Chapter 6

Explain the difference between an ionic bond and a covalent bond.

What kind of bonding (ionic or covalent) would you predict for the products resulting from the following combinations of elements? (a) \(\mathrm{Na}+\mathrm{I}_{2}\) (b) \(\mathrm{C}+\mathrm{S}_{8}\) (c) \(\mathrm{Mg}+\mathrm{Br}_{2}\) (d) \(\mathrm{P}_{4}+\mathrm{Cl}_{2}\)

What characteristics must atoms A and \(X\) have if they are able to form a covalent bond \(\mathrm{A}-\mathrm{X}\) with each other? A polar covalent bond with each other?

Indicate how alkanes, alkenes, and alkynes differ by giving the structural formula of a compound in each class that contains three carbon atoms.

While sulfur forms the compounds \(\mathrm{SF}_{4}\) and \(\mathrm{SF}_{6},\) no equivalent compounds of oxygen, \(\mathrm{OF}_{4}\) and \(\mathrm{OF}_{6}\), are known. Explain.

Which of these molecules have an odd number of valence electrons: \(\mathrm{NO}_{2}, \mathrm{SCl}_{2}, \mathrm{NH}_{3}, \mathrm{NO}_{3} ?\)

Consider a series of molecules in which the \(\mathrm{C}\) atom is bonded to atoms of second-period elements: \(\mathrm{C}-\mathrm{O}\), \(\mathrm{C}-\mathrm{F}, \mathrm{C}-\mathrm{N}, \mathrm{C}-\mathrm{C},\) and \(\mathrm{C}-\mathrm{B}\). Place these bonds in order of increasing bond length.

Describe the trends in bond length and bond energy for single, double, and triple carbon-to-oxygen bonds.

Why is cis-trans isomerism not possible for alkynes?

Explain in your own words why the energy of two \(\mathrm{H}\) atoms is lower when the atoms are \(74 \mathrm{pm}\) apart than when the atoms are (a) \(25 \mathrm{pm}\) apart; (b) \(100 \mathrm{pm}\) apart.

Write Lewis structures for these molecules or ions. (a) ClF (b) \(\mathrm{H}_{2} \mathrm{Se}\) (c) \(\mathrm{BF}_{4}^{-}\) (d) \(\mathrm{PO}_{4}^{3-}\)

Write the Lewis structures of (a) dichlorine monoxide, \(\mathrm{Cl}_{2} \mathrm{O} ;\) (b) hydrogen peroxide, \(\mathrm{H}_{2} \mathrm{O}_{2} ;\) (c) borohydride ion, \(\mathrm{BH}_{4}^{-} ;\) (d) phosphonium ion, \(\mathrm{PH}_{4}^{+}\); and (e) \(\mathrm{PCl}_{5}\).

Write Lewis structures for these molecules or ions. (a) \(\mathrm{CH}_{3} \mathrm{Cl}\) (b) \(\mathrm{SiO}_{4}^{4-}\) (c) \(\mathrm{ClF}_{4}^{+}\) (d) \(\mathrm{C}_{2} \mathrm{H}_{6}\)

Write Lewis structures for these molecules. (a) Tetrafluoroethylene, \(\mathrm{C}_{2} \mathrm{~F}_{4}\), the molecule from which Teflon is made (b) Acrylonitrile, \(\mathrm{CH}_{2} \mathrm{CHCN},\) the molecule from which Orlon is made

Write Lewis structures for these molecules. (a) Formic acid, \(\mathrm{HCOOH}\), in which atomic arrangement is (b) Acetonitrile, \(\mathrm{CH}_{3} \mathrm{CN}\) (c) Vinyl chloride, \(\mathrm{CH}_{2} \mathrm{CHCl}\), the molecule from which PVC plastics are made

Write Lewis structures for (a) \(\mathrm{N}_{2}^{+}\) (b) \(\mathrm{XeF}_{7}^{-}\) (c) tetracyanoethene, \(\mathrm{C}_{6} \mathrm{~N}_{4}\)

Write the structural formulas for all the branched-chain compounds with the molecular formula \(\mathrm{C}_{6} \mathrm{H}_{14}\).

Write structural formulas for two straight-chain alkenes with the formula \(\mathrm{C}_{5} \mathrm{H}_{10}\). Are these the only two structures that meet these specifications?

