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Midterm Review
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1. The state of matter in which a material has neither a definite shape nor a definite volume is the _____ state.
a. gaseous
c. elemental
b. liquid
d. solid
____
2. Under ordinary conditions of temperature and pressure, the particles in a gas are
a. closely packed.
c. held in fixed positions.
b. very far from each other.
d. unevenly distributed.
____
3. A chemical change occurs when
a. dissolved minerals solidify to form a crystal.
b. ethanol is purified through distillation.
c. salt deposits form from evaporated seawater.
d. a leaf changes color.
____
4. A physical change occurs when a
a. peach spoils.
b. copper bowl tarnishes.
c. bracelet turns your wrist green.
d. glue gun melts a glue stick.
____
5. Nitrogen monoxide and oxygen, both colorless gases, form a red-brown gas when mixed. Nitrogen monoxide
and oxygen are called the
a. products.
c. synthetics.
b. equilibria.
d. reactants.
____
6. Matter includes all of the following except
a. air.
b. light.
c. smoke.
d. water vapor.
7. A measure of the quantity of matter is
a. density.
b. weight.
c. volume.
d. mass.
____
____
8. A true statement about mass is that
a. mass is often measured with a spring scale.
b. mass is expressed in pounds.
c. as the force of Earth’s gravity on an object increases, the object’s mass increases.
d. mass is determined by comparing the mass of an object with a set of standard masses that
are part of a balance.
____
9. The liter is defined as
a. 1000 m3.
b. 1000 cm3.
____ 10. Quantitative observations are recorded using
a. numerical information.
b. a control.
c. 1000 g3.
d. 1000 c3.
c. non-numerical information.
d. a system.
____ 11. Which of the following does not describe a measurement standard?
a. Measurement standards avoid ambiguity.
b. Measurement standards must be unchanging.
c. A standard need not agree with a previously defined size.
d. Confusion is eliminated when the correct measurement is applied.
____ 12. Which of the following does not describe a unit?
a. A unit compares what is being measured with a previously defined size.
b. A unit is usually preceded by a number.
c. A unit is not needed to find a solution to a problem.
d. The choice of unit depends on the quantity being measured.
____ 13. All of the following are examples of units except
a. weight.
c. gram.
b. kilometer.
d. teaspoon.
____ 14. Which of the following is not an SI base unit?
a. kilogram
c. liter
b. second
d. kelvin
____ 15. The base unit for mass is the
a. gram.
b. cubic centimeter.
c. meter.
d. kilogram.
____ 16. The abbreviation that represents a volume unit is
a. mL.
c. mm.
b. mg.
d. cm.
____ 17. The SI base unit for time is the
a. day.
b. hour.
c. minute.
d. second.
____ 18. The metric unit for length that is closest to the thickness of a dime is the
a. micrometer.
c. centimeter.
b. millimeter.
d. decimeter.
____ 19. The abbreviation mm represents
a. micrometer.
b. millimeter.
c. milliliter.
d. meter.
____ 20. 1 cubic centimeter is equivalent to
a. 1 milliliter.
b. 1 gram.
c. 1 liter.
d. 10–1 cubic decimeters.
____ 21. The relationship between the mass m of a material, its volume V, and its density D is
a. V = mD.
c. DV = m.
b. Vm = D.
d. D + V = m.
____ 22. To calculate the density of an object,
a. multiply its mass and its volume.
b. divide its mass by its volume.
____ 23. When density is measured,
a. a balance is always used.
b. the units are always kg/m3.
c. the temperature should be specified.
c. divide its volume by its mass.
d. divide its mass by its area.
d. the mass and volume do not need to be measured.
____ 24. Electricity can convert oxygen into ozone. The ozone created by this process is
a. a new substance.
c. a different state of oxygen.
b. an isotope of oxygen.
d. an allotrope of oxygen.
____ 25. The only pure substance listed below is
a. bread dough.
b. vinegar (5% acetic acid).
c. vitamin C (ascorbic acid).
d. seawater.
____ 26. The homogeneous mixture in the illustration above is in container
a. a.
c. c.
b. b.
d. d.
____ 27. A chemical formula for a molecular compound represents the composition of
a. a molecule.
c. the ions that make up the compound.
b. an atom.
d. the crystal lattice.
____ 28. Which part of the illustration above shows the particles in a heterogeneous mixture?
a. a
c. c
b. b
d. d
____ 29. The energy that results from the breaking or formation of chemical bonds is
a. temperature.
c. chemical energy.
b. potential energy.
d. kinetic energy.
