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					       Sample Multiple Choice Questions for the Chemistry Final Exam 2010
 1.   One chemical property of matter is
         a) boiling point.                       b) texture.
         c) reactivity.                          d) density.

2.    The state of matter in which a material has definite shape and definite volume is the
          a) liquid state.                       b) solid state.
          c) gaseous state.                      d) vaporous state.

 3.   Under ordinary conditions of temperature and pressure, the particles in a gas are
         a) closely packed.                    b) very far from each other.
         c) held in fixed positions.           d) able to slide past each other.

 4.   A horizontal row of elements in the periodic table is called a(n)
         a) group.                              b) period.
         c) family.                             d) octet.

 5.   Elements in a group in the periodic table can be expected to have similar
         a) atomic masses.                       b) atomic numbers.
         c) numbers of neutrons.                 d) properties.

 6.   The symbol that represents the measured unit for volume is
         a) mL.                                b) mg.
         c) mm.                                d) cm.

 7.   A measure of the quantity of matter is
         a) density.                              b) weight.
         c) volume.                               d) mass.

 8.   To determine density, the quantities that must be measured are
         a) mass and weight.                     b) volume and weight.
         c) volume and concentration             d) volume and mass.

9.    The density of aluminum is 2.70 g/m3. The volume of a solid piece of aluminum is 1.50 cm3.
      Find its mass.
            a) 1.50 g                         b) 1.80 g
            c) 2.70 g                         d) 4.05 g

10.   1.06 L of water is equivalent to
            a) 0.00106 mL.                     b) 10.6 mL.
            c) 106 mL.                         d) 1060 mL.

 11. Convert -25'C to the kelvin scale.
          a) -323.15 K                         b) -248.15 K
          c) 248.15 K                          d) 323.15 K

 12. In oxides of nitrogen, such as N2O, NO, NO2, and N2O3, atoms combine in small whole-number
     ratios. This evidence supports the law of
            a) conservation of mass.       b) multiple proportion.
            c) definite composition.       d) mass action.
13.   Who was the schoolmaster who studied chemistry and proposed an atomic theory?
          a) John Dalton              b) Jons Berzehus
          c) Robert Brown             d) Dmitri Mendeleev

14.   According to Dalton’s atomic theory, atoms
           a) of different elements         b) can be divided into
           combine in simple                protons, neutrons, and
           whole-number ratios to form      electrons.
           c) of all elements are identical d) can be destroyed in
           in size and mass.                chemical reactions.

15.   The discovery of the electron resulted from experiments using
           a) gold foil.                   b) cathode rays.
           c) neutrons.                    d) alpha particles.

16.   Who discovered the nucleus by bombarding gold foil with positively charged particles and
      noting that some particles were widely deflected?
            a) Rutherford                 b) Dalton
             c) Chadwick                  d) Bohr

17. Rutherford fired positively charged particles at metal foil and concluded that most of the mass
    of an atom was
           a) in the electrons.          b) concentrated in the nucleus.
           c) evenly spread throughout d) in rings around the atom.
           the atom

18. A nuclear particle that has about the same mass as a proton, but with no electrical charge, is
called a(n)
            a) nuclide.                     b) neutron.
            c) electron.                    d) isotope.

19. An atom is electrically neutral because
          a) neutrons balance the           b) nuclear forces stabilize the
          protons and electrons.            charges.
          c) the numbers of protons and d) the numbers of protons and
          electrons are equal               neutrons are equal.

20. Atoms of the same element that have different masses are called
          a) moles.                     b) isotopes.
          c) nuclides.                  d) neutrons.

21. All atoms of the same element have the same
           a) atomic mass.               b) number of neutrons.
           c) mass number.               d) atomic number.

22. An aluminum isotope consists of 13 protons, 13 electrons, and 14 neutrons. Its mass
number is
          a) 13.                        b) 14.
          c) 27.                        d) 40.
23. Carbon-14 (atomic number 6), the radioactive nuclide used in dating fossils, has
          a) 6 neutrons.                 b) 8 neutrons.
           c) 10 neutrons.              d) 14 neutrons.

24.The number of atoms in 1 mol of carbon is
          a) 6.022 x 1022.              b) 6.022 x 1023.
          c) 5.022 x 10 .               d) 5.022 x 1023.

