid are closely spaced, but are not confined to specific positions.
3. Particles in a solid are closely spaced and are confined to specific positions.
a. 1 only b. 2 only c. 3 only d. 1, 2, and 3
____ 4. Nitrogen, N2, is a(n) ____ that is composed of two nitrogen ____.
a. atom, molecules b. compound, molecules c. element, atoms d. atom, elements
____ 5. Which of the following are chemical properties of iodine?
1. Iodine is a purple solid at 25 C.
2. Iodine reacts with sodium metal to form sodium iodide.
3. The density of iodine is 4.93 g/cm3.
a. 1 only b. 2 only c. 3 only d. 1 and 3
____ 6. Which of the following are physical properties of potassium?
1. Potassium reacts with water, producing hydrogen gas and aqueous potassium hydroxide.
2. Potassium conducts electricity.
3. Potassium is malleable at room temperature.
a. 1 only b. 2 only c. 3 only d. 2 and 3
____ 7. You can identify a metal by carefully determining its density. A 23.1 g piece of an unknown metal is 1.23 cm long, 2.11 cm wide, and 1.00 cm thick. What is the identity of the element?
a. nickel, 8.90 g/cm3 b. aluminum, 2.70 g/cm3 c. zirconium, 6.51 g/cm3 d. chromium 7.20 g/cm3
____ 8. The density of liquid mercury is 13.5 g/cm3. What mass of mercury (in kg) is required to fill a hollow cylinder having an inner diameter of 2.00 cm to a height of 25.0 cm? a.1.06 kg b. 4.24 kg c. 0.171 kg d.1.71 10-4 kg
____ 9. Which one of the following lists contains only extensive properties? a. melting point, density and color b. electrical conductivity and mass c. density, boiling point and volume d. volume and mass
____ 10. Thermostats are often set to 22 C. What is this temperature in Kelvin?a. 251 K b. 284 K c. 295 K d. 321 K
____ 11. Helium boils at 4.3 K. What is this temperature in Celsius?
a. 268.9 C b. 277.4 C c. -277.4 C d. -268.9 C
____ 12. The radius of a carbon atom is 7.7 10-11 m. What is the radius in picometers?
a. 7.7 pm b. 77 pm c. 7.7 102 pm d. 7.7 103 pm
____ 13. A typical volumetric flask holds a volume of 0.250 L. What is this volume in cubic centimeters?
a. 0.25 cm3 b. 2.50 cm3 c. 2.50 102 cm3 d. 2.50 103 cm3
____ 14. A student determines the density of a bar of silver by measuring its dimensions (2.00 cm by 1.15 cm by 1.00 cm) and determining its mass (25.3 g). If the true density of silver is 10.5 g/cm3, what is the percent error in the student's measurement? a. 1% b. 3% c. 5% d. 10%
___ 15. You and your lab partner are asked to determine the mass of a bar of silver. You use an analytical balance that measures mass to four decimal places (Method A). Your partner uses a top-loader balance that measures mass to two decimal places (Method B). The results are tabulated below.
Method A (g) Method B (g)
Measurement #1 3.3682 3.41
Measurement #2 3.3684 3.71
Measurement #3 3.3682 3.35
Measurement #4 3.3681 3.92
Average mass 3.3682 3.60
Percent Error 6.495% 0.1%
The actual mass of the silver bar is 3.6022 g. Which statements best describe the results?
a. Method A has good precision and poor accuracy. Method B has poor precision and good accuracy. b. Method A has poor precision and good accuracy. Method B has good precision and poor accuracy. c. Method A has poor precision and poor accuracy. Method B has good precision and good accuracy. d. Method A has good precision and good accuracy. Method B has good precision and poor accuracy.