____ 30. “In any chemical or physical process, energy is neither created nor destroyed” is a statement of the
a. law of conservation of energy.
c. law of physical and chemical change.
b. law of conservation of mass.
d. system.
____ 31. The specific heat of a given substance
a. depends on the mass of the substance.
b. varies with temperature.
c. is the same as that for any other substance.
d. is unique to that substance.
____ 32. A temperature of 323 K is equivalent to
a. 50°C.
b. 1.2°C.
c. 50°C.
d. 596°C.
____ 33. The temperature that is equivalent to 20°C is
a. 253 K.
b. 293 K.
c. 293 K.
d. 13.7 K.
____ 34. For every investigation, the scientific method
a. is the same set of procedures.
b. is a logical set of procedures.
c. must be abandoned if there are unexpected results.
d. helps to predict what results will be obtained.
____ 35. The reason for organizing, analyzing, and classifying data is
a. so that computers can be used.
b. to prove a law.
c. to find relationships among the data.
d. to separate qualitative and quantitative data.
____ 36. There are _____ variables represented in the weather graphs.
a. four
c. six
b. five
d. seven
____ 37. Of the items depicted, only _____ is not a model.
a. figure a
c. figure c
b. figure b
d. figure d
____ 38. A plausible explanation of a body of observed natural phenomena is a scientific
a. principle.
c. law.
b. experiment.
d. theory.
____ 39. The matter diagram represents a(n)
a. law.
b. theory.
c. assumption.
d. hypothesis.
____ 40. A theory is an accepted explanation of an observed phenomenon until
a. one study conflicts with the theory.
b. repeated data and observation conflict with the theory.
c. scientists disagree about the research method used to gather data.
d. an eminent scientist feels that it is inadequate.
____ 41. Poor precision in scientific measurement may arise from
a. the standard being too strict.
b. human error.
c. limitations of the measuring instrument.
d. both human error and the limitations of the measuring instrument.
____ 42. Precision pertains to all of the following except
a. reproducibility of measurements.
b. agreement among numerical values.
c. sameness of measurements.
d. closeness of a measurement to an accepted value.
____ 43. Five darts strike near the center of the target. Whoever threw the darts is
a. accurate.
c. both accurate and precise.
b. precise.
d. neither accurate nor precise.
____ 44. A chemist who frequently carries out a complex experiment is likely to have high
a. accuracy, but low precision.
c. precision.
b. accuracy.
d. precision, but low accuracy.
____ 45. When applied to scientific measurements, the words accuracy and precision
a. are used interchangeably.
b. have limitations.
c. can cause uncertainty in experiments.
d. have distinctly different meanings.
____ 46. When determining the number of significant digits in a measurement,
a. all zeros are significant.
b. all nonzero digits are significant.
c. all zeros between two nonzero digits are not significant.
d. all nonzero digits are not significant.
____ 47. To two significant figures, the measurement 0.0255 g should be reported as
a. 0.02 g.
c. 0.026 g.
b. 0.025 g.
d. 2.5  102 g.
____ 48. In division and multiplication, the answer must not have more significant figures than the
a. number in the calculation with the fewest significant figures.
b. number in the calculation with the most significant figures.
c. average number of significant figures in the calculation.
d. total number of significant figures in the calculation.
____ 49. The number of significant figures in the measurement 170.040 km is
a. 3.
c. 5.
b. 4.
d. 6.
____ 50. The measurement that has been expressed to four significant figures is
a. 0.0020 mm.
c. 30.00 mm.
b. 0.004 02 mm.
d. 402.10 mm.
____ 51. What is the density of 37.72 g of matter whose volume is 6.80 cm3?
a. 0.18 g/cm3
c. 30.92 g/cm3
3
b. 5.55 g/cm
d. 256.4 g/cm3
____ 52. The dimensions of a rectangular solid are measured to be 1.27 cm, 1.3 cm, and 2.5 cm. The volume should be
recorded as
a. 4.128 cm3.
c. 4.13 cm3.
3
b. 4.12 cm .
d. 4.1 cm3.
____ 53. Three samples of 0.12 g, 1.8 g, and 0.562 g are mixed together. The combined mass of all three samples,
expressed to the correct number of significant figures, should be recorded as
a. 2.4 g.
c. 2.482 g.
b. 2.48 g.
d. 2.5 g.
____ 54. Expressed in scientific notation, 0.0930 m is
a. 93  10–3 m.
b. 9.3  10–3 m.
c. 9.30  10–2 m.
d. 9.30  10–4 m.
