Multiple Choice
Identify the
choice that best completes the statement or answers the question.
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1.
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A mutual electrical attraction between the nuclei and valence electrons of
different atoms that binds the atoms together is called a(n)
a. | dipole. | c. | chemical bond. | b. | Lewis structure. | d. | London force. |
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2.
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The electrons involved in the formation of a chemical bond are called
a. | dipoles. | c. | Lewis electrons. | b. | s electrons. | d. | valence
electrons. |
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3.
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The electrostatic attraction between positively charged nuclei and negatively
charged electrons permits two atoms to be held together by a(n)
a. | chemical bond. | c. | neutron. | b. | London force. | d. | ion. |
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4.
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As independent particles, most atoms are
a. | at relatively high potential energy. | c. | very stable. | b. | at relatively low
potential energy. | d. | part of a
chemical bond. |
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5.
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Atoms naturally move
a. | toward high potential energy. | c. | toward less
stability. | b. | toward low potential energy. | d. | away from each
other. |
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6.
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As atoms bond with each other, they
a. | increase their potential energy, thus creating less-stable arrangements of
matter. | b. | decrease their potential energy, thus creating less-stable arrangements of
matter. | c. | increase their potential energy, thus creating more-stable arrangements of
matter. | d. | decrease their potential energy, thus creating more-stable arrangements of
matter. |
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7.
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If two covalently bonded atoms are identical, the bond is
a. | nonpolar covalent. | c. | dipole covalent. | b. | polar covalent. | d. | coordinate
covalent. |
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8.
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When atoms share electrons, the electrical attraction of an atom for the shared
electrons is called the atom's
a. | electron affinity. | c. | resonance. | b. | electronegativity. | d. | hybridization. |
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9.
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If the atoms that share electrons have an unequal attraction for the electrons,
the bond is called
a. | nonpolar. | c. | ionic. | b. | polar. | d. | dipolar. |
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10.
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What are shared in a covalent bond?
a. | ions | c. | electrons | b. | Lewis structures | d. | dipoles |
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11.
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Most chemical bonds are
a. | purely ionic. | c. | partly ionic and partly covalent. | b. | purely
covalent. | d. | metallic. |
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12.
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Nonpolar covalent bonds are not common because
a. | one atom usually attracts electrons more strongly than the other. | b. | ions always form
when atoms join. | c. | the electrons usually remain equally distant from both atoms. | d. | dipoles are rare in
nature. |
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13.
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The greater the electronegativity difference between two bonded atoms, the
greater the percentage of ____ in the bond.
a. | ionic character | c. | metallic character | b. | covalent character | d. | electron
sharing |
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14.
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A bond that is less than 5% ionic is considered
a. | polar covalent. | c. | nonpolar covalent. | b. | ionic. | d. | metallic. |
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15.
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The pair of elements that forms a bond with the least ionic character is
a. | Na and Cl. | c. | O and Cl. | b. | K and Cl. | d. | Mg and Cl. |
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16.
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The B—F bond in BF3 (electronegativity for B is 2.0;
electronegativity for F is 4.0) is
a. | polar covalent. | c. | nonpolar covalent. | b. | ionic. | d. | metallic. |
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17.
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In the three molecules, O2, HCl, and F2, what atom would
have a partial negative charge?
a. | oxygen | c. | chlorine | b. | hydrogen | d. | fluorine |
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18.
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The percentage ionic character and the type of bond in Br2
(electronegativity for Br is 2.8) is
a. | 0%; nonpolar covalent. | c. | 0%; pure ionic. | b. | 100%; polar covalent. | d. | 100%; pure
ionic. |
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19.
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A neutral group of atoms held together by covalent bonds is a
a. | molecular formula. | c. | polyatomic ion. | b. | chemical formula. | d. | molecule. |
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20.
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Which of the following shows the types and numbers of atoms joined in a single
molecule of a molecular compound?
a. | molecular formula | c. | covalent bond | b. | potential energy diagram | d. | ionic bond |
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21.
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Which of the following is not an example of a molecular formula?
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22.
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When a stable covalent bond forms, the potential energy of the atoms
a. | increases. | c. | remains constant. | b. | decreases. | d. | becomes zero. |
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23.
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Bond length is the average distance between two bonded atoms
a. | at which potential energy is at a minimum. | b. | at which kinetic
energy is at a maximum. | c. | at which potential energy is at a
maximum. | d. | and equal to one-half the diameter of the electron
cloud. |
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24.
