Download Final Exam Review

Final Exam Review
Chemistry 2
14-15
1. Liquid ethanol evaporates readily. Which of the following statements about this process is wrong? (Ch. 14)
a. Evaporation is a phase change.
b. The molecules become separated to relatively great distances during evaporation.
c. The molecules of vapor have a different chemical composition from those in the liquid.
d. The density of the vapor is less than that of the liquid.
e. The vapor is much more compressible than the liquid.
Once vaporized and ignited, ethanol burns readily. Balance the reaction and use it to answer question 2-5.
C2H5OH + O2  CO2 + H2O + 327 kcal
2. How many moles of CO2 are produced when 1.0 mole of ethanol burns? (Ch. 9)
a. 0.50
d. 2.0
b. 1.0
e. impossible to determine
c. 1.5
3. When one mole of C2H5OH burns completely, the volume of O2 consumed, measured at standard temperature and pressure, is
(Ch. 9)
a. Variable
d. 44.8 liters
b. 7.5 liters
e. 67.2 liters
c. 22.4 liters
4. How many grams of H2O are formed when one mole of ethanol burns? (Ch. 9)
a. 3.0 grams
d. 36 grams
b. 6.0 grams
e. 54 grams
c. 18 grams
5. If all the heat evolved from burning 1.0 mole of ethanol is transferred to 10,000 grams of water, the temperature of the water
will rise approximately (Ch. 10)
a. 8.0 C
d. 430 C
b. 33 C
e. 327 kcal
c. 43 C
6. How many moles of gas are present in a sample which occupies 2.05 liters at 3.0 atm pressure and 27 C? (Ch. 13)
a. 0.15 moles
d. 2.77 moles
b. 0.25 moles
e. 3.0 moles
c. 2.05 moles
7. If 100 moles of Mg and 100 moles of O2 are allowed to react to form MgO, the maximum mass of MgO that can be formed is
(Ch. 9)
2 Mg + O2 → 2 MgO
8. Which of the following statements about ionic and covalent bonding is false? (Ch. 12)
a. Covalent bonds are always formed between atoms having high ionization energies.
b. Ionic bonding results in substances which form conducting solutions when dissolved in water.
c. In ionic bonding, atoms gain and lose electrons to obtain the same number as the nearest noble gas.
d. Covalent bonding can only happen between identical atoms.
e. Covalent bonding may result in single, double or triple bonds.
9. Assume that each of the following reactions is initially at equilibrium. For which of them will a decrease in pressure favor
increased formation of products? (Ch. 17)
a. H2(g) + I2(g)  2 HI(g)
b. 2 NOBr(g)  2 NO(g) + Br2(g)
c. H2O(g) + CuSO4(s)  CuSO42H2O(s)
d. 2 NO(g)  N2O4(g)
10.
The number of moles of sulfate ion, SO42-, in 500 mL of a 0.2 M solution of Al 2(SO4)3 is (Ch. 15)
a. 0.1 mole
d. 0.4 mole
b. 0.2 mole
e. 0.5 mole
c. 0.3 mole
1
Questions 11-13 deal with the changes in kinetic and potential energy which occur at the molecular level as water is heated from
a temperature below 0 C to an extremely high temperature. Each question might have more than one correct answer. The
temperature intervals which are to be considered are as follows:
(1)
(2)
(3)
(4)
(5)
-20 C to 0 C, the warming of solid ice
0 C, the melting of solid water into liquid water
0 C to 100 C, the warming of liquid water
100 C, the vaporization of water
100 C to 150 C, the warming of water vapor
11. The incoming heat energy is converted mainly into potential energy. (Ch. 14 )
a. (1)
d. (4)
b. (2)
e. (5)
c. (3)
12. The incoming heat energy is being converted mainly into kinetic energy. (Ch. 14)
a. (1)
d. (4)
b. (2)
e. (5)
c. (3)
13. For each of the following, determine whether you would use Q=mCT , Hfus, or Hvap to calculate the amount of heat energy
involved in the change. (Ch. 14)
a. (1)
d. (4)
b. (2)
e. (5)
c. (3)#20
14. The equation for the dissolving of Ba(NO3)2 in water is which of the following? (Ch. 15)
a. Ba(NO3)2(s) 
Ba2+(aq) + 2 NO3-(aq)
b. Ba(NO3)2(s) 
Ba2+(aq) + (NO3)2-(aq)
c. Ba(NO3)2(s) 
Ba2+(aq) + (NO3)22-(aq)
d. Ba(NO3)2(s) 
Ba(NO3)2(aq)
e. Ba2+(aq) + (NO3)2-(aq)
 Ba(NO3)2(s)
15. Which of the following statements about the part of the electromagnetic spectrum that is visible is not true? (Ch. 11)
a. The velocity of visible light is 3.00 x 108 meters/second.
b. The frequency of visible light can be found by the relationship  = c/.
c. The higher the frequency of light, the greater its energy per photon.
d. The longer the wavelength of light, the greater its energy per photon.
e. Visible light is a very small part of the entire electromagnetic spectrum.