____ 55. When 6.02  1023 is multiplied by 9.1  10–31, the product is
a. 5.5  10–8.
c. 5.5  10–7.
54
b. 5.5  10 .
d. 5.5  10–53.
____ 56. The law of conservation of mass follows from the concept that
a. atoms are indivisible.
b. atoms of different elements have different properties.
c. matter is composed of atoms.
d. atoms can be destroyed in chemical reactions.
____ 57. The composition of the two oxides of lead, PbO and PbO2, are explained by the
a. periodic law.
c. atomic law.
b. law of multiple proportions.
d. law of conservation of mass.
____ 58. Who first proposed an atomic theory based on scientific knowledge?
a. John Dalton
c. Robert Brown
b. Jons Berzelius
d. Dmitri Mendeleev
____ 59. According to Dalton’s atomic theory, atoms
a. are destroyed in chemical reactions.
b. can be divided.
c. of each element are identical in size, mass, and other properties.
d. of different elements cannot combine.
____ 60. Which of the following is NOT part of Dalton’s atomic theory?
a. Atoms cannot be divided, created, or destroyed.
b. The number of protons in an atom is its atomic number.
c. In chemical reactions, atoms are combined, separated, or rearranged.
d. All matter is composed of extremely small particles called atoms.
____ 61. The law of definite proportions
a. contradicted Dalton’s atomic theory.
b. was explained by Dalton’s atomic theory.
c. replaced the law of conservation of mass.
d. assumes that atoms of all elements are identical.
____ 62. In a cathode tube, electrical current passes from one electrode, the _____, to the oppositely charged electrode.
a. cathode
c. negatively charged electrode
b. anode
d. Both (a) and (c)
____ 63. Experiments with cathode rays led to the discovery of the
a. proton.
c. neutron.
b. nucleus.
d. electron.
____ 64. Who explained the behavior of positively charged particles being deflected from a metal foil as the nucleus?
a. Ernest Rutherford
c. James Chadwick
b. John Dalton
d. Niels Bohr
____ 65. In the gold foil experiment, most of the particles fired at the foil
a. bounced back.
c. were absorbed by the foil.
b. passed through the foil.
d. combined with the foil.
____ 66. The gold foil experiment led to the discovery of the
a. electron.
c. nucleus.
b. cathode ray.
d. neutron.
____ 67. What did Rutherford conclude about the structure of the atom?
a. An atom is indivisible.
b. Electrons make up the center of an atom.
c. An atom carries a positive charge.
d. An atom contains a small, dense, positively charged central region.
____ 68. A nuclear particle that has about the same mass as a proton, but with no electrical charge, is called a(n)
a. nuclide.
c. electron.
b. neutron.
d. isotope.
____ 69. Which part of an atom has a mass approximately equal to
of the mass of a common hydrogen atom?
a. nucleus
c. proton
b. electron
d. electron cloud
____ 70. The mass of a neutron is
a. about the same as that of a proton.
b. about the same as that of an electron.
c. double that of a proton.
d. double that of an electron.
____ 71. The nucleus of most atoms is composed of
a. tightly packed protons.
b. tightly packed neutrons.
c. tightly packed protons and neutrons.
d. loosely connected protons and electrons.
____ 72. Protons and neutrons strongly attract when they
a. are moving fast.
c. are at high energies.
b. are very close together.
d. have opposite charges.
____ 73. An aluminum isotope consists of 13 protons, 13 electrons, and 14 neutrons. Its mass number is
a. 13.
c. 27.
b. 14.
d. 40.
____ 74. Isotopes are atoms of the same element that have different
a. principal chemical properties.
c. numbers of protons.
b. masses.
d. numbers of electrons.
____ 75. Atoms of the same element that have different masses are called
a. moles.
c. nuclides.
b. isotopes.
d. neutrons.
____ 76. Isotopes of an element contain different numbers of
a. electrons.
c. neutrons.
b. protons.
d. nuclides.
____ 77. When the light from excited atoms of an element is passed through a prism, the distinct colored lines
produced are called
a. ground states.
c. line-emission spectra.
b. excited states.
d. electromagnetic spectra.
____ 78. Bohr’s theory helped explain why
a. electrons have negative charge.
b. most of the mass of the atom is in the nucleus.
c. excited hydrogen gas gives off certain colors of light.
d. atoms combine to form molecules.
____ 79. According to Bohr’s theory, an excited atom would
a. collapse.
c. produce line-emission spectra.
b. absorb photons.
d. radiate energy.
____ 80. If electrons in an atom have the lowest possible energies, the electrons are in their
a. ground states.
b. inert states.
c. excited states.
d. radiation-emitting states.
