Solution: Write a Lewis structure for SbCl5. Does this

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

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.

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

Write Lewis dot symbols for the following ions: (a) Li1, (b) Cl2, (c) S22, (d) Sr21, (e) N32.

Write Lewis dot symbols for the following atoms and ions: (a) I, (b) I2, (c) S, (d) S22, (e) P, (f) P32, (g) Na, (h) Na1, (i) Mg, ( j) Mg21, (k) Al, (l) Al31, (m) Pb, (n) Pb21.

Explain what an ionic bond is.

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

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.

Name one ionic compound that contains only nonmetallic elements.

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

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

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

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

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.)

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 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).

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.