Draw Lewis structures and predict the molecular structures

We use differences in electronegativity to account for certain properties of bonds. What if all atoms had the same electronegativity values? How would bonding between atoms be affected? What are some differences we would notice?

Ions have different radii than their parent ions. What if ions stayed the same size as their parent ions? How would this affect ionic bonding in compounds?

You and a friend are studying for a chemistry exam. What if your friend tells you that all molecules with polar bonds are polar molecules? How would you explain to your friend that this is not correct? Provide two examples to support your answer.

Explain the electronegativity trends across a row and down a column of the periodic table. Compare these trends with those of ionization energy and atomic radii. How are they all related?

The ionic compound AB is formed. The charges on the ions may be 11, 21; 12, 22; 13, 23; or even larger. What are the factors that determine the charge for an ion in an ionic compound?

Using only the periodic table, predict the most stable ion for Na, Mg, Al, S, Cl, K, Ca, and Ga. Arrange these from largest to smallest radius, and explain why the radius varies as it does. Compare your predictions with Fig. 13.8.

The bond energy for the COH bond is about 413 kJ/mol in CH4 but 380 kJ/mol in CHBr3. Although these values are relatively close in magnitude, they are different. Explain why they are different. Does the fact that the COH bond energy in CHBr3 is lower make any sense? Why?

Consider the following statement: Because oxygen seems to prefer a negative two charge, the second electron affinity is more negative than the first. Indicate everything that is correct in this statement. Indicate everything that is incorrect. Correct the incorrect information, and explain.

Which has the greater bond lengths: NO2 2 or NO3 2? Explain.

The following ions are best described with resonance structures. Draw the resonance structures, and using formal charge arguments, predict the best Lewis structure for each ion. a. NCO2 b. CNO2

Would you expect the electronegativity of titanium to be the same in the species Ti, Ti21, Ti31, and Ti41? Explain.

The second electron affinity values for both oxygen and sulfur are unfavorable (endothermic). Explain.

Arrange the following molecules from most to least polar, and explain your order: CH4, CF2Cl2, CF2H2, CCl4, and CCl2H2.

What is meant by a chemical bond? Why do atoms form bonds with each other? Why do some elements exist as molecules in nature instead of as free atoms?

Does a Lewis structure tell which electrons come from which atoms? Explain.

Distinguish between the terms electronegativity and electron affinity, covalent bond and ionic bond, and pure covalent bond and polar covalent bond. Characterize the types of bonds in terms of electronegativity difference. Energetically, why do ionic and covalent bonds form?

The following electrostatic potential diagrams represent H2, HCl, or NaCl. Label each, and explain your choices. (a) (b)

An alternative definition of electronegativity is Electronegativity 5 constant (I.E. 2 E.A.) where I.E. is the ionization energy and E.A. is the electron affinity using the sign conventions of this book. Use data in Chapter 12 to calculate the (I.E. 2 E.A.) term for F, Cl, Br, and I. Do these values show the same trend as the electronegativity values given in this chapter? The first ionization energies of the halogens are 1678, 1255, 1138, and 1007 kJ/mol, respectively. (Hint: Choose a constant so that the electronegativity of fluorine equals 4.0. Using this constant, calculate relative electronegativities for the other halogens and compare to values given in the text.)

Use Coulombs law, V 5 Q1Q2 4pP0r 5 2.31 3 10219 J nma Q1Q2 r b to calculate the energy of interaction for the following two arrangements of charges, each having a magnitude equal to the electron charge. a. 1 3 10210 m 1 3 10210 m 11 m888n 21 m88 ` 88n 11 m888n 21 b. 1 3 10210 m 21 1 3 10210 m 11 11 1 3 10210 m 21 1 3 10210 m

Without using Fig. 13.3, predict the order of increasing electronegativity in each of the following groups of elements. a. C, N, O d. Tl, S, Ge b. S, Se, Cl e. Na, K, Rb c. Si, Ge, Sn f. B, O, Ga

Without using Fig. 13.3, predict which bond in each of the following groups is the most polar. a. COF, SiOF, GeOF b. POCl, SOCl c. SOF, SOCl, SOBr d. TiOCl, SiOCl, GeOCl e. COH, SiOH, SnOH f. AlOBr, GaOBr, InOBr, TlOBr

Repeat Exercises 17 and 18. This time use the values of the electronegativities of the elements given in Fig. 13.3. Are there any differences among your answers?