____ 81. 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.
____ 82. Most of the volume of an atom is occupied by the
a. nucleus.
c. electron cloud.
b. nuclides.
d. protons.
____ 83. The main energy levels of an atom are indicated by the
a. orbital quantum numbers.
c. spin quantum numbers.
b. magnetic quantum numbers.
d. principal quantum numbers.
____ 84. 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.
c. s: 2, p: 6, d: 10, and f: 14.
b. s: 1, p: 3, d: 5, and f: 7.
d. s: 1, p: 2, d: 3, and f: 4.
____ 85. The number of orbitals for the d sublevel is
a. 1.
b. 3.
c. 5.
d. 7.
____ 86. The statement that an electron occupies the lowest available energy orbital is
a. Hund’s rule.
c. Bohr’s law.
b. the aufbau principle.
d. the Pauli exclusion principle.
____ 87. “Orbitals of equal energy are each occupied by one electron before any is occupied by a second electron, and
all electrons in singly occupied orbitals must have the same spin” is a statement of
a. the Pauli exclusion principle.
c. the quantum effect.
b. the aufbau principle.
d. Hund’s rule.
____ 88. The statement that no more than two electrons in the same atom can occupy a single orbital is
a. the Pauli exclusion principle.
c. Bohr’s law.
b. Hund’s rule.
d. the aufbau principle.
____ 89. 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
c. the aufbau principle
b. the Pauli exclusion principle
d. the quantum rule
____ 90. Two electrons in the 1s orbital must have different spin quantum numbers to satisfy
a. Hund’s rule.
c. the Pauli exclusion principle.
b. the magnetic rule.
d. the aufbau principle.
____ 91. The sequence in which energy sublevels are filled is specified by
a. the Pauli exclusion principle.
c. Lyman’s series.
b. the orbital rule.
d. aufbau principle.
____ 92. Which is the ground-state electron configuration for
?
2
4
3
a. [Ar] 4s 3d
c. [Ar] 4s 3d3
b. [Ar] 4s13d5
d. [Ar] 4s43d2
____ 93. The carbon-12 atom is assigned a relative mass of exactly
a. 1 amu.
c. 12 amu.
b. 6 amu.
d. 100 amu.
____ 94. The abbreviation for atomic mass unit is
a. amu.
b. mu.
c. a.
d. µ.
____ 95. The idea of arranging the elements in a table according to their chemical and physical properties is attributed
to
a. Newlands.
c. Bohr.
b. Moseley.
d. Ramsay.
____ 96. From gaps in his table, Mendeleev predicted the existence of several elements and their
a. atomic numbers.
c. properties.
b. colors.
d. radioactivity.
____ 97. Mendeleev is credited with developing the first successful
a. periodic table.
b. method for determining atomic number.
c. test for radioactivity.
d. use of X rays.
____ 98. The principle that states that the physical and chemical properties of the elements are periodic functions of
their atomic numbers is
a. the periodic table.
c. the law of properties.
b. the periodic law.
d. Mendeleev’s law.
____ 99. The periodic law states that
a. no two electrons with the same spin can be found in the same place in an atom.
b. the physical and chemical properties of the elements are functions of their atomic
numbers.
c. electrons exhibit properties of both particles and waves.
d. the chemical properties of elements can be grouped according to periodicity, but physical
properties cannot.
____ 100. Elements in a group or column in the periodic table can be expected to have similar
a. atomic masses.
c. numbers of neutrons.
b. atomic numbers.
d. properties.
____ 101. The atomic number of sodium, the first element in Period 3, is 11. The atomic number of the second element
in this period is
a. 3.
c. 12.
b. 10.
d. 18.
____ 102. Elements in a group have similar
a. reactivities.
b. densities.
c. symbols.
d. electron configurations.
____ 103. An element that has an electron configuration of [He]2s22p3 is in Period _____ of the periodic table.
a. 1
c. 3
b. 2
d. 4
____ 104. An element that has an electron configuration of [Ne]3s23p3 is in Group _____ of the periodic table.
a. 2
c. 6
b. 3
d. 15
____ 105. To which group of the periodic table do lithium and potassium belong?
a. alkali metals
c. halogens
b. transition metals
d. noble gases
____ 106. To which group of the periodic table do fluorine and chlorine belong?
a. alkaline-earth metals
c. halogens
b. transition elements
d. actinides
____ 107. The outer electron configuration of an alkali metal has
a. 1 electron in the s orbital.
c. 1 electron in the p orbital.
b. 2 electrons in the s orbital.
d. 2 electrons in the p orbital.