____ 16. Which of the following statements are correct?
1. Electrons and protons have identical masses but opposite charges.
2. Most of an atom's mass is concentrated in a small, positively charged, nucleus.
3. Atoms have equal numbers of protons and neutrons.
a. 1 only b. 2 only c. 3 only d. 1 and 2
____ 17. Which of the following statements are correct?
1. Atomic number equals number of protons plus neutrons.
2. Mass number equals the number of neutrons.
3. An atomic mass unit equals 1/12 the mass of a carbon-12 atom.
a. 1 only b. 2 only c. 3 only d. 2 and 3
____ 18. Which of the following atoms has the greatest number of protons? a. 12C b. 15O c. 14C d. 15N
____ 19. Which of the following atoms contains the greatest number of neutrons? a. b. c. d.
____ 20. Silver has an average atomic mass of 107.9 u and is known to have only two naturally occurring isotopes. If 51.84% of Ag exists as Ag-107 (106.9051 u), what is the identity and the atomic mass of the other isotope?
a. Ag-110; 110.1 u b. Ag-110; 109.9 u c. Ag-108; 107.9 u d. Ag-109; 109.0 u
____ 21. An element consists of two isotopes. The abundance of one isotope is 60.40% and its atomic mass is 68.9257 u. The atomic mass of the second isotope is 70.9249 u. What is the average atomic mass of the element?
a. 69.72 u b. 69.93 u c. 70.13 u d. 139.9 u
____ 22. You have 0.500 g of the following elements: He, Ne, Ar, and Kr. Which sample contains the largest number of atoms? a. He b. Ne c. Ar d. Kr
____ 23. What is the mass of 0.442 mol Fe? a. 7.91 10-3 g b. 24.7 g c. 4.05 10-2 g d. 126 g
____ 24. Calculate the moles of Na in a 4.15 mg sample.
a. 1.81 10-4 mol b. 9.54 10-2 mol c. 1.05 10-2 mol d. 5.54 103 mol
____ 25. The molar mass of cesium is 132.9 g/mol. What is the mass of a single Cs atom?
a. 2.207 10-22 g b. 1.249 10-26 g c. 2.763 10-23 g d. 4.531 1021 g
____ 26. What is the mass of 5.8 1018 atoms of Ne?a. 9.6 10-6 g b. 1.9 10-4 g c. 4.8 10-7 g d. 5.1 103 g
____ 27. The density of silver is 10.5 g/cm3. What is the volume of a piece of Ag that contains 2.8 1022 atoms?
a. 0.48 cm3 b. 53 cm3 c. 4.8 10-2 cm3 d. 2.1 cm3
____ 28. What alkaline earth metal is located in the fourth period? a. K b. Ca c. Ga d. Ge
____ 29. Which chalcogen is located in the second period? a. P b. N c. O d. S
____ 30. Which grouping of elements is composed entirely of nonmetals? a. iodine, indium, and xenon b. aluminum, silicon, and phosphorus c. sulfur, neon, and bromine d. gallium, argon, and oxygen
____ 31. Which grouping of elements is composed entirely of metalloids?
a. B, As, and Sb b. Si, P, and Ge c. As, Ge, and Pb d. In, Sn, and Ge
____ 32. The following lists of elements are all found in the human body. Which three elements are found in the highest concentrations? a. sodium, oxygen, and magnesium b. oxygen, carbon, and hydrogen c. selenium, oxygen, and potassium d. carbon, oxygen, and iron
____ 33. The formula for acetic acid, CH3CO2H, is an example of a(n)
a. condensed formula. b. mathematical formula. c. structural formula. d. molecular formula.
____ 34. C2H6O is the formula for two possible molecules, ethanol and dimethyl ether. This type of formula is known as a(n) a. condensed formula. b. mathematical formula. c. structural formula. d. molecular formula.
