Chapter 8

(a) Use bond enthalpies to estimate the enthalpy change for the reaction of hydrogen with ethene: $$\mathrm{H}_{2}(g)+\mathrm{C}_{2} \mathrm{H}_{4}(g) \longrightarrow \mathrm{C}_{2} \mathrm{H}_{6}(g)$$ (b) Calculate the standard enthalpy change for this reaction, using heats of formation. Why does this value differ from that calculated in (a)?

Given the following bond-dissociation energies, calculate the average bond enthalpy for the \(\mathrm{Ti}-\mathrm{Cl}\) bond. $$ \begin{array}{lc} & \Delta H(\mathrm{~kJ} / \mathrm{mol}) \\ \hline \mathrm{TiCl}_{4}(g) \longrightarrow \mathrm{TiCl}_{3}(g)+\mathrm{Cl}(g) & 335 \\ \mathrm{TiCl}_{3}(g) \longrightarrow \mathrm{TiCl}_{2}(g)+\mathrm{Cl}(g) & 423 \\\ \mathrm{TiCl}_{2}(g) \longrightarrow \mathrm{TiCl}(g)+\mathrm{Cl}(g) & 444 \\ \mathrm{TiCl}(g) \longrightarrow \mathrm{Ti}(g)+\mathrm{Cl}(g) & 519 \\ \hline \end{array} $$

(a) Using average bond enthalpies, predict which of the following reactions will be most exothermic: (i) \(C(g)+2 \mathrm{~F}_{2}(g) \longrightarrow \mathrm{CF}_{4}(g)\) (ii) \(\mathrm{CO}(g)+3 \mathrm{~F}_{2} \longrightarrow \mathrm{CF}_{4}(g)+\mathrm{OF}_{2}(g)\) (iii) \(\mathrm{CO}_{2}(g)+4 \mathrm{~F}_{2} \longrightarrow \mathrm{CF}_{4}(g)+2 \mathrm{OF}_{2}(g)\) (b) Explain the trend, if any, that exists between reaction exothermicity and the extent to which the carbon atom is bonded to oxygen.

How many elements in the periodic table are represented by a Lewis symbol with a single dot? Are all these elements in the same group? Explain.

Would you expect AlN to have a lattice energy that is larger or smaller than ScN? Explain.

(a) How does a polar molecule differ from a nonpolar one? (b) Atoms \(X\) and \(Y\) have different electronegativities. Will the diatomic molecule \(X-Y\) necessarily be polar? Explain. (c) What factors affect the size of the dipole moment of a diatomic molecule?

Which of the following molecules or ions contain polar bonds: (a) \(\mathrm{P}_{4}\), (b) \(\mathrm{H}_{2} \mathrm{~S}\), (c) \(\mathrm{NO}_{2}^{-}\), (d) \(\mathrm{S}_{2}{\underline{\phantom{xx}}}^{2-}\) ?

To address energy and environmental issues, there is great interest in powering vehicles with hydrogen rather than gasoline. One of the most attractive aspects of the "hydrogen economy" is the fact that in principle the only emission would be water. However, two daunting obstacles must be overcome before this vision can become a reality. First, an economical method of producing hydrogen must be found. Second, a safe, lightweight, and compact way of storing hydrogen must be found. The hydrides of light metals are attractive for hydrogen storage because they can store a high weight percentage of hydrogen in a small volume. One of the most attractive hydrides is \(\mathrm{NaAlH}_{4}\), which can release \(5.6 \%\) of its mass as \(\mathrm{H}_{2}\) upon decomposing to \(\mathrm{NaH}(\mathrm{s})\), Al(s), and \(\mathrm{H}_{2}(\mathrm{~g})\). NaAlH \(_{4}\) possesses both covalent bonds, which hold polyatomic anions together, and ionic bonds. (a) Write a balanced equation for the decomposition of \(\mathrm{NaAlH}_{4} .\) (b) Which element in \(\mathrm{NaAlH}_{4}\) is the most electronegative? Which one is the least electronegative? (c) Based on electronegativity differences, what do you think is the identity of the polyatomic anion? Draw a Lewis structure for this ion.

For the following collection of nonmetallic elements, \(\mathrm{O}\), \(\mathrm{P}, \mathrm{Te}, \mathrm{I}, \mathrm{B}\), (a) which two would form the most polar single bond? (b) Which two would form the longest single bond? (c) Which two would be likely to form a compound of formula \(\mathrm{XY}_{2}\) ? (d) Which combinations of elements would likely yield a compound of empirical formula \(\mathrm{X}_{2} \mathrm{Y}_{3} ?\) In each case explain your answer.

