Classify the following bonds as ionic, polar cova­lent, or

What is a Lewis dot symbol? To what elements does the symbol mainly apply?

Write the Lewis structure for formic acid (HCOOH).

The molecular model shown here represents guanine, a component of a DNA molecule. Only the connections between the atoms are shown in this model.Draw a complete Lewis structure of the molecule, showing all the multiple bonds and lone pairs. (For color code, see inside back endpaper.)

Write Lewis dot symbols for the following atoms and ions: (a) I, (b) \(I^-\), (c) S, (d) \(S^{2-}\) (e) P, (f) \(P^{3-}\), (g) Na, (h) \(Na^+\), (i) Mg, (j) \(Mg^{2+}\) , (k) Al, (l) \(Al^{3+}\), (m) Pb. (n) \(Pb^{2+}\).

Draw three resonance structures for the thiocyanate ion, \(\mathrm{SCN}^?\) Rank the structures in decreasing order of importance.

Name one ionic compound that contains only non-metallic elements.

Draw the Lewis structure for \(\mathrm{BeF}_2\).

An ionic bond is formed between a cation \(A^+\) and an anion \(B^-\). How would the energy of the ionic bond [see Equation (9.2)] be affected by the following changes? (a) doubling the radius of \(A^+\), (b) tripling the charge on \(A^+\), (c) doubling the charges on \(A^+\) and \(B^-\), (d) decreasing the radii of \(A^+\) and \(B^-\) to half their original values.

Give the empirical formulas and names of the com­pounds formed from the following pairs of ions: (a) \(\mathrm{Rb^{+}}\) and \(\mathrm{I^{-}}\) (b) \(\mathrm{CS^{+}}\) and \(\mathrm{SO_4^{2-}}\) (c) \(\mathrm{Sr^{2+}}\) and \(\mathrm{N^{3-}}\) (d) \(\mathrm{Al^{3+}}\) and \(\mathrm{S^{2-}}\)

Use Lewis dot symbols to show the transfer of elec­trons between the following atoms to form cations and anions: (a) Na and F (b) K and S (c) Ba and 0 (d) Al and N.

What is Lewis's contribution to our understanding of the covalent bond?

Use an example to illustrate each of the following terms: lone pairs, Lewis structure, the octet rule, bond length.

What is the difference between a Lewis dot symbol and a Lewis structure?

Classify the following bonds as ionic, polar cova­lent, or covalent, and give your reasons: (a) the CC bond in \(\mathrm{H}_{3} \mathrm{CCH}_{3}\), (b) the KI bond in KI, (c) the NB bond in \(\mathrm{H}_{3} \mathrm{NBCl}_{3}\), (d) the CF bond in \(\mathrm{CF}_{4}\).

Classify the following bonds as ionic, polar cova­lent, or covalent, and give your reasons: (a) the SiSi bond in \(\mathrm{Cl}_{3} \mathrm{SiSiCl}_{3}\), (b) the SiCl bond in \(\mathrm{Cl}_{3} \mathrm{SiSiCl}_{3}\), (c) the CaF bond in \(\mathrm{CaF}_{2}\), (d) the NH bond in \(\mathrm{NH}_{3}\).

Summarize the essential features of the Lewis octet rule. The octet rule applies mainly to the second-period elements. Explain.

Write Lewis structures for the following species, including all resonance forms, and show formal charges: (a) \(\mathrm{HCO}_2^{-}\), (b) \(\mathrm{CH}_2 \mathrm{NO}_2^{-}\). Relative positions of the atoms are as follows:

Write three resonance structures for hydrazoic acid, \(\mathrm{HN_3}\). The atomic arrangement is HNNN. Show formal charges.

Of the noble gases, only Kr, Xe, and Rn are known to form a few compounds with O and/or F. Write Lewis structures for the following molecules: (a) \(\mathrm{XeF_2}\) (b) \(\mathrm{XeF_4}\) (c) \(\mathrm{XeF_6}\) (d) \(\mathrm{XeOF_4}\) (e) \(\mathrm{XeO_2F_4}\) In each case Xe is the central atom.

Write a Lewis structure for \(\mathrm{SbCl_5}\). Does this molecule obey the octet rule?

Write Lewis structures for \(\mathrm{ SeF_4}\) and \(\mathrm{SeF_6}\). Is the octet rule satisfied for Se?

Match each of the following energy changes with one of the processes given: ionization energy, elec­tron affinity, bond enthalpy, and standard enthalpy of formation. (a) \(\mathrm{F}(g)+e^{-} \longrightarrow \mathrm{F}^{-}(g)\) (b) \(\mathrm{F}_{2}(g) \longrightarrow 2 \mathrm{~F}(g)\) (c) \(\mathrm{Na}(g) \longrightarrow \mathrm{Na}^{+}(g)+e^{-}\) (d) \(\mathrm{Na}(s)+\frac{1}{2} \mathrm{~F}_{2}(g) \longrightarrow \mathrm{NaF}(s)\)

The formulas for the fluorides of the third-period elements are \(\mathrm{NaF}, \mathrm{MgF}_{2}, \mathrm{AlF}_{3}, \mathrm{SiF}_{4}, \mathrm{PF}_{5}, \mathrm{SF}_{6}\), and \(\mathrm{ClF}_{3}\). Classify these compounds as covalent or ionic.

Use ionization energy (see Table 8.2) and electron affinity values (see Table 8.3) to calculate the energy change (in kJ/mol) for the following reactions: (a) \(\mathrm{Li}(g)+\mathrm{I}(g) \longrightarrow \mathrm{Li}^{+}(g)+\mathrm{I}^{-}(g)\) (b) \(\mathrm{Na}(g)+\mathrm{F}(g) \longrightarrow \mathrm{Na}^{+}(g)+\mathrm{F}^{-}(g)\) (c) \(\mathrm{K}(g)+\mathrm{Cl}(g) \longrightarrow \mathrm{K}^{+}(g)+\mathrm{Cl}^{-}(g)\)

A rule for drawing plausible Lewis structures is that the central atom is invariably less electronegative than the surrounding atoms. Explain why this is so. Why does this rule not apply to compounds like \(\mathrm {H_2O}\) and \(\mathrm {NH_3}\)?

