Chapter 8

How many elements in the periodic table are represented by a Lewis symbol with a single dot? Which groups are they in?

(a) Consider the lattice energies for the following compounds: \(\mathrm{BeH}_{2}, 3205 \mathrm{~kJ} / \mathrm{mol} ; \mathrm{MgH}_{2}, 2791 \mathrm{~kJ} / \mathrm{mol} ; \mathrm{CaH}_{2}, 2410 \mathrm{~kJ} / \mathrm{mol} ;\) \(\mathrm{SrH}_{2}, 2250 \mathrm{~kJ} / \mathrm{mol} ; \mathrm{BaH}_{2}, 2121 \mathrm{~kJ} / \mathrm{mol}\). Plot lattice energy versus cation radius for these compounds. If you draw a line through your points, is the slope negative or positive? Explain. (b) The lattice energy of \(\mathrm{ZnH}_{2}\) is \(2870 \mathrm{~kJ} / \mathrm{mol}\). Based on the data given in part (a), the radius of the \(\mathrm{Zn}^{2+}\) ion is expected to be closest to that of which group \(2 \mathrm{~A}\) element?

An ionic substance of formula MX has a lattice energy of \(6 \times 10^{3} \mathrm{~kJ} / \mathrm{mol}\). Is the charge on the ion M likely to be \(1+, 2+\), or \(3+\) ? Explain.

A classmate of yours is convinced that he knows everything about electronegativity. (a) In the case of atoms \(\mathrm{X}\) and \(\mathrm{Y}\) having different electronegativities, he says, the diatomic molecule \(\mathrm{X}\) - Y must be polar. Is your classmate correct? (b) Your classmate says that the farther the two atoms are apart in a bond, the larger the dipole moment will be. Is your classmate correct?

Consider the collection of nonmetallic elements \(\mathrm{O}, \mathrm{P}, \mathrm{Te}\), I and 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 \(\mathrm{X}_{2} \mathrm{Y}_{3}\) ?

The substance chlorine monoxide, \(\mathrm{ClO}(\mathrm{g})\), is important in atmospheric processes that lead to depletion of the ozone layer. The ClO molecule has an experimental dipole moment of \(1.24 \mathrm{D}\), and the \(\mathrm{Cl}\) - \(\mathrm{O}\) bond length is \(1.60 \hat{A}\). (a) Determine the magnitude of the charges on the \(\mathrm{Cl}\) and \(\mathrm{O}\) atoms in units of the electronic charge, \(e\). (b) Based on the electronegativities of the elements, which atom would you expect to have a partial negative charge in the \(\mathrm{ClO}\) molecule? (c) Using formal charges as a guide, propose the dominant Lewis structure for the molecule. (d) The anion \(\mathrm{ClO}^{-}\)exists. What is the formal charge on the \(\mathrm{Cl}\) for the best Lewis structure for \(\mathrm{ClO}^{-}\)?

A major challenge in implementing the "hydrogen economy" is finding a safe, lightweight, and compact way of storing hydrogen for use as a fuel. The hydrides of light metals are attractive for hydrogen storage because they can store a high weight percentage of hydrogen in a small volume. For example, \(\mathrm{NaAlH}_{4}\) can release \(5.6 \%\) of its mass as \(\mathrm{H}_{2}\) upon decomposing to \(\mathrm{NaH}(s), \mathrm{Al}(s)\), and \(\mathrm{H}_{2}(g) . \mathrm{NaAlH}_{4}\) 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, predict the identity of the polyatomic anion. Draw a Lewis structure for this ion. (d) What is the formal charge on hydrogen in the polyatomic ion?

Although \(\mathrm{I}_{3}^{-}\)is known, \(\mathrm{F}_{3}^{-}\)is not. Which statement is the most correct explanation? (a) Iodine is more likely to be electron-deficient. (b) Fluorine is too small to accommodate three nonbonding electron pairs and two bonding electron pairs. (c) Fluorine is too electronegative to form anions. (d) \(\mathrm{I}_{2}\) is known but \(\mathrm{F}_{2}\) is not. (e) Iodine has a larger electron affinity than fluorine.

Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in \(\mathrm{O}_{3}\), (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 \(\mathrm{ClO}_{4}^{-}\)? (c) Perchlorate is a much stronger oxidizing agent than hypochlorite. Suggest an explanation.