16. What is the partial pressure of helium in a mixture that consists of 8.0 g helium gas, He, and 8.0 g of hydrogen gas, H 2, if the
total pressure is 1,200 mm Hg? (Ch. 13)
a. 120 mm Hg
d. 600 mm Hg
b. 240 mm Hg
e. 800 mm Hg
c. 400 mm Hg
17. A 2-cm-thick piece of cardboard placed over a radiation source would be most effective in protecting against which type of
radiation? (Ch. 19)
A) alpha
C) gamma
B) beta
D) x-ray
18. Which of the following atoms has six valence electrons? (Ch. 12)
A) magnesium (Mg)
C) sulfur (S)
B) silicon (Si)
D) argon (Ar)
19. Which of the following statements about the quantum-mechanical model for the atom is not true? (Ch. 11)
a. Electrons exist only in shells which are a specific distance from the nucleus, determined by the principal quantum number, n.
b. The principal quantum number, n, is used to describe the energy level change in which an electron may be found.
c. For every principal quantum number, n, there are n2 orbitals available for electron occupancy.
d. It is often useful to consider the electron’s wave characteristics rather than its particle characteristics.
e. The quantum-mechanical model tells only the probability of finding an electron in a given area of space.
2
20. Given the generalizations implied by the periodic table, which of the following formulas is not reasonable? (Ch. 12)
a. SiO2
d. CaF
b. Al2O3
e. BeF2
c. CO2
21. The mass of a 1.00 liter sample of gas, measured at STP, is found to be 1.96 grams. The gas could have which of the following
formulas? [HINT: Calculate the moles of gas. Molar mass has these units…. Grams / moles] (Ch. 6)
a. CH4
d. N2
b. NH3
e. CO2
c. NO2
22. Consider two identical flasks, both at 25 C and 1 atm pressure. One contains H2 gas and the other O2 gas. Which of the
following statements about the gases contained in the two flasks is not true? (Ch. 13)
a. The same number of molecules are contained in each flask.
b. The average kinetic energy of the molecules in both flasks is the same.
c. The average velocity of the molecules in both flasks is the same.
d. The flask filled with H2 gas will have a smaller mass than the flask filled with O2 gas.
e. The molecules in both flasks are in rapid motion.
23. Positive ions are usually formed by which of the following processes? (Ch. 3)
a. gain of protons by the nucleus.
b. loss of electrons from the nucleus.
c. loss of neutrons from the nucleus.
d. gain of electrons either inside or outside the nucleus.
e. loss of electrons from outside the nucleus.
24. Plutonium-238 (238Pu) undergoes radioactive decay to give ?U as one of products. The other product and the mass number of
the uranium produced are which of the following? (Ch. 19)
a. an alpha particle (4He) and 238.
2
b. a beta particle (0e-) and 239
-1
c. a gamma ray and 235
d. an alpha particle (4He) and 234
2
0 -
e. a beta particle ( e ) and 235
-1
25. Which of the following statements about the periodic table is false? (Ch. 11)
a. Metals are generally located on the left side of the periodic table, and nonmetals on the right.
b. The elements are arranged in order of increasing atomic number.
c. Elements in each vertical column of the periodic table have similar chemical reactions.
d. Elements in each horizontal row have nearly identical physical characteristics.
e. Elements on the left side of the periodic table generally form positively charged ions when they react.
26. Consider Al, As, Ga, Ge. Which of the following elements has the same Lewis dot structure as silicon? (Ch. 12)
A) germanium (Ge)
C) arsenic (As)
B) aluminum (Al)
D) gallium (Ga)
27. Given the reaction at equilibrium
N2O4(g)  2 NO2(g)
H = +14.1 kcal
Which of the following changes would not increase the concentration of NO2(g)? (Ch. 17)
a. increasing the concentration of N2O4(g)
b. increasing the volume of the container
c. decreasing the pressure on the system
d. addition of a catalyst
e. increasing the temperature at which the reaction is run
28. Which of the following is not a characteristic of an exothermic reaction? (Ch. 10)
a. The enthalpy of the products is greater than that of the reactants.
b. The reaction container becomes hot.
c. H is negative.
d. Heat is a product of the reaction.
e. The reaction will probably go farther toward completion at low temperatures than at high temperatures.
3
29. Which of the following would not cause an increase in the volume of a gas trapped in a syringe with one atm pressure upon it?