____ 108. The elements in Group 17 are known by what name?
a. halogens
c. alkaline-earth metals
b. alkali metals
d. noble gases
____ 109. Transition elements are found in which of the following group(s) of the periodic table?
a. Group 1
c. Groups 3–12
b. Groups 1 and 2
d. Group 18
____ 110. An element that has four electrons in its outermost d orbitals and one electron in its outermost s orbital is a
member of what group in the periodic table?
a. Group 1
c. Group 5
b. Groups 4
d. Group 15
____ 111. Moving from left to right across _____, electrons are being added to the 4f orbitals.
a. the noble gases
c. the lanthanides
b. the halogens
d. Period 4
____ 112. A property of all actinides is that they are
a. members of Period 7.
b. radioactive.
c. nonmetals.
d. Both (a) and (b)
____ 113. A metal is called malleable if it
a. has a shiny appearance.
b. can be hammered into sheets.
c. can be squeezed out into a wire.
d. exists naturally as an element.
____ 114. A solution of two or more metals is
a. an insulator.
b. a jelly.
c. brittle.
d. an alloy.
____ 115. Trends in the properties of elements in a group or period can be explained in terms of
a. binding energy.
c. electron configuration.
b. atomic number.
d. electron affinity.
____ 116. The effect of inner electrons on the attraction between the nucleus and the outer electrons of an atom is called
a. electron affinity.
c. electron shielding.
b. electron ionization.
d. electronegativity.
____ 117. Trends in the periodic table indicate that the element with the greatest ionization energy is in which of the
following periods and groups?
a. Period 2, Group 1
c. Period 1, Group 18
b. Period 7, Group 2
d. Period 6, Group 17
____ 118. One method of measuring the size of an atom involves calculating a value that is _____ the distance between
the nuclei of two bonded atoms.
a. twice
c. half
b. equal to
d. one-quarter
____ 119. Going down a group in the periodic table, electron shielding generally causes the effective nuclear charge to
a. increase.
c. decrease.
b. remain the same.
d. vary unpredictably.
____ 120. Going across a period in the periodic table, electron shielding generally has little effect. As a result, the
effective nuclear charge
a. increases.
c. decreases.
b. remain the same.
d. varies unpredictably.
____ 121. Which is the best reason that the atomic radius generally increases with atomic number in each group of
elements?
a. The nuclear charge increases.
c. The number of energy levels increases.
b. The number of neutrons increases.
d. A new octet forms.
____ 122. For the alkaline-earth metals, atoms with the smallest radii have the
a. largest atomic numbers.
c. most mass.
b. greatest volumes.
d. highest ionization energies.
____ 123. Trends in the periodic table indicate that an element in which of the following periods and groups will have
the smallest anion (negative ion) radius?
a. Period 2, Group 1
c. Period 4, Group 16
b. Period 7, Group 2
d. Period 1, Group 17
____ 124. Trends in the periodic table indicate that an element in which of the following periods and groups will have
the largest cation (positive ion) radius?
a. Period 7, Group 1
c. Period 4, Group 2
b. Period 2, Group 2
d. Period 1, Group 17
____ 125. Electron affinity tends to
a. decrease across a period and decrease down a group.
b. increase across a period and increase down a group.
c. decrease across a period and increase down a group.
d. increase across a period and decrease down a group.
____ 126. The process that changes one element into another different element is called
a. transfiguration.
c. transmutation.
b. transformation.
d. transgeneration.
____ 127. Which of the following electron configurations belongs to an element that is NOT chemically reactive?
a. 1s22s22p3
c. 1s22s22p5
b. 1s22s22p63s23p6
d. 1s22s22p63s23p64s1
____ 128. In the compound sodium fluoride, NaF, the sodium atom loses one electron and the fluorine atom gains one
electron to form ions that have electron configurations similar to
a. helium.
b. oxygen.
c. neon.
d. calcium.
____ 129. In many compounds, atoms of main-group elements form ions so that the number of electrons in the
outermost energy levels of each ion is
a. 2.
c. 8.
b. 6.
d. 10.
____ 130. The electron configuration of nitrogen is 1s22s22p3. How many more electrons does nitrogen need to have an
electron configuration similar to neon?
a. 1
c. 5
b. 3
d. 8
____ 131. Atoms of copper and iron
a. generally form stable bonds with transition elements.
b. have stable electron configurations.
c. tend to form cations.
d. tend to form anions.
