Ch 9 - 105AE

Problem 1E Why are bonding theories important? Provide some examples of what bonding theories can predict.

Problem 1SAQ Which pair of elements is most likely to form an ionic bond? a) nitrogen and oxygen b) carbon and hydrogen c) sulfur and oxygen d) calcium and oxygen

Which set of elements is arranged in order of increasing electronegativity?

Why do chemical bonds form? What basic forces are involved in bonding?

Problem 3E What are the three basic types of chemical bonds? What happens to electrons in the bonding atoms in each type?

Which is the correct Lewis structure for magnesium bromide?

Problem 4E How do you determine how many dots to put around the Lewis symbol of an element?

Problem 5E Describe the octet rule in the Lewis model.

Which compound is likely to have an incomplete octet? a) NH3 b) SO3 c) N2O d) BH3

Which compound has the highest magnitude of lattice energy? A) MgS b)CaS c) SrS d) BaS

Problem 6E According to the Lewis model, what is a chemical bond?

Problem 7E How do you draw an ionic Lewis structure?

Which set of compounds is arranged in order of increasing magnitude of lattice energy?

Problem 7SAQ Which pair of atoms forms the most polar bond? a) N and O b) C and O c) C and F d) N and F

Problem 8E How can Lewis structures be used to determine the formula of ionic compounds? Give an example.

Problem 8SAQ Which pair of atoms forms a nonpolar covalent bond? a) C and S b) C and O c) B and O d) Na and Cl

Problem 9E What is lattice energy?

Which is the correct Lewis structure for nitrogen trif uoride?

Problem 10E Why is the formation of solid sodium chloride from solid sodium and gaseous chlorine exothermic, even though it takes more energy to form the Na+ ion than the amount of energy released upon formation of Cl- ?

Which is the correct Lewis structure for CO32-?]

Problem 11E What is the Born–Haber cycle? List each of the steps in the cycle and show how the cycle is used to calculate lattice energy.

Determine the formal charge of nitrogen in this structure. a) +1 b) +2 c) -1 d) -2

Problem 12E How does lattice energy relate to ionic radii? To ion charge?

A Lewis structure for the acetate ion is shown here: Which structure is the best resonance structure for the acetate ion?

Problem 13E How does the ionic bonding model explain the relatively high melting points of ionic compounds?

Use formal charge to choose the best Lewis structure for CH3SOCH3.

Problem 14E How does the ionic bonding model explain the non conductivity of ionic solids, and at the same time the conductivity of ionic solutions?

Use bond energies to determine ?Hrxn for the reaction between ethanol and hydrogen chloride. a) -1549 kJ b) 1549 kJ c) -12 kJ d) 12 kJ

Problem 15E Within a covalent Lewis structure, what is the difference between lone pair and bonding pair electrons?

Consider the halogenation of ethene, where X is a generic halogen: Use bond energies to determine which halogen produces the most exothermic halogenation reaction with ethene. The C—F, C —Br, and C —I bond energies are 552 kJ/mol, 280 kJ/mol, and 209 kJ/mol, respectively. Find all other necessary bond energies in Table 9.3 . a) fluorine b) chlorine c) bromine d) iodine

Problem 16E In what ways are double and triple covalent bonds different from single covalent bonds?

Problem 17E How does the Lewis model for covalent bonding account for why certain combinations of atoms are stable while others are not?

Problem 18E How does the Lewis model for covalent bonding account for the relatively low melting and boiling points of molecular compounds (compared to ionic compounds)?

Problem 20E Explain the difference between a pure covalent bond, a polar covalent bond, and an ionic bond.

Problem 21E Explain what is meant by the percent ionic character of a bond. Do any bonds have 100% ionic character?

Problem 22E What is a dipole moment?

Problem 23E What is the magnitude of the dipole moment formed by separating a proton and an electron by 100 pm? 200 pm?

Problem 24E What is the basic procedure for writing a covalent Lewis structure?

Problem 25E How do you determine the number of electrons that go into the Lewis structure of a molecule? A polyatomic ion?

Problem 26E What are resonance structures? What is a resonance hybrid?

Problem 27E Do resonance structures always contribute equally to the overall structure of a molecule? Explain.

Problem 28E What is formal charge? How is formal charge calculated? How is it helpful?

Problem 29E Why does the octet rule have exceptions? Give the three major categories of exceptions and an example of each.

Problem 30E What elements can have expanded octets? What elements should never have expanded octets?

Problem 31E What is bond energy? How can you use average bond energies to calculate enthalpies of reaction?

Write a Lewis symbol for each atom or ion. a. Al b. Na+ c. Cl d. Cl-

Write a Lewis symbol for each atom or ion. a. S2- b. Mg c. Mg2+ d. P

Write the Lewis symbols that represent the ions in each ionic compound. a. NaF b. CaO c. SrBr2 d. K2O

Write the Lewis symbols that represent the ions in each ionic compound. a. SrO b. Li2S c. CaI2 d. RbF

Problem 41E Use Lewis symbols to determine the formula for the compound that forms between each pair of elements. a. Sr and Se b. Ba and Cl c. Na and S d. Al and O

Problem 42E Use Lewis symbols to determine the formula for the compound that forms between each pair of elements: a. Ca and N b. Mg and I c. Ca and S d. Cs and F

Explain the trend in the lattice energies of the alkaline earth metal oxides.

Problem 44E Rubidium iodide has a lattice energy of -617 kJ/mol, while potassium bromide has a lattice energy of -671 kJ/mol. Why is the lattice energy of potassium bromide more exothermic than the lattice energy of rubidium iodide?

Problem 45E The lattice energy of CsF is -744 kJ/mol, whereas that of BaO is -3029 kJ/mol. Explain this large difference in lattice energy.

Problem 46E Arrange these compounds in order of increasing magnitude of lattice energy: KCl, SrO, RbBr, CaO.

Use the Born–Haber cycle and data from Appendix IIB, and Chapter 8 and this chapter to calculate the lattice energy of KCl. (?Hsub for potassium is 89.0 kJ/mol.)

Use the Born–Haber cycle and data from Appendix IIB and Table 9.3 to calculate the lattice energy of CaO. (?Hsub for calcium is 178 kJ/mol; IE1 and IE2 for calcium are 590 kJ/mol and 1145 kJ/mol, respectively; EA1 and EA2 for O are –141 kJ/mol and 744 kJ/mol, respectively.)

Write an appropriate Lewis structure for each compound. Make certain to distinguish between ionic and molecular compounds. a. BI3 b. K2S c. HCFO d. PBr3

Write an appropriate Lewis structure for each compound. Make certain to distinguish between ionic and molecular compounds. a. Al2O3 b. ClF5 c. MgI2 d. XeO4

Each compound contains both ionic and covalent bonds. Write ionic Lewis structures for each of them, including the covalent structure for the ion in brackets. Write resonance structures if necessary. a. BaCO3 b. Ca(OH)2 c. KNO3 d. LiIO

Each compound contains both ionic and covalent bonds. Write ionic Lewis structures for each of them, including the covalent structure for the ion in brackets. Write resonance structures if necessary. a. RbIO2 b. NH4Cl c. KOH d. Sr(CN)2

Carbon ring structures are common in organic chemistry. Draw a Lewis structure for each carbon ring structure, including any necessary resonance structures. a. C4H8 b. C4H4 c. C6H12 d. C6H6

Amino acids are the building blocks of proteins. The simplest amino acid is glycine (H2NCH2COOH). Draw a Lewis structure for glycine. (Hint: the central atoms in the skeletal structure are nitrogen bonded to carbon, which is bonded to another carbon. The two oxygen atoms are bonded directly to the rightmost carbon atom.)

Problem 91E Formic acid is responsible for the sting of ant bites. By mass, formic acid is 26.10% C, 4.38% H, and 69.52% O. The molar mass of formic acid is 46.02 g/mol. Find the molecular formula of formic acid and draw its Lewis structure.

Diazomethane is a highly poisonous, explosive compound because it readily evolves N2. Diazomethane has the following composition by mass: 28.57% C; 4.80% H; and 66.64% N. The molar mass of diazomethane is 42.04 g/mol. Find the molecular formula of diazomethane, draw its Lewis structure, and assign formal charges to each atom. Why is diazomethane not very stable? Explain.