(Ch. 13)
a. changing to a different gas of higher molecular weight
b. heating the syringe
c. increasing the amount of gas in the syringe
d. a sudden, dramatic decrease in atmospheric pressure
e. placing the syringe in a vacuum
30. Which of the following electron configurations for neutral atoms in their lowest energy state is not correct? (Ch. 11)
a. fluorine
1s22s22px22py22pz1
b. rubidium
[Kr] 5s1
c. scandium
1s22s22p63s23p64s24px1
d. zinc
[Ar] 4s23d10
e. germanium
[Ar] 4s23d104px14py1
31. When a 9.81-gram sample of sulfuric acid, H2SO4, is added to 200 ml of water, the temperature of the water rose 8.9 C. On the
basis of this information, the heat of solution, H of sulfuric acid is which of the following? [HINT: Calculate the energy
absorbed by the water, using the calorimetry equation (q = mTC), where C = 1.00 cal/g ºC. Then calculate the moles of
sulfuric acid dissolved. Knowing that energy is conserved, you set the energy absorbed by the water = energy released by the
acid. heat of solution, H, has these units: kcal per mole. Convert cal to kcal and plug in your data.] (Ch. 10)
a. –1.78 kcal/mole
d. –17. 8 kcal/mole
b. –8.9 kcal/mole
e. –17,800 kcal/mole
c. –9.81 kcal/mole
32. An isotope of krypton, 89Kr, has a half-life of 3.2 minutes. If the original sample was 100 grams, how many grams
of the original sample will remain undecayed at the end of 9.6 minutes? (Ch. 19)
a. none
d. 37.5 g
b. 12.5 g
e. 50 g
c. 25 g
33. The isotope 27Al contains ______ protons in its nucleus. (Ch. 3)
a. 13
d. 40
b. 14
e. none of these
c. 27
34. Two identical containers, one containing helium gas and the other containing oxygen gas, are at the same temperature and
pressure. Which of the following statements is false? (Ch. 13)
a. The average kinetic energy of the oxygen molecules is the same as the average kinetic energy of the helium molecules.
b. The average velocity of the helium molecules is the same as the average velocity of the oxygen molecules.
c. The oxygen in one container weighs eight times more than the helium gas in the other container.
d. The number of moles oxygen in one flask is the same as the number of moles of helium in the other.
e. If both containers were heated to 100 ºC, the average kinetic energies of both gases would increase but remain equal to
each other.
35. Which of the following types of nuclear decay results in an increase in the nuclear charge? (Ch. 19)
a. alpha emission
d. positron emission
b. beta emission
e. gamma emission
c. electron capture
36. Which element is a liquid at room temperature? (Ch. 3)
a. N2
d. S
b. Fe
e. Na
c. Hg
37. Oxygen gas can be produced by the decomposition of potassium chlorate, according this equation:
2 KClO3(s)  2 KCl(s) + 3 O2 (g). The oxygen gas is collected over water at 27 C in a 2.00 liter container at a total
pressure of 760. torr. The vapor pressure of water at 27 C is 26.0 torr. Determine the moles of potassium chlorate that
reacted in this experiment. (Ch. 13)
a. 0.0790 moles
d. 0.0813 moles
b. 0.119 moles
e. 0.0345 moles
c. 0.0527 moles
4
38. Which of the following is the correct form of the equilibrium constant expression for the reaction below? (Ch. 17)
NH3(aq) + H2O(l)
 NH4+(aq) + OH-(aq)
a. K = [NH4+][OH-]/[NH3][H2O]
b. K = [NH3]
c. K = [NH4+][OH-]
d. K = [NH4+][OH-]/[NH3]
e. K = [NH3][H2O]/[NH4+][OH-]
39. Which atom has the highest ionization energy? (Ch. 11)
a. Mg
b. Be
c. Ba
d. Ca
e. Sr
40. One gram (1.00 g) of a gaseous hydrocarbon occupies 0.821 L at 1.00 atm and 147 C. The compound is [HINT: calc moles of
the gas using PV=nRT. Solve for molar mass by dividing grams by moles (g/mole)] (Ch.13)
a. CH4
d. C2H2
b. C2H4
e. C3H6
c. C4H8
41. Which type of orbital has the highest energy? (Ch. 11)
a. 3s
b. 3p
c. 4s
d. 3d
e. 4f
42. When acid is added to an active metal, ___________ a flammable gas is produced. (Think about the gas produced when you
added hydrochloric acid to magnesium metal in our lab). (Ch. 13)
a. CO b. H2
c. OH- d. CO2 e. H3O+
43. Which element in the nitrogen family is most metallic? (Ch. 11)
a. nitrogen
d. antimony
b. phosphorus
e. bismuth
c. arsenic
44. Which of the following is a chemical property for carbon? (Ch. 2)
a. it is a solid at room temperature
b. it is not magnetic
c. it has a density of 2.25 g/cm3
d. it undergoes a combustion to carbon dioxide
e. it boils at 4200 C
45. In the reaction of
CaO + 3 C  CaC2 + CO
If 52 grams of CaO are mixed with 27 grams of C, how many grams of CaC 2 can be produced? (Ch. 9)
a. 25
b. 75
c. 57
d. 48
e. 133
46. What is the percentage composition of N in the fertilizer (NH4)2SO4? [HINT: percent composition is just the mass of the
individual element divided by the mass of the total compound, x 100] (Ch. 6)
a.10.6%
b. 7.58%
c. 5.30%
d. 21.1%
e. 37.9%
47. In 0.500 moles of acetic acid, CH3COOH, there are (Ch. 6)
a. 6.02 x 1023 molecules
d. 2.4 x 1024 atoms
b. 1.2 x 1024 molecules
e. 4.8 x 1024 atoms
24
c. 1.8 x 10 atoms
48. Which one of the following statements about the periodic table is false? (Ch. 11)
a. The noble gases, the least reactive of all the elements, are located in the last vertical column to the far right of the table.
b. Each vertical column of the periodic table contains elements with similar chemical and physical properties.
c. Metals are generally located on the left side of the periodic table, and nonmetals on the right.
d. Elements in the same horizontal row of the table have similar physical characteristics.
e. The alkali metals are located in the first vertical column at the far left of the table.
