Given that the ionization energy of F2 2 is 290 kJ/mol, do the following: a. Calculate

What if the sp3 hybrid orbitals were higher in energy than the p orbitals in the free atom? How would this affect our model of bonding?

Describe the bonding in the ammonia molecule using the localized electron model.

Describe the bonding in the N2 molecule.

Describe the bonding in the triiodide ion (I3 2).

How is the xenon atom in XeF4 hybridized?

For each of the following molecules or ions, predict the hybridization of each atom, and describe the molecular structure. a. CO b. BF4 2 c. XeF2

What if p2p orbitals were lower in energy than s2p orbitals? What would you expect the B2 molecular orbital energy-level diagram to look like (without considering ps mixing)? Compare your expected diagram to Figs. 9.34 and 9.35, and state the differences from each

For the species O2, O2 1, and O2 2, give the electron configuration and the bond order for each. Which has the strongest bond?

Use the molecular orbital model to predict the bond order and magnetism of each of the following molecules. a. Ne2 b. P2

Use the molecular orbital model to predict the magnetism and bond order of the NO1 and CN2 ions

Why do we hybridize atomic orbitals to explain the bonding in covalent compounds? What type of bonds form from hybrid orbitals, s or p? Explain.

What hybridization is required for central atoms that have a tetrahedral arrangement of electron pairs? A trigonal planar arrangement of electron pairs? A linear arrangement of electron pairs? How many unhybridized p atomic orbitals are present when a central atom exhibits tetrahedral geometry? Trigonal planar geometry? Linear geometry? What are the unhybridized p atomic orbitals used for?

Describe the bonding in H2S, CH4, H2CO, and HCN using the localized electron model.

What hybridization is required for central atoms exhibiting trigonal bipyramidal geometry? Octahedral geometry? Describe the bonding of PF5, SF4, SF6, and IF5 using the localized electron model.

Electrons in s bonding molecular orbitals are most likely to be found in the region between the two bonded atoms. Why does this arrangement favor bonding? In a s antibonding orbital, where are the electrons most likely to be found in relation to the nuclei in a bond?

Show how 2s orbitals combine to form s bonding and s antibonding molecular orbitals. Show how 2p orbitals overlap to form s bonding, p bonding, p antibonding, and s antibonding molecular orbitals

What are the relationships among bond order, bond energy, and bond length? Which of these can be measured? Distinguish between the terms paramagnetic and diamagnetic. What type of experiment can be done to determine if a material is paramagnetic?

How does molecular orbital theory explain the following observations? a. H2 is stable, while He2 is unstable. b. B2 and O2 are paramagnetic, while C2, N2, and F2 are diamagnetic. c. N2 has a very large bond energy associated with it. d. NO1 is more stable than NO2.

Consider the heteronuclear diatomic molecule HF. Explain in detail how molecular orbital theory is applied to describe the bonding in HF.

What is delocalized p bonding, and what does it explain? Explain the delocalized p bonding system in C6H6 (benzene) and O3 (ozone).

What are molecular orbitals? How do they compare with atomic orbitals? Can you tell by the shape of the bonding and antibonding orbitals which is lower in energy? Explain

Explain the difference between the s and p MOs for homonuclear diatomic molecules. How are bonding and antibonding orbitals different? Why are there two p MOs and one s MO? Why are the p MOs degenerate?

Compare Figs. 9.35 and 9.37. Why are they different? Because B2 is known to be paramagnetic, the p2p and s2p molecular orbitals must be switched from the first prediction. What is the rationale for this? Why might one expect the s2p to be lower in energy than the p2p? Why cant we use diatomic oxygen to help us decide whether the s2p or p2p is lower in energy?