The reaction of Fe2O3(s) with Al(s) to form Al2O3(s) and Fe(s) is called the thermite reaction and is highly exothermic. What role does lattice energy play in the exothermicity of the reaction?

NaCl has a lattice energy of -787 kJ/mol. Consider a hypothetical salt XY. X3+ has the same radius of Na+ and Y3- has the same radius as Cl-. Estimate the lattice energy of XY.

Problem 95E Draw the Lewis structure for nitric acid (the hydrogen atom is attached to one of the oxygen atoms). Include all three resonance structures by alternating the double bond among the three oxygen atoms. Use formal charge to determine which of the resonance structures is most important to the structure of nitric acid.

Phosgene (Cl2CO) is a poisonous gas used as a chemical weapon during World War I. It is a potential agent for chemical terrorism today. Draw the Lewis structure of phosgene. Include all three resonance forms by alternating the double bond among the three terminal atoms. Which resonance structure is the best?

The cyanate ion (OCN-) and the fulminate ion (CNO-) share the same three atoms but have vastly different properties. The cyanate ion is stable, while the fulminate ion is unstable and forms explosive compounds. The resonance structures of the cyanate ion were explored in Example 9.8 . Draw Lewis structures for the fulminate ion—including possible resonance forms—and use formal charge to explain why the fulminate ion is less stable (and therefore more reactive) than the cyanate ion.

Draw the Lewis structure for each organic compound from its condensed structural formula. a. C3H8 b. CH3OCH3 c. CH3COCH3 d. CH3COOH e. CH3CHO

Draw the Lewis structure for each organic compound from its condensed structural formula. a. C2H4 b. CH3NH2 c. HCHO d. CH3CH2OH e. HCOOH

Use Lewis structures to explain why Br3- and I3- are stable, while F3- is not.

Draw the Lewis structure for HCSNH2. (The carbon and nitrogen atoms are bonded together and the sulfur atom is bonded to the carbon atom.) Label each bond in the molecule as polar or nonpolar.

Draw the Lewis structure for urea, H2NCONH2, one of the compounds responsible for the smell of urine. (The central carbon atom is bonded to both nitrogen atoms and to the oxygen atom.) Does urea contain polar bonds? Which bond in urea is most polar?

Some theories of aging suggest that free radicals cause certain diseases and perhaps aging in general. As you know from the Lewis model, such molecules are not chemically stable and will quickly react with other molecules. According to certain theories, free radicals may attack molecules within the cell, such as DNA, changing them and causing cancer or other diseases. Free radicals may also attack molecules on the surfaces of cells, making them appear foreign to the body’s immune system. The immune system then attacks the cells and destroys them, weakening the body. Draw Lewis structures for each free radical implicated in this theory of aging. a. O2- b. O- c. OH d. CH3OO (unpaired electron on terminal oxygen)

Free radicals are important in many environmentally significant reactions (see the Chemistry in the Environment box on free radicals in this chapter). For example, photochemical smog— smog that results from the action of sunlight on air pollutants— forms in part by these two steps: The product of this reaction, ozone, is a pollutant in the lower atmosphere. (Upper atmospheric ozone is a natural part of the atmosphere that protects life on Earth from ultraviolet light.) Ozone is an eye and lung irritant and also accelerates the weathering of rubber products. Rewrite the given reactions using the Lewis structure of each reactant and product. Identify the free radicals.

If hydrogen were used as a fuel, it could be burned according to this reaction: Use average bond energies to calculate rxn for this reaction and also for the combustion of methane (CH4). Which fuel yields more energy per mole? Per gram?

Calculate rxn for the combustion of octane (C8H18), a component of gasoline, by using average bond energies and then calculate it using enthalpies of formation from Appendix IIB . What is the percent difference between your results? Which result would you expect to be more accurate?

Draw Lewis structures for each compound. a. Cl2O7 (no Cl — Cl bond) b. H3PO3 (two OH bonds) c. H3AsO4

The azide ion, N3-, is a symmetrical ion, all of whose contributing resonance structures have formal charges. Draw three important contributing structures for this ion.

List the following gas-phase ion pairs in order of the quantity of energy released when they form from separated gas-phase ions. Start with the pair that releases the least energy. Na+F-, Mg2+F-, Na+O2-, Mg2+O2-, Al3+O2-.

Calculate o for the reaction H2(g) + Br2(g) 2 HBr(g) using the bond energy values. The of HBr(g) is not equal to one-half of the value calculated. Account for the difference.

Problem 93AE

Problem 96AE

Problem 97AE

Problem 100AE

Problem 101AE

Problem 104AE

Problem 105AE

Problem 106AE

Problem 107AE

Problem 108AE

In the Chemistry and the Environment box on free radicals in this chapter, we discussed the importance of the hydroxyl radical in reacting with and eliminating many atmospheric pollutants. However, the hydroxyl radical does not clean up everything. For example, chlorofluorocarbons—which destroy stratospheric ozone—are not attacked by the hydroxyl radical. Consider the hypothetical reaction by which the hydroxyl radical might react with a chlorofluorocarbon: Use bond energies to explain why this reaction is improbable.

Hydrogen, a potential future fuel, can be produced from carbon (from coal) and steam by this reaction: Use average bond energies to calculate rxn for the reaction.

Hydrogenation reactions are used to add hydrogen across double bonds in hydrocarbons and other organic compounds. Use average bond energies to calculate rxn for the hydrogenation reaction.

Write Lewis structures for each molecule or ion. Use expanded octets as necessary. a. ClF5 b. AsF6- c. Cl3PO d. IF5

Which of these compounds has the stronger nitrogen–nitrogen bond? The shorter nitrogen–nitrogen bond?

Write Lewis structures for each molecule or ion. Use expanded octets as necessary. a. PF5 b. I3- c. SF4 d. GeF4

Write Lewis structures for each molecule or ion. Include resonance structures if necessary and assign formal charges to all atoms. If necessary, expand the octet on the central atom to lower formal charge. a. SO42- b. HSO4- c. SO3 d. BrO2-

Write the Lewis structure for each ion. Include resonance structures if necessary and assign formal charges to all atoms. If necessary, expand the octet on the central atom to lower formal charge. a. PO43- b. CN- c. SO32- d. ClO2-

Problem 36E Write an electron configuration for Ne. Then write a Lewis symbol for Ne and show which electrons from the electron configuration are included in the Lewis symbol.

Problem 35E Write an electron configuration for N. Then write a Lewis symbol for N and show which electrons from the electron configuration are included in the Lewis symbol.

Problem 34E How does the electron sea model explain the conductivity of metals? The malleability and ductility of metals?

Problem 33E What is the electron sea model for bonding in metals?

Problem 32E Explain the difference between endothermic reactions and exothermic reactions with respect to the bond energies of the bonds broken and formed.

Problem 19E What is electronegativity? What are the periodic trends in electronegativity?

An ionic bond is formed between a cation A1 and an anion B2. How would the energy of the ionic bond [see Equation (9.2)] be affected by the following changes? (a) doubling the radius of A1, (b) tripling the charge on A1, (c) doubling the charges on A1 and B2, (d) decreasing the radii of A1 and B2 to half their original values.

Give the empirical formulas and names of the compounds formed from the following pairs of ions: (a) Rb1 and I2, (b) Cs1 and SO22 4 , (c) Sr21 and N32, (d) Al31 and S22

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.

Write the Lewis dot symbols of the reactants and products in the following reactions. (First balance the equations.) (a) Sr 1 Se SrSe (b) Ca 1 H2 CaH2 (c) Li 1 N2 Li3N (d) Al 1 S Al2S3

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.

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.

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) Mg3N2 or NaCl. Explain your choice.

Compare the stability (in the solid state) of the following pairs of compounds: (a) LiF and LiF2 (containing the Li21 ion), (b) Cs2O and CsO (containing the O2 ion), (c) CaBr2 and CaBr3 (containing the Ca31 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 Hf(NaCl) 5 2411 kJ/mol. Energy needed to dissociate 1 2 mole of Cl2 into Cl atoms 5 121.4 kJ.]