25. Molar mass
           a) is the mass in gram of one b) is numerically equal to the
           mole of a substance.          average atomic mass of the
            c) both a and b              d) neither a nor b

26. A sample of tin (atomic mass 118.69 amu) contains 3.01 x 1023 atoms. The mass of the sample is
           a) 3.01                       b) 59.3 g.
           c) 72.6 g.                    d) 11 g.

27. A bright line spectrum of an atom is caused by the energy released when electrons
           a) jump to a higher energy      b) fall to a lower energy level.
            c) absorb energy and jump to d) absorb energy and fall to a
           a higher energy level.          lower energy level.

28. For an electron in an atom to change from the ground state to an excited state,
            a) energy must be released.     b) energy must be absorbed.
            c) radiation must be emitted     d) the electron must make a transition from a higher
                                             to a lower energy level.

29. The set of orbitals that are dumbbell-shaped and directed along the x, y, and z axes are
            a) d orbitals.                 b) p oibitals.
            c) f orbitals.                 d) s orbitals.

30. The letter designations for the first four sublevels with the number of electrons that can be
accommodated in each sublevel are
            a) s:1, p:3, d:10, and f:14.     b) s:1, p:3, d.5, and f 7.
             c) s:2, p:6, d.10, and f 14.    d) s:1, p:2, d: 3, and f 4.

31. Which of the following rules requires that each of the p orbitals at a particular energy level
receive one electron before any of them can have two electrons?
           a) Hund’s rule                 b) the Pauli exclusion principle
           c) the Aufbau principle        d) the quantum rule

32. What is the electron configuration for nitrogen, atomic number 7?
           a) 1s2 2s2 2p3                  b) 1s2 2s3 2p2
                  2   3   l
            C) 1s 2s 2p                    d) 1s2 2s2 2p2 3s1

33. Mendeleev predicted that the spaces in his periodic table represented
          a) isotopes.                    b) radioactive elements.
          c) permanent gaps.              d) undiscovered elements.
34. The discovery of the noble gases changed Mendeleev's periodic table by adding a new
           a) period.                    b) series.
           c) group .                    d) sublevel block.

35. In the modem periodic table, elements are ordered according to
            a) decreasing atomic mass.     b) Mendeleev's original design.
            c) increasing atomic number. d) the date of their discovery.

36. Krypton, atomic number 36, is the fourth element in Group 18. What is the atomic number of
xenon, the fifth element in Group 18?
            a) 54                        b) 68
            c) 72                        d) 90

37. The electron configuration of aluminum, atomic number 13, is [Ne] 3s2 3pl. Aluminum is in
            a) 2.                        b) 3.
            c) 6.                        d) 13.

38. Calcium, atomic number 20, has the electron configuration [Ar] 4s2. In what period is calcium?
           a) Period 2                   b) Period 4
           c) Period 8                   d) Period 20

39. The energy required to remove an electron from an atom is the atom’s
           a) electron affinity.         b) electron energy.
           c) electronegativity.         d) ionization energy.

40. A measure of the ability of an atom in a chemical compound to attract electrons is called
          a) electron affinity.            b) electron configuration
          c) electronegativity.            d) ionization potential.

41. Across a period in the periodic table, atomic radii
           a) gradually decrease.           b) gradually decrease, then
                                            sharply increase.
           c) gradually increase.           d) gradually increase, then
                                            sharply decrease.

42. The number of valence electrons in Group 17 elements is
          a) 7.                          b) 8.
          c) 17.                         d) equal to the period number.

43. The electrons involved in the formation of a chemical bond are called
           a) dipoles.                    b) s electrons.
           c) Lewis electrons.            d) valence electrons.

44. The chemical bond formed when two atoms share electrons is called a(n)
           a) ionic bond.             b) orbital bond.
           c) Lewis structure.        d) covalent bond.

45. If the atoms that share electrons have an uneqa1 attraction for the electrons, the bond is called
             a) nonpolar.                   b) polar.
              c) ionic.                     d) dipolar.
46. The electron configuration of nitrogen is 1s2 2s2 2p3. How many more electrons does nitrogen
need to satisfy the octet rule?
            a) 1                          b) 3
            c) 5                          d) 8

47. A formula that shows the types and numbers of atoms combined in a single molecule is called
           a) molecular formula.         b) ionic formula.
           c) Lewis structure.           d) covalent formula.

48. VSEPR theory is a model for predicting
          a) the strength of metallic    b) the shape of molecules.
          c) lattice energy values.      d) ionization energy.