Which of the following incorrectly shows the bond polarity? Show the correct bond polarity for those that are incorrect. a. d1HOFd2 d. d1BrOBrd2 b. d1ClOId2 e. d1OOPd2 c. d1SiOSd2

Indicate the bond polarity (show the partial positive and partial negative ends) in the following bonds. a. COO c. HOCl e. SeOS b. POH d. BrOTe

Hydrogen has an electronegativity value between boron and carbon and identical to phosphorus. With this in mind, rank the following bonds in order of decreasing polarity: POH, OOH, NOH, FOH, COH.

Rank the following bonds in order of increasing ionic character: NOO, CaOO, COF, BrOBr, KOF.

List all the possible bonds that can occur between the elements P, Cs, O, and H. Predict the type of bond (ionic, covalent, or polar covalent) one would expect to form for each bond.

Some plant fertilizer compounds are (NH4)2SO4, Ca3(PO4)2, K2O, P2O5, and KCl. Which of these compounds contain both ionic and covalent bonds?

The following electrostatic potential diagrams represent CH4, NH3, or H2O. Label each, and explain your choices. a. c. b.

When an element forms an anion, what happens to the radius? When an element forms a cation, what happens to the radius? Why? Define the term isoelectronic. When comparing sizes of ions, which ion has the largest radius, and which ion has the smallest radius in an isoelectronic series? Why?

Consider the ions Sc31, Cl2, K1, Ca21, and S22. Match these ions to the following pictures that represent the relative sizes of the ions

For each of the following groups, place the atoms and ions in order of decreasing size. a. Cu, Cu1, Cu21 b. Ni21, Pd21, Pt21 c. O, O2, O22 d. La31, Eu31, Gd31, Yb31 e. Te22, I2, Cs1, Ba21, La31

Write electron configurations for each of the following. a. the cations: Mg21, Sn21, K1, Al31, Tl1, As31 b. the anions: N32, O22, F2, Te22

Write electron configurations for the most stable ion formed by each of the elements Rb, Ba, Se, and I (when in stable ionic compounds).

Give an example of an ionic compound where both the anion and the cation are isoelectronic with each of the following noble gases. a. Ne b. Ar c. Kr d. Xe

What noble gas has the same electron configuration as each of the ions in the following compounds? a. cesium sulfide c. calcium nitride b. strontium fluoride d. aluminum bromide

Which of the following ions have noble gas electron configurations? a. Fe21, Fe31, Sc31, Co31 c. Pu41, Ce41, Ti41 b. Tl1, Te22, Cr31 d. Ba21, Pt21, Mn21

Give three ions that are isoelectronic with krypton. Place these ions in order of increasing size.

Some of the important properties of ionic compounds are as follows: i. low electrical conductivity as solids and high conductivity in solution or when molten ii. relatively high melting and boiling points iii. brittleness iv. solubility in polar solvents How does the concept of ionic bonding discussed in this chapter account for these properties?

Predict the empirical formulas of the ionic compounds formed from the following pairs of elements. Name each compound. a. Al and S c. Mg and Cl b. K and N d. Cs and Br

Which compound in each of the following pairs of ionic substances has the most exothermic lattice energy? Justify your answers. a. NaCl, KCl b. LiF, LiCl c. Mg(OH)2, MgO d. Fe(OH)2, Fe(OH)3 e. NaCl, Na2O f. MgO, BaS

Use the following data to estimate DH8f for potassium chloride. K(s) 1 1 2Cl2(g) 88n KCl(s) Lattice energy 2690. kJ/mol Ionization energy for K 419 kJ/mol Electron affinity of Cl 2349 kJ/mol Bond energy of Cl2 239 kJ/mol Enthalpy of sublimation for K 90. kJ/mol

Use the following data to estimate DH8f for magnesium fluoride. Mg(s) 1 F2(g) 88n MgF2(s) Lattice energy 22913 kJ/mol First ionization energy of Mg 735 kJ/mol Second ionization energy of Mg 1445 kJ/mol Electron affinity of F 2328 kJ/mol Bond energy of F2 154 kJ/mol Enthalpy of sublimation of Mg 150. kJ/mol

Consider the following: Li(s) 1 1 2 I2(g) n LiI(s) DH 5 2292 kJ. LiI(s) has a lattice energy of 2753 kJ/mol. The ionization energy of Li(g) is 520. kJ/mol, the bond energy of I2(g) is 151 kJ/mol, and the electron affinity of I(g) is 2295 kJ/mol. Use these data to determine the heat of sublimation of Li(s).