____ 35. Which of the following statements are correct?
1. Metals generally lose electrons to become cations.
2. Nonmetals generally gain electrons to become anions.
3. Group 2A metals form ions with a 2+ charge.
a. 1 only b. 2 only c. 3 only d. 1,2, and 3
____ 36. What charge is most commonly observed for a calcium ion? a. 2- b. 1- c. 1+ d. 2+
____ 37. What charge is most commonly observed for an oxide ion? a. 3- b. 2- c. 1- d. 1+
____ 38. What charge is most commonly observed for a nickel ion? a. 4+ b. 2+ c. 1+ d. 2-
____ 39. What is the formula for magnesium oxide? a. MgO b. Mg2O c. MgO2 d. Mg2O3
____ 40. What is the formula for aluminum chloride? a. AlCl b. AlCl2 c. AlCl3 d. Al2Cl3
____ 41. What is the formula for calcium fluoride? a. CaF b. CaF2 c. Ca2F d. Ca2F3
____ 42. What is the formula for lead(IV) sulfate? a. PbS b. Pb(SO3)2 c. PbSO4 d. Pb(SO4)2
____ 43. What is the name of Mg(OH)2?
a. magnesium oxide b. magnesium dihydroxide c. magnesium hydroxide d. magnesium(II) hydroxide
____ 44. What is the name of KHCO3? a. hydrogen carbonate potassium ion b. potassium hydrogen carbon trioxide c. potassium bicarbonate d. potassium hydrocarbonate
____ 45. What is the name of (NH4)2Cr2O7? a. ammonium dichromium heptaoxide b. dichromate diammine c. diammonia dichromate d. ammonium dichromate
____ 46. What is the name of Cu2S? a. copper sulfide b. copper(I) sulfide c. copper(II) sulfide d. dicopper sulfide
____ 47. What is the name of PCl5? a. phosphorus chloride b. phosphorus pentachlorine c. phosphorus pentachloride d. phosphorus(V) chloride
____ 48. What is the molecular formula for silicon dioxide? a. SiO b. Si2O c. SiO2 d. (SiO)2
____ 49. What is the common name for NH3?
a. ammonia b. nitrogen trihydride c. trihydrogen nitride d. ammonium
____ 50. Which of the following statements are correct.
1. The attractive forces between ions of opposite charges increase with increasing separation.
2. The greater the charges on ions, the greater the attractive forces between ions.
3. Ions combine to form small molecules, such as NaCl.
a. 1 only b. 2 only c. 3 only d. 1 and 2
____ 51. Calculate the moles present in 10.5 g CaCO3. a. 0.105 mol b. 0.410 mol c. 9.53 mol d. 1.05 103 mol
____ 52. Determine the mass of 0.350 mol Na3PO4. a. 2.14 10-3 g b. 0.0174 g c. 57.4 g d. 163.9 g
____ 53. How many oxygen atoms are in 0.20 g CO2? a. 2.4 1023 oxygen atoms b. 2.7 1021 oxygen atoms c. 5.5 1021 oxygen atoms d. 1.2 1023 oxygen atoms