49. Which of these involves a chemical change? (Ch. 2)
a. burning
b. dissolving
c. melting
d. dilution
5
e. condensation
50. Hydrogen sulfide reacts with oxygen gas to produce sulfur dioxide and water. What is the sum of the coefficients for the
balanced equation? (Ch. 7)
a. 6
b. 8
c. 9
d. 11
e. 17
51. Which of the following is not a solid at room temperature? (Ch. 11)
a. sodium
b. bromine
c. phosphorous
d. boron
52. Once cubic millimeter is equal to a volume of (Ch. 5)
a. 1 mL
b. 0.01 mL
c. 100 mL
d. 0.001 mL
e. silicon
e.10 mL
53. What is the mass of one platinum atom? (Ch. 6)
a.
1.6 x 10-22 g
d.
195.08 g
b.
3.2 x 10-22 g
e. none of these
c.
5.13 x 10-3 g
54. A 6.12 g sample of silver undergoes a temperature change from 25.5 C to 55.0 C when 10.24 calories of heat are supplied.
What is the specific heat of silver in calories /g C ? (Ch. 10)
a. 0.0669
b. 0.0567
c. 0.0304
d. 1.67
e. 1.00
55. Which is not true for covalent bonds? (Ch. 12)
a. The pair of electrons is shared between the two atoms.
b. More than two electrons can be shared between the two atoms.
c. Covalent bonds usually do not break in aqueous solutions.
d. One electron in the bond must come from each atom.
e. Ionic bonds are about as strong as covalent bonds.
56. Which of the following molecules contains one triple bond? (Ch. 12)
a. NH3
d. H2CCH2
b. CO
e. HNNH
c. H2CCO
57. Which of the following diatomic molecules has the greatest bond strength? (Ch. 12)
a. H2
d. O2
b. F2
e. HF
c. N2
58. How many atoms are in 160 grams of diatomic nitrogen gas? (Ch. 6)
a. 6.9 x 1024
d. 9.6 x 1025
24
b. 3.4 x 10
e. 1.6 x 1023
25
c. 1.4 x 10
59. Which of the following is not a pure substance? (Ch. 2)
a. table salt
d. water
b. milk
e. hydrochloric acid
c. baking soda
60. A phase change from solid to gas without melting in between is called (Ch. 14)
a. sublimation
d. vaporization
b. freezing
e. fusion
c. precipitation
61. Which of the following is a metal? (Ch. 3)
a. B
d. Be
b. C
e. Ge
c. Si
62. The physical state that retains both volume and shape is (Ch. 3)
a. excited
d. solid
b. gas
e. liquid
c. plasma
6
63. What is the electron configuration of a stable magnesium atom? (Ch. 11)
a. 1s22s22p5
d. 1s22s22p63s2
b. 1s22s22p6
e. 1s22s22p63s23p1
2 2
6 1
c. 1s 2s 2p 3s
64. The following species S2-, Cl-, Ar, K+ , all have the same number of (Ch. 3)
a. protons
d. neutrons
b. nucleons
e. isotopes
c. electrons
65. Which pair of atoms is most likely to form a covalent chemical bond? (Ch. 12)
a. H and H
d. Na and Cl
b. He and Ne
e. Li and Br
c. Na and Na
66. Which of the following is likely to have the largest radius? (Ch. 11)
a. H
d. Rb
b. Mn
e. Ag
c. Cl
67. Which property is not a characteristic of the alkali metals? (Ch. 11)
a. Their chlorides are water soluble.
b. They react with water to give off hydrogen gas.
c. They can be found as the free element in nature.
d. They are all reactive, readily losing one electron to form ions with a +1 charge.
e. They all have the outer electron configuration ns1.
