Chapter 4

Draw the Lewis structures, predict the molecular structures, and describe the bonding (in terms of the hybrid orbitals for the central atom) for the following. a. \(\mathrm{XeO}_{3}\) b. \(\mathrm{XeO}_{4}\) c. \(\mathrm{XeOF}_{4}\) d. \(\mathrm{XeOF}_{2}\) e. \(\mathrm{XeO}_{3} \mathrm{F}_{2}\)

What do each of the following sets of compounds/ions have in common with each other? a. \(\mathrm{SO}_{3}, \mathrm{NO}_{3}^{-}, \mathrm{CO}_{3}^{2-}\) b. \(\mathrm{O}_{3}, \mathrm{SO}_{2}, \mathrm{NO}_{2}^{-}\)

What do each of the following sets of compounds/ions have in common with each other? a. \(\mathrm{XeCl}_{4}, \mathrm{XeCl}_{2}\) b. ICls, \(\operatorname{TeF}_{4}, \operatorname{ICl}_{3}, \mathrm{PCl}_{3}, \mathrm{SCl}_{2}, \mathrm{SeO}_{2}\)

The three most stable oxides of carbon are carbon monoxide \((\mathrm{CO}),\) carbon dioxide \(\left(\mathrm{CO}_{2}\right),\) and carbon suboxide \(\left(\mathrm{C}_{3} \mathrm{O}_{2}\right) .\) The space-filling models for these three compounds are For each oxide, draw the Lewis structure, predict the molecular structure, and describe the bonding (in terms of the hybrid orbitals for the carbon atoms).

Which of the following molecules have net dipole moments? For the molecules that are polar, indicate the polarity of each bond and the direction of the net dipole moment of the molecule. a. \(\mathrm{CH}_{2} \mathrm{Cl}_{2}, \mathrm{CHCl}_{3}, \mathrm{CCl}_{4}\) b. \(\mathrm{CO}_{2}, \mathrm{N}_{2} \mathrm{O}\) c. \(\mathrm{PH}_{3}, \mathrm{NH}_{3}\)

The structure of \(\operatorname{Te} \mathrm{F}_{5}^{-}\) is Draw a complete Lewis structure for \(\operatorname{TeF}_{5}^{-},\) and explain the distortion from the ideal square pyramidal structure. (See Exercise 26.)

Complete the following resonance structures for \(\mathrm{POCl}_{3}\) a. Would you predict the same molecular structure from each resonance structure? b. What is the hybridization of \(P\) in each structure? c. What orbitals can the \(P\) atom use to form the \(\pi\) bond in structure B? d. Which resonance structure would be favored on the basis of formal charges?

The \(\mathrm{N}_{2} \mathrm{O}\) molecule is linear and polar. a. On the basis of this experimental evidence, which arrangement, NNO or NON, is correct? Explain your answer. b. On the basis of your answer to part a, write the Lewis structure of \(\mathrm{N}_{2} \mathrm{O}\) (including resonance forms). Give the formal charge on each atom and the hybridization of the central atom. c. How would the multiple bonding in : \(\mathrm{N} \equiv \mathrm{N}-\mathrm{O}:\) be described in terms of orbitals?

Describe the bonding in the first excited state of \(\mathrm{N}_{2}\) (the one closest in energy to the ground state) using the molecular orbital model. What differences do you expect in the properties of the molecule in the ground state as compared to the first excited state? (An excited state of a molecule corresponds to an electron arrangement other than that giving the lowest possible energy.)

Using an MO energy-level diagram, would you expect \(\mathrm{F}_{2}\) to have a lower or higher first ionization energy than atomic fluorine? Why?

Show how a \(d_{x z}\) atomic orbital and a \(p_{z}\) atomic orbital combine to form a bonding molecular orbital. Assume the \(x\) -axis is the internuclear axis. Is a \(\sigma\) or a \(\pi\) molecular orbital formed? Explain.

Consider three molecules: \(\mathrm{A}, \mathrm{B},\) and \(\mathrm{C}\). Molecule A has a hybridization of \(s p^{3} .\) Molecule \(\mathrm{B}\) has two more effective pairs (electron pairs around the central atom) than molecule A. Molecule C consists of two \(\sigma\) bonds and two \(\pi\) bonds. Give the molecular structure, hybridization, bond angles, and an example for each molecule.

Predict the molecular structure, bond angles, and polarity (has a net dipole moment or has no net dipole moment) for each of the following compounds. a. \(\mathrm{SeCl}_{4}\) b. \(\mathrm{SF}_{2}\) c. \(\mathrm{KrF}_{4}\) d. \(C B r_{4}\) e. \(\mathrm{IF}_{3}\) f. \(\mathrm{ClF}_{5}\)

Draw the Lewis structures for \(\mathrm{SO}_{2}, \mathrm{PCl}_{3}, \mathrm{NNO}, \mathrm{COS},\) and \(\mathrm{PF}_{3} .\) Which of the compounds are polar? Which of the compounds exhibit at least one bond angle that is approximately 120 degrees? Which of the compounds exhibit \(s p^{3}\) hybridization by the central atom? Which of the compounds have a linear molecular structure?