____ 54. What is the mass percent of oxygen in acetic acid, CH3CO2H? a. 26.64% b. 46.71% c. 53.29% d. 73.36%
____ 55. Toluene is composed of 91.25% C and 8.75% H. Determine the empirical formula for toluene.
a. CH b. CH3 c. C4H5 d. C7H8
____ 56. Benzene, an organic solvent, has the empirical formula CH. If the molar mass of benzene is 78.11 g/mol, what is the molecular formula of benzene? a. C4H30 b. C5H18 c. C6H6 d. C7H8
____ 57. You want to determine the value of x in hydrated iron(II) sulfate, FeSO4 • x H2O. In the laboratory you weigh out 1.983 g of the hydrated salt. After thorough heating, 1.084 g of the anhydrous salt remains. What is the value of x? a. 2 b. 3 c. 6 d. 7
____ 58. Balance the following chemical equation: Na2S(s) + HCl(aq) H2S(g) + NaCl(aq)
a. Na2S(s) + HCl(aq) H2S(g) + NaCl(aq) b. Na2S(s) + 2 HCl(aq) H2S(g) + 2 NaCl(aq) c. NaS(s) + HCl(aq) HS(g) + NaCl(aq) d. Na2S(s) + H2Cl(aq) H2S(g) + Na2Cl(aq)
____ 59. Balance the following chemical equation: CH3OH() + O2(g) CO2(g) + H2O(g)
a. CH3OH() + O2(g) CO2(g) + H2O(g) b. CH3OH() + 2 O2(g) CO2(g) + 2 H2O(g) c. 2 CH3OH() + 3 O2(g) 2 CO2(g) + 4 H2O(g) d. CH3OH() + O3(g) CO2(g) + 2 H2O(g)
____ 60. Potassium metal reacts with water to form aqueous potassium hydroxide and hydrogen gas. Write a balanced chemical equation for this reaction. a. K(s) + H2O() KOH(aq) + H2(g) b. 2 K(s) + 2 H2O() 2 KOH(aq) + H2(g) c. 2 K(s) + H2O() 2 KOH(aq) + H2(g) d. 4 K(s) + 2 H2O() 4 KOH(aq) + H2(g)
____ 61. What amount of carbon dioxide (in moles) is produced from the reaction of 2.24 moles of ethanol with excess oxygen? C2H5OH(aq) + 3 O2(g) 2 CO2(g) + 3 H2O(g) a. 1.12 mol b. 2.24 mol c. 4.48 mol d. 6.72 mol
____ 62. What amount of bromine (in moles) reacts with 4.0 mol of aluminum to produce AlBr3?
2 Al(s) + 3 Br2(g) 2 AlBr3(s) a. 2.0 mol b. 4.0 mol c. 6.0 mol d. 12 mol
____ 63. What amount of sodium (in moles) reacts with 1.5 moles of fluorine gas? The unbalanced chemical equation is given below. Na(s) + F2(g) NaF(s) a. 0.75 mol b. 1.5 mol c. 3.0 mol d. 6.0 mol
____ 64. What mass of magnesium iodide is produced from the reaction of 0.555 g Mg with excess iodine?
Mg(s) + I2(s) MgI2(s) a. 0.0228 g b. 0.0485 g c. 0.555 g d. 6.35 g
____ 65. What mass of sodium will react with 1.23 g of chlorine gas to produce sodium chloride?
2 Na(s) + Cl2(g) 2 NaCl(s) a. 0.199 g b. 0.399 g c. 0.798 g d. 2.46 g
____ 66. The unbalanced equation for the process of reducing iron ore to the metal is given below. What mass of iron metal is produced from the reduction of 10.0 g Fe2O3? Fe2O3(s) + CO(g) Fe(s) + CO2(g)
a. 3.50 g b. 6.99 g c. 24.6 g d. 57.2 g
____ 67. What amount of ammonia (in moles) is produced by the reaction of 4.00 mol H2 with 3.00 mol N2?
3 H2(g) + N2(g) 2 NH3(g) a. 2.67 mol b. 2.00 mol c. 6.00 mol d. 7.00 mol
____ 68. What mass of H2S is produced by the reaction of 1.55 g Na2S with 2.11 g HCl?
Na2S(s) + 2 HCl(g) 2 NaCl(s) + H2S(g)
a. 0.677 g b. 0.986 g c. 1.97 g d. 3.66 g
____ 69. The combustion of methane produces carbon dioxide and water. If 5.00 g CH4 reacts with 25.0 g O2, what mass of water is produced? a. 2.82 g b. 5.62 g c. 11.2 g d. 14.1 g
____ 70. The decomposition of 3.32 g CaCO3 yields 1.24 g CaO. What is the percent yield of this reaction?
CaCO3(s) CaO(s) + CO2(g) a. 1.86% b. 37.3% c. 62.7% d. 66.7%
____ 71. The reaction of 23.1 g NH3 and 18.3 g O2 produces 4.10 g NO. What is the percent yield of this reaction?
4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g) a. 9.90% b. 10.1% c. 22.4% d. 29.9%
____ 72. A mixture of CaCO3 and CaO has a mass of 2.500 g. After heating, the mass of the sample is reduced to 2.016 g. What is the mass percent of CaCO3 in the mixture? CaCO3(s) CaO(s) + CO2(g)
a. 19.4% b. 44.0% c. 54.6% d. 80.6%
____ 73. Cyclooctene is a hydrocarbon containing only C and H atoms. When burned in oxygen, 1.000 g of cyclooctene produces 3.195 g CO2 and 1.144 g H2O. What is the empirical formula of cyclooctene?