____ 132. An anion
a. is an ion with a negative charge.
b. attracts ions with negative charges.
c. results when an alkaline-earth metal loses one of its two outermost electrons.
d. has more protons than electrons.
____ 133. When the octet rule is satisfied, the outermost _____ are filled.
a. d and f orbitals
c. s and d orbitals
b. s and p orbitals
d. d and p orbitals
____ 134. The elements of the _____ group satisfy the octet rule without forming compounds.
a. halide
c. alkali metal
b. noble gas
d. alkaline-earth metal
____ 135. Once an atom has full s and p orbitals in its outermost energy level,
a. it is highly reactive only with alkali metals.
b. it is highly reactive only with halogens.
c. it can be combined with most elements.
d. it has a stable octet and is unreactive.
____ 136. An ion and its parent atom have the same
a. electron configuration.
b. number of charges.
c. atomic number.
d. chemical reactivity.
____ 137. The energy released when a salt is formed from gaseous ions is called the
a. bond energy.
c. lattice energy.
b. potential energy.
d. energy of crystallization.
____ 138. The lattice energy is a measure of the
a. strength of an ionic bond.
b. strength of a metallic bond.
c. strength of a covalent bond.
d. number of ions in a crystal.
____ 139. If the lattice energy of compound A is greater than that of compound B,
a. compound A is not an ionic compound.
b. the bonds in compound A are stronger than the bonds in compound B.
c. compound B is probably a gas.
d. compound A has larger crystals than compound B.
____ 140. When ions are formed,
a. energy is always released.
b. energy is needed.
c. there is no net change in energy.
d. energy can be released or absorbed.
____ 141. When an electron is added to an atom,
a. an input of energy is required.
b. energy is usually released.
c. a stable octet is always formed.
d. ionic bonds are broken.
____ 142. Which of the following is NOT a property of an ionic compound?
a. low boiling point
c. hardness
b. brittleness
d. molten compound conducts electricity
____ 143. Compared with ionic compounds, molecular compounds
a. have higher boiling points.
c. have lower melting points.
b. are brittle.
d. are harder.
____ 144. The melting points of ionic compounds are higher than the melting points of molecular compounds because
a. ionic substances tend to vaporize at room temperature.
b. ionic substances are brittle.
c. attractive forces between ions are greater than the attractive forces between molecules.
d. None of the above
____ 145. Because ions are more strongly attracted in an ionic compound than molecules are attracted in molecular
compounds, the melting points of ionic compounds are
a. equal for all ionic compounds.
b. lower than the melting points of molecular compounds.
c. higher than the melting points of molecular compounds.
d. approximately equal to room temperature.
____ 146. Ionic compounds are brittle because the strong attractive forces
a. allow the layers to shift easily.
b. cause the compound to vaporize easily.
c. keep the surface dull.
d. hold the layers in relatively fixed positions.
____ 147. In the NaCl crystal, each Na+ and Cl– ion has clustered around it _____ of the oppositely charged ions.
a. 1
c. 4
b. 2
d. 6
____ 148. In a crystal of an ionic compound, each cation is surrounded by
a. molecules.
c. dipoles.
b. positive ions.
d. anions.
____ 149. A compound that has the same number of positive and negative charges is said to be
a. cationic.
c. electroneutral.
b. anionic.
d. isoelectronic.
____ 150. What is the ratio of cations to anions in a compound composed of sodium ions, Na+, and carbonate ions,
?
a. 1 to 1
c. 2 to 1
b. 1 to 2
d. 3 to 1
____ 151. The indium(II) ion and indium(III) ion
a. have the same charge.
b. are polyatomic ions.
c. have charges of 1+ and 2+, respectively.
d. have charges of 2+ and 3+, respectively.
____ 152. What is the metallic ion in the compound CuCl2?
a. Cu2–
c. Cu+
–
b. Cu
d. Cu2+
____ 153. What is the formula for the compound formed by calcium ions, Ca2+, and chloride ions, Cl–?
a. CaCl
c. CaCl3
b. Ca2Cl
d. CaCl2
____ 154. The chemical formula for an ionic compound represents the
a. number of atoms in each molecule.
b. number of ions in each molecule.
c. simplest ratio of ions that results in an electrically neutral compound.
d. total number of ions in the crystal lattice.
____ 155. The symbol HCO3– represents a
a. monatomic ion.
b. stable compound.
c. polyatomic ion.
d. salt.
____ 156. A comparison of calcium sulfate and calcium sulfite shows that
a. both have a monatomic cation and a polyatomic anion.
b. calcium sulfite has more oxygen atoms than calcium sulfate.
c. only calcium sulfite contains a polyatomic anion.
d. only calcium sulfate is arranged in a crystal lattice pattern.