In general, the higher the charge on the ions in an ionic compound, the more favorable is the lattice energy. Why do some stable ionic compounds have 11 charged ions even though 14, 15, and 16 charged ions would have a more favorable lattice energy?

Consider the following energy changes: DE (kJ/mol) Mg(g) 88n Mg1(g) 1 e2 735 Mg1(g) 88n Mg21(g) 1 e2 1445 O(g) 1 e2 88n O2(g) 2141 O2(g) 1 e2 88n O22(g) 878 a. Magnesium oxide exists as Mg21O22, not as Mg1O2. Explain. b. What experiment could be done to confirm that magnesium oxide does not exist as Mg1O2?

Use the following data (in kJ/mol) to estimate DH for the reaction S2(g) 1 e2 n S22(g). Include an estimate of uncertainty. DH8f Lattice Energy IE of M DHsub of M Na2S 2365 22203 495 109 K2S 2381 22052 419 90. Rb2S 2361 21949 409 82 Cs2S 2360. 21850. 382 78 S(s) 88n S(g) DH 5 277 kJ/mol S(g) 1 e2 88n S2(g) DH 5 2200. kJ/mo

Rationalize the following lattice energy values: Compound Lattice Energy (kJ/mol) CaSe 22862 Na2Se 22130 CaTe 22721 Na2Te 22095

The lattice energies of FeCl3, FeCl2, and Fe2O3 are (in no particular order) 22631 kJ/mol, 25339 kJ/mol, and 214,774 kJ/mol. Match the appropriate formula to each lattice energy

Use bond energy values in Table 13.6 to estimate DH for each of the following reactions in the gas phase. a. H2(g) 1 Cl2(g) n 2HCl(g) b. NqN(g) 1 3H2(g) n 2NH3(g) c. H C N(g) + 2H2(g) H C H H H H N(g) d. (g)  2F2(g) N N(g)  4HF(g) H H H H N N

Compare your answers from parts a and b of Exercise 47 with DH values calculated for each reaction using standard enthalpies of formation in Appendix 4. Do enthalpy changes calculated from bond energies give a reasonable estimate of the actual values?

Use bond energies to predict DH for the isomerization of methyl isocyanide to acetonitrile. CH3NqC(g) 88n CH3CqN(g

The major industrial source of hydrogen gas is by the following reaction: CH4 1g2 1 H2O1g2 h CO1g2 1 3H2 1g2 Use bond energies to predict DH for this reaction

Combustion reactions of fossil fuels provide most of the energy needs of the world. Why are combustion reactions of fossil fuels so exothermic?

Use data from Table 13.6 to estimate DH for the combustion of methane (CH4), as shown below: + +

Use bond energies to estimate DH for the combustion of 1 mole of acetylene: C2H2(g) 1 5 2O2(g) 88n 2CO2(g) 1 H2O(g)

Consider the following reaction: A2 1 B2 88n 2AB DH 5 2285 kJ The bond energy for A2 is one-half the amount of the AB bond energy. The bond energy of B2 5 432 kJ/mol. What is the bond energy of A2?

The space shuttle Orbiter uses the oxidation of methyl hydrazine by dinitrogen tetroxide for propulsion: 5N2O4(g) 1 4N2H3CH3(g) 88n 12H2O(g) 1 9N2(g) 1 4CO2(g) Use bond energies to estimate DH for this reaction. The structures for the reactants are N N O O O O JM O D E G N N H H H3C H

Following are three processes that have been used for the industrial manufacture of acrylonitrilean important chemical used in the manufacture of plastics, synthetic rubber, and fibers. Use bond energy values (Tables 13.6 and 13.7) to estimate DH for each of the reactions. a. A A A A O O O G DO CH2 CH2+ HCN HOC C C N H H H H q G G D D C HOCH2CH2CN CC + H2O N H H H q P b. 4CH2PCHCH3 1 6NO 8888n 4CH2PCHCN 1 6H2O 1 N2 The nitrogenoxygen bond energy in nitric oxide (NO) is 630. kJ/mol. c. 2CH2PCHCH3 1 2NH3 1 3O2 8888888n 2CH2PCHCN 1 6H2O

Is the elevated temperature noted in parts b and c of Exercise 56 needed to provide energy to endothermic reactions?