a. CH2 b. C2H5 c. C3H7 d. C4H7
____ 74. Benzoic acid contains C, H, and O atoms. When 1.500 g benzoic acid is burned in oxygen, 3.784 g CO2 and 0.6639 g H2O are produced. What is the empirical formula of benzoic acid?
a. C3H5O2 b. C4H6O c. C5H6O2 d. C7H6O2
____ 75. Which of the following compounds is a nonelectrolyte?
a. calcium chloride, CaCl2 b. acetic acid, CH3CO2H c. potassium iodide, KI d. ethanol, CH3CH2OH
____ 76. Which of the following statements concerning electrolytes are correct?
1. All ionic compounds are strong electrolytes.
2. Weak acids, such as acetic acid, are weak electrolytes.
3. Molecular species are weak electrolytes, provided they dissolve in water.
a. 1 only b. 2 only c. 3 only d. 1, 2, and 3
____ 77. Which compound is insoluble in water? a. BaSO4 b. K2HPO4 c. NaOH d. Ni(NO3)2
____ 78. Which compound is soluble in water? a. AgI b. Cu3(PO4)2 c. Zn(OH)2 d. (NH4)2CO3
____ 79. What ions are formed when NaHCO3 dissolves in water?
a. Na+, H+, and CO3- b. NaH+, and CO32- c. Na+ and HCO3- d. Na-, H+, and CO32-
____ 80. Which of the following acids is a weak acid? a. HF b. HCl c. HNO3 d. H2SO4
____ 81. What ions are produced from the dissolution of sodium hydroxide in water?
a. Na+ and H+ b. Na- and H+ c. Na+ and OH- d. Na+ and O2-
____ 82. Metal oxides are often referred to as basic oxides. Write a balanced chemical equation for the reaction of CaO and water. a. CaO(s) + H2O() CaO2(s) + H2(g) b. CaO(s) + H2O() Ca(OH)2(aq) c. CaO(s) + H2O() CaOH(aq) + OH-(aq) d. CaO(s) + H2O() H2CO2(aq)
____ 83. Nonmetal oxides are often referred to as acidic oxides. Write a balanced chemical equation for the reaction of SO3 and water. a. SO3(g) + H2O() H2SO4(aq) b. SO3(g) + H2O() SO2(g) + 2 OH-(aq) c. SO3(g) + H2O() SO2(g) + H2O2(aq) d. SO3(g) + H2O() H2SO3(aq) + O2-(aq)
____ 84. Write a balanced net ionic equation for the reaction of NaOH with FeCl3.
a. Na+(aq) + Cl-(aq) NaCl(s) b. Fe3+(aq) + 3 Na+(aq) Fe(Na)3(s) c. Fe3+(aq) + 3 OH-(aq) Fe(OH)3(s) d. Cl-(aq) + OH-(aq) HOCl(s)
____ 85. Write a balanced equation for the reaction of aqueous solutions of silver nitrate with calcium bromide.
a. 2 AgNO3(aq) + CaBr2(aq) CaAg2(s) + 2 BrNO3(aq) b. Ag(NO3)2(aq) + CaBr2(aq) AgBr2(aq) + Ca(NO3)2(s) c. AgNO3(aq) + CaBr2(aq) AgBr2(s) + CaNO3(aq) d. 2 AgNO3(aq) + CaBr2(aq) 2 AgBr(s) + Ca(NO3)2(aq)