49. The following molecules contain polar bonds. The only polar molecule is
            a) CCl4.                      b) CO2.
            c) NH3                        d) CH4-

50. How many atoms of fluorine are present in a molecule of carbon tetrafluoride, CF4?
         a) 1                           b) 2
         c) 4                           d) 5

51. The formula for carbon dioxide, CO2, can represent
           a) one molecule of carbon     b) 1 mol of carbon dioxide
           dioxide.                      molecules.
           c) one molar mass of carbon d) all of the above. dioxide.

52. What is the formula for aluminum sulfate?
           a) AlSO4                      b) Al2SO4
           c) Al2(SO4)3                  d) Al(SO4)3

53. Name the compound KClO3.
          a) potassium chloride            b) potassium trioxychlorite
          c) potassium chlorate            d) hypochlorite

54. What is the metallic ion in copper(II) chloride?
           a) Co2+                          b) O2-
           c) Cu2+                          d) Cl-

55. Name the compound N2O4.
          a) sodium tetroxide              b) dinitrogen tetroxide
          c) nitrous oxide                 d) binitrogen oxide

56. What is the formula for sulfur dichloride?
           a) NaCl2                        b) SCl2
           c) S2Cl                         d) S2Cl2

57. The molar mass of MgI2 is
           a) the sum of the masses of     b) the sum of the masses of
           1 mol of Mg and 2 mol of 1.     1 mol of Mg and I mol of 1.
          c) the sum of the masses of      d) impossible to calculate.
           2 mol of Mg and 2 mol of 1.
58. The molar mass of NO2 is 46.01 g/mol. How many moles of NO2 are present in 114.95 g?
           a) 0.4003 mol                b) 1.000 mol
           c) 2.498 mol                d) 114.95 mol

59. What is the percentage composition of CF4?
           a) 20% C, 80% F               b) 13.6% C, 86.4% F
            c) 16.8% C, 83.2% F          d) 81% C, 19% F

60. What is the empirical formula for a compound that is 43.6% phosphorus and 56.4% oxygen?
           a) P3O7                        b) PO3
           c) P2O3                        d) P2O5

61. The molecular formula for vitamin C is C6H8O6. What is the empirical formula?
           a) CHO                        b) CH2O
           c) C3H4O3                     d) C2H4O2

62. A compound’s empirical formula is HO. If the formula mass is 34 amu, what is the molecular
          a) H2O                       b) H2O2
          c) HO3                       d) H2O3

63. In writing an equation that produces hydrogen gas, the correct representation of hydrogen gas is
             a) H.                        b) 2H.
             c) H2.                       d) OH.

64. Which equation is NOT balanced?
           a) 2H2 + O2  2H2O             b) 4H2 + 2O2  4H2O
           c) H2 + H2 + O2               d) 2H2 + O2  H2O
              H2O + H2O

65. The reaction Mg(s) + 2HCI(aq)  H2(g) + MgCl2(aq) is a
           a) composition reaction       b) decomposition reaction.
           c) single-replacement reaction. d) double-replacement reaction.

66. The reaction Pb(NO3)2(aq) + 2KI(aq)  PbI2(S) + 2KNO3(aq) is a
           a) double-replacement reaction. b) synthesis reaction.
           c) decomposition reaction.      d) combustion reaction.

67. What is the balanced equation when aluminum reacts with copper(II) sulfate?
           a) Al + Cu2S A12S + Cu
           b) 2Al + 3CuSO4  Al2(SO4)3 + 3Cu
           c) Al + CuSO4  AlSO4 + Cu
           d) 2Al + Cu2SO4 -4 A12SO4 + 2Cu

68. In the reaction N2 + 3H22NH3, what is the mole ratio of nitrogen to ammonia?
             a) 1:1                    b) 1:2
             c) 1:3                    d) 2:3

69. In the equation 2KClO3  2KCI + 3O2, how many moles of oxygen are produced when 3.0 mol
of KClO3 decompose completely?
            a) 1.0 mol                b) 2.5 mol
             c) 3.0 mol               d) 4.5 mol
70. For the reaction 2HNO3 + Mg(OH)2  Mg(NO3)2 + 2H2O, how many grams of magnesium
nitrate are produced from 8.00 mol of nitric acid, HNO3?
             a) 148 g                     b) 445 g
              c) 592 g                     d) 818 g

71. For the reaction 3Fe + 4H2O  Fe3O4 + 4H2, how many moles of iron oxide are produced from
500 g of iron?
            a) 1 mol                   b) 3 mol
             c) 9 mol                  d) 12 mol

72. For the reaction SO3 + H2O  H2SO4, calculate the percent yield if 500. g of sulfur trioxide
react with excess water to produce 575 g of sulfuric acid.
            a) 82.7%                      b) 88.3%
             c) 91.2%                     d) 93.9%

73. An ideal gas is an imaginary gas
           a) not made of particles.    b) that conforms to all of the
                                        assumptions of the kinetic theory.
           c) whose particles have zero d) made of motionless
           mass.                        particles.