Calculate the lattice energy of calcium chloride given that the heat of sublimation of Ca is 121 kJ/mol and Hf(CaCl2) 5 2795 kJ/mol. (See Tables 8.2 and 8.3 for other data.)

What is Lewiss 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?

How many lone pairs are on the underlined atoms in these compounds? HBr, H2S, CH4

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.

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.

List the following 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-tooxygen bond in SO2, the chlorine-to-fluorine bond in ClF3.

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 5 3.8, E 5 3.3, F 5 2.8, and G 5 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.

Classify the following 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 CF bond in CF4

Classify the following 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.

Summarize the essential features of the Lewis octet rule. The octet rule applies mainly to the secondperiod elements. Explain.

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

Write Lewis structures for the following molecules and ions: (a) NCl3, (b) OCS, (c) H2O2, (d) CH3COO2, (e) CN2, (f) CH3CH2NH3 1.

Write Lewis structures for the following molecules and ions: (a) OF2, (b) N2F2, (c) Si2H6, (d) OH2, (e) CH2ClCOO2, (f) CH3NH3 1.

Write Lewis structures for the following ions: (a) O22 2 , (b) C22 2 , (c) NO1, (d) NH1 4 . Show formal charges.

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.) (a) O Q HOCPN O (b) HPCPCPH (c) O OQOSnOO OQ (d) SF G D OQ FS O Q O B A F N A F (e) HOO OPQ OFS (g) F G D SOQ OFS S Q SQS QS (f) OCOQ OFS OQ H G O D

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. H A A H O A HPCOCOOOH

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.

Write Lewis structures for the following species, including all resonance forms, and show formal charges: (a) HCO2 2, (b) CH2NO2 2. Relative positions of the atoms are as follows: OH O HC CN OH O

Draw three resonance structures for the chlorate ion, ClO3 2. Show formal charges.

Write three resonance structures for hydrazoic acid, HN3. The atomic arrangement is HNNN. Show formal charges.

Draw two resonance structures for diazomethane, CH2N2. Show formal charges. The skeletal structure of the molecule is H CNN H

Draw three resonance structures for the molecule N2O3 (atomic arrangement is ONNO2). Show formal charges.

Draw three reasonable resonance structures for the OCN2 ion. Show formal charges.

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 2p5 ), 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?

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

The AlI3 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 BeCl2 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?

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) XeF2, (b) XeF4, (c) XeF6, (d) XeOF4, (e) XeO2F2. In each case Xe is the central atom.

Write a Lewis structure for SbCl5. Does this molecule obey the octet rule?

Write Lewis structures for SeF4 and SeF6. Is the octet rule satisfied for Se?

Write Lewis structures for the reaction AlCl3 1 Cl2 AlCl4 2 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?

From the following data, calculate the average bond enthalpy for the NH bond: NH3(g) NH2(g) 1 H(g) H 5 435 kJ/mol NH2(g) NH(g) 1 H(g) H 5 381 kJ/mol NH(g) N(g) 1 H(g) H 5 360 kJ/mol

For the reaction O(g) 1 O2(g) O3(g) H 5 2107.2 kJ/mol Calculate the average bond enthalpy in O3.

The bond enthalpy of F2(g) is 156.9 kJ/mol. Calculate Hf for F(g).

For the reaction 2C2H6(g) 1 7O2(g) 4CO2(g) 1 6H2O(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: CH4, KF, CO, SiCl4, BaCl2

Which of the following are ionic compounds? Which are covalent compounds? RbCl, PF5, BrF3, KO2, CI4

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) F(g) 1 e2 F2(g) (b) F2(g) 2F(g) (c) Na(g) Na1(g) 1 e2 (d) Na(s) 1 1 2F2(g) NaF(s)

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.

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) Li(g) 1 I(g) Li1(g) 1 I 2(g) (b) Na(g) 1 F(g) Na1(g) 1 F2(g) (c) K(g) 1 Cl(g) K1(g) 1 Cl2(g)

Describe some characteristics of an ionic compound such as KF that would distinguish it from a covalent compound such as benzene (C6H6).

Write Lewis structures for BrF3, ClF5, and IF7. Identify those in which the octet rule is not obeyed.

Write three reasonable resonance structures for the azide ion N2 3 in which the atoms are arranged as NNN. Show formal charges.

The amide group plays an important role in determining the structure of proteins: A H O B ONOCO S S O 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 PO3F22 ion. The central P atom is bonded to the three O atoms and to the F atom. Show formal charges.

Attempts to prepare the compounds listed here as stable species under atmospheric conditions have failed. Suggest possible reasons for the failure. CF2, LiO2, CsCl2, PI5

Draw reasonable resonance structures for the following ions: (a) HSO2 4 , (b) PO32 4 , (c) HSO2 3 , (d) SO22 3 . (Hint: See comment on p. 396.)

Are the following statements true or false? (a) Formal charges represent actual separation of charges. (b) Hrxn 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

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?

Using the following information and the fact that the average CH bond enthalpy is 414 kJ/mol, estimate the standard enthalpy of formation of methane (CH4). C(s) C(g) Hrxn 5 716 kJ/mol 2H2(g) 4H(g) Hrxn 5 872.8 kJ/mol

Based on energy considerations, which of the following reactions will occur more readily? (a) Cl(g) 1 CH4(g) CH3Cl(g) 1 H(g) (b) Cl(g) 1 CH4(g) CH3(g) 1 HCl(g) (Hint: Refer to Table 9.4, and assume that the average bond enthalpy of the CCl bond is 338 kJ/mol.)

Which of the following molecules has the shortest nitrogen-to-nitrogen bond? Explain. N2H4, N2O, N2, N2O4

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 CH3 in acetic acid, CH3COOH.) (a) Draw a Lewis structure for the carboxyl group. (b) Upon ionization, the carboxyl group is converted to the carboxylate group, COO2. Draw resonance structures for the carboxylate group

Which of the following species are isoelectronic? NH4 1, C6H6, CO, CH4, N2, B3N3H6

The following species have been detected in interstellar space: (a) CH, (b) OH, (c) C2, (d) HNC, (e) HCO. Draw Lewis structures for these species and indicate whether they are diamagnetic or paramagnetic.

The amide ion, NH2 2, is a Brnsted base. Represent the reaction between the amide ion and water.

Draw Lewis structures for the following organic molecules: (a) tetrafluoroethylene (C2F4), (b) propane (C3H8), (c) butadiene (CH2CHCHCH2), (d) propyne (CH3CCH), (e) benzoic acid (C6H5COOH). (To draw C6H5COOH, replace a H atom in benzene with a COOH group.)

The triiodide ion (I2 3 ) in which the I atoms are arranged in a straight line is stable, but the corresponding F2 3 ion does not exist. Explain

Compare the bond enthalpy of F2 with the energy change for the following process: F2(g) F1(g) 1 F2(g) Which is the preferred dissociation for F2, energetically speaking?

Methyl isocyanate (CH3NCO) 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 CH3NCO, showing formal charges.

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 this molecule.

Several resonance structures for the molecule CO2 are shown next. Explain why some of them are likely to be of little importance in describing the bonding in this molecule. (a) OQOPCPO OQ OQ (b) SOQOOCOOQOS 2 SOqC OS SOqCOOS O (c) Q O (d) 9.10

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) C2H6, (b) C4H10, (c) C5H12. For (b) and (c), show only structures in which each C atom is bonded to no more than two other C atoms

Draw Lewis structures for the following chlorofluorocarbons (CFCs), which are partly responsible for the depletion of ozone in the stratosphere: (a) CFCl3, (b) CF2Cl2, (c) CHF2Cl, (d) CF3CHF2.

Draw Lewis structures for the following organic molecules. In each there is one CC bond, and the rest of the carbon atoms are joined by CC bonds. C2H3F, C3H6, C4H8

Calculate H for the reaction H2(g) 1 I2(g) 2HI(g) using (a) Equation (9.3) and (b) Equation (6.18), given that Hf for I2(g) is 61.0 kJ/mol.