____ 86. Write a balanced equation for the reaction of acetic acid with potassium hydroxide.
a. CH3CO2H(aq) + KOH(aq) KCH3CO2(aq) + H2O() b. CH3CO2H(aq) + KOH(aq) KCH3O(aq) + OH-(aq) + CO2(g) c. CH3CO2H(aq) + KOH(aq) KCH3(aq) + H2O() + CO2(g) d. CH3CO2H(aq) + KOH(aq) KCH3CO3H(aq)
____ 87. Write a net ionic equation for the reaction of hydrobromic acid with sodium hydroxide.
a. Na+(aq) and Br-(aq) NaBr(aq) b. HBr(aq) + NaOH(aq) H2O() + NaBr(aq) c. H+(aq) + OH-(aq) H2O() d. Na+(aq) + Br-(aq) NaBr(s)
____ 88. Write a balanced chemical equation for the reaction of zinc(II) carbonate with nitric acid.
a. ZnCO3(s) + 2 HNO3(aq) Zn(NO3)2(aq) + H2O() + CO2(g) b. ZnCO3(s) + HNO3(aq) ZnO(s) + CO2(g) + HNO3(aq) c. ZnCO3(s) + 2 HNO3(aq) Zn(OH)2(s) + N2O5(s) + CO2(g) d. ZnCO3(s) + 2 HNO3(aq) Zn(NO3)2(aq) + H2CO3(aq)
____ 89. The chemical equation below, 2 Na(s) + 2 H2O() 2 NaOH(aq) + H2(g)
is an example of a(n) ____ reaction.
a. precipitation b. oxidation-reduction c. acid-base d. both oxidation-reduction and acid-base
____ 90. The chemical equation below, AlCl3(aq) + 3 NaOH(aq) Al(OH)3(s) + 3 NaCl(aq)
is an example of a(n) ____ reaction.
a. precipitation b. oxidation-reduction c. strong acid-strong base d. gas-forming
____ 91. Determine the oxidation number of sulfur in SO42-. a. +2 b. +4 c. +6 d. +8
____ 92. Determine the oxidation number of each atom in CH2Cl2.
a. C = 0, H = +1, Cl = -1 b. C = +2, H = 0, Cl = -1 c. C = 0, H = 0, Cl = 0 d. C = -2, H = +1, Cl = 0
____ 93. Which of the following lists contains only common oxidizing agents?
a. K, Ca, and F2 b. Li, Na, and K c. F2, MnO4-, and HNO3 d. F-, Cl-, and Br-
____ 94. If 8.33 g Ca(NO3)2 is dissolved in enough water to make 0.250 L of solution, what is the molar concentration of Ca(NO3)2? a. 0.0127 M b. 0.203 M c. 3.33 M d. 4.92 M
____ 95. What volume of 0.250 M KOH contains 25.0 g KOH? a. 0.111 L b. 0.781 L c. 0.625 L d. 1.78 L
____ 96. If 45.0 mL of 0.244 M NiCl2 is diluted to 275 mL with pure water, what is the molar concentration of NiCl2 in the diluted solution? a. 0.0399 M b. 0.105 M c. 0.671 M d. 1.49 M
____ 97. If 0.40 M of acetic acid solution has a pH of 2.57, what is the H+ concentration of the solution?
a. 2.7 10-3 M b. 0.039 M c. 0.40 M d. 3.7 102 M
____ 98. What is the pH of a 0.34 M HNO3 solution? a. -0.47 b. 0.34 c. 0.47 d. 2.13
____ 99. If you combine 35.0 mL of 0.100 M AgNO3 with 45.0 mL of 0.0800 M NaBr, what mass of AgBr is produced? a. 0.537 g b. 0.657 g c. 0.676 g d. 0.818 g
____ 100. A 25.00 mL sample of sulfuric acid, H2SO4, requires 42.13 mL of 0.1533 M NaOH for titration to the equivalence point. What is the concentration of the sulfuric acid?
H2SO4(aq) + 2 NaOH Na2SO4(aq) + 2 H2O()
a. 0.04558 M b. 0.1292 M c. 0.2583 M d. 0.5167 M
