The Basics Of Chemical Bonding Full Test Bank Chapter 8

Chemistry: Molecular Nature of Matter, 8e (Jespersen)

Chapter 8 The Basics of Chemical Bonding

1) Which of the following statements is true?

A) Bond formation tends to increase the potential energy of the atoms involved.

B) Compounds with negative heats of formation tend to be unstable with respect to decomposition into their elements.

C) When ΔH°f is negative, the reaction produces a stable compound.

D) An endothermic reaction produces a fairly stable compound.

E) Compounds with positive heats of formation tend to be stable with respect to decomposition into their elements.

Diff: 2

Section: 8.1

2) Based on the ΔH°f data given, which compound is the most stable?

A) N2O4(g), +9.7 kJ/mol

B) H2S(g), -20.6 kJ/mol

C) N2H4(g), +94.5 kJ/mol

D) PH3(g), +5.4 kJ/mol

E) NH3(g), -46.4 kJ/mol

Diff: 2

Section: 8.1

3) Based on the ΔH°f data given, which of the following compounds is the most unstable?

A) SrO(s), -592.0 kJ/mol

B) CsBr(s), -396 kJ/mol

C) CaF2(s), -1215 kJ/mol

D) NaI(s), -288 kJ/mol

E) KF(s), -568.6 kJ/mol

Diff: 2

Section: 8.1

4) Based on the ΔH°f data given, which compound is the most stable?

A) CO2(g), -393.5 kJ/mol

B) NO2(g), +33.85 kJ/mol

C) NO(g), +90.4 kJ/mol

D) CO(g), -110.5 kJ/mol

E) NH3(g), -46.4 kJ/mol

Diff: 2

Section: 8.1

5) Given that the first ionization energy of cesium is +376 kJ/mol and the electron affinity of bromine is −325 kJ/mol, calculate ΔE for the reaction, Cs(g) + Br(g) s Cs+(g) + Br−(g).

A) +376 kJ/mol

B) +701 kJ/mol

C) +51 kJ/mol

D) -701 kJ/mol

E) -51 kJ/mol

Diff: 2

Section: 8.2

6) Lithium fluoride has a lattice energy of -1033 kJ/mol. In the ionic solid AB, A2+ has approximately the same radius as Li+, and B2− has approximately the same radius as F−. What is a reasonable estimate of the lattice energy of AB?

A) About 3 × (-1033 kJ/mol)

B) About -1033 kJ/mol

C) About 4 × (-1033 kJ/mol)

D) About 2 × (-1033 kJ/mol)

E) About 6 × (-1033 kJ/mol)

Diff: 2

Section: 8.2

7) Which ionic solid is likely to have the largest exothermic lattice energy?

A) KCl

B) NaCl

C) LiCl

D) CsCl

E) They all have the same value.

Diff: 2

Section: 8.2

8) Which of the following solids is likely to have the largest exothermic lattice energy?

A) LiF

B) NaCl

C) AlCl3

D) Al2O3

E) CaCl2

Diff: 2

Section: 8.2

9) Which ionic solid is likely to have the least exothermic lattice energy?

A) KCl

B) NaCl

C) LiCl

D) CsCl

E) RbBr

Diff: 2

Section: 8.2

10) Which of the following solids is likely to have the least exothermic lattice energy?

A) LiF

B) NaCl

C) AlCl3

D) Al2O3

E) CaCl2

Diff: 2

Section: 8.2

11) Which of the following solids would have the highest melting point?

A) NaF

B) NaCl

C) NaBr

D) NaI

E) They should all be the same.

Diff: 1

Section: 8.2

12) Which of the following solids would have the highest melting point?

A) LiF

B) NaCl

C) AlCl3

D) Al2O3

E) CaCl2

Diff: 2

Section: 8.2

13) Which of the following statements is true in relation to the formation of the stable ionic compound sodium chloride?

Hint: Consider why the sodium to chloride bond occurs.

A) The net absorption of 147 kJ/mol of energy for the formation of the gaseous sodium and chloride ions is mainly responsible for the formation of the stable ionic compound.

B) The lattice energy provides the necessary stabilization energy for the formation of sodium chloride.

C) Sodium chloride is stable because it is formed from the combination of isolated gaseous ions.

D) The first ionization energy of sodium contributes favorably to the overall formation of sodium chloride.

E) The release of energy as the solid sodium chloride is formed leads to an overall increase in the potential energy.

Diff: 3

Section: 8.2

14) For two ions with charges q1 and q2 separated by distance r, the potential energy can

be calculated from Coulomb's law: E =

Calculate the energy released when 1 mole NaCl is formed if the constant k = 1.11 × 10−10 C2/J·m, the charge Na+ = +1 e, and the charge on Cl− is -1 e, e = 1.602 × 10−19 C, and .

A) +308 kJ/mol

B) +8.20 kJ/mol

C) −297 kJ/mol

D) −494 kJ/mol

E) −371 kJ/mol

Diff: 2

Section: 9.2

15) Sodium forms a monatomic ion that has the electron configuration of a noble gas. What is the electron configuration of that noble gas?

A) 1s2

B) 1s2 2p6

C) 1s2 2s2 2p6

D) 1s2 2s2 2p6 3s2

E) 1s2 2s2 2p6 3s2 3p6

Diff: 1

Section: 8.3

16) Magnesium forms a monatomic ion that has the electron configuration of a noble gas. What is the electron configuration of that noble gas?