Using the following information and the fact that the average C—H bond enthalpy is 414 kJ/mol, estimate the standard enthalpy of formation of methane \(\mathrm {(CH_4)}\). \(\mathrm{C}(s) \longrightarrow \mathrm{C}(g) \quad \Delta H_{\mathrm{rxn}}^{\circ}=716 \mathrm{~kJ} / \mathrm{mol}\) \(2 \mathrm{H}_{2}(g) \longrightarrow 4 \mathrm{H}(g) \quad \Delta H_{\mathrm{rxn}}^{\circ}=872.8 \mathrm{~kJ} / \mathrm{mol}\)

Based on energy considerations, which of the fol­lowing reactions will occur more readily? (a) \(\mathrm{Cl}(g)+\mathrm{CH}_{4}(g) \longrightarrow \mathrm{CH}_{3} \mathrm{Cl}(g)+\mathrm{H}(g)\) (b) \(\mathrm{Cl}(g)+\mathrm{CH}_{4}(g) \longrightarrow \mathrm{CH}_{3}(g)+\mathrm{HCl}(g)\) (Hint: Refer to Table 9.4, and assume that the average bond enthalpy of the C—CI bond is 338 kJ/mol.)

The chlorine nitrate molecule \(\left(\mathrm{ClONO}_{2}\right)\) is believed to be involved in the destruction of ozone in the Antarctic stratosphere. Draw a plausible Lewis structure for this molecule.

Several resonance structures for the molecule \(\mathrm{CO}_{2}\) are shown next. Explain why some of them are likely to be of little importance in describing the bonding in this molecule.

For each of the following organic molecules draw a Lewis structure in which the carbon atoms are bonded to each other by single bonds: (a) \(\mathrm{C}_{2} \mathrm{H}_{6}\), (b) \(\mathrm{C}_{4} \mathrm{H}_{10}\), (c) \(\mathrm{C}_{5} \mathrm{H}_{12}\). For (b) and (c), show only structures in which each \(\mathrm{C}\) atom is bonded to no more than two other \(\mathrm{C}\) atoms.

Draw Lewis structures for the following chlorofluorocarbons (CFCs), which are partly responsible for the depletion of ozone in the stratosphere: (a) \(\mathrm{CFCl}_{3}\), (b) \(\mathrm{CF}_{2} \mathrm{Cl}_{2}\), (c) \(\mathrm{CHF}_{2} \mathrm{Cl}\), (d) \(\mathrm{CF}_{3} \mathrm{CHF}_{2}\).

Draw Lewis structures for the following organic molecules. In each there is one \(\mathrm{C}=\mathrm{C}\) bond, and the rest of the carbon atoms are joined by \(\mathrm{C}?\mathrm{C}\) bonds. \(\mathrm{C}_2\mathrm{H}_3\mathrm{F}\), \(\mathrm{C}_{3} \mathrm{H}_{6}\), \(\mathrm{C}_{4} \mathrm{H}_{8}\)

Calculate \(\Delta H^{\circ}\) for the reaction \(\mathrm{H}_{2}(g)+\mathrm{I}_{2}(g) \longrightarrow 2 \mathrm{HI}(g)\) using (a) Equation (9.3) and (b) Equation (6.18), given that \(\Delta H_{\mathrm{f}}^{\circ}\) for \(\mathrm{I}_{2}(g)\) is 61.0 kJ/mol.

The \(\mathrm{N}?\mathrm{O}\) bond distance in nitric oxide is 115 pm, which is intermediate between a triple bond (106 pm) and a double bond (120 pm). (a) Draw two resonance structures for \(\mathrm{NO}\) and comment on their relative importance. (b) Is it possible to draw a resonance structure having a triple bond between the atoms?

Write the formulas of the binary hydride for the second-period elements \(\mathrm{LiH}\) to \(\mathrm{HF}\). Comment on the change from ionic to covalent character of these compounds. Note that beryllium behaves differently from the rest of the Group 2A metals (see p. 350).

In 1998 scientists using a special type of electron microscope were able to measure the force needed to break a single chemical bond. If \(2.0\times10^{29}\mathrm{\ N}\) was needed to break a \(\mathrm{C}?\mathrm{Si}\) bond, estimate the bond enthalpy in kJ/mol. Assume that the bond had to be stretched by a distance of 2 Å \((2 \times 10^{-10} \mathrm{ m})\) before it is broken.

Although nitrogen dioxide \(\left(\mathrm{NO}_{2}\right)\) is a stable compound, there is a tendency for two such molecules to combine to form dinitrogen tetroxide \(\left(\mathrm{N}_{2} \mathrm{O}_{4}\right)\). Why? Draw four resonance structures of \(\mathrm{N}_{2} \mathrm{O}_{4}\), showing formal charges.

The American chemist Robert S. Mulliken suggested a different definition for the electronegativity (EN) of an element, given by \(\mathrm{EN}=\frac{\mathrm{IE}+\mathrm{EA}}{2}\) where IE is the first ionization energy and EA the electron affinity of the element. Calculate the electronegativities of O, F, and Cl using the above equation. Compare the electronegativities of these elements on the Mulliken and Pauling scale. (To convert to the Pauling scale, divide each EN value by 230 kJ/mol.)

Among the common inhaled anesthetics are: halothane: \(\mathrm{CF}_{3} \mathrm{CHClBr}\) enflurane: \(\mathrm{CHFClCF}_{2} \mathrm{OCHF}_{2}\) isoflurane: \(\mathrm{CF}_{3} \mathrm{CHClOCHF}_{2}\) methoxyflurane: \(\mathrm{CHCl}_{2} \mathrm{CF}_{2} \mathrm{OCH}_{3}\) Draw Lewis structures of these molecules.