Which of the following would you expect to be more favorable energetically? Explain. a. an H2 molecule in which enough energy is added to excite one electron from the bonding to the antibonding MO b. two separate H atoms

Draw the Lewis structure for HCN. Indicate the hybrid orbitals, and draw a picture showing all the bonds between the atoms, labeling each bond as s or p

Which is the more correct statement: The methane molecule (CH4) is a tetrahedral molecule because it is sp3 hybridized or The methane molecule (CH4) is sp3 hybridized because it is a tetrahedral molecule? What, if anything, is the difference between these two statements?

Compare and contrast the MO model with the LE model. When is each useful?

What are the relationships among bond order, bond energy, and bond length? Which of these quantities can be measured?

In the hybrid orbital model, compare and contrast s bonds with p bonds. What orbitals form the s bonds and what orbitals form the p bonds? Assume the z-axis is the internuclear axis

In the molecular orbital model, compare and contrast s bonds with p bonds. What orbitals form the s bonds and what orbitals form the p bonds? Assume the z-axis is the internuclear axis

Why are d orbitals sometimes used to form hybrid orbitals? Which period of elements does not use d orbitals for hybridization? If necessary, which d orbitals (3d, 4d, 5d, or 6d) would sulfur use to form hybrid orbitals requiring d atomic orbitals? Answer the same question for arsenic and for iodine.

The atoms in a single bond can rotate about the internuclear axis without breaking the bond. The atoms in a double and triple bond cannot rotate about the internuclear axis unless the bond is broken. Why?

Compare and contrast bonding molecular orbitals with antibonding molecular orbitals.

What modification to the molecular orbital model was made from the experimental evidence that B2 is paramagnetic?

Why does the molecular orbital model do a better job in explaining the bonding in NO2 and NO than the hybrid orbital model?

The three NO bonds in NO3 2 are all equivalent in length and strength. How is this explained even though any valid Lewis structure for NO3 2 has one double bond and two single bonds to nitrogen?

Use the localized electron model to describe the bonding in H2O.

Use the localized electron model to describe the bonding in CCl4.

Use the localized electron model to describe the bonding in H2CO (carbon is the central atom).

Use the localized electron model to describe the bonding in C2H2 (exists as HCCH).

The space-filling models of ethane and ethanol are shown below. Ethane (C2H6) Ethanol (C2H5OH) C O Use the localized electron model to describe the bonding in ethane and ethanol.

The space-filling models of hydrogen cyanide and phosgene are shown below. Hydrogen cyanide (HCN) Phosgene (COCl2) C Cl O N H Use the localized electron model to describe the bonding in hydrogen cyanide and phosgene.

Give the expected hybridization of the central atom for the molecules or ions in Exercises 83 and 89 from Chapter 8.

Give the expected hybridization of the central atom for the molecules or ions in Exercises 84 and 90 from Chapter 8.

Give the expected hybridization of the central atom for the molecules or ions in Exercise 87 from Chapter 8.

Give the expected hybridization of the central atom for the molecules in Exercise 88 from Chapter 8.

Give the expected hybridization of the central atom for the molecules in Exercises 113 and 114 from Chapter 8

Give the expected hybridization of the central atom for the molecules in Exercises 115 and 116 from Chapter 8.

For each of the following molecules, write the Lewis structure(s), predict the molecular structure (including bond angles), give the expected hybrid orbitals on the central atom, and predict the overall polarity. a. CF4 e. BeH2 i. KrF4 b. NF3 f. TeF4 j. SeF6 c. OF2 g. AsF5 k. IF5 d. BF3 h. KrF2 l. IF3

For each of the following molecules or ions that contain sulfur, write the Lewis structure(s), predict the molecular structure (including bond angles), and give the expected hybrid orbitals for sulfur. a. SO2 b. SO3 c. S2O3 2 2 S S O O O d. S2O8 2 O S O O O O 2 S O O O e. SO3 22 i. SF6 f. SO4 22 j. F3SOSF g. SF2 k. SF5 1 h. SF4

Why must all six atoms in C2H4 lie in the same plane?

The allene molecule has the following Lewis structure: H H C C H H C Must all hydrogen atoms lie the same plane? If not, what is their spatial relationship? Explain.