A) 1s2 2p6

B) 1s2 2s2 2p6

C) 1s2 2s2 2p6 3s1

D) 1s2 2s2 2p6 3s2

E) 1s2 2s2 2p6 3s2 3p6

Diff: 1

Section: 8.3

17) Bromine tends to form a monatomic ion that has the electron configuration of a noble gas. What is the symbol of that noble gas?

A) Ne

B) Ar

C) Kr

D) Xe

E) Rn

Diff: 2

Section: 8.3

18) Sulfur tends to form a monatomic ion that has the electron configuration of a noble gas. What is the symbol of the noble gas?

A) Ne

B) Ar

C) Kr

D) Xe

E) Rn

Diff: 2

Section: 8.3

19) The octet rule is generally followed for cations of elements located where on the periodic table?

A) post transition metals and transition metals

B) Group IA and Group IIA metals

C) Group IVA elements

D) lanthanide elements

E) actinide elements

Diff: 2

Section: 8.3

20) Which ion of uranium has a noble gas electron configuration?

A) U2+

B) U3+

C) U4+

D) U6+

E) U2−

Diff: 2

Section: 8.3

21) Which ion has a noble gas electron configuration?

A) Fe3+

B) Sn2+

C) Ni2+

D) Ti4+

E) Cr3+

Diff: 2

Section: 8.3

22) Which ion has a noble gas electron configuration?

A) Fe3+

B) Ni2+

C) Sc3+

D) V3+

E) Cr3+

Diff: 2

Section: 8.3

23) Which atom has the same electron configuration as In3+?

A) Te

B) Pd

C) Ir

D) Zn

E) Ga

Diff: 2

Section: 8.3

24) Give the ground state electron configuration and number of unpaired electrons for Sb3+.

A) [Kr] 4d1 5s1; 2 unpaired electrons

B) [Ar]; 1 unpaired electron

C) [Kr] 4d10 5s2; 0 unpaired electrons

D) [Kr] 4d7 5s2; 3 unpaired electrons

E) [Ar] 4d4 5s2; 4 unpaired electrons

Hint: Draw the orbital diagram of Sb3+ to determine unpaired electrons.

Diff: 3

Section: 8.3

25) Give the ground state electron configuration and number of unpaired electrons for V3+.

A) [Ar] 3d1; 2 unpaired electrons

B) [Ar]; 1 unpaired electron

C) [Ar] 3d2; 2 unpaired electrons

D) [Ar] 3d3; 3 unpaired electrons

E) [Ar] 3d4; 4 unpaired electrons

Hint: Draw the orbital diagram of V3+ to determine unpaired electrons.

Diff: 3

Section: 8.3

26) Which metal ion is expected to have the electron configuration [Kr]4d4?

A) Mn2+

B) Ru2+

C) Zr2+

D) Mo2+

E) Sr2+

Hint: Remember that positive ions lose electrons.

Diff: 3

Section: 8.3

27) Which element below has three valence electrons in its Lewis symbol?

A) gallium

B) fluorine

C) iron

D) nickel

E) sulfur

Diff: 1

Section: 8.4

28) Which species below has eight valence electrons in its Lewis symbol?

A) Ar+

B) F+

C) Mg+

D) S2−

E) Si

Diff: 1

Section: 8.4

29) Which species below has the lowest number of valence electrons in its Lewis symbol?

A) Ar+

B) Ga+

C) Mg2+

D) S2−

E) F−

Diff: 2

Section: 8.4

30) Which element below includes six electrons in its Lewis symbol?

A) Cr

B) Ni

C) Ba

D) Se

E) As

Diff: 1

Section: 8.4

31) Which species has the most valence electrons in its Lewis symbol?

A) Ar+

B) Ga+

C) Ba2+

D) Br

E) O2−

Diff: 2

Section: 8.4

32) The atoms in the nitrogen molecule, N2, are held together by

Hint: Draw the Lewis structure of the molecule.

A) a single covalent bond.

B) a double covalent bond.

C) a triple covalent bond.

D) an ionic bond.

E) a magnetic dipole bond.

Diff: 3

Section: 8.5

33) The atoms in the oxygen molecule, O2, are held together by

Hint: Draw the Lewis structure of the molecule.

A) a single covalent bond.

B) a double covalent bond.

C) a triple covalent bond.

D) an ionic bond.

E) a magnetic dipole bond.

Diff: 3

Section: 8.5

34) A covalent bond is characterized by two properties. These are

A) the energy of the bond and the location of the electron pair in the bond.

B) the energy of the separated atoms and the energy of the new bond.

C) the bond length and the polarity of the covalent bond.

D) the bond length and the bond energy.

E) the energy released when a bond forms, and the electronegativities of the two atoms.

Diff: 2

Section: 8.5

35) The Lewis symbol for the carbon atom has ________ valence electrons. The number of covalent bonds that carbon usually forms to complete its valence shell and obey the octet rule is ________.

A) 4, 1

B) 4, 2

C) 2, 4

D) 4, 3

E) 4, 4

Diff: 2

Section: 8.5

36) The Lewis symbol for the nitrogen atom has ________ valence electrons. The number of covalent bonds that nitrogen usually forms to complete its valence shell and obey the octet rule is ________.

A) 5, 1

B) 5, 2

C) 3, 4

D) 5, 3

E) 5, 4

Diff: 2

Section: 8.5

37) Which statement below is true?

Hint: Consider how properties of a covalent bond are measured.

A) The two electrons in a single covalent bond must be paired as required by the Pauli exclusion principle.

B) One mole of hydrogen atoms is more stable than one mole of hydrogen molecules.

C) The buildup of electron density between two atoms repels each nucleus, making them less stable.