Acetic acid is responsible for the sour taste of vinegar. It can be manufactured using the following reaction: CH3OH(g) + CqO(g) CH3COOH(l) O B Use tabulated values of bond energies (Table 13.6) to estimate DH for this reaction. Compare this result to the DH value calculated using standard enthalpies of formation in Appendix 4. Explain any discrepancies.

Use bond energies (Table 13.6), values of electron affinities (Table 12.8), and the ionization energy of hydrogen (1312 kJ/mol) to estimate DH for each of the following reactions. a. HF(g) 88n H1(g) 1 F2(g) b. HCl(g) 88n H1(g) 1 Cl2(g) c. HI(g) 88n H1(g) 1 I2(g) d. H2O(g) 88n H1(g) 1 OH2(g) (Electron affinity of OH(g) 5 2180. kJ/mol.)

The standard enthalpies of formation of S(g), F(g), SF4(g), and SF6(g) are 1278.8 kJ/mol, 179.0 kJ/mol, 2775 kJ/mol, and 21209 kJ/mol, respectively. a. Use these data to estimate the energy of an SOF bond. b. Compare the value that you calculated in part a with the value given in Table 13.6. What conclusions can you draw? c. Why are the DHf 8 values for S(g) and F(g) not equal to zero, even though sulfur and fluorine are elements?

Use the following standard enthalpies of formation to estimate the NOH bond energy in ammonia. Compare this with the value in Table 13.6. N(g) 472.7 kJ/mol H(g) 216.0 kJ/mol NH3(g) 246.1 kJ/mol

The standard enthalpy of formation for NO(g) is 90. kJ/mol. Use this and the values for the OPO and NqN bond energies to estimate the bond strength in NO.

Write Lewis structures that obey the octet rule for each of the following. Except for HCN and H2CO, the first atom listed is the central atom. For HCN and H2CO, carbon is the central atom. a. HCN d. NH4 1 g. CO2 b. PH3 e. H2CO h. O2 c. CHCl3 f. SeF2 i. HBr

Draw a Lewis structure that obeys the octet rule for each of the following molecules and ions. In each case the first atom listed is the central atom. a. POCl3, SO4 22, XeO4, PO4 32, ClO4 2 b. NF3, SO3 22, PO3 32, ClO3 2 c. ClO2 2, SCl2, PCl2

Considering your answers to Exercise 64, what conclusions can you draw concerning the structures of species containing the same number of atoms and the same number of valence electrons?

Draw Lewis structures for the following. Show all resonance structures, where applicable. Carbon is the central atom in OCN2 and SCN2. a. NO2 2, NO3 2, N2O4(N2O4 exists as O2NONO2.) b. OCN2, SCN2, N3 2

Some of the pollutants in the atmosphere are ozone, sulfur dioxide, and sulfur trioxide. Draw Lewis structures for these three molecules. Show all resonance structures.

Peroxyacetyl nitrate, or PAN, is present in photochemical smog. Draw Lewis structures (including resonance forms) for PAN. The skeletal arrangement is H H H OOO CCOO O O O O N A O A A G D the Lewis structures for methyl isocyanate, CH3NCO, including resonance forms.

Explain the terms resonance and delocalized electrons. When a substance exhibits resonance, we say that none of the individual Lewis structures accurately portrays the bonding in the substance. Why do we draw resonance structures?

Benzene (C6H6) consists of a six-membered ring of carbon atoms with one hydrogen bonded to each carbon. Draw Lewis structures for benzene, including resonance structures.

An important observation supporting the need for resonance in the LE model is that there are only three different structures of dichlorobenzene (C6H4Cl2). How does this fact support the need for the concept of resonance?

Borazine (B3N3H6) has often been called inorganic benzene. Draw Lewis structures for borazine. Borazine is a six-membered ring of alternating boron and nitrogen atoms.

Draw all the possible Lewis structures for dimethylborazine [(CH3)2B3N3H4]. (See Exercise 73.) Would there be a different number of structures if there was no resonance?

The most common type of exception to the octet rule are compounds or ions with central atoms having more than eight electrons around them. PF5, SF4, ClF3 and Br3 2 are examples of this type of exception. Draw the Lewis structure for these compounds or ions. Which elements, when they have to, can have more than eight electrons around them? How is this rationalized?

SF6, ClF5, and XeF4 are three compounds whose central atoms do not follow the octet rule. Draw Lewis structures for these compounds.