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The energy released when a covalent bond forms is the difference between zero
and the
a. | maximum potential energy. | c. | minimum potential
energy. | b. | kinetic energy of the atom. | d. | bond length expressed in nanometers. |
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25.
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In a molecule of fluorine, the two shared electrons give each fluorine atom how
many electron(s) in the outer energy level?
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26.
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The electron configuration of nitrogen is 1s2
2s2 2p3. How many more electrons does nitrogen need to satisfy
the octet rule?
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27.
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What group of elements satisfies the octet rule without forming
compounds?
a. | halogen | c. | alkali metal | b. | noble gas | d. | alkaline-earth
metal |
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28.
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In drawing a Lewis structure, each nonmetal atom except hydrogen should be
surrounded by
a. | 2 electrons. | c. | 8 electrons. | b. | 4 electrons. | d. | 10 electrons. |
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29.
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In drawing a Lewis structure, the central atom is generally the
a. | atom with the greatest mass. | b. | atom with the highest atomic
number. | c. | atom with the fewest electrons. | d. | least electronegative
atom. |
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30.
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To draw a Lewis structure, one must know the
a. | number of valence electrons in each atom. | b. | atomic mass of each
atom. | c. | bond length of each atom. | d. | ionization energy of each
atom. |
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31.
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After drawing a Lewis structure, one should
a. | determine the number of each type of atom in the molecule. | b. | add unshared pairs
of electrons around nonmetal atoms. | c. | confirm that the total number of valence
electrons used equals the number available. | d. | determine the electronegativity of each
atom. |
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32.
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Multiple covalent bonds may occur in atoms that contain carbon, nitrogen,
or
a. | chlorine. | c. | oxygen. | b. | hydrogen. | d. | helium. |
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33.
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The substance whose Lewis structure shows three covalent bonds is
a. | H2O. | c. | NH3. | b. | CH2Cl2. | d. | CCl4. |
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34.
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What is the correct Lewis structure for hydrogen chloride, HCl?  
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35.
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Bonding in molecules or ions that cannot be correctly represented by a single
Lewis structure is
a. | polyatomic. | c. | single bonding. | b. | resonance. | d. | double bonding. |
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36.
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What is placed between a molecule's resonance structures to indicate
resonance?
a. | double-headed arrow | c. | series of dots | b. | single-headed arrow | d. | Lewis structure |
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37.
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Chemists once theorized that a molecule that contains a single bond and a double
bond split its time existing as one of these two structures. This effect became known as
a. | alternation. | c. | multiple bonding. | b. | resonance. | d. | single-double
bonding. |
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38.
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The chemical formula for an ionic compound represents the
a. | number of atoms in each molecule. | b. | number of ions in each
molecule. | c. | ratio of the combined ions present in a sample. | d. | total number of ions
in the crystal lattice. |
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39.
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A formula that shows only the types and numbers of atoms combined in a single
molecule is called a(n)
a. | molecular formula. | c. | Lewis structure. | b. | ionic formula. | d. | covalent
formula. |
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40.
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The chemical formula for water, a covalent compound, is H2O. This
formula is an example of a(n)
a. | formula unit. | c. | ionic formula. | b. | Lewis structure. | d. | molecular
formula. |
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41.
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In the NaCl crystal, each Na+ and Cl– ion has how
many oppositely charged ions clustered around it?
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42.
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In an ionic compound, the orderly arrangement of ions in a crystal is the state
of
a. | maximum potential energy. | c. | average potential
energy. | b. | minimum potential energy. | d. | zero potential energy. |
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43.
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The ions in most ionic compounds are organized into a
a. | molecule. | c. | polyatomic ion. | b. | Lewis structure. | d. | crystal. |
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44.
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In a crystal, the electrons of adjacent ions
a. | repel each other. | c. | neutralize each other. | b. | attract each
other. | d. | have no effect on
each other. |
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45.
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The energy released when 1 mol of an ionic crystalline compound is formed from
gaseous ions is called the
a. | bond energy. | c. | lattice energy. | b. | potential energy. | d. | energy of
crystallization. |
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46.
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The lattice energy is a measure of the
a. | strength of an ionic bond. | c. | strength of a covalent
bond. | b. | strength of a metallic bond. | d. | net charge on a
crystal. |
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47.
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Compared with energies of neutral atoms, a crystal lattice has
a. | higher potential energy. | c. | equal potential
energy. | b. | lower potential energy. | d. | less stability. |
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48.
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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. |
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49.