Indigo is the dye used in coloring blue jeans. The term navy blue is derived from the use of indigo to dye British naval uniforms in the eighteenth century. The structure of the indigo molecule is C N H H O C C C C C C H H H C O C H H C C C C C C H H H C N a. How many s bonds and p bonds exist in the molecule? b. What hybrid orbitals are used by the carbon atoms in the indigo molecule?

Urea, a compound formed in the liver, is one of the ways humans excrete nitrogen. The Lewis structure for urea is H N H H C H O N Using hybrid orbitals for carbon, nitrogen, and oxygen, determine which orbitals overlap to form the various bonds in urea.

Biacetyl and acetoin are added to margarine to make it taste more like butter. Biacetyl Acetoin CH3 C C CH3 CH3 CH C CH3 O OH O O Complete the Lewis structures, predict values for all COCOO bond angles, and give the hybridization of the carbon atoms in these two compounds. Must the four carbon atoms and two oxygen atoms in biacetyl lie the same plane? How many s bonds and how many p bonds are there in biacetyl and acetoin?

Many important compounds in the chemical industry are derivatives of ethylene (C2H4). Two of them are acrylonitrile and methyl methacrylate. a H H C C C H N b c Acrylonitrile Methyl methacrylate H H C C C O O CH3 CH3 Complete the Lewis structures, showing all lone pairs. Give approximate values for bond angles a through f. Give the hybridization of all carbon atoms. In acrylonitrile, how many of the atoms in the molecule must lie in the same plane? How many s bonds and how many p bonds are there in methyl methacrylate and acrylonitrile?

Two molecules used in the polymer industry are azodicarbonamide and methyl cyanoacrylate. Their structures are O O N N N NH2 C C H H a b d c O O N H2C CH3 C C C e f h g Azodicarbonamide Methyl cyanoacrylate Azodicarbonamide is used in forming polystyrene. When added to the molten plastic, it decomposes to nitrogen, carbon monoxide, and ammonia gases, which are captured as bubbles in the molten polymer. Methyl cyanoacrylate is the main ingredient in super glue. As the glue sets, methyl cyanoacrylate polymerizes across the carboncarbon double bond. (See Chapter 22.) a. Complete the Lewis structures showing all lone pairs of electrons. b. Which hybrid orbitals are used by the carbon atoms in each molecule and the nitrogen atom in azodicarbonamide? c. How many p bonds are present in each molecule? d. Give approximate values for the bond angles marked a through h in the above structures

Hot and spicy foods contain molecules that stimulate paindetecting nerve endings. Two such molecules are piperine and capsaicin: Piperine H C H H O H f H CH CH CH CH C O N H H H H H H H H H H a b c e d O Capsaicin H H H H N C O g h i j k l CH3 H CH2 CH2 3CO CH CH CH HO CH3 (CH2)3 Piperine is the active compound in white and black pepper, and capsaicin is the active compound in chili peppers. The ring structures in piperine and capsaicin are shorthand notation. Each point where lines meet represents a carbon atom.a. Complete the Lewis structure for piperine and capsaicin showing all lone pairs of electrons. b. How many carbon atoms are sp, sp2 , and sp3 hybridized in each molecule? c. Which hybrid orbitals are used by the nitrogen atoms in each molecule? d. Give approximate values for the bond angles marked a through l in the above structures.

One of the first drugs to be approved for use in treatment of acquired immune deficiency syndrome (AIDS) was azidothymidine (AZT). Complete the Lewis structure for AZT. N N N H H H H H H C C O O C C C H O CH3 H H H C N C C C O N a. How many carbon atoms are sp3 hybridized? b. How many carbon atoms are sp2 hybridized? c. Which atom is sp hybridized? d. How many s bonds are in the molecule? e. How many p bonds are in the molecule? f. What is the NPNPN bond angle in the azide (ON3) group? g. What is the HOOOC bond angle in the side group attached to the five-membered ring? h. What is the hybridization of the oxygen atom in the OCH2OH group?