74. Which is an example of gas diffusion?
             a) inflating a flat tire     b) the odor of perfume spreading throughout a
           c) a cylinder of oxygen stored d) All of the above under high pressure

75. According to the kinetic-molecular theory, bow does a gas expand?
          a) Its particles become larger. b) Collisions between particles become elastic.
          c) Its temperature rises.        d) Its particles move greater

76. Which is an example of effusion?
          a) air slowly escaping from a   b) the aroma of a cooling pie
          pinhole in a tire               spreading across a room
          c) helium dispersing into a     d) oxygen and gasoline fumes
          room after a balloon pops       mixing in an automobile

77. What does the constant bombardment of gas molecules against the inside walls of a container
          a) temperature               b) density
          c) pressure                  d) diffusion

78. Convert the pressure 0.75 atm to mm Hg.
           a) 101.325 mm Hg             b) 430 mm Hg
           c) 570 mm Hg                 d) 760 mm Hg

79. Standard temperature is exactly
           a) 100 C.                      b) 273 C.
           c) 0 C.                        d) 0 K.
80. Standard pressure is exactly
           a) 1 atm                       b) 760 atm.
           c) 101.325 atm                 d) 101 atm.

81. Pressure and volume changes at a constant temperature can be calculated using
           a) Boyle's law.              b) Charles's law.
           c) Kelvin’s law.             d) Dalton's law.

83. The volume of a gas is 5.0 L when the temperature is 5.0 C. If the temperature is increased to
10.0 C without changing the pressure, what is the new volume?
           a) 2.5 L                     b) 4.8 L
           c) 5.1 L                     d) 10.0 L

84. A 150.0 L sample of gas is collected at 1.20 atm and 25 C. What volume does the gas have at
1.50 atm and 20.0 C?
           a) 94 L                        b) 120 L
           c) 143 L                       d) 183 L

85. In the equation H2(g) + Cl2(g)  2HCI(g), one volume of hydrogen yields how many volumes
of hydrogen chloride?
            a) 1                        b) 2
            c) 3                        d) 4

86. At constant temperature and pressure, gas volume is directly proportional to the
           a) molar mass of the gas.     b) number of moles of gas.
           c) density of the gas at STP. d) rate of diffusion.

87. According to Avogadro's law, 1 L of H2(g) and 1 L of O2(g) at the same temperature and
          a) have the same mass.        b) have unequal volumes.
          c) contain 1 mol of gas each. d) contain equal numbers of molecules.

88. The standard molar volume of a gas at STP is
           a) 22.4 L.                   b) g/22.4 L.
           c) g-mol wt/22.4 L.          d) 1 L.

89. A 1.00 L sample of a gas has a mass of 1.25 g at STP. What is the mass of 1 mol of this gas?
           a) a little less then 1.0 g     b) 1.25 g
            c) 22.4 g                     d) 28.0 g

90. According to Graham’s law, two gases at the same temperature and pressure will have different
rates of diffusion because they have different
             a) volumes.                   b) molar masses.
             c) kinetic energies.          d) condensation points.

91. Why are water molecules polar?
          a) They contain two kinds of b) The electrons in the covalent
          atoms.                       bonds spend more time closer
                                       to the oxygen nucleus.
          c) The hydrogen bonds are    d) They have covalent bonds.
92. What is the freezing point of water at standard pressure?
           a) -10 C                        b) 0 C
           c) 4'C                          d) 32 C

93. What is the boiling point of water at standard pressure?
           a) 100 C                        b) 112 C
           c) 212 C                        d) 200 C

94. Which of the following is a pure substance?
          a) water                       b) milk
          c) soil                        d) concrete

95. Sugar in water is an example of which solute-solvent combination?
           a) gas-liquid                b) liquid-liquid
           c) solid-liquid              d) liquid-solid

96. To conduct electricity, a solution must contain
            a) nonpolar molecules.          b) polar molecules.
            c) ions.                        d) free electrons.