D) The bond energy is the minimum energy required to bring about pairing of the electrons in a covalent bond.

E) As the distance between the nuclei decreases when forming a covalent bond, there is a corresponding decrease in the probability of finding both electrons near either nucleus.

Diff: 3

Section: 8.5

38) Which bond is the most polar?

A) H–C

B) H–Cl

C) H–P

D) H–S

E) H–Se

Diff: 1

Section: 8.6

39) Which bond is the most polar?

A) H–C

B) H–S

C) H–P

D) H–O

E) H–Se

Diff: 1

Section: 8.6

40) Which bond is the most polar?

Hint: Consider the electronegativity values of each element.

A) Al–I

B) Si–I

C) Al–Cl

D) Si–Cl

E) Si–P

Diff: 3

Section: 8.6

41) Which bond is the most polar?

Hint: Consider the electronegativity values of each element.

A) Br–Br

B) S–O

C) C–P

D) C–O

E) B–O

Diff: 3

Section: 8.6

42) Which bond is the most polar?

Hint: Consider the electronegativity values of each element.

A) B–C

B) C–N

C) C–O

D) Si–O

E) C–C

Diff: 3

Section: 8.6

43) Arrange the following in terms of increasing electronegativity values.

Ga, P, As, Ge

A) Ga < Ge < As < P

B) Ge < Ga < As < P

C) P < Ge < Ga < As

D) P < As < Ge < Ga

E) As < Ge < As <Ga

Diff: 2

Section: 8.6

44) Which bond is the least polar?

A) H–C

B) H–N

C) H–P

D) H–O

E) H–F

Diff: 2

Section: 8.6

45) Based on electronegativity considerations, which species is predicted to be the strongest oxidizing agent?

A) Ne

B) Kr

C) Br2

D) Cl2

E) S

Diff: 2

Section: 8.6

46) Arrange the following in terms of decreasing electronegativity values.

P, O, As, S

Hint: Remember to refer to your periodic table for trends.

A) P > O > As > S

B) As > P > S > O

C) S > O > P > As

D) As > S > P > O

E) O > S > P > As

Diff: 3

Section: 8.6

47) In the chlorine monoxide molecule, chlorine has a charge of +0.167 e−. If the bond length is 154.6 pm, calculate the dipole moment of the molecule in debyes.

Hint: Watch your units carefully with this problem.

A) 3.11 D

B) 2.30 D

C) 0.167 D

D) 1.24 D

E) 1.65 D

Diff: 3

Section: 8.6

48) In the nitrogen monoxide molecule, the dipole moment is 0.16 D and the bond length is 115 pm. Determine the sign and magnitude of the charge on the oxygen atom.

Hint: Watch your units carefully with this problem.

A) -1.3 e−

B) -0.029 e−

C) -0.097 e−

D) -0.72 e−

E) +1.3 e−

Diff: 3

Section: 8.6

49) Based on the "best" Lewis structure after applying formal charge considerations, how many non-bonding valence electrons are around the nitrogen atom in the nitrate ion?

A) 0

B) 2

C) 4

D) 6

E) 8

Diff: 2

Section: 8.7

50) Draw the Lewis structure for HClO2 from the skeletal structure presented below. If the valence shells are filled to the usual limit (a maximum of 8), how many nonbonding valence electrons are in the molecule?

H O Cl O

A) 14

B) 16

C) 18

D) 20

E) 26

Diff: 1

Section: 8.7

51) Draw the Lewis structure for HNO2. If the valence shells are filled to the usual limit (a maximum of 8), how many nonbonding valence electrons are in the molecule?

A) 0

B) 2

C) 6

D) 10

E) 14

Diff: 2

Section: 8.7

52) Draw a Lewis structure for CH2Cl2. Based on this Lewis structure, the calculated value for the formal charge on the carbon atom is

A) 0.

B) 4+.

C) 2+.

D) 2−.

E) 4− .

Diff: 2

Section: 8.7

53) Draw a Lewis structure for NH4+. Based on this Lewis structure, the calculated value for the formal charge on the nitrogen atom is

A) 0.

B) 4+.

C) 1+.

D) 1−.

E) 4− .

Diff: 2

Section: 8.7

54) When the fluoride ion reacts with a BF3 molecule (a molecule in which there are no multiple bonds), an ion is formed in which the boron atom is the central atom. The bond between the boron trifluoride and the fluoride ion is

A) an ionic bond.

B) a regular covalent bond, where both species contribute 1 electron to the bond.

C) a coordinate covalent bond.

D) a resonance hybrid bond.

E) a bond where two atoms share one electron instead of two.

Diff: 1

Section: 8.7

55) How many coordinate covalent bonds are formed in the reaction between H+ and NH3 and what is the formula of the only product?

A) 0, NH3

B) 1, NH4+

C) 2, NH4+

D) 3, N2H4+

E) 4, N2H5+

Diff: 2

Section: 8.7

56) How many coordinate covalent bonds are formed in this product in the reaction between F− and BF3 and what is the formula of the product?

A) BF3, 0

B) BF4−, 1

C) BF4−, 2

D) BF52−, 3

E) B2F5−, 4

Diff: 2

Section: 8.7

57) When the hydrogen ion combines with a water molecule, a coordinate covalent bond is formed because

A) a monatomic species reacts with a molecular substance.

B) both electrons are supplied by the oxygen atom of the water molecule.

C) a nonpolar bond is formed.

D) a polyatomic species is formed.

E) a bond is formed where two atoms share one electron instead of two.