Give a brief description of the medical uses of the following ionic compounds: \(\mathrm{AgNO}_{3}\), \(\mathrm{BaSO}_{4}\) \(\mathrm{CaSO}_{4}\), \(\mathrm{KI}\), \(\mathrm{Li}_{2} \mathrm{CO}_{3}\), \(\mathrm{Mg}(\mathrm{OH})_{2}\), \(\mathrm{MgSO}_{4}\), \(\mathrm{NaHCO}_{3}\), \(\mathrm{Na}_{2} \mathrm{CO}_{3}\), \(\mathrm{NaF}\), \(\mathrm{TiO}_{2}\), \(\mathrm{ZnO}\). You would need to do a Web search of some of these compounds.

Use Table 9.4 to estimate the bond enthalpy of the \(\mathrm{C}?\mathrm{C}\), \(\mathrm{N}?\mathrm{N}\), and \(\mathrm{O}?\mathrm{O}\) bonds in \(\mathrm{C}_{2} \mathrm{H}_{6}\), \(\mathrm{N}_{2} \mathrm{H}_{4}\), and \(\mathrm{H}_{2} \mathrm{O}_{2}\), respectively. What effect do lone pairs on adjacent atoms have on the strength of the particular bonds?

The isolated \(\mathrm{O}^{2-}\) ion is unstable so it is not possible to measure the electron affinity of the \(\mathrm{O}^{-}\) ion directly. Show how you can calculate its value by using the lattice energy of \(\mathrm {MgO}\) and the Born-Haber cycle. [Useful information: \(Mg(s)\rightarrow Mg(g)\ \Delta H^{\circ}=148\mathrm{\ kJ}/\mathrm{mol}\).]

Write the Lewis structure for carbon disulfide \(\mathrm {(CS_2)}\).

Identify the electrostatic potential maps shown here with HCI and LiH. In both diagrams, the H atom is on the left.

Draw the most reasonable Lewis structure of a molecule that contains a N atom, a C atom, and a H atom.

Why does \(\Delta H_\mathrm{rxn}^\circ\) calculated using bond enthalpies not always agree with that calculated using \(\Delta H_\mathrm{f}^\circ\) values?

Name five metals and five nonmetals that are very likely to form ionic compounds. Write formulas for compounds that might result from the combination of these metals and nonmetals. Name these compounds.

In which of the following states would NaCl be elec­trically conducting? (a) solid (b) molten (that is, melted) (c) dissolved in water. Explain your answers.

Calculate the enthalpy of the reaction \(\mathrm{H}_{2}(g)+\mathrm{F}_{2}(g) \longrightarrow 2 \mathrm{HF}(g)\) using (a) Equation (9.3) and (b) Equation (6.18).

Beryllium forms a compound with chlorine that has the empirical formula \(\mathrm{BeCl_2}\). How would you deter­mine whether it is an ionic compound? (The com­pound is not soluble in water.)

Compare the stability (in the solid state) of the fol­lowing pairs of compounds: (a) LiF and \(\mathrm{LiF_2}\) (containing the \(\mathrm{Li^{2+}}\) ion) (b) \(\mathrm{Cs_2O}\) and CsO (containing the \(\mathrm{O^-}\) ion) (c) \(\mathrm{CaBr_2}\) and \(\mathrm{CaBr_3}\), (containing the \(\mathrm{Ca^{3+}}\) ion).

Use the Born-Haber cycle outlined in Section 9.3 for LiF to calculate the lattice energy of NaCl. [The heat of sublimation of Na is 108 kJ/mol and \(\Delta H_\mathrm f^\circ(\mathrm{NaCl})=-411~\mathrm{ kj/mol}\). Energy needed to dissociate \(\frac{1}{2}\) mole of \(\mathrm{Cl_2}\) into CI atoms = 121.4 kJ.]

Calculate the lattice energy of calcium chloride given that the heat of sublimation of Ca is 121 kJ/mol and \(\Delta H_{\mathrm{f}}^{\circ}\left(\mathrm{CaCl}_{2}\right)=-795 \mathrm{~kJ} / \mathrm{mol}\). (See Tables 8.2 and 8.3 for other data.)

Arrange the following bonds in order of increasing ionic character: carbon to hydrogen, fluorine to hy­drogen, bromine to hydrogen, sodium to chlorine, potassium to fluorine, lithium to chlorine.

Four atoms are arbitrarily labeled D, E, F. and G. Their electronegativities are as follows: D = 3.8, E = 3.3, F = 2.8, and G = 1.3. If the atoms of these elements form the molecules DE, DG, EG, and DF. how would you arrange these molecules in order of increasing covalent bond character?

List the following bonds in order of increasing ionic character: cesium to fluorine, chlorine to chlorine, bromine to chlorine, silicon to carbon.

The skeletal structure of acetic acid shown below is correct, but some of the bonds are wrong. (a) Identify the incorrect bonds and explain what is wrong with them. (b) Write the correct Lewis structure for acetic acid.

Define bond length, resonance, and resonance structure. What are the rules for writing resonance structures?

Is it possible to "trap" a resonance structure of a compound for study? Explain.

What is a coordinate covalent bond? Is it different from a normal covalent bond?

The \(\mathrm{AlI_3}\) molecule has an incomplete octet around Al. Draw three resonance structures of the molecule in which the octet rule is satisfied for both the Al and the I atoms. Show formal charges.

In the vapor phase, beryllium chloride consists of discrete \(\mathrm{BeCl_2}\) molecules. Is the octet rule satisfied for Be in this compound? If not, can you form an octet around Be by drawing another resonance structure? How plausible is this structure?

For the reaction \(2 \mathrm{C}_{2} \mathrm{H}_{6}(g)+7 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{CO}_{2}(g)+6 \mathrm{H}_{2} \mathrm{O}(g)\) (a) Predict the enthalpy of reaction from the average bond enthalpies in Table 9.4. (b) Calculate the enthalpy of reaction from the standard enthalpies of formation (see Appendix 3) of the reactant and product molecules, and compare the result with your answer for part (a).