From their molecular formulas, classify each of these straight-chain hydrocarbons as an alkane, an alkene, or an alkyne. (a) \(\mathrm{C}_{5} \mathrm{H}_{8}\) (b) \(\mathrm{C}_{24} \mathrm{H}_{50}\) (c) \(\mathrm{C}_{7} \mathrm{H}_{14}\)

From their molecular formulas, classify each of these straight-chain hydrocarbons as an alkane, an alkene, or an alkyne. (a) \(\mathrm{C}_{21} \mathrm{H}_{44}\) (b) \(\mathrm{C}_{4} \mathrm{H}_{6}\) (c) \(\mathrm{C}_{8} \mathrm{H}_{16}\)

In each case, tell whether cis and trans isomers exist. If they do, write structural formulas for the two isomers and label each cis or trans. For those that cannot have cistrans isomers, explain why. (a) \(\mathrm{Br}_{2} \mathrm{CH}_{2}\) (b) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}=\mathrm{CHCH}_{2} \mathrm{CH}_{3}\) (c) \(\mathrm{CH}_{3} \mathrm{CH}=\mathrm{CHCH}_{3}\) (d) \(\mathrm{CH}_{2}=\mathrm{CHCH}_{2} \mathrm{CH}_{3}\)

Which of these molecules can have cis and trans iso- mers? For those that do, write the structural formulas of the two isomers and label each cis or trans. For those that cannot have these isomers, explain why. (a) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{BrC}=\mathrm{CBrCH}_{3}\) (b) \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{C}=\mathrm{C}\left(\mathrm{CH}_{3}\right)_{2}\) (c) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{IC}=\mathrm{CICH}_{2} \mathrm{CH}_{3}\) (d) \(\mathrm{CH}_{3} \mathrm{ClC}=\mathrm{CHCH}_{3}\) (e) \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{C}=\mathrm{CHCH}_{3}\)

For each pair of bonds, predict which is the shorter. (a) \(\mathrm{B}-\mathrm{Cl}\) or \(\mathrm{Ga}-\mathrm{Cl}\) (b) \(\mathrm{C}-\mathrm{O}\) or \(\mathrm{Sn}-\mathrm{O}\) (c) \(\mathrm{P}-\mathrm{S}\) or \(\mathrm{P}-\mathrm{O}\) (d) The \(\mathrm{C}=\mathrm{C}\) or the \(\mathrm{C}=\mathrm{O}\) bond in acrolein

For each pair of bonds, predict which is the shorter. (a) \(\mathrm{Si}-\mathrm{N}\) or \(\mathrm{P}-\mathrm{O}\) (b) \(\mathrm{Si}-\mathrm{O}\) or \(\mathrm{C}-\mathrm{O}\) (c) \(\mathrm{C}-\mathrm{F}\) or \(\mathrm{C}-\mathrm{Br}\) (d) The \(\mathrm{C}=\mathrm{C}\) or the \(\mathrm{C} \equiv \mathrm{N}\) bond in acrylonitrile, \(\mathrm{H}_{2} \mathrm{C}=\mathrm{CH}-\mathrm{C} \equiv \mathrm{N}\)

Using only a periodic table (not a table of bond energies), predict which is the strongest bond. (a) \(\mathrm{Si}-\mathrm{F}\) (b) \(\mathrm{P}-\mathrm{S}\) (c) \(\mathrm{Si}-\mathrm{O}\)

When sulfur is heated to above \(720^{\circ} \mathrm{C}\), the major component is \(\mathrm{S}_{2}\) in which the bond distance is \(189 \mathrm{pm}\). This is significantly shorter than the sulfur-to-sulfur distance in \(S_{8},\) which has a ring structure. Propose a reason for this difference.

Estimate \(\Delta_{1} H^{\circ}\) for the conversion of \(1 \mathrm{~mol}\) carbon monoxide to carbon dioxide by combination with molecular oxygen. Is this reaction exothermic or endothermic?

Light of appropriate wavelength can break chemical bonds. Light having \(\lambda<240 \mathrm{nm}\) can dissociate gaseous \(\mathrm{O}_{2}\). It requires light with \(\lambda<819 \mathrm{nm}\) to dissociate gaseous \(\mathrm{H}_{2} \mathrm{O}_{2}\) to \(2 \mathrm{OH}\). Assume that all of the photon energy is used solely for these dissociations. (a) Calculate the energy required to dissociate (i) \(\mathrm{O}_{2}\) and (ii) \(\mathrm{H}_{2} \mathrm{O}_{2}\). (b) Consider the results of part (a). How well do they correlate with the Lewis structures of \(\mathrm{O}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}_{2}\) ? Explain your answer.

For each pair of bonds, indicate the more polar bond and use \(\delta+\) or \(\delta-\) to show the partial charge on each atom. (a) \(\mathrm{C}-\mathrm{O}\) and \(\mathrm{C}-\mathrm{N}\) (b) \(\mathrm{B}-\mathrm{O}\) and \(\mathrm{P}-\mathrm{S}\) (c) \(\mathrm{P}-\mathrm{H}\) and \(\mathrm{P}-\mathrm{N}\) (d) \(\mathrm{B}-\mathrm{H}\) and \(\mathrm{B}-\mathrm{Cl}\)

For each pair of bonds, identify the more polar one and use \(\delta+\) or \(\delta-\) to indicate the partial charge on each atom. (a) \(\mathrm{B}-\mathrm{Cl}\) and \(\mathrm{B}-\mathrm{O}\) (b) \(\mathrm{O}-\mathrm{F}\) and \(\mathrm{O}-\mathrm{Se}\) (c) \(\mathrm{S}-\mathrm{Cl}\) and \(\mathrm{B}-\mathrm{F}\) (d) \(\mathrm{N}-\mathrm{H}\) and \(\mathrm{N}-\mathrm{F}\)

Urea is used in plastics and fertilizers. (a) Which bonds in the molecule are polar and which are nonpolar? (b) Which is the most polar bond in the molecule? Which atom is the partial negative end of this bond?