Draw Lewis structures for the following organic molecules: (a) methanol (CH3OH); (b) ethanol (CH3CH2OH); (c) tetraethyllead [Pb(CH2CH3)4], which was used in leaded gasoline; (d) methylamine (CH3NH2), which is used in tanning; (e) mustard gas (ClCH2CH2SCH2CH2Cl), a poisonous gas used in World War I; (f) urea [(NH2)2CO], a fertilizer; and (g) glycine (NH2CH2COOH), an amino acid.

Write Lewis structures for the following four isoelectronic species: (a) CO, (b) NO1, (c) CN2, (d) N2. Show formal charges.

Oxygen forms three types of ionic compounds in which the anions are oxide (O22), peroxide (O2 22), and superoxide (O2 2 ). 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 (NCO2) and (b) the isocyanate ion (CNO2). In each case, rank the resonance structures in order of increasing importance

(a) From the following data calculate the bond enthalpy of the F2 2 ion. F2(g) 2F(g) Hrxn 5 156.9 kJ/mol F2(g) F(g) 1 e2 Hrxn 5 333 kJ/mol F2 2(g) F2(g) 1 e2 Hrxn 5 290 kJ/mol (b) Explain the difference between the bond enthalpies of F2 and F2 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. 397, 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.

The NO 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 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 LiH to 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. 348)

Hydrazine borane, NH2NH2BH3, has been proposed as a hydrogen storage material. When reacted with lithium hydride (LiH), hydrogen gas is released NH2NH2BH3 1 LiH LiNH2NHBH3 1 H2 Write Lewis structures for NH2NH2BH3 and NH2NHBH3 2 and assign all formal charges

Although nitrogen dioxide (NO2) is a stable compound, there is a tendency for two such molecules to combine to form dinitrogen tetroxide (N2O4). Why? Draw four resonance structures of N2O4, showing formal charges.

Another possible skeletal structure for the CO22 3 (carbonate) ion besides the one presented in Example 9.5 is O C O O. Why would we not use this structure to represent CO22 3 ?

Draw a Lewis structure for nitrogen pentoxide (N2O5) in which each N is bonded to three O atoms.

In the gas phase, aluminum chloride exists as a dimer (a unit of two) with the formula Al2Cl6. Its skeletal structure is given by AlAl D Cl G Cl D Cl G G Cl D Cl G Cl D Complete the Lewis structure and indicate the coordinate covalent bonds in the molecule.

in atmospheric chemistry. It is highly reactive and has a tendency to combine with a H atom from other compounds, causing them to break up. Thus, 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 H atoms. (c) Estimate the enthalpy change for the following reaction: OH(g) 1 CH4(g) CH3(g) 1 H2O(g) (d) The radical is generated when sunlight hits water vapor. Calculate the maximum wavelength (in nanometers) required to break an OH bond in H2O.

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 CC bond.

Calculate Hrxn at 25C of the reaction between carbon monoxide and hydrogen shown here using both bond enthalpy and Hf values. 1 8n

Calculate Hrxn at 25C of the reaction between ethylene and chlorine shown here using both bond enthalpy and Hf values. (Hf for C2H4Cl2 is 2132 kJ/mol.) 1 8n

Draw three resonance structures of sulfur dioxide (SO2). Indicate the most plausible structure(s).

Vinyl chloride (C2H3Cl) differs from ethylene (C2H4) in that one of the H atoms is replaced with a Cl 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 CH2CHCl. Draw a portion of the molecule showing three such repeating units. (c) Calculate the enthalpy change when 1.0 3 103 kg of vinyl chloride forms poly(vinyl chloride).

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 3 1029 N was needed to break a CSi bond, estimate the bond enthalpy in kJ/mol. Assume that the bond had to be stretched by a distance of 2 (2 3 10210 m) before it is broken.

The American chemist Robert S. Mulliken suggested a different definition for the electronegativity (EN) of an element, given by EN 5 IE 1 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: CF3CHClBr enflurane: CHFClCF2OCHF2 isoflurane: CF3CHClOCHF2 methoxyflurane: CHCl2CF2OCH3 Draw Lewis structures of these molecules.

A student in your class claims that magnesium oxide actually consists of Mg1 and O2 ions, not Mg21 and O22 ions. Suggest some experiments one could do to show that your classmate is wrong

Shown here is a skeletal structure of borazine (B3N3H6). Draw two resonance structures of the molecule, showing all the bonds and formal charges. Compare its properties with the isoelectronic molecule benzene. N H H N B N B B H H H H

Calculate the wavelength of light needed to carry out the reaction H2 H1 1 H2

Sulfuric acid (H2SO4), 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 H2SO4 with water vapor that is hard to condense. Instead, sulfur trioxide is first dissolved in 98 percent sulfuric acid to form oleum (H2S2O7). 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 KCl in Table 9.1 and the ionization energy of K and electron affinity of Cl in Tables 8.2 and 8.3, calculate the H for the reaction K(g) 1 Cl(g) KCl(s)

The species H3 1 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 2H 1 H1 H3 1 H 5 2849 kJ/mol and H2 2H H 5 436.4 kJ/mol calculate H for the reaction H1 1 H2 H3

The bond enthalpy of the CN bond in the amide group of proteins (see Problem 9.81) can be treated as an average of CN and CN bonds. Calculate the maximum wavelength of light needed to break the bond.

In 1999 an unusual cation containing only nitrogen (N5 1) was prepared. Draw three resonance structures of the ion, showing formal charges. (Hint: The N atoms are joined in a linear fashion.)

Nitroglycerin, one of the most commonly used explosives, has the following structure CH2ONO2 CHONO2 CH2ONO2 The decomposition reaction is 4C3H5N3O9(l) 12CO2(g) 1 10H2O(g) 1 6N2(g) 1 O2(g) The explosive action is the result of the heat released and the large increase in gaseous volume. (a) Calculate the H for the decomposition of one mole of nitroglycerin using both standard enthalpy of formation values and bond enthalpies. Assume that the two O atoms in the NO2 groups are attached to N with one single bond and one double bond. (b) Calculate the combined volume of the gases at 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 2371.1 kJ/mol.)

Give a brief description of the medical uses of the following ionic compounds: AgNO3, BaSO4, CaSO4, KI, Li2CO3, Mg(OH)2, MgSO4, NaHCO3, Na2CO3, NaF, TiO2, 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 CC, NN, and OO bonds in C2H6, N2H4, and H2O2, respectively. What effect do lone pairs on adjacent atoms have on the strength of the particular bonds?

The isolated O22 ion is unstable so it is not possible to measure the electron affinity of the O2 ion directly. Show how you can calculate its value by using the lattice energy of MgO and the BornHaber cycle. [Useful information: Mg(s) S Mg(g) H 5 148 kJ/mol.]

When irradiated with light of wavelength 471.7 nm, the chlorine molecule dissociates into chlorine atoms. One 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 Cl2 molecule?

Recall from Chapter 8 that the product of the reaction between Xe(g) and PtF6(g) was originally thought to be an ionic compound composed of Xe1 cations and PtF6 2 anions (see Figure 8.22). This prediction was based on the theoretical enthalpy of formation of XePtF6 calculated using a Born- Haber cycle. (a) The lattice energy for XePtF6 was estimated to be 460 kJ/mol. Explain whether or not this value is consistent with the lattice energies in Table 9.1. (b) Calculate Hf for XePtF6 given IE1 for Xe(g) is 1170 kJ/mol and EA1 for PtF6(g) is 770 kJ/mol. Comment on the expected stability of XePtF6 based on your calculation.

The reaction between fluorine (F2) with ethane (C2H6) produces predominantly CF4 rather than C2F6 molecules. Explain.

A new allotrope of oxygen, O4, has been reported. The exact structure of O4 is unknown, but the simplest possible structure would be a four-member ring consisting of oxygen-oxygen single bonds. The report speculated that the O4 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 O4 and write a balanced chemical equation for the reaction between ethane, C2H6(g), and O4(g) to give carbon dioxide and water vapor. (b) Estimate H for the reaction. (c) Write a chemical equation illustrating the standard enthalpy of formation of O4(g) and estimate Hf . (d) Assuming the oxygen allotropes are in excess, which will release more energy when reacted with ethane (or any other fuel): O2(g) or O4(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 H2(g), what would be the increase in velocity of a third hydrogen atom that absorbs the energy released from this process?