Draw the Lewis structures for \(\mathrm{TeCl}_{4}, \mathrm{ICl}_{5}, \mathrm{PCl}_{5}, \mathrm{KrCl}_{4},\) and \(\mathrm{XeCl}_{2} .\) Which of the compounds exhibit at least one bond angle that is approximately 120 degrees? Which of the compounds exhibit \(d^{2} s p^{3}\) hybridization? Which of the compounds have a square planar molecular structure? Which of the compounds are polar?

A variety of chlorine oxide fluorides and related cations and anions are known. They tend to be powerful oxidizing and fluorinating agents. \(\mathrm{FClO}_{3}\) is the most stable of this group of compounds and has been studied as an oxidizing component in rocket propellants. Draw a Lewis structure for \(\mathrm{F}_{3} \mathrm{ClO}\) \(\mathrm{F}_{2} \mathrm{ClO}_{2}^{+},\) and \(\mathrm{F}_{3} \mathrm{ClO}_{2}\). What is the molecular structure for each species, and what is the expected hybridization of the central chlorine atom in each compound or ion?

Complete a Lewis structure for the compound shown below, then answer the following questions. What are the predicted bond angles about the carbon and nitrogen atoms? How many lone pairs of electrons are present in the Lewis structure? How many double bonds are present?

Which of the following statements concerning \(\mathrm{SO}_{2}\) is(are) true? a. The central sulfur atom is \(s p^{2}\) hybridized. b. One of the sulfur-oxygen bonds is longer than the other(s). c. The bond angles about the central sulfur atom are about 120 degrees. d. There are two \(\sigma\) bonds in \(\mathrm{SO}_{2}\) e. There are no resonance structures for \(\mathrm{SO}_{2}\)

Consider the molecular orbital electron configurations for \(\mathrm{N}_{2}\) \(\mathrm{N}_{2}^{+},\) and \(\mathrm{N}_{2}^{-} .\) For each compound or ion, fill in the table below with the correct number of electrons in each molecular orbital.

Place the species \(\mathrm{B}_{2}^{+}, \mathrm{B}_{2},\) and \(\mathrm{B}_{2}^{-}\) in order of increasing bond length and increasing bond energy.

The compound \(\mathrm{NF}_{3}\) is quite stable, but \(\mathrm{NCl}_{3}\) is very unstable (NCl \(_{3}\) was first synthesized in 1811 by P. L. Dulong, who lost three fingers and an eye studying its properties). The compounds \(\mathrm{NBr}_{3}\) and \(\mathrm{NI}_{3}\) are unknown, although the explosive compound \(\mathrm{NI}_{3} \cdot \mathrm{NH}_{3}\) is known. Account for the instability of these halides of nitrogen.

Predict the molecular structure for each of the following. (See Exercises 25 and \(26 .\) ) a. \(\mathrm{BrFI}_{2}\) b. \(\mathrm{XeO}_{2} \mathrm{F}_{2}\) c. \(\operatorname{TeF}_{2} \mathrm{Cl}_{3}^{-}\) For each formula there are at least two different structures that can be drawn using the same central atom. Draw all possible structures for each formula.

Cyanamide \(\left(\mathrm{H}_{2} \mathrm{NCN}\right),\) an important industrial chemical, is produced by the following steps: Calcium cyanamide (CaNCN) is used as a direct-application fertilizer, weed killer, and cotton defoliant. It is also used to make cyanamide, dicyandiamide, and melamine plastics: a. Write Lewis structures for \(\mathrm{NCN}^{2-}, \mathrm{H}_{2} \mathrm{NCN}\), dicyandiamide, and melamine, including resonance structures where appropriate. b. Give the hybridization of the \(\mathrm{C}\) and \(\mathrm{N}\) atoms in each species. c. How many \(\sigma\) bonds and how many \(\pi\) bonds are in each species? d. Is the ring in melamine planar? e. There are three different \(\mathrm{C}-\mathrm{N}\) bond distances in dicyandiamide, NCNC(NH_)_2, and the molecule is nonlinear. Of all the resonance structures you drew for this molecule, predict which should be the most important.

As compared with \(\mathrm{CO}\) and \(\mathrm{O}_{2}, \mathrm{CS}\) and \(\mathrm{S}_{2}\) are very unstable molecules. Give an explanation based on the relative abilities of the sulfur and oxygen atoms to form \(\pi\) bonds.