The antibiotic thiarubin-A was discovered by studying the feeding habits of wild chimpanzees in Tanzania. The structure for thiarubin-A is H3C C C C C SS HH C C C C C C CH CH2 a. Complete the Lewis structure showing all lone pairs of electrons. b. Indicate the hybrid orbitals used by the carbon and sulfur atoms in thiarubin-A. c. How many s and p bonds are present in this molecule?

Consider the following molecular orbitals formed from the combination of two hydrogen 1s orbitals: a. Which is the bonding molecular orbital and which is the antibonding molecular orbital? Explain how you can tell by looking at their shapes. b. Which of the two molecular orbitals is lower in energy? Why is this true?

Sketch the molecular orbital and label its type (s or p ; bonding or antibonding) that would be formed when the following atomic orbitals overlap. Explain your labels. a. + + b. + + c. + + d. + +

Which of the following are predicted by the molecular orbital model to be stable diatomic species? a. H2 1, H2, H2 2, H2 22 b. He2 21, He2 1, He2

Which of the following are predicted by the molecular orbital model to be stable diatomic species? a. N2 22, O2 22, F2 22 b. Be2, B2, Ne2

Using the molecular orbital model, write electron configurations for the following diatomic species and calculate the bond orders. Which ones are paramagnetic? a. Li2 b. C2 c. S2

Consider the following electron configuration: 1s3s 2 2 1s3s*2 2 1s3p2 2 1p3p2 4 1p3p*2 4 Give four species that, in theory, would have this electron configuration.

Using molecular orbital theory, explain why the removal of one electron in O2 strengthens bonding, while the removal of one electron in N2 weakens bonding.

Using the molecular orbital model to describe the bonding in F2 1, F2, and F2 2, predict the bond orders and the relative bond lengths for these three species. How many unpaired electrons are present in each species?

The transport of O2 in the blood is carried out by hemoglobin. Carbon monoxide can interfere with oxygen transport because hemoglobin has a stronger affinity for CO than for O2. If CO is present, normal uptake of O2 is prevented, depriving the body of needed oxygen. Using the molecular orbital model, write the electron configurations for CO and for O2. From your configurations, give two property differences between CO and O2.

A Lewis structure obeying the octet rule can be drawn for O2 as follows: O O Use the molecular orbital energy-level diagram for O2 to show that the above Lewis structure corresponds to an excited state.

Using the molecular orbital model, write electron configurations for the following diatomic species and calculate the bond orders. Which ones are paramagnetic? Place the species in order of increasing bond length and bond energy. a. CO b. CO1 c. CO21

Using the molecular orbital model, write electron configurations for the following diatomic species and calculate the bond orders. Which ones are paramagnetic? Place the species in order of increasing bond length and bond energy. a. CN1 b. CN c. CN2

In which of the following diatomic molecules would the bond strength be expected to weaken as an electron is removed? a. H2 c. C2 22 b. B2 d. OF

In terms of the molecular orbital model, which species in each of the following two pairs will most likely be the one to gain an electron? Explain. a. CN or NO b. O2 21 or N2 21

Show how two 2p atomic orbitals can combine to form a s or a p molecular orbital.

Show how a hydrogen 1s atomic orbital and a fluorine 2p atomic orbital overlap to form bonding and antibonding molecular orbitals in the hydrogen fluoride molecule. Are these molecular orbitals s or p molecular orbitals?