Estimate Hf for sodium astatide (NaAt) according to the equation Na(s) 1 1 2At2(s) NaAt(s) The information in Problem 8.147 may be useful.

Problem 49E Use covalent Lewis structures to explain why each element (or family of elements) occurs as diatomic molecules. a. hydrogen b. the halogens c. oxygen d. nitrogen

Use covalent Lewis structures to explain why the compound that forms between nitrogen and hydrogen has the formula NH3. Show why NH2 and NH4 are not stable.

Write the Lewis structure for each molecule. a. PH3 b. SCl2 c. HI d. CH4

Write the Lewis structure for each molecule. a. NF3 b. HBr c. SBr2 d. CCl4

Write the Lewis structure for each molecule. a. SF2 b. SiH4 c. HCOOH (both O bonded to C) d. CH3SH (C and S central)

Write the Lewis structure for each molecule. a. CH2O b. C2Cl4 c. CH3NH2 d. CFCl3 (C central)

Problem 55E Determine whether a bond between each pair of atoms would be pure covalent, polar covalent, or ionic. a. Br and Br b. C and Cl c. C and S d. Sr and O

Determine whether a bond between each pair of atoms would be pure covalent, polar covalent, or ionic. a. C and N b. N and S c. K and F d. N and N

Draw the Lewis structure for BrF with an arrow representing the dipole moment. Use Figure 9.10 to estimate the percent ionic character of the BrF bond.

Draw the Lewis structure for CO with an arrow representing the dipole moment. Use Figure 9.10 to estimate the percent ionic character of the CO bond.

Write the Lewis structure for each molecule or ion. a. CI4 b. N2O c. SiH4 d. Cl2CO

Write the Lewis structure for each molecule or ion. a. H3COH b. OH- c. BrO- d. O22+

Write the Lewis structure for each molecule or ion. a. N2H2 b. N2H4 c. C2H2 d. C2H4

Write the Lewis structure for each molecule or ion. a. H3COCH3 b. CN- c. NO2- d. ClO-

Write a Lewis structure that obeys the octet rule for each molecule or ion. Include resonance structures if necessary and assign formal charges to each atom. a. SeO2 b. CO32- c. ClO- d. NO2-

Write a Lewis structure that obeys the octet rule for each ion. Include resonance structures if necessary and assign formal charges to each atom. a. ClO3- b. ClO4- c. NO3- d. NH4+

Use formal charge to determine which Lewis structure is better:

Use formal charge to determine which Lewis structure is better:

How important is this resonance structure to the overall structure of carbon dioxide? Explain.

In N2O, nitrogen is the central atom and the oxygen atom is terminal. In OF2, however, oxygen is the central atom. Use formal charges to explain why.

Draw the Lewis structure (including resonance structures) for the acetate ion (CH3COO-). For each resonance structure, assign formal charges to all atoms that have formal charge.

Draw the Lewis structure (including resonance structures) for methyl azide (CH3N3). For each resonance structure, assign formal charges to all atoms that have formal charge.

What are the formal charges of the atoms shown in red?

What are the formal charges of the atoms shown in red?

Write the Lewis structure for each molecule (octet rule not followed). a. BCl3 b. NO2 c. BH3

Write the Lewis structure for each molecule (octet rule not followed). a. \(\mathrm{BBr}_{3}\) b. NO c. \(\mathrm{ClO}_{2}\)

Order these compounds in order of increasing carbon–carbon bond strength and in order of decreasing carbon–carbon bond length: HCCH, \(\mathrm{H}_2\mathrm{CCH}_2,\mathrm{\ H}_3\mathrm{CCH}_3\).

Ethanol is a possible fuel. Use average bond energies to calculate \(\Delta H_{\mathrm{rxn}}\) for the combustion of ethanol. \(\mathrm{CH}_3 \mathrm{CH}_2 \mathrm{OH}(g)+3 \mathrm{O}_2(g) \longrightarrow 2 \mathrm{CO}_2(g)+3 \mathrm{H}_2 \mathrm{O}(g)\)

The heat of atomization is the heat required to convert a molecule in the gas phase into its constituent atoms in the gas phase. The heat of atomization is used to calculate average bond energies. Without using any tabulated bond energies, calculate the average C — Cl bond energy from the following data: the heat of atomization of \(\mathrm{CH}_{4}\) is 1660 kJ/mol and of \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\) is 1495 kJ/mol.

Problem 113E A compound composed of only carbon and hydrogen is 7.743% hydrogen by mass. Propose a Lewis structure for the compound.

Calculate the heat of atomization (see previous problem) of C2H3Cl, using the average bond energies in Table 9.3 .

A compound composed of only carbon and chlorine is 85.5% chlorine by mass. Propose a Lewis structure for the compound.

The main component of acid rain (H2SO4) forms from SO2 pollutant in the atmosphere via these steps: Draw the Lewis structure for each of the species in these steps and use bond energies and Hess’s law to estimate rxn for the overall process. (Use 265 kJ/mol for the S — O single bond energy.)

A 0.167 g sample of an unknown acid requires 27.8 mL of 0.100 M NaOH to titrate to the equivalence point. Elemental analysis of the acid gives the following percentages by mass: 40.00% C, 6.71% H, 53.29% O. Determine the molecular formula, molar mass, and Lewis structure of the unknown acid.

Use the dipole moments of HF and HCl (given at the end of the problem) together with the percent ionic character of each bond ( Figure 9.10 ) to estimate the bond length in each molecule. How well does your estimated bond length agree with the bond length in Table 9.4 ?

Use average bond energies together with the standard enthalpy of formation of C(g) (718.4 kJ/mol) to estimate the standard enthalpy of formation of gaseous benzene, C6H6(g). (Remember that average bond energies apply to the gas phase only.) Compare the value you obtain using average bond energies to the actual standard enthalpy of formation of gaseous benzene, 82.9 kJ/mol. What does the difference between these two values tell you about the stability of benzene?

The standard state of phosphorus at \(25 ^\circ C\) is \(P_4\). This molecule has four equivalent P atoms, no double or triple bonds, and no expanded octets. Draw its Lewis structure.

The standard heat of formation of CaBr2 is -675 kJ/mol. The first ionization energy of Ca is 590 kJ/mol and its second ionization energy is 1145 kJ/mol. The heat of sublimation of Ca[Ca(s) Ca(g)] is 178 kJ/mol. The bond energy of Br2 is 193 kJ/mol, the heat of vaporization of Br2() is 31 kJ/mol, and the electron affinity of Br is -325 kJ/mol. Calculate the lattice energy of CaBr2.

The standard heat of formation of PI3(s) is -24.7 kJ/mol and the PI bond energy in this molecule is 184 kJ/mol. The standard heat of formation of P(g) is 334 kJ/mol and that of I2(g) is 62 kJ/mol. The I2 bond energy is 151 kJ/mol. Calculate the heat of sublimation of PI3[PI3(s) PI3(g)].

A compound has the formula C8H8 and does not contain any double or triple bonds. All the carbon atoms are chemically identical and all the hydrogen atoms are chemically identical. Draw the Lewis structure for this molecule.

Find the oxidation number of each sulfur in the molecule \(\mathrm{H}_2\mathrm{S}_4\), which has a linear arrangement of its atoms.

Ionic solids of the O- and O3- anions do not exist, while ionic solids of the O2- anion are common. Explain.

The standard state of sulfur is solid rhombic sulfur. Use the appropriate standard heats of formation given in Appendix II to find the average bond energy of the S = O in SO2.

Problem 126E Which statement is true of an endothermic reaction? a. Strong bonds break and weak bonds form. b. Weak bonds break and strong bonds form. c. The bonds that break and those that form are of approximately the same strength.

Problem 127E When a firecracker explodes, energy is obviously released. The compounds in the firecracker can be viewed as being “energy rich.” What does this mean? Explain the source of the energy in terms of chemical bonds.

Problem 128E A fundamental difference between compounds containing ionic bonds and those containing covalent bonds is the existence of molecules. Explain why molecules exist in solid covalent compounds but do not exist in solid ionic compounds.

In the very first chapter of this book, we described the scientific approach and put a special emphasis on scientific models or theories. In this chapter, we looked carefully at a model for chemical bonding (the Lewis model). Why is this theory successful? What are some of the limitations of the theory?