Classify the following substances as ionic com­pounds or covalent compounds containing discrete molecules: \(\mathrm{CH}_{4}, \mathrm{KF}, \mathrm{CO}, \mathrm{SiCl}_{4}, \mathrm{BaCl}_{2}\).

Which of the following are ionic compounds? Which are covalent compounds? \(\mathrm{RbCl}, \mathrm{PF}_{5}, \mathrm{BrF}_{3}, \mathrm{KO}_{2}, \mathrm{CI}_{4}\).

Attempts to prepare the compounds listed here as stable species under atmospheric conditions have failed. Suggest possible reasons for the failure. \(\mathrm {CF_2, LiO_2, CsCl2, PI_5}\)

Draw reasonable resonance structures for the follow­ing ions: (a) \(\mathrm{HSO}_{4}^{-}\) (b) \(\mathrm{PO}_{4}^{3-}\) (c) \(\mathrm{HSO}_{3}^{-}\) (d) \(\mathrm{SO}_{3}^{2-}\) (Hint: See comment on p. 398.)

Are the following statements true or false? (a) Formal charges represent actual separation of charges. (b) \(\Delta H_\mathrm{rxn}^\circ\) can be estimated from the bond enthalpies of reactants and products, (c) All second-period ele­ments obey the octet rule in their compounds. (d) The resonance structures of a molecule can be separated from one another.

The triiodide ion \(\mathrm{(I^-_3)}\) in which the I atoms are arranged in a straight line is stable, but the corre­sponding \(\mathrm{(F^-_3)}\) ion does not exist. Explain.

Compare the bond enthalpy of \(\mathrm{ F_2}\) with the energy change for the following process: \(\mathrm{F}_{2}(g) \longrightarrow \mathrm{F}^{+}(g)+\mathrm{F}^{-}(g)\) Which is the preferred dissociation for \(\mathrm{ F_2}\), energeti­cally speaking?

Methyl isocyanate \(\left(\mathrm{CH}_{3} \mathrm{NCO}\right)\) is used to make certain pesticides. In December 1984, water leaked into a tank containing this substance at a chemical plant, producing a toxic cloud that killed thousands of people in Bhopal, India. Draw Lewis structures for \(\mathrm{CH}_3\mathrm{NCO}\), showing formal charges.

Which of the following bonds is covalent, which is polar covalent and which is ionic? (a) the bond in CSCI (b) the bond in \(\mathrm {H_2S}\) (c) the NN bond in \(\mathrm {H_2NNH_2}\).

Why is it not possible for hydrogen to form double or triple bonds in covalent compounds?

Without referring to Figure 9.1, write Lewis dot sym­bols for atoms of the following elements: (a) Be (b) K (c) Ca (d) Ga (e) O (f) Br (g) N (h) I (i) As (j) F

Write formal charges for the nitrite ion \(\mathrm {(NO_2^-)}\).

The molecular model shown here represents acetamide, which is used as an organic solvent. Only the connections between the atoms are shown in this model. Draw two resonance structures for the molecule. showing the positions of multiple bonds and formal charges . (For color code, see inside back endpaper.)

Explain how ionization energy and electron affinity determine whether atoms of elements will combine to form ionic compounds.

Draw reasonable Lewis structures of sulfuric acid \(\mathrm{(H_2SO_4)}\).

The term "molar mass" was introduced in Chapter 3. What is the advantage of using the term “molar mass" when we discuss ionic compounds?

Write the Lewis structure of sulfuric tetrafluoride \(\mathrm{(SF_4)}\).

What is lattice energy and what role does it play in the stability of ionic compounds?

Explain how the lattice energy of an ionic compound such as KCl can be determined using the Born-Haber cycle. On what law is this procedure based?

Specify which compound in the following pairs of ionic compounds has the higher lattice energy: (a) KCl or MgO (b) LiF or LiBr (c) \(\mathrm{Mg_3N_2}\) or NaCl Explain your choice.

Define electronegativity, and explain the difference between electronegativity and electron affinity. Describe in general how the electronegativities of the elements change according to position in the periodic table.

What is a polar covalent bond? Name two compounds that contain one or more polar covalent bonds.

Write Lewis structures for the following molecules: (a) ICl (b) \(\mathrm{PH_3}\) (c) \(\mathrm{P_4}\) (each P is bonded to three other P atoms), (d) \(\mathrm{H_2S}\) (e) \(\mathrm{N_2H_4}\) (f) \(\mathrm{HClO_3}\) (g) \(\mathrm{COBr_2}\) (C is bonded to O and Br atoms).

List the following bonds in order of increasing ionic character: the lithium-to-fluorine bond in LiF, the potassium-to-oxygen bond in \(\mathrm{K_2O}\), the nitrogen-to-nitrogen bond in \(\mathrm{N_2}\), the sulfur-to-oxygen bond in \(\mathrm{SO_2}\), the chlorine-to-fluorine bond in \(\mathrm{ClF_3}\).

Write Lewis structures for the following ions: (a) \(\mathrm{O_2^{2-}}\) (b) \(\mathrm{C_2^{2-}}\) (c) \(\mathrm{NO^+}\) (d) \(\mathrm{NH_4^+}\) Show formal charges.

The following Lewis structures for (a) HCN, (b) \(\mathrm{C_2H_2}\), (c) \(\mathrm{SnO_2}\), (d) \(\mathrm{BF_3}\) (e) HOF, (f) HCOF. and (g) \(\mathrm{NF_3}\) are incorrect. Explain what is wrong with each one and give a correct structure for the molecule. (Relative positions of atoms are shown correctly.)

Why does the octet rule not hold for many com­pounds containing elements in the third period of the periodic table and beyond?

Give three examples of compounds that do not satisfy the octet rule. Write a Lewis structure for each.

Because fluorine has seven valence electrons \((2s^2 p^5)\), seven covalent bonds in principle could form around the atom. Such a compound might be \(\mathrm{FH_7}\) or \(\mathrm{FCl_7}\). These compounds have never been prepared. Why?