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Compared with ionic compounds, molecular compounds
a. | have higher boiling points. | c. | have lower melting
points. | b. | are brittle. | d. | are harder. |
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50.
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The forces of attraction between molecules in a molecular compound are
a. | stronger than the forces among formula units in ionic bonding. | b. | weaker than the
forces among formula units in ionic bonding. | c. | approximately equal to the forces among formula
units in ionic bonding. | d. | zero. |
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51.
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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. |
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52.
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The properties of both ionic and molecular compounds are related to the
a. | lattice energies of the compounds. | b. | strengths of attraction between the particles
in the compounds. | c. | number of covalent bonds each
contains. | d. | mobile electrons that they contain. |
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53.
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The Lewis structure for the ammonium ion, NH4, has
a. | nonpolar covalent bond. | c. | polar covalent
bond. | b. | ionic bond. | d. | metallic bond. |
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54.
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How many extra electrons are in the Lewis structure of the phosphate ion,
PO43–?
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55.
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How many electrons must be shown in the Lewis structure of the hydroxide ion,
OH–?
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56.
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A chemical bond formed by the attraction between positive ions and surrounding
mobile electrons is a(n)
a. | nonpolar covalent bond. | c. | polar covalent
bond. | b. | ionic bond. | d. | metallic bond. |
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57.
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Compared with nonmetals, the number of valence electrons in metals is
generally
a. | smaller. | c. | about the same. | b. | greater. | d. | almost triple. |
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58.
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In metals, the valence electrons
a. | are attached to particular positive ions. | c. | are immobile. | b. | are shared by all of
the atoms. | d. | form covalent
bonds. |
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59.
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In metallic bonds, the mobile electrons surrounding the positive ions are called
a(n)
a. | Lewis structure. | c. | electron cloud. | b. | electron sea. | d. | dipole. |
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60.
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To appear shiny, a material must be able to
a. | form crystals. | b. | absorb and re-emit light of many
wavelengths. | c. | absorb light and change it all to energy as heat. | d. | change light to
electricity. |
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61.
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The shiny appearance of a metal is most closely related to the
metal's
a. | highly mobile valence electrons. | c. | brittle crystalline
structure. | b. | covalent bonds. | d. | positive ions. |
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62.
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As light strikes the surface of a metal, the electrons in the electron
sea
a. | allow the light to pass through. | b. | become attached to particular positive
ions. | c. | fall to lower energy levels. | d. | absorb and re-emit the
light. |
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63.
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If a material can be shaped or extended by physical pressure, such as hammering,
which property does the material have?
a. | conductivity | c. | ductility | b. | malleability | d. | luster |
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64.
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Metals are malleable because the metallic bonding
a. | holds the layers of ions in rigid positions. | b. | maximizes the
repulsive forces within the metal. | c. | allows one plane of ions to slide past
another. | d. | is easily broken. |
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65.
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Which best explains the observation that metals are malleable and ionic crystals
are brittle?
a. | their chemical bonds | c. | their enthalpies of vaporization | b. | their London
forces | d. | their net
change |
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66.
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Malleability and ductility are characteristic of substances with
a. | covalent bonds. | c. | Lewis structures. | b. | ionic bonds. | d. | metallic bonds. |
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67.
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Shifting the layers of an ionic crystal causes the crystal to
a. | be drawn into a wire. | c. | become metallic. | b. | shatter. | d. | emit light. |
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68.
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According to VSEPR theory, an AB2 molecule is
a. | trigonal-planar. | c. | linear. | b. | tetrahedral. | d. | octahedral. |
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69.
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VSEPR theory is a model for predicting
a. | the strength of metallic bonds. | c. | lattice energy
values. | b. | the shape of molecules. | d. | ionization energy. |
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70.
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The concept that electrostatic repulsion between electron pairs surrounding an
atom causes these pairs to be separated as far as possible is the foundation of
a. | VSEPR theory. | c. | the electron-sea model. | b. | the hybridization
model. | d. | Lewis
theory. |
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71.
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According to VSEPR theory, the shape of an AB3 molecule is
a. | trigonal-planar. | c. | linear. | b. | tetrahedral. | d. | bent. |
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72.
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According to VSEPR theory, the structure of the ammonia molecule,
NH3, is
a. | trigonal-planar. | c. | trigonal-pyramidal. | b. | bent. | d. | tetrahedral. |
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73.
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Use VSEPR theory to predict the shape of the hydrogen chloride molecule,
HCl.
a. | tetrahedral | c. | bent | b. | linear | d. | trigonal-planar |
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74.
