Chapter 10

Write Lewis symbols for the following atoms. (a) \(\mathrm{Kr}\); (b) Ge; (c) \(\mathrm{N} ;\) (d) Ga; (e) As; (f) Rb.

Write Lewis symbols for the following ions. (a) \(\mathrm{H}^{-}\) (b) \(\operatorname{Sn}^{2+} ;\) (c) \(\mathrm{K}^{+} ;\) (d) \(\mathrm{Br}^{-} ;\) (e) \(\mathrm{Se}^{2-} ;\) (f) \(\mathrm{Sc}^{3+}\).

Write plausible Lewis structures for the following molecules that contain only single covalent bonds. (a) \(\mathrm{FCl} ;\) (b) \(\mathrm{I}_{2} ;\) (c) \(\mathrm{SF}_{2} ;\) (d) \(\mathrm{NF}_{3} ;\) (e) \(\mathrm{H}_{2} \mathrm{Te}\).

Each of the following molecules contains at least one multiple (double or triple) covalent bond. Give a plausible Lewis structure for (a) \(\mathrm{OCS} ;\) (b) \(\mathrm{CH}_{3} \mathrm{CHO}\) (c) \(\mathrm{F}_{2} \mathrm{CO} ;\) (d) \(\mathrm{Cl}_{2} \mathrm{SO} ;\) (e) \(\mathrm{C}_{2} \mathrm{H}_{2}\).

By means of Lewis structures, represent bonding between the following pairs of elements: (a) Cs and \(\mathrm{Br} ;\) (b) \(\mathrm{H}\) and \(\mathrm{Sb} ;\) (c) \(\mathrm{B}\) and \(\mathrm{Cl} ;\) (d) \(\mathrm{Cs}\) and \(\mathrm{Cl}\); (e) Li and O; (f) Cl and L. Your structures should show whether the bonding is essentially ionic or covalent.

Which of the following have Lewis structures that \(d o\) not obey the octet rule: \(\mathrm{NF}_{3}, \mathrm{AlCl}_{3}, \mathrm{SiF}_{6}^{2-}, \mathrm{SO}_{3}, \mathrm{PH}_{4}^{+}\) \(\mathrm{PO}_{4}^{3-}, \mathrm{ClO}_{2} ?\)

Give several examples for which the following statement proves to be incorrect. "All atoms in a Lewis structure have an octet of electrons in their valence shells."

Suggest reasons why the following do not exist as stable molecules: (a) \(\mathrm{H}_{3} ;\) (b) \(\mathrm{HHe} ;\) (c) \(\mathrm{He}_{2}\); (d) \(\mathrm{H}_{3} \mathrm{O}\).

Write Lewis structures for the following ionic compounds: (a) calcium chloride; (b) barium sulfide; (c) lithium oxide; (d) sodium fluoride.

Under appropriate conditions, both hydrogen and nitrogen can form monatomic anions. What are the Lewis symbols for these ions? What are the Lewis structures of the compounds (a) lithium hydride; (b) calcium hydride; (c) magnesium nitride?

Derive the correct formulas for the following ionic compounds by writing Lewis structures. (a) lithium sulfide; (b) sodium fluoride; (c) calcium iodide; (d) scandium chloride.

Each of the following ionic compounds consists of a combination of monatomic and polyatomic ions. Represent these compounds with Lewis structures. (a) \(\mathrm{Al}(\mathrm{OH})_{3}\); (c) \(\mathrm{NH}_{4} \mathrm{F}\); (d) \(\mathrm{KClO}_{3}\); (b) \(\mathrm{Ca}(\mathrm{CN})_{2}\); (e) \(\mathrm{Ba}_{3}\left(\mathrm{PO}_{4}\right)_{2}\).

Both oxidation state and formal charge involve conventions for assigning valence electrons to bonded atoms in compounds, but clearly they are not the same. Describe several ways in which these concepts differ.

Although the notion that a Lewis structure in which formal charges are zero or held to a minimum seems to apply in most instances, describe several significant situations in which this appears not to be the case.