From the following data, calculate the average bond enthalpy for the N—H bond: \(\mathrm{NH}_{3}(g) \longrightarrow \mathrm{NH}_{2}(g)+\mathrm{H}(g) \Delta H^{\circ} =435 \mathrm{~kJ} / \mathrm{mol}\) \(\mathrm{NH}_{2}(g) \longrightarrow \mathrm{NH}(g)+\mathrm{H}(g) \Delta H^{\circ} =381 \mathrm{~kJ} / \mathrm{mol}\) \(\mathrm{NH}(g) \longrightarrow \mathrm{N}(g)+\mathrm{H}(g) \Delta H^{\circ} =360 \mathrm{~kJ} / \mathrm{mol}\)

For the reaction \(\mathrm{O}(g)+\mathrm{O}_{2}(g) \longrightarrow \mathrm{O}_{3}(g) \quad \Delta H^{\circ}=-107.2 \mathrm{~kJ} / \mathrm{mol}\) Calculate the average bond enthalpy in \(\mathrm O_3\).

The bond enthalpy of \(\mathrm {F_2}(g)\) is 156.9 kJ/mol. Calculate \(\Delta H_\mathrm f^\circ\) for F(g).

The amide group plays an important role in deter­mining the structure of proteins: Draw another resonance structure for this group. Show formal charges.

Give an example of an ion or molecule containing Al that (a) obeys the octet rule, (b) has an expanded octet, and (c) has an incomplete octet.

Draw four reasonable resonance structures for the \(\mathrm {PO_3F^{2-}}\) ion. The central P atom is bonded to the three O atoms and to the F atom. Show formal charges.

The following species have been detected in interstellar space: (a) CH (b) OH (c) \(\mathrm{C_2}\) (d) HNC (e) HCO Draw Lewis structures for these species and indicate whether they are diamagnetic or paramagnetic.

The amide ion, \(\mathrm{NH_2^-}\), is a Brønsted base. Represent the reaction between the amide ion and water.

Draw Lewis structures for the following organic mol­ecules: (a) tetrafluoroethylene \(\mathrm{(C_2F_4)}\), (b) propane \(\mathrm{(C_3H_8)}\), (C) butadiene \(\mathrm{(CH_2CHCHCH_2)}\), (d) propyne \(\mathrm{(CH_3CCH)}\), (e) benzoic acid \(\mathrm{(C_6H_5COOH)}\). (To draw \(\mathrm{C_6H_5COOH}\), replace a H atom in benzene with a COOH group.)

Use Lewis dot symbols to represent the formation of barium hydride.

Which of the following compounds has a larger lattice energy, LiCI or CsBr?

Use the second member of each group from Group 1A to Group 7A to show that the number of valence electrons on an atom of the element is the same as its group number.

Write the Lewis structure for the nitrite ion \(\mathrm {(NO_2^-)}\).

Consider three possible atomic arrangements for cyanamide \((CH_2N_2)\) : (a) \(H_2CNN\) (b) \(H_2NCN\) (c) \(HNNCH\) Using formal charges as a guide, determine which is the most plausible arrangement.

Explain what an ionic bond is.

Name one ionic compound that contains a polyatomic cation and a polyatomic anion (see Table 2.3).

Draw the Lewis structure for arsenic pentafluoride \(\mathrm{(AsF_5)}\).

Explain why ions with charges greater than 3 are seldom found in ionic compounds.

Write the Lewis dot symbols of the reactants and products in the following reactions. (First balance the equations.) (a) \(\mathrm{Sr}+\mathrm{Se} \longrightarrow \mathrm{SrSe}\) (b) \(\mathrm{Ca}+\mathrm{H}_{2} \longrightarrow \mathrm{CaH}_{2}\) (c) \(\mathrm{Li}+\mathrm{N}_{2} \longrightarrow \mathrm{Li}_{3} \mathrm{~N}\) (d) \(\mathrm{Al}+\mathrm{S} \longrightarrow \mathrm{Al}_{2} \mathrm{~S}_{3}\)

For each of the following pairs of elements, state whether the binary compound they form is likely to be ionic or covalent. Write the empirical formula and name of the compound: (a) I and CI (b) Mg and F

For each of the following pairs of elements, state whether the binary compound they form is likely to be ionic or covalent. Write the empirical formula and name of the compound: (a) B and F (b) K and Br

How many lone pairs are on the underlined atoms in these compounds? HBr, \(\mathrm{H_2\underline S}\), \(\mathrm{\underline CH_4}\)

Compare single, double, and triple bonds in a mol­ecule, and give an example of each. For the same bonding atoms, how does the bond length change from single bond to triple bond?

Compare the properties of ionic compounds and covalent compounds.

Explain the concept of formal charge. Do formal charges represent actual separation of charges?

Write Lewis structures for the following molecules and ions: (a) \(\mathrm{NC_{13}}\) (b) OCS (c) \(\mathrm{H_2O_2}\) (d) \(\mathrm{CH_3COO^-}\) (e) \(\mathrm{CN^-}\) (f) \(\mathrm{CH_3CH_2NH_3^+}\)

Write Lewis structures for the following molecules and ions: (a) \(\mathrm{OF_2}\) (b) \(\mathrm{N_2F_2}\) (c) \(\mathrm{Si_2H_6}\) (d) \(\mathrm{OH^-}\) (e) \(\mathrm{CH_2ClCOO^-}\) (f) \(\mathrm{CH_3NH_3^+}\)

Draw two resonance structures for diazomethane, \(\mathrm{CH}_2 \mathrm{~N}_2\). Show formal charges. The skeletal structure of the molecule is

Draw three resonance structures for the molecule \(\mathrm{N_2O_3}\) (atomic arrangement is \(\mathrm{ONNO_2}\)). Show formal charges.

Draw three reasonable resonance structures for the \(\mathrm{OCN^-}\) ion. Show formal charges.