Values of measured bond energies may vary greatly depending on the molecule studied. Consider the following reactions: $$ \begin{aligned} & \mathrm{NCl}_{3}(g) \longrightarrow \mathrm{NCl}_{2}(g)+\mathrm{Cl}(g) & \Delta E &=375 \mathrm{kJ} / \mathrm{mol} \\ \mathrm{ONCl}(g) & \longrightarrow \mathrm{NO}(g)+\mathrm{Cl}(g) & & \Delta E=158 \mathrm{kJ} / \mathrm{mol} \end{aligned} $$ Rationalize the difference in the values of \(\Delta E\) for these reactions, even though each reaction appears to involve only the breaking of one \(\mathrm{N}-\mathrm{Cl}\) bond. (Hint: Consider the bond order of the NO bond in ONCl and in NO.)

Bond energy has been defined in the text as the amount of energy required to break a chemical bond, so we have come to think of the addition of energy as breaking bonds. However, in some cases the addition of energy can cause the formation of bonds. For example, in a sample of helium gas subjected to a high-energy source, some He_ molecules exist momentarily and then dissociate. Use MO theory (and diagrams) to explain why \(\mathrm{He}_{2}\) molecules can come to exist and why they dissociate.

Arrange the following from lowest to highest ionization energy: \(\mathbf{O}, \mathbf{O}_{2}, \mathbf{O}_{2}^{-}, \mathbf{O}_{2}^{+} .\) Explain your answer.

Use the MO model to determine which of the following has the smallest ionization energy: \(\mathrm{N}_{2}, \mathrm{O}_{2}, \mathrm{N}_{2}^{2-}, \mathrm{N}_{2}^{-}, \mathrm{O}_{2}^{+} .\) Explain your answer.

Carbon monoxide (CO) forms bonds to a variety of metals and metal ions. Its ability to bond to iron in hemoglobin is the reason that \(\mathrm{CO}\) is so toxic. The bond carbon monoxide forms to metals is through the carbon atom: $$\mathbf{M}-\mathbf{C} \equiv \mathbf{O}$$ a. On the basis of electronegativities, would you expect the carbon atom or the oxygen atom to form bonds to metals? b. Assign formal charges to the atoms in CO. Which atom would you expect to bond to a metal on this basis? c. In the MO model, bonding MOs place more electron density near the more electronegative atom. (See the HF molecule in Figs. \(4-54\) and \(4-55 .\) ) Antibonding MOs place more electron density near the less electronegative atom in the diatomic molecule. Use the MO model to predict which atom of carbon monoxide should form bonds to metals.

As the head engineer of your starship in charge of the warp drive, you notice that the supply of dilithium is critically low. While searching for a replacement fuel, you discover some diboron, \(\mathbf{B}_{2}\) a. What is the bond order in \(\mathrm{Li}_{2}\) and \(\mathrm{B}_{2} ?\) b. How many electrons must be removed from \(\mathrm{B}_{2}\) to make it isoelectronic with \(\mathrm{Li}_{2}\) so that it might be used in the warp drive? c. The reaction to make \(\mathrm{B}_{2}\) isoelectronic with \(\mathrm{Li}_{2}\) is generalized (where \(n=\) number of electrons determined in part b) as follows: $$\mathrm{B}_{2} \longrightarrow \mathrm{B}_{2}^{n+}+n \mathrm{e}^{-} \quad \Delta E=6455 \mathrm{kJ} / \mathrm{mol}$$ How much energy is needed to ionize \(1.5 \mathrm{kg} \mathrm{B}_{2}\) to the desired isoelectronic species?

A flask containing gaseous \(\mathrm{N}_{2}\) is irradiated with 25 -nm light. a. Using the following information, indicate what species can form in the flask during irradiation. $$\begin{array}{ll} \mathrm{N}_{2}(g) \longrightarrow 2 \mathrm{N}(g) & \Delta E=941 \mathrm{kJ} / \mathrm{mol} \\ \mathrm{N}_{2}(g) \longrightarrow \mathrm{N}_{2}^{+}(g)+\mathrm{e}^{-} & \Delta E=1501 \mathrm{kJ} / \mathrm{mol} \\ \mathrm{N}(g) \longrightarrow \mathrm{N}^{+}(g)+\mathrm{e}^{-} & \Delta E=1402 \mathrm{kJ} / \mathrm{mol} \end{array}$$ b. What range of wavelengths will produce atomic nitrogen in the flask but will not produce any ions? c. Explain why the first ionization energy of \(\mathrm{N}_{2}(1501 \mathrm{kJ} /\) mol) is greater than the first ionization energy of atomic nitrogen (1402 kJ/mol).

Determine the molecular structure and hybridization of the central atom \(X\) in the polyatomic ion \(X Y_{3}^{+}\) given the following information: A neutral atom of X contains 36 electrons, and the element Y makes an anion with a \(1-\) charge, which has the electron configuration \(1 s^{2} 2 s^{2} 2 p^{6}\).