What does Talese tell us about the position of sports in American popular culture?

Talese describes DiMaggio as a kind of male Garbo, referring to the legendary, reclusive film star Greta Garbo (para. 85). The comparison suggests that DiMaggios detachment was a masculine ideal. Does this ideal still resonate? How does it square with the image of todays superstar athletes? Does our media- crazy era demand more engagement from our heroes? Are we still capable of being moved by the mystique (para. 85)?

Talese has said that his work is a highly personal response to the world as an Italian American outsider. What evidence do you find that it is an outsiders voice? How does the authorial voice of an outsider add nuance to the profile of Joe DiMaggio?

Talese suggests but doesnt come out and say that Joe DiMaggio and Marilyn Monroe were competitive about their celebrity. How does Talese use Monroes mythical status to develop his portrait of DiMaggio? Is he sympathetic to her? Does his profile of DiMaggio deepen our understanding of Monroe, or does she remain as tantalizingly out of reach as she seems to have been to DiMaggio? Explain your responses.

Several parts of the essay espe cially paragraphs 3 and 41 evoke Hemingway. How do the images and language of those allusions create another level of meaning? How do they add to the portrait of DiMaggio?

What is the overall tone of the essay? How does Talese achieve his tone? How does the tone add a layer of meaning?

Talese notes that sportswriters have called DiMaggio an immortal (para. 124). How does the essay both support and debunk that myth?

Explain the assumptions Talese makes about his audience based on his portrait of DiMaggio as an aging hero.

Do your children have fears about going away to camp?

Antigone was very protective of Oedipus in Oedipus at Colonus.

Identify the verbs and verbals in the following passages. Discuss how these verbs affect the tone of the passages. His vets warned us all along that the odds were against him, but we didnt really believe them. They had hope, too. How could a horse who appeared so full of life break his leg and be so suddenly close to death? His head was fine. His back was fine. His lungs and heart and chest were fine. In fact, after a while, his broken leg was fairly fine. It was another leg that was so worrisome, since the weight of his body constantly bearing down on the delicate structures inside his foot eventually damaged and destroyed them. Jane Smiley, Barbaro, The Heart in the Winners Circle She maintains her composure until we get in the car, then crumples in humiliation. Once were on the highway, forty miles to home, freezing rain coats our windows, but I cant see well enough to find a safe place to pull off. I drive with trepidation over the slippery road, through the foggy darkness, while Emily cries so hard it sounds as though she will break apart. Mommy, Im such a failure! she weeps. At first my attempts to comfort her only increase her misery, so I shut up. Im left to listen and worry about the road and think my resentful thoughts. I remember all the years in elementary school when she was benched in the classroom left to do bulletin boards for the teacher because shed already mastered what was being taught. I think about the studies that suggest that girls who compete in athletics are far less likely to drink or take drugs or become pregnant. Kris Vervaecke, A Spectators Notebook

Analyze the verbs in the opening paragraph of The Cruelest Sport (p. 622). How would you describe the verbs Joyce Carol Oates uses? How do they mirror the subject she is writing about? Do the verbs she uses tip you off that this is a piece that is more complex and aca demic than usual for sportswriting? Cite specific examples to support your view. Professional boxing is the only major American sport whose primary, and often murderous, energies are not coyly deflected by such artifacts as balls and pucks. Though highly ritualized, and as rigidly bound by rules, traditions, and taboos as any religious ceremony, it survives as the most primitive and terrifying of contests: two men, near- naked, fight each other in a brightly lit, elevated space roped in like an animal pen (though the ropes were originally to keep rowdy spectators out); two men climb into the ring from which only one, symbolically, will climb out. (Draws do occur in boxing, but are rare, and unpopular.)

Count the verbs in one of the passages in Exercise 2. Then categorize them into linking verbs and more vivid action verbs, and calculate the ratio. Do the same for several paragraphs of your own writing. Are you relying more on linking verbs, or are most of your verbs direct and precise action verbs?

How does the ionic bonding model explain the relatively high melting points of ionic compounds?

How does the ionic bonding model explain the nonconductivity of ionic solids, and at the same time the conductivity of ionic solutions?

Within a covalent Lewis structure, what is the difference between lone pair and bonding pair electrons?

In what ways are double and triple covalent bonds different from single covalent bonds?

How does the Lewis model for covalent bonding account for why certain combinations of atoms are stable while others are not?

How does the Lewis model for covalent bonding account for the relatively low melting and boiling points of molecular compounds (compared to ionic compounds)?

What is electronegativity? What are the periodic trends in electronegativity?

Explain the difference between a pure covalent bond, a polar covalent bond, and an ionic bond.

Explain what is meant by the percent ionic character of a bond. Do any bonds have 100% ionic character?

What is a dipole moment?

What is the magnitude of the dipole moment formed by separating a proton and an electron by 100 pm? 200 pm?

What is the basic procedure for writing a covalent Lewis structure?

How do you determine the number of electrons that go into the Lewis structure of a molecule? A polyatomic ion?

What are resonance structures? What is a resonance hybrid?

Do resonance structures always contribute equally to the overall structure of a molecule? Explain.

What is formal charge? How is formal charge calculated? How is it helpful?

Why does the octet rule have exceptions? Give the three major categories of exceptions and an example of each.

What elements can have expanded octets? What elements should never have expanded octets?

What is bond energy? How can you use average bond energies to calculate enthalpies of reaction?

Explain the difference between endothermic reactions and exothermic reactions with respect to the bond energies of the bonds broken and formed.

What is the electron sea model for bonding in metals?

How does the electron sea model explain the conductivity of metals? The malleability and ductility of metals?

Write an electron configuration for N. Then write a Lewis symbol for N and show which electrons from the electron configuration are included in the Lewis symbol.

Write an electron configuration for Ne. Then write a Lewis symbol for Ne and show which electrons from the electron configuration are included in the Lewis symbol.

Write a Lewis symbol for each atom or ion. a. Al b. Na+ c. Cl d. Cl

Write a Lewis symbol for each atom or ion. a. S2- b. Mg c. Mg2+ d. P

Write the Lewis symbols that represent the ions in each ionic compound. a. NaF b. CaO c. SrBr2 d. K2O

Write the Lewis symbols that represent the ions in each ionic compound. a. SrO b. Li2S c. CaI2 d. RbF

Use Lewis symbols to determine the formula for the compound that forms between each pair of elements. a. Sr and Se b. Ba and Cl c. Na and S d. Al and O

Use Lewis symbols to determine the formula for the compound that forms between each pair of elements: a. Ca and N b. Mg and I c. Ca and S d. Cs and F

Explain the trend in the lattice energies of the alkaline earth metal oxides.

Rubidium iodide has a lattice energy of -617 kJ/mol, while potassium bromide has a lattice energy of -671 kJ/mol. Why is the lattice energy of potassium bromide more exothermic than the lattice energy of rubidium iodide?

The lattice energy of CsF is -744 kJ/mol, whereas that of BaO is -3029 kJ/mol. Explain this large difference in lattice energy.

Arrange these compounds in order of increasing magnitude of lattice energy: KCl, SrO, RbBr, CaO.

Use the BornHaber cycle and data from Appendix IIB, and Chapter 8 and this chapter to calculate the lattice energy of KCl. (\(\Delta H_{sub}\) for potassium is 89.0 kJ/mol.)

Use the BornHaber cycle and data from Appendix IIBand Table 9.3 to calculate the lattice energy of CaO. ( H sub for calcium is 178 kJ/mol; IE1 and IE2 for calcium are 590 kJ/mol and 1145 kJ/mol, respectively; EA1 and EA2 for O are 141 kJ/mol and 744 kJ/mol, respectively.)

Use covalent Lewis structures to explain why each element (or family of elements) occurs as diatomic molecules. a. hydrogen b. the halogens c. oxygen d. nitrogen

Use covalent Lewis structures to explain why the compound that forms between nitrogen and hydrogen has the formula NH3. Show why NH2 and NH4 are not stable.

