Chapter 6 Chemistry Review
Multiple Choice
Identify the choice that best completes the statement or answers the question. Put the LETTER of the correct answer in the
blank.
____
1. The electrons involved in the formation of a chemical bond are called
a. dipoles.
c. Lewis electrons.
b. s electrons.
d. valence electrons.
____
2. 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.
____
3. As independent particles, most atoms are
a. at relatively high potential energy.
b. at relatively low potential energy.
c. very stable.
d. part of a chemical bond.
____
4. 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.
____
5. What are shared in a covalent bond?
a. ions
b. Lewis structures
c. electrons
d. dipoles
6. Most chemical bonds are
a. purely ionic.
b. purely covalent.
c. partly ionic and partly covalent.
d. metallic.
____
____
7. 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
____
8. A neutral group of atoms held together by covalent bonds is a
a. molecular formula.
c. polyatomic ion.
b. chemical formula.
d. molecule.
____
9. 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
____ 10. 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.
____ 11. In a molecule of fluorine, the two shared electrons give each fluorine atom how many electron(s) in the outer
energy level?
a. 1
b. 2
c. 8
d. 32
____ 12. What group of elements satisfies the octet rule without forming compounds?
a. halogen
c. alkali metal
b. noble gas
d. alkaline-earth metal
____ 13. 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.
____ 14. 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.
____ 15. 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.
____ 16. Multiple covalent bonds may occur in atoms that contain carbon, nitrogen, or
a. chlorine.
c. oxygen.
b. hydrogen.
d. helium.
____ 17. 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.
____ 18. 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
____ 19. 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.
____ 20. 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.
____ 21. In the NaCl crystal, each Na+ and Cl– ion has how many oppositely charged ions clustered around it?
a. 1
c. 4
b. 2
d. 6
____ 22. 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.
____ 23. 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.
____ 24. 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.
____ 25. 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.
____ 26. How many extra electrons are in the Lewis structure of the phosphate ion, PO43–?
a. 0
c. 3
b. 2
d. 4
____ 27. 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.
____ 28. 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
____ 29. Malleability and ductility are characteristic of substances with
a. covalent bonds.
c. Lewis structures.
b. ionic bonds.
d. metallic bonds.
____ 30. 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.
____ 31. According to VSEPR theory, the shape of an AB3 molecule is
a. trigonal-planar.
c. linear.
b. tetrahedral.
d. bent.
____ 32. Use VSEPR theory to predict the shape of the hydrogen chloride molecule, HCl.
a. tetrahedral
c. bent
b. linear
d. trigonal-planar
____ 33. Use VSEPR theory to predict the shape of the chlorate ion, ClO3–.
a. trigonal-planar
c. trigonal-pyramidal
b. octahedral
d. bent
____ 34. Use VSEPR theory to predict the shape of carbon dioxide, CO2.
a. tetrahedral
c. bent
b. linear
d. octahedral
____ 35. The hybridized orbitals responsible for the bent shape of the water molecule are
a. 1s2 2s2.
c. sp3.
1
b. ps .
d. 2s2 sp2.
____ 36. 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.
____ 37. 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.
____ 38. 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.
____ 39. 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.
____ 40. 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.
____ 41. The following molecules contain polar bonds. The only polar molecule is
a. CCl4.
c. NH3.
b. CO2.
d. CH4.
Short Answer
42. Why do most atoms form chemical bonds?
43. Explain why scientists use resonance structures to represent some molecules.
44. Differentiate between an ionic compound and a molecular compound.
45. Explain why metals are good conductors of electricity.
Problem
46. Draw a Lewis structure for the oxalate ion, C2O42–.
47. Draw a Lewis structure for carbon disulfide, CS2.
48. Draw a Lewis structure for the nitrate ion,
. Use VSEPR theory to predict its molecular geometry.
49. 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.
Essay
50. How many different kinds of covalent bonds can a nitrogen atom form? Explain.
Chapter 6 Chemistry Review
Answer Section
MULTIPLE CHOICE
1.
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41.
D
A
A
D
C
C
A
D
A
C
C
B
C
D
A
C
B
A
A
D
D
B
C
B
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C
A
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B
A
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B
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C
SHORT ANSWER
42. Atoms form chemical bonds to establish a more-stable arrangement. As independent particles, they are at high
potential energy. By bonding, they decrease their potential energy, thus becoming more stable.
43. Resonance structures represent the bonding in molecules that cannot be adequately represented with a single
Lewis structure.
44. Atoms in a molecular compound share electrons to achieve stability. Atoms in an ionic compound gain or lose
electrons to form ions, which combine so that the number of positive and negative charges is equal.
45. The valence electrons in a metal's structure are delocalized, so they can move freely and carry an electric
charge throughout the metal.
PROBLEM
46.
47.
48.
trigonal-planar
O
49.
H
δ-
H
δ+
ESSAY
50. A nitrogen atom has five valence electrons. To have a complete octet of electrons, the nitrogen atom forms
three covalent bonds. It could form three single bonds, one single, and one double bond, or one triple bond.