The following three Lewis structures can be drawn for \(\mathrm{N}_{2} \mathrm{O}\) : \(: \mathrm{N} \equiv \mathrm{N}-\ddot{\mathrm{O}}: \longleftrightarrow: \ddot{\mathrm{N}}-\mathrm{N} \equiv \mathrm{O}: \longleftrightarrow: \ddot{\mathrm{N}}=\mathrm{N}=\ddot{\mathrm{O}}:\) (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 \(1.19 \AA\) Ä, slightly shorter than a typical bond (see Table 8.5). Based on these data, which resonance structure best represents \(\mathrm{N}_{2} \mathrm{O}\) ?

(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.

The \(\mathrm{Ti}^{2+}\) ion is isoelectronic with the Ca atom. (a) Writ the electron configurations of \(\mathrm{Ti}^{2+}\) and Ca. (b) Calculate th number of unpaired electrons for \(\mathrm{Ca}\) and for \(\mathrm{Ti}^{2+}\). (c) Wha charge would Ti have to be isoelectronic with \(\mathrm{Ca}^{2+}\) ?

You and a partner are asked to complete a lab entitled " \(\mathrm{Ox}\) ides of Ruthenium" that is scheduled to extend over two lal periods. The first lab, which is to be completed by your part ner, is devoted to carrying out compositional analysis. In th second lab, you are to determine melting points. Upon going to lab you find two unlabeled vials, one containing a soft yel low substance and the other a black powder. You also find the following notes in your partner's notebook-Compounc \(1: 76.0 \% \mathrm{Ru}\) and \(24.0 \% \mathrm{O}\) (by mass), Compound \(2: 61.2 \% \mathrm{R}\) 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? 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 o. your apparatus, \(1200^{\circ} \mathrm{C}\). (c) What is the identity of the yellow compound? (d) What is the identity of the black compound? (e) Which compound is molecular? (f) Which compound is ionic?

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 \(=k(I-E A)\), where \(k\) is a proportionality constant. (a) How does this definition explain why the electronegativity of \(\mathrm{F}\) is greater than that of \(\mathrm{Cl}\) even though \(\mathrm{Cl}\) 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 \(F\) under this definition. (d) Use your result from part (c) to determine the electronegativities of \(\mathrm{Cl}\) and \(\mathrm{O}\) using this scale. (e) Another scale for electronegativity defines electronegativity as the average of an atom's first ionization energy and its electron affinity. Using this scale, calculate the electronegativities for the halogens, and scale them so fluorine has an electronegativity of 4.0. On this scale, what is Br's electronegativity?

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 \(13.35 \% \mathrm{O}\) 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 \(\mathrm{C}\) atom and that there are \(\mathrm{C}-\mathrm{C}\) 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 \(30.4 \% \mathrm{~N}\). Measurements of its molecular mass yield a value of \(184.3 \mathrm{~g} / \mathrm{mol}\). 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 \(\mathrm{S}(g)\) 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) Are there any lone pairs of electrons in the \(\mathrm{P}_{4}\) molecule? (b) How many \(\mathrm{P}\) - P bonds are there in the molecule? (c) Draw a Lewis structure for a linear \(\mathrm{P}_{4}\) molecule that satisfies the octet rule. Does this molecule have resonance structures? (d) On the basis of formal charges, which is more stable, the linear molecule or the tetrahedral molecule?

Silicon, the element, is the heart of integrated circuits and computer chips in almost all of our electronic devices. Si has the same structure as diamond; each atom is singly bonded to four neighbors. Unlike diamond, silicon has a tendency to oxidize (to \(\mathrm{SiO}_{2}\), another extended solid) if exposed to air. (a) Estimate the enthalpy of reaction for the conversion of \(1 \mathrm{~cm}^{3}\) of silicon into \(\mathrm{SiO}_{2}\). (b) Unlike carbon, silicon rarely forms multiple bonds. Estimate the bond enthalpy of the \(\mathrm{Si}=\mathrm{Si}\) bond, assuming that the ratio of the \(\mathrm{Si}=\mathrm{Si}\) double bond enthalpy to that of the \(\mathrm{Si}-\mathrm{Si}\) single bond is the same as that for carbon-carbon bonds.