Use Figs. 9.42 and 9.43 to answer the following questions. a. Would the bonding molecular orbital in HF place greater electron density near the H or the F atom? Why? b. Would the bonding molecular orbital have greater fluorine 2p character, greater hydrogen 1s character, or an equal contribution from both? Why? c. Answer the previous two questions for the antibonding molecular orbital in HF

The diatomic molecule OH exists in the gas phase. The bond length and bond energy have been measured to be 97.06 pm and 424.7 kJ/mol, respectively. Assume that the OH molecule is analogous to the HF molecule discussed in the chapter and that molecular orbitals result from the overlap of a lowerenergy pz orbital from oxygen with the higher-energy 1s orbital of hydrogen (the OOH bond lies along the z-axis). a. Which of the two molecular orbitals will have the greater hydrogen 1s character? b. Can the 2px orbital of oxygen form molecular orbitals with the 1s orbital of hydrogen? Explain. c. Knowing that only the 2p orbitals of oxygen will interact significantly with the 1s orbital of hydrogen, complete the molecular orbital energy-level diagram for OH. Place the correct number of electrons in the energy levels. d. Estimate the bond order for OH. e. Predict whether the bond order of OH1 will be greater than, less than, or the same as that of OH. Explain.

Acetylene (C2H2) can be produced from the reaction of calcium carbide (CaC2) with water. Use both the localized electron and molecular orbital models to describe the bonding in the acetylide anion (C2 22).

Describe the bonding in NO1, NO2, and NO using both the localized electron and molecular orbital models. Account for any discrepancies between the two models

Describe the bonding in the O3 molecule and the NO2 2 ion using the localized electron model. How would the molecular orbital model describe the p bonding in these two species?

Describe the bonding in the CO3 22 ion using the localized electron model. How would the molecular orbital model describe the p bonding in this species?

Draw the Lewis structures, predict the molecular structures, and describe the bonding (in terms of the hybrid orbitals for the central atom) for the following. a. XeO3 d. XeOF2 b. XeO4 e. XeO3F2 c. XeOF4

FClO2 and F3ClO can both gain a fluoride ion to form stable anions. F3ClO and F3ClO2 will both lose a fluoride ion to form stable cations. Draw the Lewis structures and describe the hybrid orbitals used by chlorine in these ions.

Two structures can be drawn for cyanuric acid: C O N C N N H C H O O H C N C N N C O O H H O H a. Are these two structures the same molecule? Explain. b. Give the hybridization of the carbon and nitrogen atoms in each structure. c. Use bond energies (Table 8.4) to predict which form is more stable; that is, which contains the strongest bonds?

Give the expected hybridization for the molecular structures illustrated below. a. d. b. c. e

Vitamin B6 is an organic compound whose deficiency in the human body can cause apathy, irritability, and an increased susceptibility to infections. An incomplete Lewis structure for vitamin B6 is shown below. Complete the Lewis structure and answer the following questions. Hint: Vitamin B6 can be classified as an organic compound (a compound based on carbon atoms). The majority of Lewis structures for simple organic compounds have all atoms with a formal charge of zero. Therefore, add lone pairs and multiple bonds to the structure below to give each atom a formal charge of zero. H H H O C C O H C O C C C N C H CH H H H d f c b a e g a. How many s bonds and p bonds exist in vitamin B6? b. Give approximate values for the bond angles marked a through g in the structure. c. How many carbon atoms are sp2 hybridized? d. How many carbon, oxygen, and nitrogen atoms are sp3 hybridized? e. Does vitamin B6 exhibit delocalized p bonding? Explain

Aspartame is an artificial sweetener marketed under the name NutraSweet. A partial Lewis structure for aspartame is shown below. C C OH O O C O H2N CH NH OCH3 CHCH2 CH2 Aspartame can be classified as an organic compound (a compound based on carbon atoms). The majority of Lewis structures for simple organic compounds have all atoms with a formal charge of zero. Therefore, add lone pairs and multiple bonds to the structure above to give each atom a formal charge of zero when drawing the Lewis structure. Also note that the six-sided ring is shorthand notation for a benzene ring (OC6H5). Benzene is discussed in Section 9.5. Complete the Lewis structure for aspartame. How many C and N atoms exhibit sp2 hybridization? How many C and O atoms exhibit sp3 hybridization? How many s and p bonds are in aspartame?