Write Lewis structures for the reaction \(\mathrm{AlCl_3+Cl^- \longrightarrow AlCl_4^-}\) What kind of bond joins Al and Cl in the product?

What is bond enthalpy? Bond enthalpies of polyatomic molecules are average values, whereas those of diatomic molecules can be accurately determined Why?

Explain why the bond enthalpy of a molecule is usu­ally defined in terms of a gas-phase reaction. Why are bond-breaking processes always endothermic and bond-forming processes always exothermic?

Describe some characteristics of an ionic compound such as KF that would distinguish it from a covalent compound such as benzene \(\mathrm {(C_6H_6)}\).

Write Lewis structures for \(\mathrm {BrF_3}\), \(\mathrm {ClF_5}\), and \(\mathrm {IF_7}\). Identify those in which the octet rule is not obeyed.

Write three reasonable resonance structures for the azide ion \(\mathrm {N^-_3}\) in which the atoms are arranged as NNN. Show formal charges

Which of the following molecules has the shortest nitrogen-to-nitrogen bond? Explain. \(\mathrm{N_2H_4, N_2O, N_2, N_2O_4}\)

Most organic acids can be represented as RCOOH, where COOH is the carboxyl group and R is the rest of the molecule. (For example, R is \(\mathrm{CH_3}\) in acetic acid, \(\mathrm{CH_3COOH}\).) (a) Draw a Lewis structure for the carboxyl group. (b) Upon ionization, the carboxyl group is converted to the carboxylate group, \(\mathrm{COO^-}\). Draw resonance structures for the carboxyl­ate group.

Which of the following species are isoelectronic? \(\mathrm{NH_4^+}\) \(\mathrm{C_6H_6}\) CO \(\mathrm{CH_4}\) \(\mathrm{N_2}\) \(\mathrm{B_3N_3H_6}\)

Draw Lewis structures for the following organic molecules: (a) methanol \(\left(\mathrm{CH}_{3} \mathrm{OH}\right)\); (b) ethanol \(\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\right)\); (c) tetraethyllead \(\left[\mathrm{Pb}\left(\mathrm{CH}_{2} \mathrm{CH}_{3}\right)_{4}\right]\), which was used in "leaded gasoline”; (d) methylamine \(\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right)\), which is used in tanning; (e) mustard gas \(\left(\mathrm{ClCH}_{2} \mathrm{CH}_{2} \mathrm{SCH}_{2} \mathrm{CH}_{2} \mathrm{Cl}\right)\), a poisonous gas used in World War I; (f) urea \(\left[\left(\mathrm{NH}_{2}\right)_{2} \mathrm{CO}\right]\), a fertilizer; and (g) glycine \(\left(\mathrm{NH}_{2} \mathrm{CH}_{2} \mathrm{COOH}\right)\), an amino acid.

Write Lewis structures for the following four isoelectronic species: (a) \(\mathrm{CO}\), (b) \(\mathrm{NO}^{+}\), (c) \(\mathrm{CN}^{-}\), (d) \(\mathrm{N}_2\). Show formal charges.

Oxygen forms three types of ionic compounds in which the anions are oxide \(\left(\mathrm{O}^{2-}\right)\), peroxide \(\left(\mathrm{O}_2^{2-}\right)\), and superoxide \(\left(\mathrm{O}_2^{-}\right)\). Draw Lewis structures of these ions.

Comment on the correctness of the statement, "All compounds containing a noble gas atom violate the octet rule."

Write three resonance structures for (a) the cyanate ion \(\left(\mathrm{NCO}^{-}\right)\) and (b) the isocyanate ion \(\left(\mathrm{CNO}^{-}\right)\) . In each case, rank the resonance structures in order of increasing importance.

(a) From the following data calculate the bond enthalpy of the \(\mathrm{F}_{2}^{-}\) ion. \(\mathrm{F}_2(g)\longrightarrow2\mathrm{F}(g)\quad\ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \Delta H_{\mathrm{rxn}}^{\circ}=156.9\mathrm{\ kJ}/\mathrm{mol}\) \(\mathrm{F}^-(g)\longrightarrow\mathrm{F}(g)+e^-\quad\ \ \ \ \ \ \ \ \ \Delta H_{\mathrm{rxn}}^{\circ}=333\mathrm{\ kJ}/\mathrm{mol}\) \(\mathrm{F}_2^-(g)\longrightarrow\mathrm{F}_2(g)+e^-\ \ \ \ \ \ \ \quad\Delta H_{\mathrm{rxn}}^{\circ}=290\mathrm{\ kJ}/\mathrm{mol}\) (b) Explain the difference between the bond enthalpies of \(\mathrm{F}_{2}\) and \(\mathrm{F}_{2}^{-}\).

The resonance concept is sometimes described by analogy to a mule, which is a cross between a horse and a donkey. Compare this analogy with the one used in this chapter, that is, the description of a rhinoceros as a cross between a griffin and a unicorn. Which description is more appropriate? Why?

What are the other two reasons for choosing (b) in Example 9.7?

In the Chemistry in Action essay on p. 399, nitric oxide is said to be one of about 10 of the smallest stable molecules known. Based on what you have learned in the course so far, write all the diatomic molecules you know, give their names, and show their Lewis structures.

Another possible skeletal structure for the \(\mathrm{CO}_{3}^{2-}\) (carbonate) ion besides the one presented in Example 9.5 is \(\mathrm{O}\mathrm{\ C}\mathrm{\ O}\mathrm{\ O}\). Why would we not use this structure to represent \(\mathrm{CO}_{3}^{2-}\)?

Draw a Lewis structure for nitrogen pentoxide \(\left(\mathrm{N}_{2} \mathrm{O}_{5}\right)\) in which each \(\mathrm{N}\) is bonded to three \(\mathrm{O}\) atoms.

In the gas phase, aluminum chloride exists as a dimer (a unit of two) with the formula \(\mathrm{Al}_{2} \mathrm{Cl}_{6}\). Its skeletal structure is given by Complete the Lewis structure and indicate the coordinate covalent bonds in the molecule.