Diff: 1

Section: 8.7

58) Draw the Lewis structure for hydrogen peroxide, H2O2. Based on this structure, how many polar bonds and non-polar bonds are present?

A) 3 polar bonds and no non-polar bonds

B) 2 polar bonds and 1 non-polar bond

C) 1 polar bond and 2 non-polar bonds

D) no polar bonds and 3 non-polar bonds

E) 2 polar bonds and 2 non-polar bonds

Diff: 2

Section: 8.7

59) How many lone pairs of electrons are in the nitric oxide molecule?

A) 2

B) 1

C) 0

D) 3

E) 4

Diff: 2

Section: 8.7

60) The formal charge on the nitrogen atom in the nitrate ion (NO3−) is

A) 3−

B) 0

C) 1+

D) 3+

E) 5+

Diff: 1

Section: 8.7

61) The formal charge on the carbon atom in the carbonate (CO32−) ion is

A) 2−.

B) 0.

C) 1+.

D) 2+.

E) 4+.

Diff: 1

Section: 8.7

62) Based on the "best" Lewis structure after applying formal charge considerations, how many non-bonding valence electrons are around the carbon atom in the CO molecule?

A) 0

B) 2

C) 4

D) 6

E) 8

Diff: 2

Section: 8.7

63) Based on the "best" Lewis structure after applying formal charge considerations, how many non-bonding valence electrons are around the nitrogen atom in the nitrite ion (NO2−)?

A) 0

B) 2

C) 4

D) 6

E) 8

Diff: 2

Section: 8.7

64) Based on the "best" Lewis structure after applying formal charge considerations, how many non-bonding valence electrons are around the nitrogen atom in the HCN molecule?

A) 0

B) 1

C) 2

D) 4

E) 6

Diff: 2

Section: 8.7

65) Draw the Lewis structure for HClO3. If the valence shells are filled to the usual limit (a maximum of 8), what is the formal charge on the chlorine atom?

Hint: Do not forget about double or triple bonds potentially being in the Lewis structure.

A) −1

B) 0

C) +1

D) +2

E) +3

Diff: 3

Section: 8.7

66) Draw the Lewis structure for H2SeO3. If the valence shells are filled to the usual limit (a maximum of 8), how many nonbonding valence electrons are there in the molecule?

Hint: Do not forget about double or triple bonds potentially being in the Lewis structure.

A) 14

B) 16

C) 18

D) 20

E) 28

Diff: 3

Section: 8.7

67) Draw a Lewis structure for H3C–NH2. Based on this Lewis structure, the calculated value for the formal charge on the nitrogen atom is

Hint: Do not forget about double or triple bonds potentially being in the Lewis structure.

A) 2−.

B) 3+.

C) 3−.

D) 2+.

E) 0.

Diff: 3

Section: 8.7

68) Draw the Lewis structure for the H2CO molecule. Based on this structure, how many polar bonds and non-polar bonds are present?

Hint: Do not forget about double or triple bonds potentially being in the Lewis structure.

A) 3 polar bonds and no non-polar bonds

B) 2 polar bonds and 1 non-polar bond

C) 1 polar bond and 2 non-polar bonds

D) no polar bonds and 3 non-polar bonds

E) 2 polar bonds and 2 non-polar bonds

Diff: 3

Section: 8.7

69) Draw the Lewis structure for the F2CO molecule. Based on this structure, how many polar bonds and non-polar bonds are present?

Hint: Do not forget about double or triple bonds potentially being in the Lewis structure.

A) 3 polar bonds and no non-polar bonds

B) 2 polar bonds and 1 non-polar bond

C) 1 polar bond and 2 non-polar bonds

D) no polar bonds and 3 non-polar bonds

E) 2 polar bonds and 2 non-polar bonds

Diff: 3

Section: 8.7

70) Which species has the same bond order as the carbon monoxide molecule?

A) ClF

B) CO32−

C) CO2

D) NO+

E) O3

Diff: 2

Section: 8.7

71) The metaphosphate ion, PO3−, is the structural analog of the NO3− ion with respect to the arrangement of the atoms in the ion. After drawing the "best" Lewis structure for the metaphosphate ion based on formal charge considerations, what is the formal charge on the phosphorus atom?

Hint: Do not forget about double or triple bonds potentially being in the Lewis structure.

A) 1−

B) 0

C) 1+

D) 2+

E) 5+

Diff: 3

Section: 8.7

72) Draw the Lewis structure for H2SeO4. If the valence shells are filled to the usual limit (a maximum of 8), what is the formal charge on the Se?

Hint: Do not forget about double or triple bonds potentially being in the Lewis structure.

A) 2−

B) 1−

C) 0

D) 1+

E) 2+

Diff: 3

Section: 8.7

73) Draw the Lewis structure for H2SO4. Experimentally, it is known that the S-O bonds are shorter than the S-OH bonds. The likely explanation for this is

Hint: Consider the number of bonded electrons in each bond.

A) the electrons in the S-OH bonds are not equally shared.

B) the sulfur atom is not as electronegative as the oxygen atom.

C) the hydrogen atom is weakly bonded to the other S atom.

D) there are no lone pairs on the hydrogen atom.

E) the S-O bond has a greater bond order than the S-OH bond.

Diff: 3

Section: 8.7

74) Draw the Lewis structure for HClO3. If the valence shells are filled to the usual limit (a maximum of 8), what is the sum of the absolute values of all the formal charges in the molecule?

Hint: Minimize formal charge to ensure you have the most stable or "best" Lewis structure.