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Use VSEPR theory to predict the shape of the magnesium hydride molecule,
MgH2.
a. | tetrahedral | c. | bent | b. | linear | d. | octahedral |
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75.
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Use VSEPR theory to predict the shape of the carbon tetraiodide molecule,
CI4.
a. | tetrahedral | c. | bent | b. | linear | d. | trigonal-planar |
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76.
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Use VSEPR theory to predict the shape of the chlorate ion,
ClO3–.
a. | trigonal-planar | c. | trigonal-pyramidal | b. | octahedral | d. | bent |
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77.
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Use VSEPR theory to predict the shape of carbon dioxide, CO2.
a. | tetrahedral | c. | bent | b. | linear | d. | octahedral |
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78.
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The hybridized orbitals responsible for the bent shape of the water molecule
are
a. | 1s2 2s2. | c. | sp3. | b. | ps1. | d. | 2s2
sp2. |
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79.
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The mixing of two or more atomic orbitals of similar energies on the same atom
to produce new orbitals of equal energies is called
a. | VSEPR theory. | c. | hybridization. | b. | malleability. | d. | dipole-dipole
interaction. |
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80.
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Which hybrid orbitals help explain the bonding in methane,
CH4?
a. | sp3 orbitals | c. | pd3
orbitals | b. | sp orbitals | d. | df3 orbitals |
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81.
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Four hybrid sp3 orbitals are formed from
a. | two s orbitals and two p orbitals. | b. | an s orbital
and a p orbital. | c. | three s orbitals and one p
orbital. | d. | one s orbital and three p orbitals. |
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82.
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Dipole-dipole forces are considered the most important forces in polar
substances because the London dispersion forces present in polar substances
a. | are no longer present. | b. | are usually much weaker than the dipole-dipole
forces. | c. | are too unpredictable. | d. | act only in
solids. |
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83.
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The strength of London dispersion forces between molecules depends on
a. | only the number of electrons in the molecule. | b. | only the number of
protons in the molecule. | c. | both the number of electrons in the molecule
and the mass of the molecule. | d. | both the number of electrons and the number of
neutrons in the molecule. |
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84.
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The strong forces of attraction between the positive and negative regions of
molecules are called
a. | dipole-dipole forces. | c. | lattice forces. | b. | London forces. | d. | orbital forces. |
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85.
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Compared with molecular bonds, the strength of intermolecular forces is
a. | weaker. | c. | about the same. | b. | stronger. | d. | too variable to
compare. |
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86.
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The equal but opposite charges present in the two regions of a polar molecule
create a(n)
a. | electron sea. | c. | crystal lattice. | b. | dipole. | d. | ionic bond. |
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87.
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The reason the boiling point of water (H2O) is higher than the
boiling point of hydrogen sulfide (H2S) is partially explained by
a. | London forces. | c. | ionic bonding. | b. | covalent bonding. | d. | hydrogen
bonding. |
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88.
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The following molecules contain polar bonds. The only polar molecule is
a. | CCl4. | c. | NH3. | b. | CO2. | d. | CH4. |
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89.
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A polar molecule contains
a. | ions. | b. | a region of positive charge and a region of
negative charge. | c. | only London forces. | d. | no bonds. |
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90.
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When a polar molecule attracts the electron in a nonpolar molecule,
a. | a dipole is induced. | c. | an ionic bond forms. | b. | a crystal lattice forms. | d. | a Lewis structure
forms. |
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91.
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Iodine monochloride (ICl) has a higher boiling point than bromine
(Br2) partly because iodine monochloride is a(n)
a. | nonpolar molecule. | c. | metal. | b. | polyatomic ion. | d. | polar molecule. |
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Short Answer
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92.
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Why do most atoms form chemical bonds?
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93.
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Explain why scientists use resonance structures to represent some
molecules.
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94.
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Differentiate between an ionic compound and a molecular compound.
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95.
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Explain why metals are good conductors of electricity.
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Problem
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96.
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Draw a Lewis structure for the oxalate ion,
C2O42–.
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97.
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Draw a Lewis structure for carbon disulfide, CS2.
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98.
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Draw a Lewis structure for the nitrate ion,  . Use VSEPR
theory to predict its molecular geometry.
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99.
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Draw a ball-and-stick model of a water molecule. Label that atoms, include the
polarities of the bonds using arrows, and indicate net molecular dipole.
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Essay
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100.
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How many different kinds of covalent bonds can a nitrogen atom form?
Explain.
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