Which of the following statements is(are) true? Correct the false statements. a. It is impossible to satisfy the octet rule for all atoms in XeF2. b. Because SF4 exists, OF4 should also exist because oxygen is in the same family as sulfur. c. The bond in NO1 should be stronger than the bond in NO2. d. As predicted from the two Lewis structures for ozone, one oxygenoxygen bond is stronger than the other oxygenoxygen bond.

Lewis structures can be used to understand why some molecules react in certain ways. Write the Lewis structures for the reactants and products in the reactions described below. a. Nitrogen dioxide dimerizes to produce dinitrogen tetroxide. b. Boron trihydride accepts a pair of electrons from ammonia, forming BH3NH3. Give a possible explanation for why these two reactions occu

Consider the following bond lengths: COO 1.43 CPO 1.23 CqO 1.09 In the CO3 22 ion, all three COO bonds have identical bond lengths of 1.36 . Why?

Order the following species with respect to the carbon oxygen bond length (longest to shortest): CO, CO2, CO3 22, CH3OH What is the order from the weakest to the strongest carbonoxygen bond?

Place the species below in order of the shortest to the longest nitrogenoxygen bond. H2NOH, N2O, NO1, NO2 2, NO3 2 (H2NOH exists as H2NOOH.)

Use the formal charge arguments to rationalize why BF3 would not follow the octet rule.

Use formal charge arguments to explain why CO has a much smaller dipole moment than would be expected on the basis of electronegativity.

Nitrous oxide (N2O) has three possible Lewis structures: N N P PO NqNOO  N N O qO Given the following bond lengths, NON 167 pm NPO 115 pm NPN 120 pm NOO 147 pm NqN 110 pm rationalize the observations that the NON bond length in N2O is 112 pm and that the NOO bond length is 119 pm. Assign formal charges to the resonance structures for N2O. Can you eliminate any of the resonance structures on the basis of formal charges? Is this consistent with observation?

Draw Lewis structures that obey the octet rule for the following species. Assign the formal charge to each central atom. a. POCl3 c. ClO4 2 e. SO2Cl2 g. ClO3 2 b. SO4 22 d. PO4 32 f. XeO4 h. NO4 32

Draw the Lewis structures that involve minimum formal charges for the species in Exercise 85.

When molten sulfur reacts with chlorine gas, a vilesmelling orange liquid forms that has an empirical formula of SCl. The structure of this compound has a formal charge of zero on all elements in the compound. Draw the Lewis structure for the vile-smelling orange liquid

Oxidation of the cyanide ion produces the stable cyanate ion (OCN2). The fulminate ion (CNO2), on the other hand, is very unstable. Fulminate salts explode when struck; Hg(CNO)2 is used in blasting caps. Write the Lewis structures and assign formal charges for the cyanate and fulminate ions. Why is the fulminate ion so unstable? (C is the central atom in OCN2, and N is the central atom in CNO2.)

Write the Lewis Structure for O2F2 (O2F2 exists as FOOOOOF). Assign oxidation states and formal charges to the atoms in O2F2. This compound is a vigorous and potent oxidizing and fluorinating agent. Are oxidation states or formal charges more useful in accounting for these properties of O2F2?

Three resonance structures can be drawn for CO2. Which resonance structure is best from a formal charge standpoint?

A common trait of simple organic compounds is to have Lewis structures where all atoms have a formal charge of zero. Consider the following incomplete Lewis structure for an organic compound called methyl cyanoacrylate, the main ingredient in Super Glue. H C O H C C N O C H H C H Draw a complete Lewis structure for methyl cyanoacrylate in which all atoms have a formal charge of zero.

Benzoic acid is a food preservative. The space-filling model for benzoic acid is shown below. Benzoic acid (C6H5CO2H) H O C Draw the Lewis structure for benzoic acid, including all resonance structures in which all atoms have a formal charge of zero

Write the name of each of the following molecular structures. a. d. b. e. c

State whether or not each of the following has a permanent dipole moment. a. b. c. d. e.

Predict the molecular structure and the bond angles for each molecule or ion in Exercises 63, 64, and 66.

Predict the molecular structure and the bond angles for each of the following. a. SeO3 b. SeO2 c. PCl3 d. SCl2 e. SiF4

There are several molecular structures based on the trigonal bipyramid geometry. Three such structures are A A B 180 Linear A A A B 90 90 T-shaped A A A A B 90 90 120 See-saw Which of the compounds or ions in Exercises 75 and 76 have these molecular structures?