____ 157. What is the formula for the compound formed by lead(II) ions and chromate ions, CrO42–?
a. PbCrO4
c. Pb2(CrO4)3
b. Pb2CrO4
d. Pb(CrO4)2
____ 158. What is the formula for the compound formed by aluminum(III) and the sulfate ion, SO42–?
a. AlSO4
c. Al2(SO4)3
b. Al2SO4
d. Al(SO4)3
____ 159. What is the formula for the compound formed by tin(IV) and the chromate ion, CrO42–?
a. Sn(CrO4)4
c. Sn2(CrO4)4
b. Sn2(CrO4)2
d. Sn(CrO4)2
____ 160. What is the formula for the compound formed by the barium ion, Ba2+, and the hydroxide ion, OH–?
a. BaOH
c. Ba(OH)2
b. BaOH2
d. Ba(OH)
____ 161. Name the compound Ni(ClO3)2.
a. nickel chlorate
b. nickel chloride
c. nickel chlorite
d. nickel peroxide
____ 162. Name the compound Zn3(PO4)2.
a. zinc potassium oxide
c. zinc phosphate
b. trizinc polyoxide
d. zinc phosphite
____ 163. Name the compound Hg2(NO3)2.
a. mercury(II) nitrate
b. dimercury dinitrate
c. mercury(I) nitrate
d. mercuric nitrate
____ 164. Name the compound KClO3.
a. potassium chloride
b. potassium trioxychlorite
c. potassium chlorate
d. hypochlorite
____ 165. Name the compound Fe(NO2)2.
a. iron(II) nitrate
b. iron(II) nitrite
c. ferric nitrate
d. ferrous nitride
____ 166. Name the compound CuCO3.
a. Copper(I) carbonate
b. cupric trioxycarbide
c. cuprous carbide
d. copper(II) carbonate
____ 167. What is the name of Sn3(PO4)4?
a. stannous phosphate
b. tin(IV) phosphate
c. tin(III) phosphate
d. tin(II) phosphate
____ 168. The name of a polyatomic ion that contains hydrogen begins with the term
a. hypo-.
c. hydrogen.
b. thio-.
d. per-.
____ 169. The electrons involved in the formation of a covalent bond are
a. transferred from one atom to another.
c. valence electrons.
b. found only in the s-orbitals.
d. in filled orbitals.
____ 170. The chemical bond formed when two atoms share one or more pairs of electrons is a(n)
a. ionic bond.
c. polar bond.
b. orbital bond.
d. covalent bond.
____ 171. If two covalently bonded atoms move farther than a distance of the bond length, the potential energy of the
atoms
a. becomes negative.
c. increases.
b. decreases.
d. remains constant.
____ 172. A covalent bond forms when the attraction between two atoms is balanced by repulsion and the potential
energy is
a. at a maximum.
c. at a minimum.
b. zero.
d. equal to the kinetic energy.
____ 173. Which of the following compounds most likely has the least bond energy?
a. Cl2; Cl–Cl bond length = 199 pm
c. HF; H–F bond length = 92 pm
b. HCl; H–Cl bond length = 127 pm
d. I2; I–I bond length = 266 pm
____ 174. A nonpolar covalent bond is most likely to form between two elements that have a difference in
electronegativity values of
a. 0.1.
c. 3.0.
b. 1.5.
d. Both (a) and (b)
____ 175. An ionic bond is most likely to form between two elements that have a difference in electronegativity values
of
a. 0.1.
c. 3.0.
b. 1.5.
d. Both (a) and (b)
____ 176. A polar covalent bond is most likely to form between two elements that have a difference in electronegativity
values of
a. 0.1.
c. 3.0.
b. 1.5.
d. Both (a) and (b)
____ 177. Which of the following molecular formulas show the polar nature of the HBr molecule?
a. H+Br+
c. H-Br+
b. H+Br
d. H-Br
____ 178. Which of the following substances most likely has the lowest boiling point?
a. Cl2
c. MgCl2
b. HF
d. Cu
____ 179. The correct Lewis structure for the oxygen atom has
a. one pair of valence electrons and one single valence electron.
b. two pairs of valence electrons and one single valence electron.
c. two pairs of valence electrons and two single valence electrons.
d. three pairs of valence electrons.
____ 180. The correct Lewis structure for the boron atom has
a. one single valence electron.
b. two single valence electrons.
c. three single valence electrons.
d. one pair of valence electrons and one single valence electron.