68. How many valence electrons are in the outer shell of a Pb atom? (Ch. 11)
a. 2
d. 5
b. 3
e. 6
c. 4
69. Atmospheric pressure is most often measured with a (Ch. 13)
a. manometer
d. hydrometer
b. barometer
e. spectrophotometer
c. conductivity apparatus
70. A given mass of gas in a rigid container is heated from 200 C to 600 C. Which of the following responses best describes what
will happen to the pressure of the gas? (Ch. 13)
a. P will stay the same
b. P will decrease by a factor of 3
c. P will increase by a factor of 3
d. P will decrease by a factor of less than 3
e. P will increase by a factor of less than 3
71. For the reaction
2 C2H6(g) + O2(g)
 4 CO2(g) + 6 H2O(g)
what total volume of products would be formed from 7 liters of C 2H6 and excess oxygen at STP? (Ch. 9)
a. 10 L
b. 20 L
c. 25 L
d. 30 L
e. 35 L
72. What is the density of nitrogen gas at 227 C and 5.00 atm pressure? [HINT: use PV=nRT to calculate the volume of one mole
of diatomic nitrogen gas. Density is mass divided by volume, so you need to divide the mass of one mole of the gas by its
volume] (Ch. 13)
a. 2.93 g/L
d. 3.41 g/L
b. 0.293 g/L
e. 1.25 g/L
c. 2.30 g/L
73. A gas has a density of 1.45 g/L at 741 mm Hg and 26.0 C. What is the molar mass of the gas? (Use PV=nRT and a volume of
1 L to calculate moles of the gas.). (Ch. 13)
a. 36.5 grams/mole
d. 44.0 grams/mole
b. 48.0 grams/mole
e. cannot be determined
c. 57.6 grams/mole
7
74. A container with a volume of 10.0 L contains 2.80g nitrogen gas, 0.403 g of hydrogen gas, and 79.9 g argon gas. Calculate the
total pressure inside the container at 25 ºC. (Ch. 13)
a. 0.471 atm
d. 5.62 atm
b. 6.43 atm
e. 2.38 atm
c. 3.20 atm
75. When the following equation is balanced
Al2(CO3)3 (s)
+
HCl(aq) 
The sum of the coefficients is (Ch. 7)
a. 7
d. 30
b. 10
e. none of these
c. 15
AlCl3
+
CO2
+ H2O
76. An element E has the electron configuration [Kr]4d 105s25p2. The formula for the fluoride of E is most likely (Ch. 11)
a. EF14
d. EF4
b. EF8
e. EF
c. EF6
77. Of the following elements, which needs three electrons to complete its ns2np6 valence shell? (Ch. 11)
a. Ba
d. Ca
b. Si
e. P
c. Cl
78. Which of the following groups of compounds contains no ionic compounds? (Ch. 12)
a. HCN
NO2
Ca(NO3)2
b. PCl5
LiBr
Zn(OH)2
c. KOH
CCl4
SF4
d. NaH
CaF2
NaNH2
e. H2O
H2S
NH3
79. A reaction has an equilibrium constant = 4.0 x 10 8 at 25º. When the system is at equilibrium, which of these represents the
“favored” situation? (Ch. 17)
a. Product concentration will be favored.
b. Reactant concentration will be favored.
c. Product concentration will equal reactant concentration. Neither is favored.
d.. Cannot be determined with this information.
80. Aluminum has a density of 2.7 g/mL. How much volume will be occupied by 4.0 moles of aluminum? (Ch. 6)
a. 10. mL
d. 40. mL
b. 20. mL
e. 80. mL
c. 30. mL
81. What volume is occupied by 4.00 grams of carbon dioxide gas at a pressure of 0.976 atm and a temperature of 25.0 C? (Ch.
13)
a. 0.191 L
d. 22.8 L
b. 19.1 L
e. 2.03 L
c. 2.28 L
82. What is the answer to the correct number of significant figures? (Ch. 5)
(0.0821)(0.023)(298)
1.5
a. 0.38
b. 0.4
c. 0.3751
d. 0.375
e. 0.37514227
83. When ammonium chloride is heated, two gases are produced.
NH4Cl(s)  NH3(g) + HCl(g)
What is the total volume of gas produced at 793 K and 2.50 atm pressure when 164 grams of ammonium chloride decomposes
completely? (Ch. 13)
a. 160 L
d. 173.4 L
b. 27.3 L
e. 54.6 L
c. 63.8 L
8
84. According to collision theory, which one of the factors listed below will change the rate of a reaction? (Ch. 17)
a. The number and strength of the chemical bonds that must be broken in the reaction.
b. The number of collisions with enough energy to break bonds per unit of time.
c. The kinetic energy of the colliding molecules.
d. The relative amount of surface area in contact between reactants (increased by stirring or grinding).
e. All of these factors affect the rate of a reaction.
85. Which electron configuration is impossible? (Ch. 11)
a. 1s22s22p63s2
b. 1s22s22p62d2
c. 1s22s22p63s23p6
d. 1s22s22p63s1
e. 1s22s22p63s23p64s1
86. Which is a chemical property of chlorine? (Ch. 2)
a. It is yellowish green.
b. It burns in sodium vapor.
c. It has a density of 3.2 g/L at STP.
d. It dissolves in carbon tetrachloride.
e. It boils at -34 C.
87. Which pair is geometrically similar? (Ch. 12)
a. SO2 and CO2
b. PH3 and BF3
c. CO2 and OF2
d. SO2 and O3
e. CO32- and NH3
88. A metal M, forms an oxide of formula M2O3. The ground state valence shell electron configuration of the M atom is (Ch. 11):
a. ns2np1
d. ns2
1
10
b. 4s 3d
e. ns1
2
3
c. ns np
89. Which ion is smallest in size? (Ch. 11)
a. Al3+
b. Fc. O2-
d. Na+
e. N3-
90. When a metal is heated in a flame, the flame has a distinctive color. This information was eventually extended to the study of
stars because (Ch. 11)
a. the color spectra of stars indicate which elements are present.
b. a red shift in star color indicates stars are moving away.
c. star color indicates absolute distance.
d. it allows the observer to determine the size of stars.