97. If the amount of solute present in a solution at a given temperature is less than the maximum
amount that can dissolve at that temperature, the solution is said to be
            a) saturated.                  b) unsaturated.
            c) supersaturated.             d) concentrated.

98. Which of the following is likely to produce crystals if disturbed?
          a) an unsaturated solution       b) a supersaturated solution
          c) a saturated solution          d) a concentrated solution

99. What is the molarity of a solution that contains 0.202 mol KC1 in 7.98 L solution?
          a) 0.0132 M                      b) 0.0253 M
           c) 0.459 M                      d) 1.363 M

100. How many moles of HCI are present in 0.70 L of a 0.33 M HCI solution?
          a) 0.23 mol                 b) 0.28 mol
          c) 0.38 mol                 d) 0.47 mol

101. An NaOH solution contains 1.90 mol of NaOH, and its concentration is 0.555 M. What is its
          a) 0.623 L                    b) 0.911 L
          c) 1.05 L                     d) 3.42 L

102. How many milliliters water are needed to make a 0.171 M solution that contains 1.00 g of
          a) 100 mL                      b) 1000 mL
          c) 171 mL                      d) 17.1 mL

103. How many moles of ions are produced by the dissociation of 1 mol of MgCl2?
         a) 0 mol                      b) 1 mol
          c) 2 mol                     d) 3 mol
104. Colligative properties depend on
           a) the identity of the solute   b) the concentration of the
          particles.                       solute particles.
          c) the physical properties of    d) the boiling point and
          the solute particles.            freezing point of the solution.

105. Compared with a 0.01 m sugar solution, a 0.01 m KCI solution has
          a) the same freezing-point    b) about twice the
          depression.                   freezing-point depression.
          c) the same freezing-point    d) about six times the
          elevation.                    freezing-point elevation.

106. Electrolytes affect colligative properties differently than do nonelectrolytes because
            a) are volatile.                b) have lower boiling points.
             c) produce fewer moles of      d) produce more moles of
             solute particles per mole of solute particles per mole of
             solute.                        solute.

107. Acids taste
           a) sweet.                       b) sour.
           c) bitter.                      d) salty.

108. Acids react with
           a) bases to produce salts and   b) salts to produce bases and
           water.                          water.
           c) water to produce bases and   d) neither bases, salts, nor
           salts.                          water.

109. Bases taste
           a) soapy.                       b) sour.
           c) sweet.                       d) bitter.

110. A binary acid contains
           a) two hydrogen atoms.          b) hydrogen and one other
            c) hydrogen and two other      d) hydrogen and three other
            elements.                      elements.

111. According to the traditional definition an acid contains
          a) hydrogen and does not           b) hydrogen and ionizes to
          ionize.                            form hydrogen ions.
          c) oxygen and ionizes to form d) oxygen and ionizes to form
           hydroxide ions.                  oxygen ions.

112. A substance that ionizes nearly completely in aqueous solutions and produces H3O+ is a
           a) weak base.                 b) strong base.
            c) weak acid.                d) strong acid.

113. A Bronsted-Lowry acid is
          a) an electron-pair acceptor.    b) an electron-pair donor.
          c) a proton acceptor.            d) a proton donor.
114. What is neutralization?
           a) an acid-base reaction that     b) a reaction of hydronium
           does not include dissocation      of ions and hydroxide ions to
           ions                              form a salt
           c) a reaction of hydronium        d) a reaction of hydronium
           ions and hydroxide ions to        ions and hydroxide ions to
           form water molecules              form water molecules and a salt

115. Pure water contains
           a) water molecules only.         b) hydronium ions only.
           c) hydroxide ions only.          d) water molecules, hydronium
                                            ions, and hydroxide ions.

116. What is the concentration of H3O+ in pure water?
          a) 10-7 M                      b) 0.7 M
          c) 55.4 M                       d) 107 M

117. Which expression represents the pH of a solution?
          a) log[H3O+1]                 b) -log[H3O+1]
          c) log[OH-]                   d) -log[OH-]

118. If [H3O+1] of a solution is less than [OH-1], the solution
           a) is always acidic.             b) is always basic.
           c) is always neutral.            d) might be acidic, basic, or neutral.

119. What is the pH of a neutral solution at 25 C?
           a) 0                            b) 1
           c) 7                            d) 14

120. The pH scale in general use ranges from
          a) 0 to 1.                      b) - 1 to 1.
          c) 0 to 7.                      d) 0 to 14.