Two variations of the octahedral geometry are illustrated below. B AA AA  90 90 Square planar B A A A A A 90 90 90 Square pyramid Which of the compounds or ions in Exercises 75 and 76 have these molecular structures?

Predict the molecular structure and the bond angles for each of the following. (See Exercises 97 and 98.) a. XeCl2 b. ICl3 c. TeF4 d. PCl5

Predict the molecular structure and the bond angles for each of the following. (See Exercises 97 and 98.) a. ICl5 b. XeCl4 c. SeCl6

Which of the molecules in Exercise 96 have net dipole moments (are polar)?

Which of the molecules in Exercises 99 and 100 have net dipole moments (are polar)?

Give two requirements that should be satisfied for a molecule to be polar. Explain why CF4 and KrF4 are nonpolar compounds (have no net dipole moments), whereas SF4 is polar (has a net dipole moment). Is CO2 polar? What about COS? Explain

What do each of the following sets of compounds/ions have in common with each other? Reference your Lewis structures for Exercises 96, 99, and 100. a. XeCl4, XeCl2 b. ICl5, TeF4, ICl3, PCl3, SCl2, SeO2

Which of the following statements is(are) true? Correct the false statements. a. The molecules SeS3, SeS2, PCl5, TeCl4, ICl3, and XeCl2 all exhibit at least one bond angle which is approximately 120 degrees. b. The bond angle in SO2 should be similar to the bond angle in CS2 or SCl2. c. Central atoms in a molecule adopt a geometry of the bonded atoms and lone pairs about the central atom in order to maximize electron repulsions.

Consider the following Lewis structure, where E is an unknown element: O E O O OO A     What are some possible identities for element E? Predict the molecular structure (including bond angles) for this ion.

Consider the following Lewis structure, where E is an unknown element: F D G O 2 k F E O What are some possible identities for element E? Predict the molecular structure (including bond angles) for this ion. (See Exercises 97 and 98.)

Although the VSEPR model is correct in predicting that CH4 is tetrahedral, NH3 is pyramidal, and H2O is bent, the model in its simplest form does not account for the fact that these molecules do not have exactly the same bond angles (, HCH is 109.5 degrees, as expected for a tetrahedron, but , HNH is 107.3 degrees and , HOH is 104.5 degrees). Explain these deviations from the tetrahedral angle

Draw Lewis structures and predict the molecular structures of the following. (See Exercises 97 and 98.) a. OCl2, KrF2, BeH2, SO2 c. CF4, SeF4, KrF4 b. SO3, NF3, IF3 d. IF5, AsF5 Which of the above compounds have net dipole moments (are polar)?

Which of the following molecules have net dipole moments? For the molecules that are polar, indicate the polarity of each bond and the direction of the net dipole moment of the molecule. a. CH2Cl2, CHCl3, CCl4 b. CO2, N2O c. PH3, NH3

The molecules BF3, CF4, CO2, PF5, and SF6 are all nonpolar, even though they contain polar bonds. Why?

Although both the Br3 2 and I3 2 ions are known, the F3 2 ion does not exist. Explain

When acid is added to an aqueous solution containing carbonate or bicarbonate ions, carbon dioxide gas is formed. We generally say that carbonic acid (H2CO3) is unstable. Use bond energies to estimate DH for the reaction (in the gas phase): H2CO3 88n CO2 1 H2O Specify a possible cause for the instability of carbonic acid.

The structure of TeF5 2 is F F F F F Te 79 Draw a complete Lewis structure for TeF5 2, and explain the distortion from the ideal square pyramidal structure. (See Exercise 98.)

The compound NF3 is quite stable, but NCl3 is very unstable (NCl3 was first synthesized in 1811 by P. L. Dulong, who lost three fingers and an eye studying its properties). The compounds NBr3 and NI3 are unknown, although the explosive compound NI3 ? NH3 is known. Account for the instability of these halides of nitrogen.

There are two possible structures of XeF2Cl2, where Xe is the central atom. Draw them, and describe how measurements of dipole moments might be used to distinguish among them.