91. How many total moles of aluminum and sulfate ions (added together) are there per mole of Al2(SO4)3? (Ch. 6)
a. 1
d. 4
b. 2
e. 5
c. 3
92. The molecule CCl2F2 is expected to have what kind of geometry? (Ch. 12)
a. bent
d. trigonal planar
b. tetrahedral
e. trigonal pyramidal
c. linear
93. An atom has in its nucleus 47 protons and 60 neutrons. What is the isotope symbol for this element? (Ch. 3)
a. 60Ag
b. 107Nd
c. 13Al
d. 107Ag
9
94. If the following pairs of elements form ions, which pair would form an ionic compound of the general formula C2A (where C =
cation and A = anion)? (Ch. 12)
a. sodium and sulfur
b. aluminum and oxygen
c. calcium and bromine
d. magnesium and sulfur
95. Reading from left to right across a row of the periodic table, the atomic radius becomes (Ch. 11):
a. larger, owing to increased nuclear charge that is only partially shielded by additional valence electrons.
b. smaller, owing to increased nuclear charge that is only partially shielded by additional valence electrons.
c. smaller, owing to placement of additional electrons in closer orbitals.
d. larger, owing to placement of additional electrons in closer orbitals.
96.
Which of the following shows atoms in expected order of increasing ionization energy? (Ch. 11)
I. Be, Mg, Ca
II. B, C, N
III. K, Na, Li
a. I only
b. II only
c. III only
d. II and III only
97. A gas mixture is made up of 0.18 moles of hydrogen gas and 0.72 moles of helium gas. If the total mixture is 4.40 atm, what is
the partial pressure of the helium? (Ch. 13)
a. 0.16 atm
b. 1.60 atm
c. 3.5 atm
d. 0.88 atm
98. A solution containing dissolved solute is placed between the electrodes of a conductivity apparatus. The bulb is dimly lit,
showing low conductivity. The best explanation for this result is that (Ch. 15)
a. the solution is either a very dilute solution of a strong electrolyte, or the solute is a weak electrolyte.
b. the solution must be very dilute.
c. the solution must contain a weak electrolyte.
d. the solute is a non-electrolyte.
e. the solution must not be saturated.
99. How many grams of alcohol are contained in 467 grams of a mixture of alcohol and water that is 23.0 % alcohol by mass?
(Ch. 5)
a. 20.30 grams
b. 23.0 grams
c. 359.6 grams
d. 107.4 grams
100.
How many atoms of hydrogen are represented in the following formula? (Ch. 6)
(NH4)2CO3
a. 4
b. 8
c. 24 x 1023
d. 48 x 1023
In an experiment similar to your silver - copper Lab, a student suspended a silver (Ag) wire in a
solution of gold nitrate, (Au(NO3)3, until the reaction was complete. By the end of the experiment, the
mass of the silver wire had decreased by 2.70 +/- 0.02 g, and 1.64 +/- 0.02 g of goldhad precipitated.
(Atomic wts: Ag=108, Au=197). Use this information to answer the next 3 questions.
101. How many moles of silver is reacted? (Ch. 6)
a. 0.0250 mole
d.
b. 0.00250 mole
e.
c. 2.70 mole
0.0400 mole
0.400 mole
102. How many moles of gold precipitated? (Ch. 6)
a. 0.0832 mole
d.
b. 0.00832 mole
e.
c. 12.0 moles
1.64 moles
1.20 moles
10
103. Use mole ratios to determine the most likely equation for the reaction (Ch. 9):
a. 3 Ag + Au(NO3)3  Au + 3 AgNO3
d. Ag + Au(NO3)3  Au + Ag(NO3)3
b. Ag + 3 AuNO3  AgNO3 + 3 Au
e. 2 Ag + Au(NO3)3  Au + 2 AgNO3
c. AgNO3 + Au  AuNO3 + Ag
104.
A cube measures 3.00 cm on an edge and has a mass of 16.23 g. The correct density is (Ch. 5):
a. 5.41 g/cm3
c. 0.601 g/cm3
3
b. 2.71 g/cm
d. 1.80 g/cm3
105. How many zeros in 0.00052010 are significant? (Ch. 5)
a. 1
c. 5
b. 2
d. 6
106.
How many atoms are there in 4.00 moles of sodium? (Ch. 6)
a. 6.02 x 1023 atoms
c. 6.02 x 1026 atoms
24
b. 1.20 x 10 atoms
d. 2.41 x 1024 atoms
107. The correct name for Cu2O is (Ch. 4)
a. copper oxide
b. copper (I) oxide
c.
d.
copper (II) oxide
dicopper monoxide
108. Which of the following has the largest radius? (Ch. 11)
a. S
d. K
b.
Cl-
c.
Ar
109.
a.
S
b.
Ca
c.
Na
e.
Ca
is the electron configuration for which one of the following ions? (Ch.11)
d. F
e.
none of these
110. How many lone pairs of electrons are in the Lewis structure for ammonia, NH 3 ? (Ch. 12)
a. 0
d. 3
b. 1
e. 4
c. 2
111. The electron dot structure of an atom of phosphorus is (Ch. 12):
a.
c.
b.
d.
112. The electron dot structure for water is (Ch. 12):
a.
c.
b.
d.
113. The electron dot structure for Cl is (Ch. 12):
a.
c.
b.
d.