The hydroxyl radical \((\mathrm {OH})\) plays an important role in atmospheric chemistry. It is highly reactive and has a tendency to combine with a \(\mathrm {H}\) atom from other compounds, causing them to break up. Thus, \(\mathrm {OH}\) is sometimes called a "detergent" radical because it helps to clean up the atmosphere. (a) Write the Lewis structure for the radical. (b) Refer to Table 9.4 and explain why the radical has a high affinity for \(\mathrm {H}\) atoms. (c) Estimate the enthalpy change for the following reaction: \(\mathrm{OH}(g)+\mathrm{CH}_{4}(g) \longrightarrow \mathrm{CH}_{3}(g)+\mathrm{H}_{2} \mathrm{O}(g)\) (d) The radical is generated when sunlight hits water vapor. Calculate the maximum wavelength (in nanometers) required to break an \(\mathrm{O}?\mathrm{H}\) bond in \(\mathrm{H}_{2} \mathrm{O}\).

Experiments show that it takes 1656 kJ/mol to break all the bonds in methane \(\left(\mathrm{CH}_{4}\right)\) and 4006 kJ/mol to break all the bonds in propane \(\left(\mathrm{C}_{3} \mathrm{H}_{8}\right)\). Based on these data, calculate the average bond enthalpy of the \(\mathrm{C}?\mathrm{C}\) bond.

Calculate \(\Delta H_{\mathrm{rxn}}^{\circ}\) at \(25^{\circ} \mathrm{C}\) of the reaction between carbon monoxide and hydrogen shown here using both bond enthalpy and \(\Delta H_{\mathrm{f}}^{\circ}\) values.

Calculate \(\Delta H_{\mathrm{rxn}}^{\circ}\) at \(25^{\circ} \mathrm{C}\) of the reaction between ethylene and chlorine shown here using both bond enthalpy and \(\Delta H_{\mathrm{f}}^{\circ}\) values. (\(\Delta H_{\mathrm{f}}^{\circ}\) for \(\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{Cl}_{2}\), is -132 kJ/mol.)

Vinyl chloride \(\left(\mathrm{C}_2 \mathrm{H}_3 \mathrm{Cl}\right)\) differs from ethylene \(\left(\mathrm{C}_2 \mathrm{H}_4\right)\) in that one of the H atoms is replaced with a CI atom. Vinyl chloride is used to prepare poly(vinyl chloride), which is an important polymer used in pipes. (a) Draw the Lewis structure of vinyl chloride. (b) The repeating unit in poly(vinyl chloride) is - \(\mathrm{CH}{ }_2-\mathrm{CHCl}_{-}\). Draw a portion of the molecule showing three such repeating units. (c) Calculate the enthalpy change when \(1.0 \times 10^3 \mathrm{~kg}\) of vinyl chloride forms poly(vinyl chloride).

Draw three resonance structures of sulfur dioxide \(\left(\mathrm{SO}_{2}\right)\). Indicate the most plausible structure(s).

A student in your class claims that magnesium oxide actually consists of \(\mathrm{Mg}^{+}\) and \(\mathrm{O}^{-}\) ions, not \(\mathrm{Mg}^{2+}\) and \(\mathrm{O}^{2-}\) ions. Suggest some experiments one could do to show that your classmate is wrong.

Shown here is a skeletal structure of borazine \(\left(\mathrm{B}_3\mathrm{N}_3\mathrm{H}_6\right)\). Draw two resonance structures of the molecule, showing all the bonds and formal charges. Compare its properties with the isoelectronic molecule benzene.

Calculate the wavelength of light needed to carry out the reaction \(\mathrm{H}_{2} \longrightarrow \mathrm{H}^{+}+\mathrm{H}^{-}\)

Sulfuric acid \(\left(\mathrm{H}_{2} \mathrm{SO}_{4}\right)\), the most important industrial chemical in the world, is prepared by oxidizing sulfur to sulfur dioxide and then to sulfur trioxide. Although sulfur trioxide reacts with water to form sulfuric acid, it forms a mist of fine droplets of \(\mathrm{H}_{2} \mathrm{SO}_{4}\), with water vapor that is hard to condense. Instead, sulfur trioxide is first dissolved in 98 percent sulfuric acid to form oleum \(\left(\mathrm{H}_2\mathrm{S}_2\mathrm{O}_7\right)\). On treatment with water, concentrated sulfuric acid can be generated. Write equations for all the steps and draw Lewis structures of oleum based on the discussion in Example 9.11.

From the lattice energy of \(\mathrm {KCl}\) in Table 9.1 and the ionization energy of \(\mathrm {K}\) and electron affinity of \(\mathrm {Cl}\) in Tables 8.2 and 8.3, calculate the \(\Delta H^{\circ}\) for the reaction \(\mathrm{K}(g)+\mathrm{Cl}(g) \longrightarrow \mathrm{KCl}(s)\)

The species \(\mathrm {H}\) is the simplest polyatomic ion. The geometry of the ion is that of an equilateral triangle. (a) Draw three resonance structures to represent the ion. (b) Given the following information \(2\mathrm{H}+\mathrm{H}^+\longrightarrow\mathrm{H}_3^+\quad\ \ \ \ \ \Delta H^{\circ}=-849\mathrm{\ kJ}/\mathrm{mol}\) and \(\mathrm{H}_2\longrightarrow2\mathrm{H}\quad\ \ \ \ \ \Delta H^{\circ}=436.4\mathrm{\ kJ}/\mathrm{mol}\) calculate \(\Delta H^{\circ}\) for the reaction \(\mathrm{H}^{+}+\mathrm{H}_{2} \longrightarrow \mathrm{H}_{3}^{+}\)

The bond enthalpy of the \(\mathrm{C}-\mathrm{N}\) bond in the amide group of proteins (see Problem 9.81) can be treated as an average of \(\mathrm{C}-\mathrm{N}\) and \(\mathrm{C}=\mathrm{N}\) bonds. Calculate the maximum wavelength of light needed to break the bond.