What is the formal charge of the indicated atom in each of the following structures? (a) the central \(\mathrm{O}\) atom in \(\mathrm{O}_{3}\) (b) \(\mathrm{Al}\) in \(\mathrm{AlH}_{4}^{-}\) (c) \(\mathrm{Cl}\) in \(\mathrm{ClO}_{3}^{-}\) (d) \(\sin \sin ^{2} \theta^{-}\) (e) \(\mathrm{Cl}\) in \(\mathrm{Cl} \mathrm{F}_{3}\)

Assign formal charges to the atoms in the following species, and then select the more likely skeletal structure. (a) \(\mathrm{H}_{2} \mathrm{NOH}\) or \(\mathrm{H}_{2} \mathrm{ONH}\) (b) SCS or CSS (c) NFO or FNO (d) \(\mathrm{SOCl}_{2}\) or \(\mathrm{OSCl}_{2}\) or \(\mathrm{OCl}_{2} \mathrm{S}\) (e) \(\mathrm{F}_{3} \mathrm{SN}\) and \(\mathrm{F}_{3} \mathrm{NS}\)

The concept of formal charge helped us to choose the more plausible of the I ewis structures for \(\mathrm{NO}_{2}^{+}\) given in expressions (10.14) and \((10.15) .\) Can it similarly help us to choose a single Lewis structure as most plausible for \(\mathrm{CO}_{2} \mathrm{H}^{+} ?\) Explain.

Show that the idea of minimizing the formal charges in a structure is at times in conflict with the observation that compact, symmetrical structures are more commonly observed than elongated ones with many central atoms. Use \(\mathrm{ClO}_{4}^{-}\) as an illustrative example.

Write acceptable Lewis structures for the following molecules: (a) \(\mathrm{H}_{2} \mathrm{NNH}_{2}\); (b) HOClO; (c) (HO) \(_{2}\) SO; (d) HOOH (e) SO \(_{4}^{2-}\).

The following polyatomic anions involve covalent bonds between O atoms and the central nonmetal atom. Propose an acceptable Lewis structure for each. (a) \(\mathrm{SO}_{3}^{2-} ;\) (b) \(\mathrm{NO}_{2}^{-} ;\) (c) \(\mathrm{CO}_{3}^{2-} ;\) (d) \(\mathrm{HO}_{2}^{-}\).

Represent the following ionic compounds by Lewis structures: (a) barium hydroxide; (b) sodium nitrite; (c) magnesium iodate; (d) aluminum sulfate.

Write a plausible Lewis structure for crotonaldehyde, \(\mathrm{CH}_{3} \mathrm{CHCHCHO},\) a substance used in tear gas and insecticides.

Write a plausible Lewis structure for \(\mathrm{C}_{3} \mathrm{O}_{2},\) a substance known as carbon suboxide.

Which of the following molecules would you expect to have a resultant dipole moment \((\mu) ;\) (a) \(\mathrm{F}_{2}\), (b) \(\mathrm{NO}_{2}\) (c) \(\mathrm{BF}_{3},\) (d) \(\mathrm{HBr},\) (e) \(\mathrm{H}_{2} \mathrm{CCl}_{2}\), (f) \(\mathrm{SiF}_{4},\) (g) OCS? Explain.

What is the percent ionic character of each of the following bonds: (a) \(\mathrm{S}-\mathrm{H} ;\) (b) \(\mathrm{O}-\mathrm{Cl} ;\) (c) \(\mathrm{Al}-\mathrm{O}\) (d) As - O?

Through appropriate Lewis structures, show that the phenomenon of resonance is involved in the nitrite ion.

Which of the following species requires a resonance hybrid for its Lewis structure: (a) \(\mathrm{CO}_{2},\) (b) \(\mathrm{OCl}^{-}\) (c) \(\mathrm{CO}_{3}^{2-},\) or \((\mathrm{d}) \mathrm{OH}^{-} ?\) Explain.