Write correct Lewis structures and assign a formal charge to each atom. (a) \(\mathrm{CH}_{3} \mathrm{CHO}\) (b) \(\mathrm{N}_{3}^{-}\) (c) \(\mathrm{CH}_{3} \mathrm{CN}\)

Write correct Lewis structures and assign a formal charge to each atom. (a) \(\mathrm{KrF}_{4}\) (b) \(\mathrm{ClO}_{3}^{-}\) (c) \(\mathrm{SO}_{2} \mathrm{Cl}_{2}\)

Write the correct Lewis structure and assign a formal charge to each atom in fulminate ion, \(\mathrm{CNO}^{-}\).

Peroxydisulfate ion, \(\mathrm{S}_{2} \mathrm{O}_{8}^{2-},\) contains an \(-\mathrm{O}-\mathrm{O}-\) linkage between two sulfate ions, \(\mathrm{SO}_{4}^{2-}\). Write the correct Lewis structures and assign a formal charge to each atom in the peroxydisulfate and sulfate ions.

Two Lewis structures can be written for nitrosyl chloride, which contains one nitrogen, one oxygen, and one chlorine atom per molecule. Write the two Lewis structures and assign a formal charge to each atom.

Two Lewis structures can be written for nitrosyl fluoride, which contains one nitrogen, one oxygen, and one fluorine atom per molecule. Write the two Lewis structures and assign a formal charge to each atom.

Use Lewis structures and formal charges to determine the bond type (single, double, or triple) for each bond in (a) \(\mathrm{POF}_{3}\) (b) \(\mathrm{SOF}_{4}\)

Consider the \(\mathrm{SCO}_{2}^{2-}\) ion in which each \(\mathrm{S}\) and \(\mathrm{O}\) atom is bonded to a central \(\mathrm{C}\) atom. Use Lewis structures and formal charges to write resonance structures and to determine which is the most plausible resonance form. Explain your choice.

Write all resonance structures for (a) nitric acid (b) nitrate ion, \(\mathrm{NO}_{3}^{-}\)

Write all the resonance structures for (a) \(\mathrm{SO}_{3}\) (b) \(\mathrm{SCN}^{-}\)

Several Lewis structures can be written for perbromate ion, \(\mathrm{BrO}_{4}^{-},\) the central \(\mathrm{Br}\) with all single \(\mathrm{Br}-\mathrm{O}\) bonds, or with one, two, or three \(\mathrm{Br}=\mathrm{O}\) double bonds. Draw the Lewis structures of these possible resonance structures, and use formal charges to predict which makes the greatest contribution to the resonance hybrid.

Several Lewis structures can be written for thiosulfate ion, \(\mathrm{S}_{2} \mathrm{O}_{3}^{2-}\). Write the Lewis structures of these possible resonance structures. Predict which structure makes the most important contribution to the resonance hybrid.

Compare the carbon-oxygen bond lengths in the formate ion, \(\mathrm{HCO}_{2}^{-},\) and in the carbonate ion, \(\mathrm{CO}_{3}^{2-} .\) In which ion is the bond longer? Explain briefly.

Compare the nitrogen-oxygen bond lengths in \(\mathrm{NO}_{2}^{+}\) and in \(\mathrm{NO}_{3}^{-}\). In which ion are the bonds longer? Explain briefly.

Draw resonance structures for each of these ions: \(\mathrm{NSO}^{-}\) and \(\mathrm{SNO}^{-}\). (The atoms are bonded in the order given in each case, that is, \(\mathrm{S}\) is the central atom in \(\mathrm{NSO}^{-} .\).) (a) Use formal charges to determine which ion is likely to be more stable. (b) Explain why the two ions cannot be considered resonance structures of each other.

Three known isomers exist of \(\mathrm{N}_{2} \mathrm{CO},\) with the atoms in these sequences: \(\mathrm{NOCN} ; \mathrm{ONNC} ;\) and ONCN. Write resonance structures for each isomer and use formal charge to predict which isomer is the most stable.

Write the Lewis structure for (a) \(\mathrm{BrF}_{5}\) (b) \(\mathrm{IF}_{5}\) (c) \(\mathrm{IBr}_{2}^{-}\)

Write the Lewis structure for (a) \(\mathrm{BrF}_{3}\) (b) \(\mathrm{I}_{3}^{-}\) (c) \(\mathrm{XeF}_{4}\)

Which of these molecules have Lewis structures that involve exceptions to the octet rule? Classify each exception. (a) \(\mathrm{PCl}_{3}\) (b) \(\mathrm{SnF}_{4}\) (c) \(\mathrm{BCl}_{3}\) (d) \(\mathrm{NO}\)