9.1 What is a Lewis dot symbol? To what elements does the symbol mainly apply?
9.2 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.
9.3 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.
9.4
Write Lewis dot symbols for the following ions: (a) Li+, (b) Cl−, (c) S2−, (d) Sr2+, (e) N3−.9.5
Write Lewis dot symbols for the following atoms and ions: (a) I, (b) I−, (c) S, (d) S2−, (e) P, (f) P3−, (g) Na, (h) Na+, (i) Mg, (j) Mg2+, (k) Al, (l) Al3+, (m) Pb, (n) Pb2+.
9.6 Explain what an ionic bond is.
9.7 Explain how ionization energy and electron affinity determine whether atoms of elements will combine to form ionic compounds.
9.8 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.
9.9 Name one ionic compound that contains only nonmetallic elements.
9.10 Name one ionic compound that contains a polyatomic cation and a polyatomic anion (see Table 2.3).
9.11 Explain why ions with charges greater than 3 are seldom found in ionic compounds.
9.12 The term “molar mass” was introduced in Chapter 3. What is the advantage of using the term “molar mass” when we discuss ionic compounds?
9.13
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.9.14
Beryllium forms a compound with chlorine that has the empirical formula BeCl2. How would you determine whether it is an ionic compound? (The compound is not soluble in water.)
9.15
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.9.16
Give the empirical formulas and names of the compounds formed from the following pairs of ions: (a) Rb+ and I−, (b) Cs+ and 
, (c) Sr2+ and N3−, (d) Al3+ and S2−.9.17 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 O, (d) Al and N.
9.18
Write the Lewis dot symbols of the reactants and products in the following reactions. (First balance the equations.)Sr + Se → SrSe
Ca + H2 → CaH2
Li + N2 → Li3N
Al + S → Al2S3
9.19
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 Cl, (b) Mg and F.9.20
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.
9.21 What is lattice energy and what role does it play in the stability of ionic compounds?
9.22 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?
9.23
Specify which compound in the following pairs of ionic compounds has the higher lattice energy: (a) KCl or MgO, (b) LiF or LiBr, (c) Mg3N2 or NaCl. Explain your choice.9.24
Compare the stability (in the solid state) of the following pairs of compounds: (a) LiF and LiF2 (containing the Li2+ ion), (b) Cs2O and CsO (containing the O− ion), (c) CaBr2 and CaBr3 (containing the Ca3+ion).
9.25
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 
(NaCl) = −411 kJ/mol. Energy needed to dissociate 
mole of Cl2 into Cl atoms = 121.4 kJ].9.26
Calculate the lattice energy of calcium chloride given that the heat of sublimation of Ca is 121 kJ/mol and (CaCl2) = −795 kJ/mol. (See Tables 8.2 and 8.3 for other data.)
9.27 What is Lewis's contribution to our understanding of the covalent bond?
9.28 What is the difference between a Lewis dot symbol and a Lewis structure?
9.29
How many lone pairs are on the underlined atoms in these compounds? HBr, H2S, CH4.9.30 Distinguish among single, double, and triple bonds in a molecule, and give an example of each.
9.31 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.
9.32 What is a polar covalent bond? Name two compounds that contain one or more polar covalent bonds.
9.33
List these bonds in order of increasing ionic character: the lithium-to-fluorine bond in LiF, the potassium-to-oxygen bond in K2O, the nitrogen-to-nitrogen bond in N2, the sulfur-to-oxygen bond in SO2, the chlorine-to-fluorine bond in ClF3.9.34
Arrange these 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.9.35
Four atoms are arbitrarily labeled D, E, F, and G. Their electronegativities are: 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?9.36 List these bonds in order of increasing ionic character: cesium to fluorine, chlorine to chlorine, bromine to chlorine, silicon to carbon.
9.37
Classify these bonds as ionic, polar covalent, or covalent, and give your reasons: (a) the CC bond in H3CCH3, (b) the KI bond in KI, (c) the NB bond in H3NBCl3, (d) the ClO bond in ClO2.9.38
Classify these bonds as ionic, polar covalent, or covalent, and give your reasons: (a) the SiSi bond in Cl3SiSiCl3, (b) the SiCl bond in Cl3SiSiCl3, (c) the CaF bond in CaF2, (d) the NH bond in NH3.
9.39 Summarize the essential features of the Lewis octet rule. The octet rule applies mainly to the second-period elements. Explain.
9.40 Explain the concept of formal charge. Do formal charges on a molecule represent actual separation of charges?
9.41
Write Lewis structures for these molecules: (a) ICl, (b) PH3, (c) P4 (each P is bonded to three other P atoms), (d) H2S, (e) N2H4, (f) HClO3, (g) COBr2 (C is bonded to O and Br atoms).9.42
Write Lewis structures for these ions: (a) 
, (b) 
, (c) NO+, (d) 
. Show formal charges.9.43
The following Lewis structures for (a) HCN, (b) C2H2, (c) SnO2, (d) BF3, (e) HOF, (f) HCOF, and (g) NF3 are incorrect. Explain what is wrong with each one and give a correct structure for the molecule. (Relative positions of atoms are shown correctly.)