A) 0

B) 1

C) 2

D) 3

E) 4

Diff: 3

Section: 8.7

75) Draw a Lewis structure for CH3NO2. Based on the Lewis structure, what is the formal charge on the nitrogen atom for the most favorable structure?

Hint: Minimize formal charge to ensure you have the most stable or "best" Lewis structure.

A) −2

B) −1

C) 0

D) +1

E) +2

Diff: 3

Section: 9.8

76) Draw the "best" Lewis structure for sulfur trioxide based on formal charge considerations. The number of resonance structures is

A) 1 (no resonance structures).

B) 2.

C) 3.

D) 4.

E) 5.

Diff: 2

Section: 8.8

77) How many resonance structures, if any, can be drawn for the nitrate ion?

A) 1 (no resonance structures).

B) 2.

C) 3.

D) 4.

E) 5.

Diff: 2

Section: 8.8

78) How many resonance structures, if any, can be drawn for the nitrite ion?

A) 1 (no resonance structures)

B) 2

C) 3

D) 4

E) 5

Diff: 2

Section: 8.8

79) Based on the "best" Lewis structure from formal charge considerations, how many resonance structures, if any, can be drawn for the PO43− ion?

A) 1 (no resonance structures)

B) 2

C) 3

D) 4

E) 5

Diff: 2

Section: 8.8

80) The metaphosphate ion, PO3−, is the structural analog of the NO3− ion with respect to the arrangement of the atoms in the ion. After drawing the "best" Lewis structure for the metaphosphate ion based on formal charge considerations, what is the number of resonance structures for this "best" Lewis structure?

Hint: Use formal charge to help determine how many realistic resonance structures there are for the metaphosphate ion.

A) 1 (no resonance structures)

B) 2

C) 3

D) 4

E) 5

Diff: 3

Section: 8.8

81) How many resonance structures, if any, can be drawn for the N2O5 molecule?

Hint: Consider formal charge when determining if the resonance structure is valid.

A) 1 (no resonance structures)

B) 2

C) 3

D) 4

E) 5

Diff: 3

Section: 8.8

82) The chlorite ion has a Lewis structure that is based on the skeletal structure shown below:

O=Cl -O

Draw the Lewis structure minimizing the formal charges present in the structure. Based on this structure, the chlorite ion has

Hint: Consider all possible resonance structures when working with formal charge.

A) two single bonds, the sum of absolute values of formal charges = 3, and no resonance hybrids.

B) one single and one double bond, sum of absolute values of formal charges = 5, and two contributing resonance hybrids.

C) one single and one double bond, sum of absolute values of formal charges = 3, and two contributing resonance hybrids.

D) two double bonds, sum of absolute values of formal charges = 1, and no resonance hybrids.

E) one single and one double bond, sum of absolute values of formal charges = 1, and two contributing resonance hybrids.

Diff: 3

Section: 8.8

83) The carbonate ion has the skeletal structure as shown below. Complete the Lewis structure by filling in the bonds and the remaining valence electrons which are not involved in bonds. Which statement made about the carbonate ion is true?

O

|

O - C - O

Based on its Lewis structure, the carbonate ion should have ________ resonance hybrids and ________ atoms with formal charges on them in its structure.

Hint: Consider the possible places the double bond could go when determining resonance of carbonate.

A) 2, 2

B) 3, 1

C) 3, 2

D) 2, 1

E) no, 2

Diff: 3

Section: 8.8

84) Draw the Lewis structure of CH3NO2.

Based on its structure, the nitromethane molecule should have ________ resonance hybrids and ________ atoms with formal charges on them.

Hint: Consider possible double and triple bonds in the structure and be sure to minimize formal charge.

A) no, 2

B) no, 3

C) 2, 2

D) 2, 3

E) 3, 2

Diff: 3

Section: 8.8

85) Alcohols belong to a class of organic compounds in which a hydrogen atom in a hydrocarbon is replaced by which of the following?

A) an Ar atom

B) a NO group

C) an OH group

D) an F atom

E) a metal ion

Diff: 2

Section: 8.9

86) An organic compound has the formula CH3CH2OH. Draw out the Lewis structure. What functional group is present in this structure?

A) a ketone

B) an alcohol

C) an ether

D) a halogen

E) an aldehyde

Diff: 2

Section: 8.9

87) An organic compound has the formula CH3CH2NH2. Draw out the Lewis structure. What functional group is present in this structure?

A) an amine

B) a ketone

C) an ether

D) a halogen

E) an aldehyde

Diff: 2

Section: 8.9

88) Ketones belong to a class of organic compounds in which a hydrogen atom in a hydrocarbon is replaced by which of the following?

A) an argon atom

B) a doubly bonded sulfur atom

C) a doubly bonded carbon atom

D) a triply bonded nitrogen atom

E) a doubly bonded oxygen atom

Diff: 2

Section: 8.9

89) Organic acids are characterized by the presence of which particular group of atoms?

A) a hydronium ion, H3O+

B) an extra —CH2 group

C) a doubly bonded carbon atom

D) a carboxyl group, —CO2H

E) a doubly bonded hydrogen atom

Diff: 2

Section: 8.9

90) How many lone pairs of electrons are on the group of atoms that characterize organic acids?

A) 6

B) 4

C) 3

D) 2

E) 1

Diff: 2

Section: 8.9

91) Amines are weak organic bases and can be considered derived from which compound?

A) H2O

B) SO2

C) NH3

D) CH3C1

E) CO2

Diff: 2

Section: 8.9

92) Based on the ∆H°f data given below, which compound is the least stable?