The Lewis structure of nitric acid, \(\mathrm{HONO}_{2}\), is a resonance hybrid. How important do you think the contribution of the following structure is to the resonance hybrid? Explain.

Draw Lewis structures for the following species, indicating formal charges and resonance where applicable: (a) \(\mathrm{HCO}_{2}=\) (b) \(\mathrm{HCO}_{3}^{-}\) (c) \(\mathrm{FSO}_{3}^{-}\) (d) \(\mathrm{N}_{2} \mathrm{O}_{3}^{2-}\) (the nitrogen atoms are joined centrally with one oxygen atom on one \(\mathrm{N}\) and two on the other)

Draw Lewis structures for the following species, indicating formal charges and resonance where applicable: (a) \(\mathrm{HOSO}_{3}\) (b) \(\mathrm{H}_{2} \mathrm{NCN}\) (c) \(\mathrm{FCO}_{2}^{-}\) (d) \(\mathrm{S}_{2} \mathrm{N}_{2}\) (a cyclic structure with \(\mathrm{S}\) and \(\mathrm{N}\) alternating)

Write plausible Lewis structures for the following odd-electron species: (a) \(\mathrm{CH}_{3} ;\) (b) \(\mathrm{ClO}_{2} ;\) (c) \(\mathrm{NO}_{3}\).

Write plausible Lewis structures for the following free radicals: (a) \(\cdot \mathrm{C}_{2} \mathrm{H}_{5} ;\) (b) \(\mathrm{HO}_{2}\) "; (c) ClO?.

Which of the following species would you expect to be diamagnetic and which paramagnetic: (a) \(\mathrm{OH}^{-} ;\) (b) \(\mathrm{OH} ;(\mathrm{c}) \mathrm{NO}_{3} ;(\mathrm{d}) \mathrm{SO}_{3} ;(\mathrm{e}) \mathrm{SO}_{3}^{2-} ;(\mathrm{f}) \mathrm{HO}_{2} ?\)

Write a plausible Lewis structure for \(\mathrm{NO}_{2}\), and indicate whether the molecule is diamagnetic or paramagnetic. Two \(\mathrm{NO}_{2}\) molecules can join together (dimerize) to form \(\mathrm{N}_{2} \mathrm{O}_{4}\). Write a plausible Lewis structure for \(\mathrm{N}_{2} \mathrm{O}_{4},\) and comment on the magnetic properties of the molecule.

In which of the following species is it necessary to employ an expanded valence shell to represent the Lewis structure: \(\mathrm{PO}_{4}^{3-}, \mathrm{PI}_{3}, \mathrm{ICl}_{3}, \mathrm{OSCl}_{2}, \mathrm{SF}_{4}, \mathrm{ClO}_{4} ?\) Explain your choices.

Describe the carbon-to-sulfur bond in \(\mathrm{H}_{2} \mathrm{CSF}_{4}\). That is, is it most likely a single, double, or triple bond?

Use VSEPR theory to predict the geometric shapes of the following molecules and ions: (a) \(\mathrm{N}_{2}\); (b) HCN; (c) \(\mathrm{NH}_{4}^{+} ;\) (d) \(\mathrm{NO}_{3}^{-} ;\) (e) NSF.

Use VSEPR theory to predict the geometric shapes of the following molecules and ions: (a) \(\mathrm{PCl}_{3} ;\) (b) \(\mathrm{SO}_{4}^{2-}\); (c) \(\mathrm{SOCl}_{2} ;\) (d) \(\mathrm{SO}_{3} ;\) (e) \(\mathrm{BrF}_{4}^{+}\).