9.44 The skeletal structure of acetic acid in this structure 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.
9.45 Define bond length, resonance, and resonance structure.
9.46 Is it possible to “trap” a resonance structure of a compound for study? Explain.
9.47
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 that 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?9.48 What are the other two reasons for choosing (b) in Example 9.7?
9.49
Write Lewis structures for these species, including all resonance forms, and show formal charges: (a) 
(b) 
Relative positions of the atoms are as follows:
9.50
Draw three resonance structures for the chlorate ion, 
Show formal charges.9.51
Write three resonance structures for hydrazoic acid, HN3. The atomic arrangement is HNNN. Show formal charges.9.52
Draw two resonance structures for diazomethane, CH2N2. Show formal charges. The skeletal structure of the molecule is
9.53
Draw three reasonable resonance structures for the OCN− ion. Show formal charges.9.54 Draw three resonance structures for the molecule OCS in which the atoms are arranged in the order OCS. Indicate formal charges.
9.55
Draw three resonance structures for the molecule N2O3 (atomic arrangement is ONNO2). Show formal charges.
9.56 Why does the octet rule not hold for many compounds containing elements in the third period of the periodic table and beyond?
9.57 Because fluorine has seven valence electrons (2s22p5), seven covalent bonds in principle could form around the atom. Such a compound might be FH7 or FCl7. These compounds have never been prepared. Why?
9.58 What is a coordinate covalent bond? Is it different from a normal covalent bond?
9.59
The BCl3 molecule has an incomplete octet around B. Draw three resonance structures of the molecule in which the octet rule is satisfied for both the B and the Cl atoms. Show formal charges.9.60 In the vapor phase, beryllium chloride consists of discrete molecular units BeCl2. 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?
9.61
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 these molecules: (a) XeF2, (b) XeF4, (c) XeF6, (d) XeOF4, (e) XeO2F2. In each case Xe is the central atom.9.62 Write a Lewis structure for SbCl5. Is the octet rule obeyed in this molecule?
9.63
Write Lewis structures for SeF4 and SeF6. Is the octet rule satisfied for Se?9.64
Write Lewis structures for the reaction
What kind of bond is between Al and Cl in the product?
9.65 Define bond enthalpy. Bond enthalpies of polyatomic molecules are average values. Why?
9.66 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?
9.67
From these data, calculate the average bond enthalpy for the N—H bond:
9.68 For the reaction
calculate the average bond enthalpy in O3.
9.69
The bond enthalpy of F2(g) is 156.9 kJ/mol. Calculate for F(g).9.70 (a) For the reaction
predict the enthalpy of reaction from the average bond enthalpies in Table 9.3. (b) Calculate the enthalpy of reaction from the standard enthalpies of formation (see Appendix 2) of the reactant and product molecules, and compare the result with your answer for part (a).
9.71
Match each of these energy changes with one of the processes given: ionization energy, electron affinity, bond enthalpy, standard enthalpy of formation.F(g) + e− → F−(g)
F2(g) → 2F(g)
Na(g) → Na+(g) + e−
9.72 The formulas for the fluorides of the third-period elements are NaF, MgF2, AlF3, SiF4, PF5, SF6, and ClF3. Classify these compounds as covalent or ionic.
9.73
Use the ionization energy (see Table 8.2) and electron affinity values (see Table 8.3) to calculate the energy change (in kilojoules) for these reactions:
Li(g) + I(g) → Li+(g) + I−(g)
Na(g) + F(g) → Na+(g) + F−(g)
K(g) + Cl(g) → K+(g) + Cl−(g)
9.74 Describe some characteristics of an ionic compound such as KF that would distinguish it from a covalent compound such as CO2.
9.75
Write Lewis structures for BrF3, ClF5, and IF7. Identify those in which the octet rule is not obeyed.9.76 Write three reasonable resonance structures of the azide ion
in which the atoms are arranged as NNN. Show formal charges.9.77
The amide group plays an important role in determining the structure of proteins:
Draw another resonance structure of this group. Show formal charges.
9.78 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.
9.79
Draw four reasonable resonance structures for the PO3F2− ion. The central P atom is bonded to the three O atoms and to the F atom. Show formal charges.9.80 Attempts to prepare these as stable species under atmospheric conditions have failed. Suggest reasons for the failure.
9.81
Draw reasonable resonance structures for these sulfur-containing ions: (a) 
, (b) 
, (c) 
, (d) 
.9.82 True or false: (a) Formal charges represent actual separation of charges; (b)
can be estimated from 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.9.83 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 H2O and NH3?
9.84 Using this information:
and the fact that the average C—H bond enthalpy is 414 kJ/mol, estimate the standard enthalpy of formation of methane (CH4).
9.85
Based on energy considerations, which of these two reactions will occur more readily?Cl(g) + CH4(g) → CH3Cl(g) + H(g)
Cl(g) + CH4(g) → CH3(g) + HCl(g)
(Hint: Refer to Table 9.3, and assume that the average bond enthalpy of the C—Cl bond is 338 kJ/mol.)