Write the Lewis structure for each molecule. a. PH3 b. SCl2 c. HI d. CH4

Write the Lewis structure for each molecule. a. NF3 b. HBr c. SBr2 d. CCl4

Write the Lewis structure for each molecule. a. SF 2 b. SiH 4 c. HCOOH (both O bonded to C) d. CH 3 SH (C and S central)

Write the Lewis structure for each molecule. a. CH 2 O b. C 2 Cl 4 c. CH 3 NH 2 d. CFCl 3 (C central)

Determine whether a bond between each pair of atoms would be pure covalent, polar covalent, or ionic. a. Br and Br b. C and Cl c. C and S d. Sr and O

Determine whether a bond between each pair of atoms would be pure covalent, polar covalent, or ionic. a. C and N b. N and S c. K and F d. N and N

Draw the Lewis structure for CO with an arrow representing the dipole moment. Use Figure 9.10 to estimate the percent ionic character of the CO bond.

Draw the Lewis structure for BrF with an arrow representing the dipole moment. Use Figure 9.10 to estimate the percent ionic character of the BrF bond.

Write the Lewis structure for each molecule or ion. a. CI4 b. N2O c. SiH4 d. Cl2CO

Write the Lewis structure for each molecule or ion. a. H3COH b. OH- c. BrO- d. O2 2-

Write the Lewis structure for each molecule or ion. a. N2H2 b. N2H4 c. C2H2 d. C2H4

Write the Lewis structure for each molecule or ion. a. H3COCH3 b. CNc. NO2 - d. ClO

Write a Lewis structure that obeys the octet rule for each molecule or ion. Include resonance structures if necessary and assign formal charges to each atom. a. SeO2 b. CO3 2- c. ClO- d. NO2 -

Write a Lewis structure that obeys the octet rule for each ion. Include resonance structures if necessary and assign formal charges to each atom. a. ClO3 - b. ClO4 - c. NO3 - d. NH4 +

Use formal charge to determine which Lewis structure is better:

Use formal charge to determine which Lewis structure is better:

How important is this resonance structure to the overall structure of carbon dioxide? Explain.

In N2O, nitrogen is the central atom and the oxygen atom is terminal. In OF2, however, oxygen is the central atom. Use formal charges to explain why

Draw the Lewis structure (including resonance structures) for the acetate ion (CH 3 COO- ). For each resonance structure, assign formal charges to all atoms that have formal charge.

Draw the Lewis structure (including resonance structures) for methyl azide (CH 3 N 3 ). For each resonance structure, assign formal charges to all atoms that have formal charge.

What are the formal charges of the atoms shown in red?

What are the formal charges of the atoms shown in red?

Write the Lewis structure for each molecule (octet rule not followed). a. BCl3 b. NO2 c. BH3

Write the Lewis structure for each molecule (octet rule not followed). a. BBr3 b. NO c. ClO2

Write the Lewis structure for each ion. Include resonance structures if necessary and assign formal charges to all atoms. If necessary, expand the octet on the central atom to lower formal charge. a. PO4 3- b. CN- c. SO3 2- d. ClO2

Write Lewis structures for each molecule or ion. Include resonance structures if necessary and assign formal charges to all atoms. If necessary, expand the octet on the central atom to lower formal charge. a. SO4 2- b. HSO4 - c. SO3 d. BrO2 -

Write Lewis structures for each molecule or ion. Use expanded octets as necessary. a. PF5 b. I3 - c. SF4 d. GeF4

Write Lewis structures for each molecule or ion. Use expanded octets as necessary. a. ClF5 b. AsF6 - c. Cl3PO d. IF5

Order these compounds in order of increasing carbon-carbon bond strength and in order of decreasing carbon-carbon bond length: \(\mathrm{HCCH}, \mathrm{H}_2 \mathrm{CCH}_2, \mathrm{H}_3 \mathrm{CCH}_3\).

Which of these compounds has the stronger nitrogennitrogen bond? The shorter nitrogennitrogen bond? H2NNH2, HNNH

Hydrogenation reactions are used to add hydrogen across double bonds in hydrocarbons and other organic compounds. Use average bond energies to calculate Hrxn for the hydrogenation reaction. H2CCH2(g) + H2(g) h H3CiCH3(g)

Ethanol is a possible fuel. Use average bond energies to calculate Hrxn for the combustion of ethanol.

Hydrogen, a potential future fuel, can be produced from carbon (from coal) and steam by this reaction: C(s) + 2 H2O(g) h 2 H2(g) + CO2(g) Use average bond energies to calculate Hrxn for the reaction.

In the Chemistry and the Environment box on free radicals in this chapter, we discussed the importance of the hydroxyl radical in reacting with and eliminating many atmospheric pollutants. However, the hydroxyl radical does not clean up everything. For example, chlorofluorocarbonswhich destroy stratospheric ozoneare not attacked by the hydroxyl radical. Consider the hypothetical reaction by which the hydroxyl radical might react with a chlorofluorocarbon: OH(g) + CF2Cl2(g) h HOF(g) + CFCl2(g) Use bond energies to explain why this reaction is improbable.

Write an appropriate Lewis structure for each compound. Make certain to distinguish between ionic and molecular compounds. a. BI3 b. K2S c. HCFO d. PBr3

Write an appropriate Lewis structure for each compound. Make certain to distinguish between ionic and molecular compounds. a. Al2O3 b. ClF5 c. MgI2 d. XeO4

Each compound contains both ionic and covalent bonds. Write ionic Lewis structures for each of them, including the covalent structure for the ion in brackets. Write resonance structures if necessary. a. BaCO3 b. Ca(OH)2 c. KNO3 d. LiIO

Each compound contains both ionic and covalent bonds. Write ionic Lewis structures for each of them, including the covalent structure for the ion in brackets. Write resonance structures if necessary. a. RbIO2 b. NH4Cl c. KOH d. Sr(CN)2

Carbon ring structures are common in organic chemistry. Draw a Lewis structure for each carbon ring structure, including any necessary resonance structures. a. C4H8 b. C4H4 c. C6H12 d. C6H6

Amino acids are the building blocks of proteins. The simplest amino acid is glycine (H2NCH2COOH). Draw a Lewis structure for glycine. (Hint: the central atoms in the skeletal structure are nitrogen bonded to carbon, which is bonded to another carbon. The two oxygen atoms are bonded directly to the rightmost carbon atom.)

Formic acid is responsible for the sting of ant bites. By mass, formic acid is 26.10% C, 4.38% H, and 69.52% O. The molar mass of formic acid is 46.02 g/mol. Find the molecular formula of formic acid and draw its Lewis structure.

Diazomethane is a highly poisonous, explosive compound because it readily evolves N2. Diazomethane has the following composition by mass: 28.57% C; 4.80% H; and 66.64% N. The molar mass of diazomethane is 42.04 g/mol. Find the molecular formula of diazomethane, draw its Lewis structure, and assign formal charges to each atom. Why is diazomethane not very stable? Explain.

The reaction of Fe2O3(s) with Al(s) to form Al2O3(s) and Fe(s) is called the thermite reaction and is highly exothermic. What role does lattice energy play in the exothermicity of the reaction?

NaCl has a lattice energy of -787 kJ/mol. Consider a hypothetical salt XY. X3+ has the same radius of Na+ and Y3- has the same radius as Cl-. Estimate the lattice energy of XY.

Draw the Lewis structure for nitric acid (the hydrogen atom is attached to one of the oxygen atoms). Include all three resonance structures by alternating the double bond among the three oxygen atoms. Use formal charge to determine which of the resonance structures is most important to the structure of nitric acid.

Phosgene (Cl2CO) is a poisonous gas used as a chemical weapon during World War I. It is a potential agent for chemical terrorism today. Draw the Lewis structure of phosgene. Include all three resonance forms by alternating the double bond among the three terminal atoms. Which resonance structure is the best?

The cyanate ion (OCN-) and the fulminate ion (CNO-) share the same three atoms but have vastly different properties. The cyanate ion is stable, while the fulminate ion is unstable and forms explosive compounds. The resonance structures of the cyanate ion were explored in Example 9.8 . Draw Lewis structures for the fulminate ionincluding possible resonance formsand use formal charge to explain why the fulminate ion is less stable (and therefore more reactive) than the cyanate ion.