H2S(g), −20.6 kJ/mol; N2H4(g), +94.5 kJ/mol;

N2O(g), +81.6 kJ/mol; Na2O2(g), −46.4 kJ/mol

Diff: 2

Section: 8.1

93) The attraction between positive and negative ions in an ionic compound is referred to as the ________.

Diff: 1

Section: 8.2

94) The Born-Haber cycle, which is an enthalpy diagram that can be used to calculate the lattice energy of an ionic compound, has how many endothermic steps?

Diff: 2

Section: 8.2

95) Of the following solids, KI, KBr, KCl, and KF, the solid with the highest melting point is ________.

Diff: 1

Section: 8.2

96) Of the following solids, NaI, NaF, MgO, MgCl2, and KF, the solid with the highest melting point is ________.

Diff: 2

Section: 8.2

97) Which of the following cations, Ca3+, Ca2+, Ca+, is the most likely to form a stable ionic compound?

Diff: 1

Section: 8.3

98) What is the electron configuration of the metal ion in Cr2(SO4)3?

Diff: 2

Section: 8.3

99) When the Sn2+ ion is formed, the electrons are removed from which set of orbital(s)?

Diff: 2

Section: 8.3

100) For the transition elements, the first electrons lost are the s electrons of the outer shell. If additional electrons are lost, they come from the ________.

Diff: 2

Section: 8.3

101) Which of these ions shown, Fe3+, P3− , B−, Ca2+, has the same number of valence electrons as a germanium atom?

Diff: 2

Section: 8.4

102) Which of these species, Fe3+, P3−, B−, Ca2+, has the most valence electrons?

Diff: 2

Section: 8.4

103) Draw the Lewis dot symbol for the S−2 ion.

Diff: 2

Section: 8.4

104) Draw the Lewis dot symbol for the Ca2+ ion.

Diff: 2

Section: 8.4

105) The amount of energy released when a covalent bond is formed is called its ________.

Diff: 1

Section: 8.5

106) The Lewis symbol for the selenium atom shows ________ valence shell electrons. The number of covalent bonds that selenium normally forms to complete its valence shell and obey the octet rule is ________.

Diff: 2

Section: 8.5

107) The Lewis symbol for the nitrogen atom shows ________ valence shell electrons. The number of covalent bonds that nitrogen normally forms to complete its valence shell and obey the octet rule is ________.

Diff: 2

Section: 8.5

108) The Lewis symbol for the oxygen atom shows ________ valence shell electrons. The number of covalent bonds that oxygen normally forms to complete its valence shell and obey the octet rule is ________.

Diff: 2

Section: 8.5

109) The Lewis symbol for the carbon atom shows ________ valence shell electrons. The number of covalent bonds that carbon normally forms to complete its valence shell and obey the octet rule is ________.

Diff: 2

Section: 8.5

110) The Lewis representation of the PH3 molecule shows ________ covalent bonds, and ________ lone pair on the ________atom.

Diff: 2

Section: 8.5

111) In a covalent bond between S and Cl, which atom carries the partial negative charge?

Diff: 1

Section: 8.6

112) For the HCl molecule, the dipole moment is calculated from q × r. Calculations show that q equals 0.17 electronic charge units. This means the hydrogen carries a charge of ________ and the chlorine a charge of ________.

Diff: 2

Section: 8.6

113) In the C-S molecule, the dipole moment is 1.96 D and the bond length is 153 pm. Determine the sign and magnitude of the charge on the carbon atom.

Hint: Carefully watch your units and signs when completing this problem.

Diff: 3

Section: 8.6

114) In the S-Si molecule, the silicon atom has a charge of 0.187 e− and the bond length is 193 pm. Determine the dipole moment of this molecule.

Hint: Carefully watch your units and signs when completing this problem.

Diff: 3

Section: 8.6

115) A covalent bond has about 50 % ionic character when the electronegativity difference between the atoms is approximately ________.

Diff: 2

Section: 8.6

116) The total number of bonds in the thiocyanate ion, SCN−.

Diff: 2

Section: 8.7

117) The total number of non-bonding electrons in the NF3 molecule is ________.

Diff: 2

Section: 8.7

118) The total number of non-bonding electrons in molecular nitrogen is ________.

Diff: 2

Section: 8.7

119) The total number of non-bonding electrons in molecular oxygen is ________.

Diff: 2

Section: 8.7

120) The formal charge on the most stable Lewis structure for the central oxygen in ozone, O3, is ________.

Diff: 2

Section: 8.7

121) Draw the most favorable Lewis structure for the NCO molecule.

Hint: Remember to minimize formal charge for your most favorable Lewis structure.

: N≡ C —:

Diff: 3

Section: 8.7

122) Draw the most favorable Lewis structure for the NCO molecule. The formal charge on the oxygen molecule is ________.

Hint: Consider double and triple bonds for your structure.

Diff: 3

Section: 8.7

123) In PCl5, the P atom violates the octet rule by sharing ________electrons. This is because the valence shell for elements in which n = 3 can hold a maximum of _______electrons.

Diff: 2

Section: 8.7

124) For bonds formed between the same elements, as the bond order increases, the bond length ________ and the bond energy ________.

Diff: 2

Section: 8.7

125) Draw the Lewis structure for the sulfite ion. When formal charge considerations are fully considered and adjustments made, if necessary, how many resonance structures, if any, can be drawn for this ion ?

Diff: 2

Section: 8.8

126) Draw the Lewis structure for the sulfur dioxide molecule. When formal charge considerations are fully considered and adjustments made, if necessary, how many resonance structures, if any, can be drawn?