Using bond energies from Table 8.4, estimate the barrier to rotation about a carboncarbon double bond. To do this, consider what must happen to go from Cl H Cl C H C to Cl Cl H C H C in terms of making and breaking chemical bonds; that is, what must happen in terms of the p bond?

The three most stable oxides of carbon are carbon monoxide (CO), carbon dioxide (CO2), and carbon suboxide (C3O2). The space-filling models for these three compounds are For each oxide, draw the Lewis structure, predict the molecular structure, and describe the bonding (in terms of the hybrid orbitals for the carbon atoms)

Complete the Lewis structures of the following molecules. Predict the molecular structure, polarity, bond angles, and hybrid orbitals used by the atoms marked by asterisks for each molecule. a. BH3 H H B* H b. N2F2 FiN*iN*iF c. C4H6 H C C H H C H C H * ** * H

Complete the following resonance structures for POCl3. Cl ClP O Cl Cl ClP O Cl (A) (B) a. Would you predict the same molecular structure from each resonance structure?b. What is the hybridization of P in each structure? c. What orbitals can the P atom use to form the p bond in structure B? d. Which resonance structure would be favored on the basis of formal charges?

The N2O molecule is linear and polar. a. On the basis of this experimental evidence, which arrangement, NNO or NON, is correct? Explain your answer. b. On the basis of your answer to part a, write the Lewis structure of N2O (including resonance forms). Give the formal charge on each atom and the hybridization of the central atom. c. How would the multiple bonding in N N O be described in terms of orbitals?

Describe the bonding in the first excited state of N2 (the one closest in energy to the ground state) using the molecular orbital model. What differences do you expect in the properties of the molecule in the ground state as compared to the first excited state? (An excited state of a molecule corresponds to an electron arrangement other than that giving the lowest possible energy.)

Using an MO energy-level diagram, would you expect F2 to have a lower or higher first ionization energy than atomic fluorine? Why?

Show how a dxz atomic orbital and a pz atomic orbital combine to form a bonding molecular orbital. Assume the x-axis is the internuclear axis. Is a s or a p molecular orbital formed? Explain

What type of molecular orbital would result from the in-phase combination of two dxz atomic orbitals shown below? Assume the x-axis is the internuclear axis

Consider three molecules: A, B, and C. Molecule A has a hybridization of sp3 . Molecule B has two more effective pairs (electron pairs around the central atom) than molecule A. Molecule C consists of two s bonds and two p bonds. Give the molecular structure, hybridization, bond angles, and an example for each molecule.

Draw the Lewis structures for SO2, PCl3, NNO, COS, and PF3. Which of the compounds are polar? Which of the compounds exhibit at least one bond angle that is approximately 120 degrees? Which of the compounds exhibit sp3 hybridization by the central atom? Which of the compounds have a linear molecular structure?

Draw the Lewis structures for TeCl4, ICl5, PCl5, KrCl4, and XeCl2. Which of the compounds exhibit at least one bond angle that is approximately 120 degrees? Which of the compounds exhibit d2 sp3 hybridization? Which of the compounds have a square planar molecular structure? Which of the compounds are polar?

A variety of chlorine oxide fluorides and related cations and anions are known. They tend to be powerful oxidizing and fluorinating agents. FClO3 is the most stable of this group of compounds and has been studied as an oxidizing component in rocket propellants. Draw a Lewis structure for F3ClO, F2ClO2 1, and F3ClO2. What is the molecular structure for each species, and what is the expected hybridization of the central chlorine atom in each compound or ion?

Pelargondin is the molecule responsible for the red color of the geranium flower. It also contributes to the color of ripe strawberries and raspberries. The structure of pelargondin is: HO HC HC C C C OH CH CH OH OH H C C C C H H C O + C C C 1 4 2 3 How many s and p bonds exist in pelargondin? What is the hybridization of the carbon atoms marked 14?