____ 181. The correct Lewis structure for a Group 18 atom has
a. one pair of valence electrons and one single valence electron.
b. two pairs of valence electrons and one single valence electron.
c. three pairs of valence electrons and one single valence electron.
d. four pairs of valence electrons.
____ 182. The correct Lewis structure for a fluorine atom in a molecule of F2 shows
a. three unshared pairs of electrons.
c. one shared pair of electrons.
b. an octet of valence electrons.
d. All of the above
____ 183. To draw the Lewis structure of the polyatomic ion,
you would have to _____ those in the structures of
Cl, O, O, and O.
a. add one electron to
c. add a proton to
b. subtract one electron from
d. subtract a proton from
____ 184. In a double bond, two atoms share a total of _____ electrons.
a. two
c. four
b. three
d. six
____ 185. The correct Lewis structure for a molecule of the compound C2H2 contains
a. three single bonds.
c. three double bonds.
b. two double bonds.
d. one triple bond.
____ 186. To indicate resonance, a _____ is placed between a molecule’s resonance structures.
a. single-headed arrow
c. long dash
b. series of three raised dots
d. double-headed arrow
____ 187. Which of the following number of atoms can form a tetrahedral molecule?
a. two
c. four
b. three
d. five
____ 188. According to VSEPR theory, which of the following shapes is possible for a molecule with the molecular
formula of AB2?
a. linear
c. trigonal pyramidal
b. bent
d. Both (a) and (b)
____ 189. According to VSEPR theory, which of the following shapes is possible for a molecule with the molecular
formula of AB3?
a. linear
c. trigonal pyramidal
b. trigonal planar
d. Both (b) and (c)
____ 190. According to VSEPR theory, a molecule with the molecular formula of A2 is
a. linear in shape.
c. bent.
b. a dipole.
d. Both (a) and (b)
____ 191. According to VSEPR theory, the molecular shape of CO2 is classified as
a. linear.
c. trigonal planar.
b. bent.
d. trigonal pyramidal.
____ 192. According to VSEPR theory, the molecular shape of H2O is classified as
a. linear.
c. trigonal planar.
b. bent.
d. trigonal pyramidal.
____ 193. Iodine monochloride, ICl, has a higher boiling point than bromine, Br2, partly because ICl is a(n)
a. nonpolar molecular substance.
c. metallic substance.
b. ionic substance.
d. polar molecular substance.
Midterm Review
Answer Section
MULTIPLE CHOICE
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
A
B
D
D
D
B
D
D
B
A
C
C
A
C
D
A
D
B
B
A
C
B
C
D
C
A
A
A
C
A
D
A
B
B
C
A
D
D
B
B
D
42.
43.
44.
45.
46.
47.
48.
49.
50.
51.
52.
53.
54.
55.
56.
57.
58.
59.
60.
61.
62.
63.
64.
65.
66.
67.
68.
69.
70.
71.
72.
73.
74.
75.
76.
77.
78.
79.
80.
81.
82.
83.
84.
85.
86.
87.
88.
D
C
C
D
B
C
A
D
C
B
D
D
C
C
A
B
A
C
B
B
A
D
A
B
C
D
B
B
A
C
B
C
B
B
C
C
C
D
A
B
C
D
C
C
B
D
A
89.
90.
91.
92.
93.
94.
95.
96.
97.
98.
99.
100.
101.
102.
103.
104.
105.
106.
107.
108.
109.
110.
111.
112.
113.
114.
115.
116.
117.
118.
119.
120.
121.
122.
123.
124.
125.
126.
127.
128.
129.
130.
131.
132.
133.
134.
A
C
D
B
C
A
A
C
A
B
B
D
C
D
B
D
A
C
A
A
C
C
C
D
B
D
C
C
C
C
C
A
C
D
D
A
D
C
B
C
C
B
C
A
B
B
135.
136.
137.
138.
139.
140.
141.
142.
143.
144.
145.
146.
147.
148.
149.
150.
151.
152.
153.
154.
155.
156.
157.
158.
159.
160.
161.
162.
163.
164.
165.
166.
167.
168.
169.
170.
171.
172.
173.
174.
175.
176.
177.
178.
179.
180.
181.
D
C
C
A
B
D
B
A
C
C
C
D
D
D
C
C
D
D
D
C
C
A
A
C
D
C
A
C
C
C
B
D
B
C
C
D
C
C
D
A
C
B
B
A
C
C
D
182.
183.
184.
185.
186.
187.
188.
189.
190.
191.
192.
193.
D
A
C
D
D
D
D
D
A
A
B
D