Diff: 2

Section: 8.8

127) Draw the Lewis structure for SOCl2. Considering formal charges, there will be ________ bonds and ________ resonance structures in the molecule.

Diff: 2

Section: 8.8

128) The total number of valence electrons used in the C4H10 molecule is ________ so there is a total of ________ covalent bonds in this hydrocarbon molecule ________.

Diff: 2

Section: 8.9

129) When (CH3)2NH is dissolved in water, the resulting solution is slightly ________. (acidic or basic)

Diff: 2

Section: 8.9

130) In the family of compounds called aldehydes, the carbon in C=O group is directly attached to another ________ atom and a ________ atom.

Diff: 2

Section: 8.9

131) Compounds with negative heats of formation tend to be unstable with respect to decomposition into its elements.

Diff: 1

Section: 8.1

132) Bond formation tends to decrease the potential energy of the two atoms involved in forming the bond.

Diff: 2

Section: 8.1

133) The lattice energy is the energy released if gaseous ions combine to form one mole of a solid ionic compound.

Diff: 2

Section: 8.2

134) The lattice energy of Al2O3 is greater than that for AlCl3, which implies that the melting point of Al2O3 should be greater than that of AlCl3.

Diff: 2

Section: 8.2

135) When electrons are removed from a given shell to form a cation, they come from the highest energy occupied subshell first before any are removed from a lower-energy subshell.

Diff: 2

Section: 8.3

136) The octet rule does not work well for ions such as Br− and O2−.

Diff: 2

Section: 8.3

137) The electron configuration for the Sb3+ ion is [Kr] 4d10 5s2.

Diff: 2

Section: 8.3

138) In barium sulfide, the number of electrons in the Lewis symbol for the barium is zero, while for the sulfide ion, it is 2.

Diff: 2

Section: 8.4

139) There are covalent bonds in the Lewis structure for magnesium fluoride.

Diff: 1

Section: 8.4

140) The energy of the bond holding two covalently bound atoms together is independent of the distance between the two atoms.

Diff: 2

Section: 8.5

141) The maximum number of electrons that can be shared by a given pair of atoms to form bonds in a molecule that has two or more atoms is 4.

Diff: 2

Section: 8.5

142) The F–O bond is more polar than the H–O bond.

Diff: 2

Section: 8.6

143) Br2 will oxidize NaI, but not NaF or NaCl______.

Hint: Consider the polarity of the bonds.

Diff: 3

Section: 8.6

144) When oxygen reacts with CuS(s) the gas that is formed is SO2.

Diff: 2

Section: 8.6

145) When the Lewis structure for the nitrite ion is drawn, based on formal charge considerations, there are four bonds in its structure.

Diff: 2

Section: 8.7

146) Based on the skeleton structure, N—N—O, the magnitude of the sum of the partial charges for N2O has to be 4.

Hint: Write out the full Lewis structure and remember your double and triple bonds.

Diff: 3

Section: 8.7

147) There are two resonance structures for the ion, HCOO−.

Diff: 2

Section: 8.8

148) There are three resonance structures for the hydrogen carbonate ion.

Hint: Draw the Lewis structure and look at the double bonds. How could they be re-arranged?

Diff: 3

Section: 8.8

149) The sulfite ion has three resonance structures based on lowest formal charges.

Hint: Draw the Lewis structure and look at the double bond. Where else could it be in the structure?

Diff: 3

Section: 8.8

150) A layer of ozone (O3) in the stratosphere, a region of the atmosphere extending from about 45 to 55 km altitude, absorbs most of the incoming UV radiation from sunlight, protecting life on the surface of the earth.

Diff: 1

Section: Chemistry and Current Affairs 8.1

151) Absorption of UV radiation causes constituents of DNA called pyrimidine bases to undergo certain types of undesirable chemical reactions.

Diff: 1

Section: Chemistry and Current Affairs 8.1

152) The UV radiation that passes through the ozone layer and is most damaging to structures like DNA, has wavelength ranging from 280 to 320 nm and is also called "UV-B" radiation.

Diff: 2

Section: Chemistry and Current Affairs 8.1

153) Gases like chlorofluorocarbons (CFCs) are believed to be responsible for protecting the ozone layer from being destroyed by chemicals that produce free radicals.

Diff: 2

Section: Chemistry and Current Affairs 8.1

154) A compound that has the formula, C4H8O, should have 12 bonds in its Lewis structure.

Diff: 2

Section: 8.9

155) There are five lone pairs of electrons in the acetate ion, CH3COO−.

Diff: 2

Section: 8.9

156) Use the data to calculate the lattice energy of sodium chloride.

Na(s) → Na(g) ΔH1 = +108 kJ

½ Cl2(g) → Cl(g) ΔH2 = +120 kJ

Na(g) → Na+(g) + e− ΔH3 = +496 kJ

Cl(g) + e− → Cl−(g) ΔH4 = −349 kJ

Na(s) + ½ Cl2(g) → NaCl(s) ΔH°f = −411 kJ

Diff: 2

Section: 8.2

157) Use the data to calculate the lattice energy of lithium chloride.

Li(s) → Li(g) ΔH1 = +155.2 kJ

½ Cl2(g) → Cl(g) ΔH2 = +120 kJ

Li(g) → Li+(g) + e− ΔH3 = +520 kJ

Cl(g) + e− → Cl−(g) ΔH4 = −349 kJ

Li(s) + ½ Cl2(g) → LiCl(s) ΔH°f = −408.8 kJ

Diff: 2

Section: 8.2

158) Use the data to calculate the lattice energy of sodium bromide.