114. How does a neutral atom of calcium (Ca) become a Ca2+ ion? (Ch. 3)
a. by gaining 2 protons
c. by losing 2 protons
b. by gaining 2 electrons
d. by losing 2 electrons
115. Acetylene gas, C2H2, is produced as a result of the following unbalanced equation. What mass of acetylene gas is produced
from the reaction of 10.0 grams of calcium carbide, CaC2? (Ch. 9)
CaC2 (s) +
H2O (l)  C2H2 (g) + Ca(OH)2 (s)
a. 24.6
d. 2.03
b. 10.0
e. 0.156
c. 4.06
11
116. How many grams of K2SO4 (molar mass= 174g/mol) would be needed to prepare 4.00 L of a 0.0510 M solution? (Ch. 15)
a. 17.8g
d. 43.5g
b. 35.5g
e. 63.8g
c. 71.0g
117. A solution is prepared by dissolving 200.1 g of NaOH (molar mass = 40.0 g/mol) in enough water to make 851 mL of
solution. Calculate the molarity of the solution. (Ch. 15)
a. 11.8M
d. 3.65 M
b. 4.78 M
e. 4.25 M
c. 5.88 M
118. A solution is prepared by dissolving 7.31g of Na 2SO4 in enough water to make 225 mL of solution. Calculate the solution
molarity (Ch. 15)
a. 30.8M
d. 0.229M
b. 1.64M
e. 3.11M
c. 0.136M
119. What are the concentrations of Na+ ions and S-2 ions in a 0.209 M Na2S solution? (Ch. 15)
a. Na+ = 0.209 M
S-2 = 0.1045 M
d. Na+ = 0.1045 M S-2 = 0.1045 M
+
-2
b. Na = 0.1045 M S = 0.209 M
e. Na+ = 0.418M S-2 = 0.209 M
+
-2
c. Na = 0.209 M S = 0.209 M
120. Which one of the following statements about the nuclear model of the atom is false? (Ch. 3)
a. The protons and neutrons in the nucleus are
d. The atom has no definite boundary.
very tightly packed.
b. The electrons occupy a very large volume
e. Almost all the mass of the atom is
compared to the nucleus.
concentrated in the nucleus.
c. The number of protons and neutrons are
always the same in a neutral atom.
121. If you disregard energy considerations, which one of the following reactions is not possible? (Ch. 3)
a. Na(g)  Na+(g) + electron
d. O2-(g) + 2 electrons  O(g)
+
2+
b. Ca (g)  Ca (g) + electron
e. Cl (g) + electron  Cl- (g)
3+
c. Al (g) )  Al (g) + 3 electrons
122. What volume of 18.0 M sulfuric acid is required to prepare 16.5 L of 0.126 M H2SO4? (Ch. 15)
a. 11.6 mL
d. 0.264 L
b. 0.116 L
e. 1.16 L
c. 232 mL
123. Adding a catalyst generally speeds up the rate of a reaction by (Ch. 17)
a. raising the activation energy
c. raising the heat of the reaction
b. lowering the activation energy
d. lowering the heat of the reaction
124. If K= 4.5 x 10-11 for a reaction, which of the following would be true? (Ch. 17)
a. reactants are favored at equilibrium
b. products are favored at equilibrium
c. reactants and products are equally favored at equilibrium
d. the reaction proceeds very rapidly
125. For the system SO2 (g) + Cl2 (g)  SO2Cl2 (g) at equilibrium, adding Cl2 (g) to the reaction container will have which of the
following effects? (Ch. 17)
a. the reaction will shift to the right
c. the concentration of SO2(g) will decrease
b. the concentration of SO2Cl2(g) will increase
d. all of the above will occur
126. If the equilibrium concentrations of A(g) + 2 B(g)  C(g) are found to be [A] = 0.015 M, [B] = 2.0 x 10-4 M, and [C] = 3.0 x
10-9 M, the equilibrium constant would be (Ch. 17):
a. 5.0
c. 1.0 x 10-3
b. 0.20
d. 9.0 x 10-15
127. For the endothermic reaction A(g) + 2 B(g)  C(g) , which of the following will shift the reaction towards the products?
(Ch. 17)
a. lowering the temperature
c. adding more C(g)
b. decreasing the volume
d. removing some A(g)
12
128. The quantum mechanical model of atoms explains or enables predictions of all except one of the following characteristics.
Identify the exception. (Ch. 11)
a. The probability of an electron being at a given d. The specific energy levels the electrons can
location at any instant.
occupy.
b. The general shape of the electron orbitals
e. The frequencies of light absorbed or emitted
by gaseous atoms.
c. The path or trajectory of the electrons.
129. Calculate the total amount of energy required to change 25.0 g of ice at 0 C into steam at 100 C. (Hfus = 6.02 kJ/mole,
Hvap = 40.6 kJ/mole) (Ch. 14 )
a. 581 J
d. 56, 326 J
b. 10,460 J
e. 75, 210 J
c. 8,931 J
130. Given the reaction A(g) + B(g)  C(g) + D(g). You have the gases A, B, C, and D at equilibrium. Upon adding gas A, the
value of Keq (Ch. 17)
a. increases because by adding A, more products are made, increasing the product-reactant ratio
b. decreases because A is a reactant so the product-to-reactant ratio decreases
c. does not change because A does not figure into the product-to-reactant ratio
d. does not change as long as the temperature is constant
e. depends on whether the reaction is endothermic or exothermic
131. The sharing of electrons in bond formation always involves (Ch. 12)
a. the formation of positive and negative ions
d. Shared electrons being attracted more by one
atom than another.
b. Formation of polar molecules.
e. Lower energy content for the bonded than for
the unbonded atoms.
c. Shared electrons being located equally near the
nuclei involved.
