CHEMISTRY 161
HW CHAPTER 8B
90, 94, 96, 98, 100, 108, 114, 118, 120, 126
8-90
For formamide, HCONH2 there are 18 valence electrons and the Lewis structures
with formal charges assigned to the atoms are
Notice that the double bond can technically be C O or C N and still satisfy the
octet rule, but that formal charge predicts the structure with C O will dominate.
For methylformate, HCO2CH3 there are 24 valence electrons, and the Lewis
structures with formal charges assigned to the atoms are
Formal charge predicts that the double bond is between the C and the terminal O
atom.
8-94
In order to expand beyond 8 electrons, the central atom must have available empty d
orbitals. In addition, expansion of the valence shell is typical for these elements
when forming bonds with highly electronegative elements such as fluorine, oxygen,
and chlorine. Finally, the expansion of the valence shell reduces formal charges on
the atoms in the structure.
8-96
As we have seen in the textbook, we draw Lewis structures with expanded octets
when it is required to connect all the atoms in the molecules and to minimize formal
charges on the atoms. It is possible that an atom that has the ability to expand its
octet would not need to do so when forming a covalently bonded compound.
8-98
8-100
(a) POCl3 has 32 valence electrons, and its Lewis structure with formal charges
assigned to its atoms is
We can reduce the formal charges on P and O through a double bond.
There are 10 e– in five covalent bonds around phosphorus in POCl3.
(b) H3PO4 has 32 valence electrons, and its Lewis structure with formal charges
assigned to its atoms is
To reduce the formal charges on P and O, we form a double bond between them.
This gives 10 e– in five covalent bonds around phosphorus in H3PO4.
(c) H3PO3 has 26 valence electrons, and its Lewis structure with formal charges
assigned to its atoms is
All the formal charges on the atoms in this structure are zero. There are 6 e– in three
covalent bonds around phosphorus in H3PO3.
(d) PF6– has 48 valence electrons, and its Lewis structure with formal charges
assigned to its atoms is
All the formal charges on the F atoms in this structure are zero. There are 12 e – in
six covalent bonds around phosphorus in PF6–.
8-108
(a) NO has (1N  5 e–) + (1O  6 e–) = 11 e–
(b) NO2 has (1N  5 e–) + (2O  6 e–) = 17 e–
(c) NO3 has (1N  5 e–) + (3O  6 e–) = 23 e–
(d) N2O4 has (2N  5 e–) + (4O  6 e–) = 34 e–
(e) N2O5 has (2N  5 e–) + (5O  6 e–) = 40 e–
The odd-electron molecules are NO, NO2, and NO3 (a–c).
8-114
The bond in O2 is a double bond.
The bond in O3 is 1.5 due to resonance.
Because of resonance the oxygen– oxygen bond in ozone (O3) is longer than the bond
8-118
For CO, the bond order for the C — O bond is 3.0.
For CO2, the bond order for the C — O bond is 2.0.
For CO32–, the bond order for the C — O bond is 1.33 due to resonance.
In order of increasing bond length: CO < CO2 < CO32–.
8-120
For CO, the bond order for the C — O bond is 3.0.
For CO2, the bond order for the C — O bond is 2.0.
For CO32–, the bond order for the C — O bond is 1.33 due to resonance.
In order of increasing bond energy: CO32– < CO2 < CO.
8-126
(a)
∆Hrxn = [(2  799 kJ/mol) + (1  436 kJ/mol)] + [– (2  463 kJ/mol) – (1  1072 kJ/mol)] = 36 kJ
(b)
∆Hrxn = [(1  941 kJ/mol) + (1  495 kJ/mol)] + [– (2  607 kJ/mol)] = 222 kJ
(c)
∆Hrxn = [(1  719 kJ/mol) + (2  799 kJ/mol)] + [– (2  1072 kJ/mol)] = 173 kJ