121. The pH of an acidic solution is
          a) less than 0.                  b) less than 7.
          c) greater than 7.               d) greater than 14.

122. The pH of a basic solution is
          a) less than 0.                  b) less than 7.
          c) greater than 7.               d) greater than 14.

123. If [H3O+] = 1.7 x 10-3 M, what is the pH of the solution?
           a) 1.81                         b) 2.13
           c) 2.42                         d) 2.77

124. What is the pH of a 0.027 M KOH solution?
           a) 6.47                     b) 12.43
           c) 12.92                    d) 14.11

125. What is the hydronium ion concentration of a solution whose pH is 4.12?
          a) 4.4 x 10-8 M               b) 5.1 X 10-6 M
          c) 6.4 x 10-5 M               d) 7.6 x 10-5 M
126. Dyes with pH-sensitive colors are used as
          a) primary standards.          b) indicators.
          c) titrants.                   d) None of the above

127.In an acid-base titration, equivalent quantities of hydronium ions and hydroxide ions are present
           a) at the beginning point.      b) at the midpoint.
           c) at the endpoint.             d) throughout the titration.

128. What is the molarity of an HCL solution if 125 mL is neutralized in a titration by 76.0 mL of
1.22 M KOH?
           a) 0.371 M                    b) 0.455 M
            c) 0.617 M                  d) 0.742 M

129. What is the molarity of a Ba(OH)2 solution if 93.9 mL is completely titrated by 15.3 mL of
0.247 M H2SO4?
           a) 0.0101 M                   b) 0.0201 M
           c) 0.0402 M                   d) 0.0805 M

130. How is a Celsius temperature reading converted to a Kelvin temperature reading?
          a) by adding 273               b) by subtracting 273
          c) by dividing by 273          d) by multiplying by 273

131. The pH of a solution is 9. What is its H3O+ concentration?
           a) 10-9 M                       b) 10-7 M
           c) 10 M                          d) 9 M

132. After balancing the equation FeCl3 + Zn  ZnCI2 + Fe, the coefficients, in order
from left to right, are
            a) 2, 2, 1, 2.                b) 1, 1, 1, 1.
            c) 4, 3, 3, 4.                d) 2, 3, 3, 2.

133. Which of the following is the electron configuration of carbon in the ground state?
           a) 1s2 2s2 2p2                 b) 2s2 2sl 2p3
                2   2
          c) 1s 2s 2p3                    d) 1s2 2s2 2p6

134. How many valence electrons does a carbon atom have?
         a) 3                           b) 4
         c) 5                           d) 6

135. When a carbon atom forms four covalent bonds, the bonds are directed toward the comers of a
          a) triangle.                 b) pyramid.
          c) square.                   d) tetrahedron

136. Which formula is most useful in distinguishing isomers?
          a) molecular formula            b) structural formula
          c) empirical formula            d) ionic formula

137. Isomers are compounds that have
          a) the same molecular formula     b) the same molecular formula
          but different structure            and the same structure.
          c) different molecular formulas   d) different molecular formulas
          and different structures.         but the same structure.
 138. How many structural isomers does C4H10 have?
          a) one                        b) two
          c) three                      d) five

 139. Which hydrocarbons have double covalent bonds?
           a) alkanes                   b) alkenes
            c) alkynes                  d) aromatic hydrocarbons

140. Which of the following is the functional group in alcohols?
          a) -COOH                        b) -OH
          c) -CO                          d) -O-

 141. The systematic names of alcohols end in
            a) -al.                      b) -one.
            c) -ol.                      d) -oic.

 142. What is the general formula for ethers?
            a) R- O- R'                    b) R-COOH
            e) R-COO- R'                   d) R-CHO

 143. What is the systematic name for the two carbon alcohol?
            a) ethanol                     b) ethanal
            c) 2-propanol                  d) 2-propanal

 144. What is the systematic name of the compound CCl3F?
            a) trichlorofluoromethane     b) 3-chloro-l-fluoromethane
            c) chloromediylfluolide       d) chlorofluorocarbon

Name the following organic molecules

 145. CH3-CH2-CH-CH3

 146.         CH2 – CH2
              |     |
              CH2 – CH2

 147.         CH2 – CH2
              |      |
              F     Cl

 148.        CH3 – CH2 – CH – CH2 – CH3

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