Why does the molecular orbital model do a better job in

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 and MOs for homonuclear diatomic molecules. How are bonding and antibonding orbitals different? Why are there two MOs and one MO? Why are the MOs degenerate? 3.

Compare Figs. 9.35 and 9.37. Why are they different? Because B2 is known to be paramagnetic, the 2p and 2p molecular orbitals must be switched from the first prediction. What is the rationale for this? Why might one expect the 2p to be lower in energy than the 2p? Why cant we use diatomic oxygen to help us decide whether the 2p or 2p is lower in energy? 4. W

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 or .

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 bonds with bonds. What orbitals form the bonds and what orbitals form the bonds? Assume the z-axis is the internuclear axis. 10

In the molecular orbital model, compare and contrast bonds with bonds. What orbitals form the bonds and what orbitals form the bonds? Assume the z-axis is the internuclear axis. 1

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 NO and NO than the hybrid orbital model?

The three NO bonds in NO3 are all equivalent in length and strength. How is this explained even though any valid Lewis structure for NO3 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

The space-filling models of hydrogen cyanide and phosgene are shown below.

Give the expected hybridization of the central atom for the molecules or ions in Exercises 81 and 87 from Chapter 8.

Give the expected hybridization of the central atom for the molecules or ions in Exercises 82 and 88 from Chapter 8.

Give the expected hybridization of the central atom for the molecules or ions in Exercise 85 from Chapter 8

Give the expected hybridization of the central atom for the molecules in Exercise 86 from Chapter 8.

Give the expected hybridization of the central atom for the molecules in Exercises 107 and 108 from Chapter 8

Give the expected hybridization of the central atom for the molecules in Exercises 109 and 110 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. 2 S2O8 2 OS S O O S O O O O O d. e. SO3 2 f. SO4 2 g. SF2 h. SF4 i. SF6 j. F3S SF k. SF5 

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

The allene molecule has the following Lewis structure: 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 a. How many bonds and 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 Using hybrid orbital theory, which orbitals overlap to form the various bonds in urea?

Biacetyl and acetoin are added to margarine to make it taste more like butter. Complete the Lewis structures, predict values for all CCO 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 bonds and how many 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. 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

Two molecules used in the polymer industry are azodicarbonamide and methyl cyanoacrylate. Their structures are O O O O N N N NH2 N H2C CH3 C C C C C H H a b e f h d g c 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 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 pain-detecting nerve endings. Two such molecules are piperine and capsaicin: G O D P B H C H H O H f H CH CH CHPCHO O C O N H H H H H H H H H H Piperine a b c e d O D G B H HO (CH2)3 H3CO H H G A G D N H C O CH CH2 CH3 CH2 G CH D J G CH CH3 Capsaicin g h i j k l 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.

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

Which of the following are predicted by the molecular orbital model to be stable diatomic species? a. N2 2, O2 2, F2 2 b. Be2, B2, Ne2 4

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: 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 , F2, and F2 , predict the bond orders and the relative bond lengths for these three species. How many unpaired electrons are present in each species?

Which charge(s) for the N2 molecule would give a bond order of 2.5?

A Lewis structure obeying the octet rule can be drawn for O2 as follows: 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. CO c. CO2

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. CN b. CN c. CN

In which of the following diatomic molecules would the bond strength be expected to weaken as an electron is removed? a. H2 c. C2 2 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 2 or N2 2

Show how two 2p atomic orbitals can combine to form a or a 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 or 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 lower-energy pz orbital from oxygen with the higher-energy 1s orbital of hydrogen (the OH 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 OH 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 2).

Describe the bonding in NO, NO, 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 ion using the localized electron model. How would the molecular orbital model describe the bonding in these two species?

Describe the bonding in the CO3 2 ion using the localized electron model. How would the molecular orbital model describe the bonding in this species?

Vitamin B6 is an organic compound whose deficiency in the human body can cause apathy, irritability, and an increased susceptibility to infections. On the next page is an incomplete Lewis structure for vitamin B6. 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. a. How many bonds and 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 bonding? Explain.

The antibiotic thiarubin-A was discovered by studying the feeding habits of wild chimpanzees in Tanzania. The structure for thiarubin-A is 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 and bonds are present in this molecule?

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. 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 bonds are in the molecule? e. How many bonds are in the molecule? f. What is the NNN bond angle in the azide (N3) group? g. What is the HOC bond angle in the side group attached to the five-membered ring? h. What is the hybridization of the oxygen atom in the CH2OH group?

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

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: 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. Describe the bonding in benzoic acid using the localized electron model combined with the molecular orbital model.

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: 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?

Aspartame is an artificial sweetener marketed under the name NutraSweet. A partial Lewis structure for aspartame is shown below. 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 (C6H5). 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 and 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 to in terms of making and breaking chemical bonds; that is, what must happen in terms of the 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.

Complete the following resonance structures for POCl3. 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 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 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 or a 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 bonds and two bonds. Give the molecular structure, hybridization, bond angles, and an example for each molecule.

Consider the following computer-generated model of caffeine: 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 and bonds are there?

Cholesterol (C27H46O) has the following structure: 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: 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:

In Exercise 89 in Chapter 8, the Lewis structures for benzene (C6H6) were drawn. Using one of the Lewis structures, estimate Hf  for C6H6(g) using bond energies and given that the standard enthalpy of formation of C(g) is 717 kJ/mol. The experimental Hf  value of C6H6(g) is 83 kJ/mol. Explain the discrepancy between the experimental value and the calculated Hf  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. 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 bonds.

Values of measured bond energies may vary greatly depending on the molecule studied. Consider the following reactions: Rationalize the difference in the values of H for these reactions, even though each reaction appears to involve only the breaking of one NCl 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 , O2 . Explain your answer.

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

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

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 number of electrons determined in part b) as follows: 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 de- fined 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][SbF6] : 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][SbF6] can be prepared when 2.93 mL anhydrous HF (density 0.975 g/mL) and 10.0 mL SbF5 (density 3.10 g/mL) are allowed to react? 93

Determine the molecular structure and hybridization of the central atom X in the polyatomic ion XY3  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 1s 2 2s 2 2p6 .