Which member of the following pairs would you expect to be more energetically stable? Justify each choice. a. NaBr or NaBr2 c. SO4 or XeO4 b. ClO4 or ClO4 2 d. OF4 or SeF4

Many times, extra stability is characteristic of a molecule or ion in which resonance is possible. How could this feature be used to explain the acidities of the following compounds? (The acidic hydrogen is marked by an asterisk.) Part c shows resonance in the phenyl ring (C6H5). a. HO OC OH* B O b. CH3O OC C CH OCH3 O OH* P AB c. OH

Arrange the following in order of increasing radius and increasing ionization energy. a. N1, N, N2 b. Se, Se2, Cl, Cl1 c. Br2, Rb1, Sr21

Draw a Lewis structure for the N,N-dimethylformamide molecule. The skeletal structure is HO OC NOCH3 CH3 A A O Various types of evidence lead to the conclusion that there is some double-bond character to one of the CON bonds. Draw one or more resonance structures that support this observation

A compound, XF5, is 42.81% fluorine by mass. Identify the element X. What is the molecular structure of XF5?

The study of carbon-containing compounds and their properties is called organic chemistry. Besides carbon atoms, organic compounds also can contain hydrogen, oxygen, and nitrogen atoms (as well as other types of atoms). A common trait of simple organic compounds is to have Lewis structures in which all atoms have a formal charge of zero. Consider the following incomplete Lewis structure for an organic compound called histidine (one of the amino acids, which are the building blocks of proteins found in human bodies): C C H H O N C C O H HH N H H C H H C N 2 1 Draw a complete Lewis structure for histidine in which all atoms have a formal charge of zero. What are the approximate bond angles about the carbon atom labeled 1 and the nitrogen atom labeled 2?

Do the Lewis structures obtained in Exercises 85 and 86 predict the same molecular structure for each case?

Predict the molecular structure for each of the following. (See Exercises 97 and 98.) a. BrFI2 b. XeO2F2 c. TeF2Cl3 2 For each formula, there are at least two different structures that can be drawn using the same central atom. Draw the possible structures for each formula.

Classify the bonding in each of the following molecules as ionic, polar covalent, or nonpolar covalent. a. H2 e. HF b. K3P f. CCl4 c. NaI g. CF4 d. SO2 h. K2S 126. List the bonds POCl, POF, OOF, and

Arrange the atoms and/or ions in the following groups in order of decreasing size. a. O, O2, O22 b. Fe21, Ni21, Zn21 c. Ca21, K1, Cl2

Use the following data to estimate DHf 8 for barium bromide. Ba1s2 1 Br2 1g2 h BaBr2 1s2 Lattice energy 21985 kJ/mol First ionization energy of Ba 503 kJ/mol Second ionization energy of Ba 965 kJ/mol Electron affinity of Br 2325 kJ/mol Bond energy of Br2 193 kJ/mol Enthalpy of sublimation of Ba 178 kJ/mol

Use bond energy values to estimate DH for the following gas phase reaction: C2H4 1 H2O2 h CH2OHCH2OH

Which of the following compounds or ions exhibit resonance? a. O3 d. CO3 22 b. CNO2 e. AsF3 c. AsI3

The formulas of several chemical substances are given in the table below. For each substance in the table, give its chemical name and predict its molecular structure. Formula Name Molecular Structure CO2 _______________ _______________ NH3 _______________ _______________ SO3 _______________ _______________ H2O _______________ _______________ ClO4 2 _______________ _______________

Predict the molecular structure, bond angles, and polarity (has a net dipole moment or has no net dipole moment) for each of the following compounds. a. SeCl4 d. CBr4 b. SO2 e. IF3 c. KrF4 f. ClF5

Predict the molecular structure of KrF2. Using hyperconjugation, draw the Lewis structures for KrF2 that obey the octet rule. Show all resonance forms

Consider the following computer-generated model of caffeine. H O N C Draw a Lewis structure for caffeine in which all atoms have a formal charge of zero.

Given the following information: Heat of sublimation of Li(s) 5 166 kJ/mol Bond energy of HCl 5 427 kJ/mol Ionization energy of Li(g) 5 520. kJ/mol Electron affinity of Cl(g) 5 2349 kJ/mol Lattice energy of LiCl(s) 5 2829 kJ/mol Bond energy of H2 5 432 kJ/mol Calculate the net change in energy for the following reaction: 2Li(s) 1 2HCl(g) 88n 2LiCl(s) 1 H2(g

Use data in this chapter and Chapter 12 to discuss why MgO is an ionic compound but CO is not an ionic compound.

A promising new material with great potential as a fuel in solid rocket motors is ammonium dinitramide [NH4N(NO2)2]. a. Draw Lewis structures (including resonance forms) for the dinitramide ion [N(NO2)2 2]. b. Predict the bond angles around each nitrogen in the dinitramide ion. c. Ammonium dinitramide can decompose explosively to nitrogen, water, and oxygen. Write a balanced equation for this reaction, and use bond energies to estimate DH for the explosive decomposition of this compound. d. To estimate DH from bond energies, you made several assumptions. What are some of your assumptions?