Draw the Lewis structure for each organic compound from its condensed structural formula. a. C 3 H 8 b. CH 3 OCH 3 c. CH 3 COCH 3 d. CH 3 COOH e. CH 3 CHO

Draw the Lewis structure for each organic compound from its condensed structural formula. a. C 2 H 4 b. CH 3 NH 2 c. HCHO d. CH 3 CH 2 OH e. HCOOH

Use Lewis structures to explain why Br3 - and I3 - are stable, while F3 - is not.

Draw the Lewis structure for HCSNH2. (The carbon and nitrogen atoms are bonded together and the sulfur atom is bonded to the carbon atom.) Label each bond in the molecule as polar or nonpolar.

Draw the Lewis structure for urea, H2NCONH2, one of the compounds responsible for the smell of urine. (The central carbon atom is bonded to both nitrogen atoms and to the oxygen atom.) Does urea contain polar bonds? Which bond in urea is most polar?

Some theories of aging suggest that free radicals cause certain diseases and perhaps aging in general. As you know from the Lewis model, such molecules are not chemically stable and will quickly react with other molecules. According to certain theories, free radicals may attack molecules within the cell, such as DNA, changing them and causing cancer or other diseases. Free radicals may also attack molecules on the surfaces of cells, making them appear foreign to the body’s immune system. The immune system then attacks the cells and destroys them, weakening the body. Draw Lewis structures for each free radical implicated in this theory of aging. a. \(\mathrm{O}_2^{-}\) b. \(\mathrm{O}^{-}\) c. \(\mathrm{OH}\) d. \(\mathrm{CH}_3 \mathrm{OO}\) (unpaired electron on terminal oxygen)

Free radicals are important in many environmentally significant reactions (see the Chemistry in the Environment box on free radicals in this chapter). For example, photochemical smog smog that results from the action of sunlight on air pollutants forms in part by these two steps: NO2 UV light NO + O O + O2 O3 The product of this reaction, ozone, is a pollutant in the lower atmosphere. (Upper atmospheric ozone is a natural part of the atmosphere that protects life on Earth from ultraviolet light.) Ozone is an eye and lung irritant and also accelerates the weathering of rubber products. Rewrite the given reactions using the Lewis structure of each reactant and product. Identify the free radicals.

If hydrogen were used as a fuel, it could be burned according to this reaction: H2(g) + 1>2 O2(g) h H2O(g) Use average bond energies to calculate Hrxn for this reaction and also for the combustion of methane (CH4). Which fuel yields more energy per mole? Per gram?

Calculate \(\Delta H_{\mathrm{rxn}}\) for the combustion of octane \(\left(\mathrm{C}_{8} \mathrm{H}_{18}\right)\), a component of gasoline, by using average bond energies and then calculate it using enthalpies of formation from Appendix IIB . What is the percent difference between your results? Which result would you expect to be more accurate?

Draw Lewis structures for each compound. a. Cl2O7 (no CliCl bond) b. H3PO3 (two OH bonds) c. H3AsO4

The azide ion, N3 - , is a symmetrical ion, all of whose contributing resonance structures have formal charges. Draw three important contributing structures for this ion.

List the following gas-phase ion pairs in order of the quantity of energy released when they form from separated gas-phase ions. Start with the pair that releases the least energy. Na+F-, Mg2+F-, Na+O2-, Mg2+O2-, Al3+O2-.

Calculate H for the reaction H2(g) + Br2(g)h 2 HBr(g) using the bond energy values. The H f of HBr(g) is not equal to one-half of the value calculated. Account for the difference. 111

The heat of atomization is the heat required to convert a molecule in the gas phase into its constituent atoms in the gas phase. The heat of atomization is used to calculate average bond energies. Without using any tabulated bond energies, calculate the average CiCl bond energy from the following data: the heat of atomization of CH4 is 1660 kJ/mol and of CH2Cl2 is 1495 kJ/mol.

Calculate the heat of atomization (see previous problem) of \(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{Cl}\), using the average bond energies in Table 9.3 .

A compound composed of only carbon and hydrogen is 7.743% hydrogen by mass. Propose a Lewis structure for the compound.

A compound composed of only carbon and chlorine is 85.5% chlorine by mass. Propose a Lewis structure for the compound.

The main component of acid rain (H2SO4) forms from SO2 pollutant in the atmosphere via these steps: SO2 + OH # h HSO3 # HSO3 # + O2 h SO3 + HOO # SO3 + H2O h H2SO4 Draw the Lewis structure for each of the species in these steps and use bond energies and Hesss law to estimate Hrxn for the overall process. (Use 265 kJ/mol for the SiO single bond energy.)

A 0.167 g sample of an unknown acid requires 27.8 mL of 0.100 M NaOH to titrate to the equivalence point. Elemental analysis of the acid gives the following percentages by mass: 40.00% C, 6.71% H, 53.29% O. Determine the molecular formula, molar mass, and Lewis structure of the unknown acid.

Use the dipole moments of HF and HCl (given at the end of the problem) together with the percent ionic character of each bond ( Figure 9.10 ) to estimate the bond length in each molecule. How well does your estimated bond length agree with the bond length in Table 9.4 ? HCl m = 1.08 D HF m = 1.82 D

Use average bond energies together with the standard enthalpy of formation of C(g) ( 718.4 kJ/mol ) to estimate the standard enthalpy of formation of gaseous benzene, C6H6(g). (Remember that average bond energies apply to the gas phase only.) Compare the value you obtain using average bond energies to the actual standard enthalpy of formation of gaseous benzene, 82.9 kJ/mol. What does the difference between these two values tell you about the stability of benzene?

The standard state of phosphorus at 25C is P4. This molecule has four equivalent P atoms, no double or triple bonds, and no expanded octets. Draw its Lewis structure.

The standard heat of formation of CaBr2 is -675 kJ/mol. The first ionization energy of Ca is 590 kJ/mol and its second ionization energy is 1145 kJ/mol. The heat of sublimation of Ca[Ca(s) h Ca(g)] is 178 kJ/mol. The bond energy of Br2 is 193 kJ/mol, the heat of vaporization of Br2(l) is 31 kJ/mol, and the electron affinity of Br is -325 kJ/mol. Calculate the lattice energy of CaBr2.

The standard heat of formation of \(\mathrm{PI}_{3}(s)\) is -24.7 kJ/mol and the PI bond energy in this molecule is 184 kJ/mol. The standard heat of formation of P(g) is 334 kJ/mol and that of \(\mathrm{I}_{2}(\mathrm{g})\) is 62 kJ/mol. The \(\mathrm{I}_{2}\) bond energy is 151 kJ/mol. Calculate the heat of sublimation of \(\mathrm{PI}_3\left[\mathrm{Pl}_3(\mathrm{s})\longrightarrow\mathrm{Pl}_3(\mathrm{g})\right]\)

A compound has the formula C8H8 and does not contain any double or triple bonds. All the carbon atoms are chemically identical and all the hydrogen atoms are chemically identical. Draw the Lewis structure for this molecule.

Find the oxidation number of each sulfur in the molecule \(\mathrm{H}_2 \mathrm{~S}_4\), which has a linear arrangement of its atoms.

Ionic solids of the O- and O3- anions do not exist, while ionic solids of the O2- anion are common. Explain.

The standard state of sulfur is solid rhombic sulfur. Use the appropriate standard heats of formation given in Appendix II to find the average bond energy of the S = O in \(\mathrm{SO}_{2}\).

Which statement is true of an endothermic reaction? a. Strong bonds break and weak bonds form. b. Weak bonds break and strong bonds form. c. The bonds that break and those that form are of approximately the same strength

When a firecracker explodes, energy is obviously released. The compounds in the firecracker can be viewed as being energy rich. What does this mean? Explain the source of the energy in terms of chemical bonds

A fundamental difference between compounds containing ionic bonds and those containing covalent bonds is the existence of molecules. Explain why molecules exist in solid covalent compounds but do not exist in solid ionic compounds.

In the very first chapter of this book, we described the scientific approach and put a special emphasis on scientific models or theories. In this chapter, we looked carefully at a model for chemical bonding (the Lewis model). Why is this theory successful? What are some of the limitations of the theory?