132. In which reaction is entropy expected to be increasing? (Ch. 10)
a. I2(g)  I2(s)
c. 2O2(g) + 2 SO(g)  SO3(g)
b. H2O(l)  H2O(s)
d. none of these
133. Which of the following shows an increase in entropy? 10)
a. Br2(g)  Br2(l)
c. H2(g) Cl2(g)  2 HCl(g)
b. NaBr(s)  Na+(aq) + Br-(aq)
d. 2 NO2(g)  N2O4(g)
134. Which of the following statements would not be included in our kinetic molecular theory of gases? (Ch. 13)
a. Molecules are in constant motion at all
c. The temperature is a measure of the potential
temperatures above absolute zero.
energy of the molecules.
b. Kinetic energy is conserved during molecular d. The mass of a gas is the sum of all the masses
collisions.
of the individual molecules.
135. In neutral atoms of two different isotopes of the same element, which one of the following properties is different
in the two isotopes? (Ch. 3)
a. Atomic number
d. General chemical reactions
b. Number of electrons
e. Number of protons
c. Mass
136. A solution is saturated when the maximum amount of solute is dissolved in the solvent. If the solubility of NaCl at
25 ºC is 36.2 g/100 g H2O, what mass of NaCl can be dissolved in 50.0 g of H 2O? (Ch. 5)
a. 18.1 g
c. 72.4 g
b. 36.2 g
d. 86.2 g
137. A chemical reaction is most likely to be spontaneous if (Ch. 10)
a. the potential energy of the system is lowered c. potential energy of the system is lowered and
and entropy increases.
entropy decreases.
b. potential energy of the system is increased
d. potential energy of the system is increased and
and entropy increases.
entropy decreases.
138. A box measures 3.50 cm x 2.915 cm. The product of these numbers = 10.2025 cm2. What is the proper way to report the
area of the box? (Ch. 5)
a. 10.20 cm2
c. 10 cm2
b. 10.2 cm2
d. 10. cm2
13
139. The result of 2.350 x (4.0 + 6.311) is, (Ch. 5)
a. 24
c.
24.21
b.
24.205
24.2
d.
140. The term that refers to the reproducibility of a laboratory measurement is (Ch. 5)
a. precision
c. repeatability
b. accuracy
d. exactness
141. The marks on the following target represent someone who is: (Ch. 5)
a.
b.
accurate, but not precise.
precise, but not accurate.
c.
d.
both accurate and precise.
neither accurate nor precise.
142. How many 1 mg salt tablets can be made from 1 kg of salt? (Ch. 5)
a. 1000
d. 1,000,000
b. 10,000
e. 10,000,000
c. 100,000
143. The number of significant figures in 30500 is (Ch. 5)
a. 1
d. 4
b. 2
e. 5
c. 3
144. The number of cubic centimeters (cm3) in 43.0 mL is (Ch. 5)
a. 43.0 cm3
c. 0.0403 cm3
3
b. 4.30 cm
d. cannot convert between cm3 and mL
145. How many significant figures should be in the measurement of the piece of wood below? (Ch. 5)
1
a.
b.
3
2
2
3
4
cm
c.
d.
1
0
146. The volume of 400 mL of chlorine gas at 400 mm Hg is decreased to 200 mL at constant temperature. What is the new
gas pressure? (Ch. 13)
a. 400 mm Hg
d. 600 mm Hg
b. 300 mm Hg
e. 650 mm Hg
c. 800 mm Hg
147. What is the equivalent of 423 Kelvin in degrees Celsius? (Ch. 13)
a. –223 ºC
d. 696 ºC
b. –23 ºC
e. 423 ºC
c. 150 ºC
148. Order these gases in order of increasing velocity: (Ch. 13)
F2, Cl2, NO, NO2, CH4
a. Cl2 < NO2 < F2 < NO < CH4
d. CH4 < NO < F2 < NO2 < Cl2
b. Cl2 < F2 < NO2 < CH4 < NO
e. F2 < NO < Cl2 < NO2 < CH4
c. CH4 < NO2 < NO < F2 < Cl2
14
149. Titanium (IV) oxide has the formula (Ch. 4)
a. Ti4O
b. TiO4
150. The correct name for P2O5 is (Ch. 4)
a. phosphorus(II) oxide
b. phosphorus(V) oxide
c. diphosphorus oxide
c.
d.
Ti(IV)O
TiO2
d.
e.
diphosphorus pentoxide
phosporus pentoxide
151. Calculate the pH of a solution with [OH-] = 1.0 x 10-9 M. (Ch. 16)
a. 9
d. 1
b. 5
e. Not enough information
c. 14
152. Is the solution in #151 acidic, basic, or neutral? (Ch. 16)
a. acidic
c. neutral
b. basic
d. Not enough information
15