You and a partner are asked to complete a lab entitled "Oxides of Ruthenium" that is scheduled to extend over two lab periods. The first lab, which is to be completed by your partner, is devoted to carrying out compositional analysis. In the second lab, you are to determine melting points. Upon going to lab you find two unlabeled vials, one containing a soft yellow substance and the other a black powder. You also find the following notes in your partner's notebook-Compound 1: \(76.0 \%\) Ru and \(24.0 \%\) O (by mass), Compound 2: \(61.2 \%\) Ru and \(38.8 \% \mathrm{O}\) (by mass). (a) What is the empirical formula for Compound \(1 ?\) (b) What is the empirical formula for Compound 2? (c) Upon determining the melting points of these two compounds, you find that the yellow compound melts at \(25^{\circ} \mathrm{C}\), while the black powder does not melt up to the maximum temperature of your apparatus, \(1200^{\circ} \mathrm{C}\). What is the identity of the yellow compound? What is the identity of the black compound? Be sure to use the appropriate naming convention depending upon whether the compound is better described as a molecular or ionic compound.

You and a partner are asked to complete a lab entitled "Fluorides of Croup \(6 \mathrm{~B}\) Metals" that is scheduled to extend over two lab periods. The first lab, which is to be completed by your partner, is devoted to carrying out compositional analysis. In the second lab, you are to determine melting points. Upon going to lab you find two unlabeled vials, one containing a colorless liquid and the other a green powder. You also find the following notes in your partner's notebook-Compound 1: \(47.7 \% \mathrm{Cr}\) and \(52.3 \% \mathrm{~F}\) (by mass), Compound 2: \(45.7 \% \mathrm{Mo}\) and \(54.3 \% \mathrm{~F}\) (by mass). (a) What is the empirical formula for Compound \(1 ?\) (b) What is the empirical formula for Compound 2? (c) Upon determining the melting points of these two compounds you find that the colorless liquid solidifies at \(18^{\circ} \mathrm{C}\), while the green powder does not melt up to the maximum temperature of your apparatus, \(1200^{\circ} \mathrm{C}\). What is the identity of the colorless liquid? What is the identity of the green powder? Be sure to use the appropriate naming convention depending upon whether the compound is better described as a molecular or ionic compound.

(a) Triazine, \(\mathrm{C}_{3} \mathrm{H}_{3} \mathrm{~N}_{3}\), is like benzene except that in triazine every other \(\mathrm{C}-\mathrm{H}\) group is replaced by a nitrogen atom. Draw the Lewis structure(s) for the triazine molecule. (b) Estimate the carbon-nitrogen bond distances in the ring.

Although \(\mathrm{I}_{3}^{-}\) is known, \(\mathrm{F}_{3}^{-}\) is not. Using Lewis structures, explain why \(\mathrm{F}_{3}^{-}\) does not form.

Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in \(\mathrm{O}_{3},(\mathrm{~b})\) phosphorus in \(\mathrm{PF}_{6}^{-}\), (c) nitrogen in \(\mathrm{NO}_{2}\), (d) iodine in \(\mathrm{ICl}_{3}\), (e) chlorine in \(\mathrm{HClO}_{4}\) (hydrogen is bonded to \(\mathrm{O}\) ).

(a) Determine the formal charge on the chlorine atom in the hypochlorite ion, \(\mathrm{ClO}^{-}\), and the perchlorate ion, \(\mathrm{ClO}_{4}^{-}\), using resonance structures where the \(\mathrm{Cl}\) atom has an octet. (b) What are the oxidation numbers of chlorine in \(\mathrm{ClO}^{-}\) and in \(\mathrm{ClO}_{4}^{-} ?(\mathrm{c})\) Is it uncommon for the formal charge and the oxidation state to be different? Explain. (d) Perchlorate is a much stronger oxidizing agent than hypochlorite. Would you expect there to be any relationship between the oxidizing power of the oxyanion and either the oxidation state or the formal charge of chlorine?

The following three Lewis structures can be drawn for $$\begin{aligned} &\mathrm{N}_{2} \mathrm{O}: \\ &: \mathrm{N} \equiv \mathrm{N}-\ddot{O}: \longleftrightarrow: \ddot{\mathrm{N}}-\mathrm{N} \equiv \mathrm{O}: \longleftrightarrow: \dot{\mathrm{N}}=\mathrm{N}=\ddot{\mathrm{O}}: \end{aligned} $$ (a) Using formal charges, which of these three resonance forms is likely to be the most important? (b) The \(\mathrm{N}-\mathrm{N}\) bond length in \(\mathrm{N}_{2} \mathrm{O}\) is \(1.12 \AA\), slightly longer than a typical \(\mathrm{N} \equiv \mathrm{N}\) bond; and the \(\mathrm{N}-\mathrm{O}\) bond length is \(1.19 \AA\), slightly shorter than a typical \(\mathrm{N}=\mathrm{O}\) bond. (See Table 8.5.) Rationalize these observations in terms of the resonance structures shown previously and your conclusion for (a).