Complete a Lewis structure for the compound shown below, then answer the following questions. How many carbon atoms are sp2 hybridized? How many CON bonds are formed by the overlap of an sp3 hybridized carbon with an sp3 hybridized nitrogen? How many lone pairs of electrons are in the Lewis structure of your molecule? How many p bonds are present? N N C C O O C C C C N H H H H N C C H H H H H H

Which of the following statements concerning SO2 is(are) true? a. The central sulfur atom is sp2 hybridized. b. One of the sulfuroxygen bonds is longer than the other(s). c. The bond angles about the central sulfur atom are about 120 degrees. d. There are two s bonds in SO2. e. There are no resonance structures for SO2

Consider the molecular orbital electron configurations for N2, N2 1, and N2 2. For each compound or ion, fill in the table below with the correct number of electrons in each molecular orbital. MO N2 N2 1 N2 2 s2p* p2p* s2p p2p s2s* s2s

Place the species B2 1, B2, and B2 2 in order of increasing bond length and increasing bond energy

Consider the following computer-generated model of caffeine: C O N H Complete a Lewis structure for caffeine in which all atoms have a formal charge of zero (as is typical with most organic compounds). How many C and N atoms are sp2 hybridized? How many C and N atoms are sp3 hybridized? sp hybridized? How many s and p bonds are there?

Cholesterol (C27H46O) has the following structure: CH3 CH3 H H H HO CH3 CH3 H3C In such shorthand structures, each point where lines meet represents a carbon atom and most H atoms are not shown. Draw the complete structure showing all carbon and hydrogen atoms. (There will be four bonds to each carbon atom.) Indicate which carbon atoms use sp2 or sp3 hybrid orbitals. Are all carbon atoms in the same plane, as implied by the structure?

Cyanamide (H2NCN), an important industrial chemical, is produced by the following steps: Cyanamide Acid CaC2 + N2 H2NCN CaNCN + C Calcium cyanamide (CaNCN) is used as a direct-application fertilizer, weed killer, and cotton defoliant. It is also used to make cyanamide, dicyandiamide, and melamine plastics: H2NCN NCNC(NH2)2 Dicyandiamide Acid NCNC(NH2)2 Heat NH3 Melamine ( bonds not shown) H2N NH2 NH2 N C C C N N a. Write Lewis structures for NCN22, H2NCN, dicyandiamide, and melamine, including resonance structures where appropriate. b. Give the hybridization of the C and N atoms in each species. c. How many s bonds and how many p bonds are in each species? d. Is the ring in melamine planar? e. There are three different CON bond distances in dicyandiamide, NCNC(NH2)2, and the molecule is nonlinear. Of all the resonance structures you drew for this molecule, predict which should be the most important

In Exercise 91 in Chapter 8, the Lewis structures for benzene (C6H6) were drawn. Using one of the Lewis structures, estimate DHf 8 for C6H6(g) using bond energies and given that the standard enthalpy of formation of C(g) is 717 kJ/mol. The experimental DHf 8 value of C6H6(g) is 83 kJ/mol. Explain the discrepancy between the experimental value and the calculated DHf 8 value for C6H6(g).

A flask containing gaseous N2 is irradiated with 25-nm light. a. Using the following information, indicate what species can form in the flask during irradiation. N2 1g2 h 2N1g2 DH 5 941 kJ/mol N2 1g2 h N2 1 1g2 1 e2 DH 5 1501 kJ/mol N1g2 h N1 1g2 1 e2 DH 5 1402 kJ/mol b. What range of wavelengths will produce atomic nitrogen in the flask but will not produce any ions? c. Explain why the first ionization energy of N2 (1501 kJ/mol) is greater than the first ionization energy of atomic nitrogen (1402 kJ/mol).

As compared with CO and O2, CS and S2 are very unstable molecules. Give an explanation based on the relative abilities of the sulfur and oxygen atoms to form p bonds.