In 1999 an unusual cation containing only nitrogen \(\left(\mathrm{N}_{5}^{+}\right)\) was prepared. Draw three resonance structures of the ion, showing formal charges. (Hint: The \(\mathrm {N}\) atoms are joined in a linear fashion.)

Nitroglycerin, one of the most commonly used explosives, has the following structure \(\mathrm{CH}_{2} \mathrm{ONO}_{2}\) ? \(\mathrm{CHONO}_{2}\) ? \(\mathrm{CH}_{2} \mathrm{ONO}_{2}\) The decomposition reaction is \(4\mathrm{C}_3\mathrm{H}_5\mathrm{N}_3\mathrm{O}_9(l)\longrightarrow12\mathrm{CO}_2(g)+10\mathrm{H}_2\mathrm{O}(g)+6\mathrm{N}_2(g)+\mathrm{O}_2(g)\) The explosive action is the result of the heat released and the large increase in gaseous volume. (a) Calculate the \(\Delta H^{\circ}\) for the decomposition of one mole of nitroglycerin using both standard enthalpy of formation values and bond enthalpies. Assume that the two \(\mathrm {O}\) atoms in the \(\mathrm {NO}_2\) groups are attached to \(\mathrm {N}\) with one single bond and one double bond. (b) Calculate the combined volume of the gases at \(\mathrm {STP}\). (c) Assuming an initial explosion temperature of 3000 K, estimate the pressure exerted by the gases using the result from (b). (The standard enthalpy of formation of nitroglycerin is -371.1 kJ/mol.)

When irradiated with light of wavelength 471.7 nm, the chlorine molecule dissociates into chlorine atoms. One \(\mathrm {Cl}\) atom is formed in its ground electronic state while the other is in an excited state that is 10.5 kJ/mol above the ground state. What is the bond enthalpy of the \(\mathrm {Cl}_2\) molecule?

The reaction between fluorine \(\left(\mathrm{F}_{2}\right)\) with ethane \(\left(\mathrm{C}_{2} \mathrm{H}_{6}\right)\) produces predominantly \(\left(\mathrm{CF}_{4}\right)\) rather than \(\mathrm{C}_{2} \mathrm{F}_{6}\) molecules. Explain.

A new allotrope of oxygen, \(\mathrm{O}_{4}\), has been reported. The exact structure of \(\mathrm{O}_{4}\), is unknown, but the simplest possible structure would be a four-member ring consisting of oxygen-oxygen single bonds. The report speculated that the \(\mathrm{O}_{4}\) molecule might be useful as a fuel "because it packs a lot of oxygen in a small space, so it might be even more energy-dense than the liquefied ordinary oxygen used in rocket fuel." (a) Draw a Lewis structure for \(\mathrm{O}_{4}\), and write a balanced chemical equation for the reaction between ethane, \(\mathrm{C}_{2} \mathrm{H}_{6}(g)\), and \(\mathrm{O}_{4}(g)\) to give carbon dioxide and water vapor. (b) Estimate \(\Delta H^{\circ}\) for the reaction. (c) Write a chemical equation illustrating the standard enthalpy of formation of \(\mathrm{O}_{4}(g)\) and estimate \(\Delta H_{\mathrm{f}}^{\circ}\). (d) Assuming the oxygen allotropes are in excess, which will release more energy when reacted with ethane (or any other fuel): \(\mathrm{O}_{2}(g)\) or \(\mathrm{O}_{4}(g)\)? Explain using your answers to parts (a)-c).

Because bond formation is exothermic, when two gas-phase atoms come together to form a diatomic molecule it is necessary for a third atom or molecule to absorb the energy that is released. Otherwise the molecule will undergo dissociation. If two atoms of hydrogen combine to form \(\mathrm{H}_{2}(g)\), what would be the increase in velocity of a third hydrogen atom that absorbs the energy released from this process?

Estimate \(\Delta H_{\mathrm{f}}^{\circ}\) for sodium astatide \((\mathrm {NaAt})\) according to the equation \(\mathrm{Na}(s)+\frac{1}{2} \mathrm{At}_{2}(s) \longrightarrow \mathrm {NaAt}(s)\)

For the reaction \(\mathrm{H}_{2}(g)+\mathrm{C}_{2} \mathrm{H}_{4}(g) \longrightarrow \mathrm{C}_{2} \mathrm{H}_{6}(g)\) (a) Estimate the enthalpy of reaction, using the bond enthalpy values in table 9.4. (b) Calculate the enthalpy of reaction using standard enthalpies of formation. (\(\Delta H_\mathrm f^\circ\) for \(\mathrm H_2\), \(\mathrm {C_2H_4}\), and \(\mathrm {C_2H_6}\) are 0, 52.3 kJ/mol, and –84.7 kJ/mol, respectively.)

Write Lewis dot symbols for the following ions: (a) \(\mathrm{Li^+}\) (b) \(\mathrm{Cl^-}\) (c) \(\mathrm{S^{2-}}\) (d) \(\mathrm{Sr^{2+}}\) (e) \(\mathrm{N^{3-}}\)

Draw three resonance structures for the chlorate ion, \(\mathrm{ClO_3^-}\). Show formal charges.

Hydrazine borane, \(\mathrm{NH}_{2} \mathrm{NH}_{2} \mathrm{BH}_{3}\), has been proposed as a hydrogen storage material. When reacted with lithium hydride \((\mathrm {LiH})\), hydrogen gas is released \(\mathrm{NH}_{2} \mathrm{NH}_{2} \mathrm{BH}_{3}+\mathrm{LiH} \longrightarrow \mathrm{LiNH}_{2} \mathrm{NHBH}_{3}+\mathrm{H}_{2}\) Write Lewis structures for \(\mathrm{NH}_{2} \mathrm{NH}_{2} \mathrm{BH}_{3}\); and \(\mathrm{NH}_{2} \mathrm{NHBH}_{3}^{-}\), and assign all formal charges.