9.86
Which of these molecules has the shortest nitrogen-to-nitrogen bond? Explain.
9.87
Most organic acids can be represented as RCOOH, in which COOH is the carboxyl group and R is the rest of the molecule. (For example, R is CH3 in acetic acid, CH3COOH.) (a) Draw a Lewis structure of the carboxyl group. (b) Upon ionization, the carboxyl group is converted to the carboxylate group, COO−. Draw resonance structures of the carboxylate group.9.88 Which of these molecules or ions are isoelectronic:
, C6H6, CO, CH4, N2, B3N3H6?9.89
These species have been detected in interstellar space: (a) CH, (b) OH, (c) C2, (d) HNC, (e) HCO. Draw Lewis structures of these species and indicate whether they are diamagnetic or paramagnetic.9.90 The amide ion,
, is a Brønsted base. Represent the reaction between the amide ion and water in terms of Lewis structures.9.91
Draw Lewis structures of these organic molecules: (a) tetrafluoroethylene (C2F4), (b) propane (C3H8), (c) butadiene (CH2CHCHCH2), (d) propyne (CH3CCH), (e) benzoic acid (C6H5COOH). (Hint: To draw C6H5COOH, replace an H atom in benzene with a COOH group.)9.92 The triiodide ion
in which the I atoms are arranged as III is stable, but the corresponding 
ion does not exist. Explain.9.93 Compare the bond enthalpy of F2 with the energy change for this process:
Which is the preferred dissociation for F2, energetically speaking?
9.94 Methyl isocyanate, CH3NCO, is used to make certain pesticides. In December 1984, water leaked into a tank containing this substance at a chemical plant to produce a toxic cloud that killed thousands of people in Bhopal, India. Draw Lewis structures for this compound, showing formal charges.
9.95
The chlorine nitrate molecule (ClONO2) is believed to be involved in the destruction of ozone in the Antarctic stratosphere. Draw a plausible Lewis structure for the molecule.9.96 Several resonance structures of the molecule CO2 are given here. Explain why some of them are likely to be of little importance in describing the bonding in this molecule.
9.97
Draw a Lewis structure for each of these organic molecules in which the carbon atoms are bonded to each other by single bonds: C2H6, C4H10, C5H12.9.98 Draw Lewis structures for these chlorofluorocarbons (CFCs), which are partly responsible for the depletion of ozone in the stratosphere: CFCl3, CF2Cl2, CHF2Cl, CF3CHF2.
9.99
Draw Lewis structures for these organic molecules, in each of which there is one C==C bond and the rest of the carbon atoms are joined by C—C bonds: C2H3F, C3H6, C4H8.9.100 Calculate ΔH° for the reaction
using (a) Equation (9.3) and (b) Equation (6.18), given that
for I2(g) is 61.0 kJ/mol.9.101
Draw Lewis structures of these organic molecules: (a) methanol (CH3OH); (b) ethanol (CH3CH2OH); (c) tetraethyllead [Pb(CH2CH3)4], which was used in “leaded” gasoline; (d) methylamine (CH3NH2); (e) mustard gas (ClCH2CH2SCH2CH2Cl), a poisonous gas used in World War I; (f) urea [(NH2)2CO], a fertilizer; (g) glycine (NH2CH2COOH), an amino acid.9.102 Write Lewis structures for these four isoelectronic species: (a) CO, (b) NO+, (c) CN−, (d) N2. Show formal charges.
9.103
Oxygen forms three types of ionic compounds in which the anions are oxide (O2−), peroxide (
), and superoxide (
). Draw Lewis structures of these ions.9.104 Comment on the correctness of this statement: All compounds containing a noble gas atom violate the octet rule.
9.105
(a) From these data:
calculate the bond enthalpy of the
ion. (b) Explain the difference between the bond enthalpies of F2 and .9.106 Write three resonance structures for the isocyanate ion (CNO−). Rank them in importance.
9.107
The only known argon-containing compound is HArF, which was prepared in 2000. Draw a Lewis structure of the compound.9.108
Experiments show that it takes 1656 kJ/mol to break all the bonds in methane (CH4) and 4006 kJ/mol to break all the bonds in propane (C3H8). Based on these data, calculate the average bond enthalpy of the C—C bond.9.109
Among the common inhaled anesthetics arehalothane: CF3CHClBr
enflurane: CHFClCF2OCHF2
isoflurane: CF3CHClOCHF2
methoxyflurane: CHCl2CF2OCH3
Draw Lewis structures of these molecules.
9.110 Industrially, ammonia is synthesized by the Haber process at high pressures and temperatures:
Calculate the enthalpy change for the reaction using (a) bond enthalpies and Equation (9.3) and (b) the
values in Appendix 2.
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
) was prepared. Draw three resonance structures of the ion, showing formal charges. (Hint: The N atoms are joined in a linear fashion.)