Na(s) → Na(g) ΔH1 = +108 kJ

½ Br2(g) → Br(g) ΔH2 = +193 kJ

Na(g) → Na+(g) + e− ΔH3 = +496 kJ

Br(g) + e− → Br−(g) ΔH4 = −325 kJ

Na(s) + ½ Br2(g) → NaBr(s) ΔH°f = −361 kJ

Diff: 2

Section: 8.2

159) Which of the following electron configuration is correct for the element or ion shown with it?

Hint: Consider your exceptions to the orbital filling rules.

A) Se: 1s2 2s2 2p6 3s2 3p6 3d10 4s3 4p3

B) Zn: 1s2 2s2 2p6 3s2 3p6 4s3 3d9

C) Br−: 1s2 2s2 2p6 3s2 3p6 4s1 3d11 4p6

D) Cu2+: 1s2 2s2 2p5 3s3 3p6 3d9

E) Sr2+: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

Diff: 3

Section: 8.3

160) Arrange the following in terms of increasing (from left to right) number of electrons in the valence shell.

Bi3+, Sn2+, Te2−, I, Rb+

A) Rb+, Sn2+, Te2−, I, Bi3+

B) Bi3+, Sn2+, Te2−, I, Rb+

C) I, Rb+, Bi3+, Te2−, Sn2+

D) Sn2+, Bi3+, Te2−, I, Rb+

E) Rb+, Sn2+, Bi3+, I, Te2−

Diff: 2

Section: 8.4

161) Arrange the following in order of increasing bond polarity:

H-Se, N-Cl, P-Cl, C-Br, Si-O, As-Br.

Hint: Consider your electronegativity values when determining bond polarity between elements.

A) H-Se, N-Cl, P-Cl, C-Br, Si-O, As-Br

B) As-Br, N-Cl, P-Cl, C-Br, Si-O, H-Se

C) P-Cl, N-Cl, H-Se, C-Br, Si-O, As-Br

D) Si-O, N-Cl, P-Cl, C-Br, H-Se, As-Br

E) N-Cl, H-Se, C-Br, As-Br, P-Cl, Si-O

Diff: 3

Section: 8.6

162) A student drew four possible Lewis structures for HBrO4

Complete these Lewis structures presented above by filling in the remaining valence electrons that are not in the bonds. Based on these structures, the preferred structure would be the structure shown as ________ in which the sum of the absolute values of the formal charges on all the atoms is ________.

Hint: When determining remaining valence electrons remember to account for the electrons already present in the bonds.

A) A, 4

B) B, 2

C) C, 0

D) D, 6

E) D, 0

Diff: 3

Section: 8.7

163) The nitrogen monoxide molecule is a unique molecule as it forms a stable radical under normal atmospheric conditions. Using formal charges develop the best Lewis structure for the nitrogen monoxide molecule. How many non-bonding electrons does the nitrogen atom have on it?

Hint: Some molecules are free radicals and have an unpaired electron present.

Diff: 3

Section: 8.7

164) Based on electronegativity considerations, which species should be the strongest oxidizing agent?

A) Xe

B) As

C) Br2

D) I2

E) Sb

Diff: 2

Section: 8.6

165) Based on electronegativity considerations, which species should be the strongest oxidizing agent?

A) O2

B) F2

C) N2

D) Cl2

E) S

Diff: 2

Section: 8.6

166) Based on electronegativity considerations, which species should be the weakest oxidizing agent?

A) F2

B) Cl2

C) Br2

D) I2

E) At2

Diff: 2

Section: 8.6

167) Complete the Lewis structures for COCl2 and SOCl2 using the skeletal structure shown below, being sure to follow the procedure for minimizing the sum of the absolute values for the formal charges where the octet rule need not be followed. Based on the complete structures, which statement below is true?

Hint: Consider the potential for multiple bonds on the molecules and be sure to account for all valence electrons.

A) The COCl2 exhibits formal charges but no resonance hybrids, while the SOCl2 exhibits both formal charges and resonance hybrids.

B) The COCl2 exhibits both residual formal charges and resonance hybrids, while the SOCl2 exhibits formal charges but no resonance hybrids.

C) The SOCl2 exhibits formal charges but no resonance hybrids, while the COCl2 exhibits resonance hybrids but no formal charges.

D) The SOCl2 exhibits both formal charges and resonance hybrids, while the COCl2 exhibits resonance hybrids but no formal charges.

E) The formal charges on all the atoms are zero, and the molecules have no resonance hybrids.

Diff: 3

Section: 8.8

168) Draw the Lewis structure for ClSO3−, optimized for formal charge considerations, showing all

resonance structures.

Diff: 2

Section: 8.8

169) The fluorophosphonate ion, FPO32−, has a structure similar to that of the orthophosphate ion, the difference being that one of the oxygen atoms has been replaced by a fluorine atom. Draw the Lewis structure for this ion, optimized for formal charge considerations, showing all resonance structures.

Diff: 2

Section: 8.8

170) Which of the following compounds is expected to produce an acidic solution when dissolved in water?

A) CH3CH2CH2CH3

B) CH3CH2OH

C) CH3COCH3

D) CH3NHCH3

E) CH3CH2CO2H

Diff: 2

Section: 8.9

171) Which of the following compounds is expected to produce a basic solution when dissolved in water?

A) CH3CH2CH2CH3

B) CH3CH2OH

C) CH3COCH3

D) CH3CH2NH2

E) CH3CH2CO2H

Diff: 2

Section: 8.9

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