Each of the following is either linear, angular (bent), planar, tetrahedral, or octahedral. Indicate the correct shape of (a) \(\mathrm{H}_{2} \mathrm{S} ;\) (b) \(\mathrm{N}_{2} \mathrm{O}_{4} ;\) (c) \(\mathrm{HCN} ;\) (d) \(\mathrm{SbCl}_{6}^{-}\); (e) \(\mathrm{BF}_{4}^{-}\)

Predict the geometric shapes of (a) \(\mathrm{CO} ;\) (b) \(\mathrm{SiCl}_{4}\); (c) \(\mathrm{PH}_{3} ;\) (d) \(\mathrm{ICl}_{3} ;\) (e) \(\mathrm{SbCl}_{5} ;\) (f) \(\mathrm{SO}_{2} ;\) (g) \(\mathrm{AlF}_{6}^{3-}\).

One of the following ions has a trigonal-planar shape: \(\mathrm{SO}_{3}^{2-} ; \mathrm{PO}_{4}^{3-} ; \mathrm{PF}_{6}^{-} ; \mathrm{CO}_{3}^{2-} .\) Which ion is it? Explain.

Each of the following molecules contains one or more multiple covalent bonds. Draw plausible Lewis structures to represent this fact, and predict the shape of each molecule. (a) \(\mathrm{CO}_{2} ;\) (b) \(\mathrm{Cl}_{2} \mathrm{CO} ;\) (c) CINO \(_{2}\).

Sketch the probable geometric shape of a molecule of (a) \(\mathrm{N}_{2} \mathrm{O}_{4}\left(\mathrm{O}_{2} \mathrm{NNO}_{2}\right) ;\) (b) \(\mathrm{C}_{2} \mathrm{N}_{2}(\mathrm{NCCN}) ;\) (c) \(\mathrm{C}_{2} \mathrm{H}_{6}\) \(\left(\mathrm{H}_{3} \mathrm{CCH}_{3}\right) ;(\mathrm{d}) \mathrm{C}_{2} \mathrm{H}_{6} \mathrm{O}\left(\mathrm{H}_{3} \mathrm{COCH}_{3}\right)\).

Use the VSEPR theory to predict the shapes of the anions (a) \(\mathrm{ClO}_{4}^{-} ;\) (b) \(\mathrm{S}_{2} \mathrm{O}_{3}^{2-}\left(\text { that is, } \mathrm{SSO}_{3}^{2-}\right) ;\) (c) \(\mathrm{PF}_{6}^{-}\); (d) I \(_{3}^{-}\).

Use the VSEPR theory to predict the shape of (a) the molecule OSF \(_{2} ;\) (b) the molecule \(\mathrm{O}_{2} \mathrm{SF}_{2} ;\) (c) the ion \(\mathrm{SF}_{5}^{-} ;\) (d) the ion \(\mathrm{ClO}_{4}^{-} ;\) (e) the ion \(\mathrm{ClO}_{3}^{-}\).

The molecular shape of \(\mathrm{BF}_{3}\) is planar (see Table 10.1 ). If a fluoride ion is attached to the \(B\) atom of \(B F_{3}\) through a coordinate covalent bond, the ion \(\mathrm{BF}_{4}^{-}\) results. What is the shape of this ion?

Explain why it is not necessary to find the Lewis structure with the smallest formal charges to make a successful prediction of molecular geometry in the VSEPR theory. For example, write Lewis structures for \(S O_{2}\) having different formal charges, and predict the molecular geometry based on these structures.

Comment on the similarities and differences in the molecular structure of the following triatomic species: \(\mathrm{CO}_{2}, \mathrm{NO}_{2}^{-}, \mathrm{O}_{3},\) and \(\mathrm{ClO}_{2}^{-}\).

Comment on the similarities and differences in the molecular structure of the following four-atom species: \(\mathrm{NO}_{3}^{-}, \mathrm{CO}_{3}^{2-}, \mathrm{SO}_{3}^{2-},\) and \(\mathrm{ClO}_{3}^{-}\).

Draw a plausible Lewis structure for the following series of molecules and ions: (a) \(\mathrm{ClF}_{2}^{-} ;\) (b) \(\mathrm{ClF}_{3}\); (c) \(\mathrm{ClF}_{4}^{-} ;\) (d) ClF \(_{5} .\) Describe the electron group geometry and molecular structure of these species.