The reaction of indium, In, with sulfur leads to three binary compounds, which we will assume to be purely ionic. The three compounds have the following properties: $$ \begin{array}{lll} \hline \text { Compound } & \text { Mass \% In } & \text { Melting Point }\left({ }^{\circ} \mathrm{C}\right) \\ \hline \text { A } & 87.7 & 653 \\ \text { B } & 78.2 & 692 \\ \text { C } & 70.5 & 1050 \\ \hline \end{array} $$ (a) Determine the empirical formulas of compounds A, B, and C. (b) Give the oxidation state of In in each of the three compounds. (c) Write the electron configuration for the In ion in each compound. Do any of these configurations correspond to a noble-gas configuration? (d) In which compound is the ionic radius of In expected to be smallest? Explain. (e) The melting point of ionic compounds often correlates with the lattice energy. Explain the trends in the melting points of compounds \(A, B\), and C in these terms.

One scale for electronegativity is based on the concept that the electronegativity of any atom is proportional to the ionization energy of the atom minus its electron affinity: electronegativity \(=k(\mathrm{IE}-\mathrm{EA})\), where \(k\) is a proportionality constant. (a) How does this definition explain why the electronegativity of \(\mathrm{F}\) is greater than that of \(C l\) even though \(C l\) has the greater electron affinity? (b) Why are both ionization energy and electron affinity relevant to the notion of electronegativity? (c) By using data in Chapter 7 , determine the value of \(k\) that would lead to an electronegativity of \(4.0\) for \(\mathrm{F}\) under this definition. (d) Use your result from part (c) to determine the electronegativities of \(\mathrm{Cl}\) and \(\mathrm{O}\) using this scale. Do these values follow the trend shown in Figure 8.6?

The compound chloral hydrate, known in detective stories as knockout drops, is composed of \(14.52 \% \mathrm{C}\), \(1.83 \% \mathrm{H}, 64.30 \% \mathrm{Cl}\), and \(19.35 \%\) O by mass and has a molar mass of \(165.4 \mathrm{~g} / \mathrm{mol}\). (a) What is the empirical formula of this substance? (b) What is the molecular formula of this substance? (c) Draw the Lewis structure of the molecule, assuming that the \(\mathrm{Cl}\) atoms bond to a single \(C\) atom and that there are a \(C-C\) bond and two \(\mathrm{C}-\mathrm{O}\) bonds in the compound.

Barium azide is \(62.04 \%\) Ba and \(37.96 \%\) N. Each azide ion has a net charge of \(1-\). (a) Determine the chemical formula of the azide ion. (b) Write three resonance structures for the azide ion. (c) Which structure is most important? (d) Predict the bond lengths in the ion.

Under special conditions, sulfur reacts with anhydrous liquid ammonia to form a binary compound of sulfur and nitrogen. The compound is found to consist of \(69.6 \% \mathrm{~S}\) and \(30.4 \%\) N. Measurements of its molecular mass yield a value of \(184.3 \mathrm{~g} \mathrm{~mol}^{-1}\). The compound occasionally detonates on being struck or when heated rapidly. The sulfur and nitrogen atoms of the molecule are joined in a ring. All the bonds in the ring are of the same length. (a) Calculate the empirical and molecular formulas for the substance. (b) Write Lewis structures for the molecule, based on the information you are given. (Hint: You should find a relatively small number of dominant Lewis structures.) (c) Predict the bond distances between the atoms in the ring. (Note: The \(\mathrm{S}-\mathrm{S}\) distance in the \(\mathrm{S}_{8}\) ring is \(2.05 \AA\).) (d) The enthalpy of formation of the compound is estimated to be \(480 \mathrm{~kJ} \mathrm{~mol}^{-1}\). \(\Delta H_{f}^{\circ}\) of \(S(g)\) is \(222.8 \mathrm{~kJ} \mathrm{~mol}^{-1}\). Estimate the average bond enthalpy in the compound.

A common form of elemental phosphorus is the tetrahedral \(\mathrm{P}_{4}\) molecule, where all four phosphorus atoms are equivalent: At room temperature phosphorus is a solid. (a) Do you think there are any unshared pairs of electrons in the \(\mathrm{P}_{4}\) molecule? (b) How many \(\mathrm{P}-\mathrm{P}\) bonds are there in the molecule? (c) Can you draw a Lewis structure for a linear \(\mathrm{P}_{4}\) molecule that satisfies the octet rule? (d) Using formal charges, what can you say about the stability of the linear molecule vs. that of the tetrahedral molecule?

Average bond enthalpies are generally defined for gasphase molecules. Many substances are liquids in their standard state. \(\mathrm{OB}\) (Section 5.7) By using appropriate thermochemical data from Appendix C, calculate average bond enthalpies in the liquid state for the following bonds, and compare these values to the gas-phase values given in Table 8.4: (a) \(\mathrm{Br}-\mathrm{Br}\), from \(\mathrm{Br}_{2}(l) ;\) (b) \(\mathrm{C}-\mathrm{Cl}\), from \(\mathrm{CCl}_{4}(l) ;\) (c) \(\mathrm{O}-\mathrm{O}\), from \(\mathrm{H}_{2} \mathrm{O}_{2}(l)\) (assume that the \(\mathrm{O}-\mathrm{H}\) bond enthalpy is the same as in the gas phase). (d) What can you conclude about the process of breaking bonds in the liquid as compared to the gas phase? Explain the difference in the \(\Delta H\) values between the two phases.