Think of forming an ionic compound as three steps (this is a simplification, as with all models): (1) removing an electron from the metal, (2) adding an electron to the nonmetal, and (3) allowing the metal cation and nonmetal anion to come together. a. What is the sign of the energy change for each of these three processes?b. In general, what is the sign of the sum of the first two processes? Use examples to support your answer. c. What must be the sign of the sum of the three processes? d. Given your answer to part c, why do ionic bonds occur? e. Given your explanations to part d, why is NaCl stable but not Na2Cl2 and NaCl2? What about MgO compared to MgO2 and Mg2O?

The compound hexaazaisowurtzitane is one of the highest-energy explosives known (C & E News, p. 26, Jan. 17, 1994). The compound, also known as CL-20, was first synthesized in 1987. The method of synthesis and detailed performance data are still classified information because of CL-20s potential military application in rocket boosters and in warheads of smart weapons. The structure of CL-20 is N N N N N N G G G NO2 NO2 NO2 O2N O2N O2N D D D In such shorthand structures, each point where lines meet represents a carbon atom. In addition, the hydrogens attached to the carbon atoms are omitted. Each of the six carbon atoms has one hydrogen atom attached. Three possible reactions for the explosive decomposition of CL-20 are i. C6H6N12O12(s) 88n 6CO(g) 1 6N2(g) 1 3H2O(g) 1 3 2O2(g) ii. C6H6N12O12(s) 88n 3CO(g) 1 3CO2(g) 1 6N2(g) 1 3H2O(g) iii. C6H6N12O12(s) 88n 6CO2(g) 1 6N2(g) 1 3H2(g) a. Use bond energies to estimate DH for these three reactions. b. Which of the three reactions releases the largest amount of energy per kilogram of CL-20?

In molecules of the type XOOOH, as the electronegativity of X increases, the acid strength increases. In addition, if the electronegativity of X is a very small value, the molecule acts like a base. Explain these observations and provide examples.

Calculate the standard heat of formation of the compound ICl(g) at 258C. (Hint: Use Table 13.6 and Appendix 4 data.)

An ionic compound made from the metal M and the diatomic gas X2 has the formula MaXb, in which a 5 1 or 2 and b 5 1 or 2. Use the data provided to determine the most likely values for a and b, along with the most likely charges for each of the ions in the ionic compound. Data (in units of kJ/mol) Successive ionization energies of M: 480., 4750. Successive electron affinity values for X: 2175, 920. Enthalpy of sublimation for M(s) n M(g): 110. Bond energy of X2: 250. Lattice energy for MX (M1 and X2): 21200. kJ/mol Lattice energy for MX2 (M21 and X2): 23500. kJ/mol Lattice energy for M2X (M1 and X22): 23600. kJ/mol Lattice energy for MX (M21 and X22): 24800. kJ/mol

Identify the following five compounds of H, N, and O. For each compound, write a Lewis structure that is consistent with the information given. a. All the compounds are electrolytes, although not all are strong electrolytes. Compounds C and D are ionic and compound B is covalent. b. Nitrogen occurs in its highest possible oxidation state in compounds A and C; nitrogen occurs in its lowest possible oxidation state in compounds C, D, and E. The formal charge on both nitrogens in compound C is 11; the formal charge on the only nitrogen in compound B is 0. c. Compounds A and E exist in solution. Both solutions give off gases. Commercially available concentrated solutions of compound A are normally 16 M. The commercial, concentrated solution of compound E is 15 M. d. Commercial solutions of compound E are labeled with a misnomer that implies that a binary, gaseous compound of nitrogen and hydrogen has reacted with water to produce ammonium ions and hydroxide ions. Actually, this reaction occurs to only a slight extent. e. Compound D is 43.7% N and 50.0% O by mass. If compound D were a gas at STP, it would have a density of 2.86 g/L. f. A formula unit of compound C has one more oxygen than a formula unit of compound D. Compounds C and A have one ion in common when compound A is acting as a strong electrolyte. g. Solutions of compound c are weakly acidic; solutions of compound A are strongly acidic; solutions of compounds B and E are basic. The titration of 0.726 g of compound B requires 21.98 mL of 1.000 M HCl for complete neutralization.