Values of measured bond energies may vary greatly depending on the molecule studied. Consider the following reactions: NCl3 1g2 h NCl2 1g2 1 Cl1g2 DH 5 375 kJ/mol ONCl1g2 h NO1g2 1 Cl1g2 DH 5 158 kJ/mol Rationalize the difference in the values of DH for these reactions, even though each reaction appears to involve only the breaking of one NOCl bond. (Hint: Consider the bond order of the NO bond in ONCl and in NO.)

Use the MO model to explain the bonding in BeH2. When constructing the MO energy-level diagram, assume that the Bes 1s electrons are not involved in bond formation

Bond energy has been defined in the text as the amount of energy required to break a chemical bond, so we have come to think of the addition of energy as breaking bonds. However, in some cases the addition of energy can cause the formation of bonds. For example, in a sample of helium gas subjected to a high-energy source, some He2 molecules exist momentarily and then dissociate. Use MO theory (and diagrams) to explain why He2 molecules can come to exist and why they dissociate.

Arrange the following from lowest to highest ionization energy: O, O2, O2 2, O2 1. Explain your answer.

Use the MO model to determine which of the following has the smallest ionization energy: N2, O2, N2 22, N2 2, O2 1. Explain your answer

Given that the ionization energy of F2 2 is 290 kJ/mol, do the following: a. Calculate the bond energy of F2 2. You will need to look up the bond energy of F2 and ionization energy of F2. b. Explain the difference in bond energy between F2 2 and F2 using MO theory.

Carbon monoxide (CO) forms bonds to a variety of metals and metal ions. Its ability to bond to iron in hemoglobin is the reason that CO is so toxic. The bond carbon monoxide forms to metals is through the carbon atom: MiC{O a. On the basis of electronegativities, would you expect the carbon atom or the oxygen atom to form bonds to metals? b. Assign formal charges to the atoms in CO. Which atom would you expect to bond to a metal on this basis? c. In the MO model, bonding MOs place more electron density near the more electronegative atom. (See the HF molecule in Figs. 9.42 and 9.43.) Antibonding MOs place more electron density near the less electronegative atom in the diatomic molecule. Use the MO model to predict which atom of carbon monoxide should form bonds to metals.

The space-filling model for benzoic acid, a food preservative, is shown below. Benzoic acid (C6H5CO2H) C O H Describe the bonding in benzoic acid using the localized electron model combined with the molecular orbital model.

As the head engineer of your starship in charge of the warp drive, you notice that the supply of dilithium is critically low. While searching for a replacement fuel, you discover some diboron, B2. a. What is the bond order in Li2 and B2? b. How many electrons must be removed from B2 to make it isoelectronic with Li2 so that it might be used in the warp drive? c. The reaction to make B2 isoelectronic with Li2 is generalized (where n 5 number of electrons determined in part b) as follows: B2 h B2 n1 1 ne2 DH 5 6455 kJ/mol How much energy is needed to ionize 1.5 kg B2 to the desired isoelectronic species?

An unusual category of acids known as superacids, which are defined as any acid stronger than 100% sulfuric acid, can be prepared by seemingly simple reactions similar to the one below. In this example, the reaction of anhydrous HF with SbF5 produces the superacid [H2F]1[SbF6]2: 2HF1l2 1 SbF5 1l2 h 3H2F4 1 3SbF6 4 2 1l2 a. What are the molecular structures of all species in this reaction? What are the hybridizations of the central atoms in each species? b. What mass of [H2F]1[SbF6]2 can be prepared when 2.93 mL anhydrous HF (density 5 0.975 g/mL) and 10.0 mL SbF5 (density 5 3.10 g/mL) are allowed to react?

Determine the molecular structure and hybridization of the central atom X in the polyatomic ion XY3 1 given the following information: A neutral atom of X contains 36 electrons, and the element Y makes an anion with a 1 charge, which has the electron configuration 1s2 2s2 2p6 .