In Chapter 10, we saw that electronegativity differences

Indicate several ways in which the valence-bond method is superior to Lewis structures in describing covalent bonds.

Explain why it is necessary to hybridize atomic orbitals when applying the valence-bond method that is, why are there so few molecules that can be described by the overlap of pure atomic orbitals only?

Describe the molecular geometry of suggested by each of the following methods: (a) Lewis theory; (b) valence-bond method using simple atomic orbitals; (c) VSEPR theory; (d) valence-bond method using hybridized atomic orbitals.

Describe the molecular geometry of suggested by each of the following methods: (a) Lewis theory; (b) valence-bond method using simple atomic orbitals; (c) VSEPR theory; (d) valence-bond method using hybridized atomic orbitals.

In which of the following, would you expect to find hybridization of the central atom? Explain.

In the manner of Example 11-1, describe the probable structure and bonding in (a) HI; (b) BrCl; (c) (d)

For each of the following species, identify the central atom(s) and propose a hybridization scheme for those atom(s): (a) (b) (c) (d)

Propose a plausible Lewis structure, geometric structure, and hybridization scheme for the NSF molecule.

Describe a hybridization scheme for the central Cl atom in the molecule that is consistent with the geometric shape pictured in Table 10.1. Which orbitals of the Cl atom are involved in overlaps, and which are occupied by lone-pair electrons?

Describe a hybridization scheme for the central S atom in the molecule that is consistent with the geometric shape pictured in Table 10.1. Which orbitals of the S atom are involved in overlaps, and which are occupied by lone-pair electrons?

Match each of the following species with one of these hybridization schemes: (a) (b) COS; (c) (d) (e)

Propose a hybridization scheme to account for bonds formed by the central carbon atom in each of the following molecules: (a) hydrogen cyanide, HCN; NO3 AsF5. -SiCl4 ; ; PF6 -sp, sp2, sp3, sp3d, sp3d2. ; SF4 ClF3 BF4 -ClO3 . -CO2 ; HONO2 ; ; OCl2. H2Se; NO2 sp2 -, CO3 SO2, CCl4, CO, 2-, CCl4 H2O (b) methyl alcohol, (c) acetone, (d) carbamic acid, CH3OH; 1CH322CO; Valence-Bond Method H 2 NCOH O

Indicate which of the following molecules and ions are linear, which are planar, and which are neither. Then propose hybridization schemes for the central atoms. (a) (b) (c) (d)

In the manner of Figure 11-17, indicate the structures of the following molecules in terms of the overlap of simple atomic orbitals and hybrid orbitals: (a) (b) (c)

Write Lewis structures for the following molecules, and then label each and bond. (a) HCN; (b) (c) (d) HONO.

Represent bonding in the carbon dioxide molecule, by (a) a Lewis structure and (b) the valencebond method. Identify and bonds, the necessary hybridization scheme, and orbital overlap.

Use the method of Figure 11-18 to represent bonding in each of the following molecules: (a) (b) ONCl; (c) HONO; (d)

Use the method of Figure 11-18 to represent bonding in each of the following ions: (a) (b) (c) (d)

The molecular model below represents citric acid, an acidic component of citrus juices. Represent bonding in the citric acid molecule using the method of Figure 11-18 to indicate hybridization schemes and orbital overlaps.

Malic acid is a common organic acid found in unripe apples and other fruit. With the help of the molecular model shown below, represent bonding in the malic acid molecule, using the method of Figure 11-18 to indicate hybridization schemes and orbital overlaps.

Shown below are ball-and-stick models. Describe hybridization and orbital-overlap schemes consistent with these structures.

Shown below are ball-and-stick models. Describe hybridization and orbital-overlap schemes consistent with these structures.

Propose a bonding scheme that is consistent with the structure for propynal. [Hint: Consult Table 10.2 to assess the multiple-bond character in some of the bonds.]

The structure of the molecule allene, is shown here. Propose hybridization schemes for the C atoms in this molecule.

Angelic acid, shown below, occurs in sumbol root, a herb used as a stimulant Represent the bonding in the angelic acid molecule by using the method in Figure 11-18 to indicate hybridization schemes and orbital overlaps. What is the maximum number of atoms that can lie in the same plane?

Dimethylolpropionic acid, shown below, is used in the preparation of resins. Represent the bonding in the dimethylolpropionic acid molecule by using the method in Figure 11-18 to indicate hybridization schemes and orbital overlaps. What is the maximum number of atoms that can lie in the same plane?

Explain the essential difference in how the valencebond method and molecular orbital theory describe a covalent bond.

Describe the bond order of diatomic carbon, with Lewis theory and molecular orbital theory, and explain why the results are different.

has an exceptionally high N21g2 bond energy. Would C2,

The paramagnetism of gaseous has been established. Explain how this observation confirms that the orbitals are at a lower energy than the orbital for B2.

In our discussion of bonding, we have not encountered a bond order higher than triple. Use the energylevel diagrams of Figure 11-26 to show why this is to be expected.

Is it correct to say that when a diatomic molecule loses an electron, the bond energy always decreases (that is, that the bond is always weakened)? Explain.

For the following pairs of molecular orbitals, indicate the one you expect to have the lower energy, and state the reason for your choice. (a) or (b) or (c) or (d) or s*2ps .

For each of the species and (a) Write the molecular orbital diagram (as in Example 11-6). (b) Determine the bond order, and state whether you expect the species to be stable or unstable. (c) Determine if the species is diamagnetic or paramagnetic; and if paramagnetic, indicate the number of unpaired electrons.

Write plausible molecular orbital diagrams for the following heteronuclear diatomic species: (a) NO; (b) (c) CO; (d) CN; (e) (f) (g) BN.

We have used the term isoelectronic to refer to atoms with identical electron configurations. In molecular orbital theory, this term can be applied to molecular species as well. Which of the species in Exercise 35 are isoelectronic?

Consider the molecules and and use molecular orbital theory to answer the following: (a) Write the molecular orbital configuration of each ion (ignore the electrons). (b) Predict the bond order of each ion. 1s N2 + NO+ CN CN+; -; NO+; NO+C2 , +, O2 -, F2 +, (c) Which of these ions is paramagnetic? Which is diamagnetic? (d) Which of these ions do you think has the greater bond length? Explain.

Consider the molecules and and use molecular orbital theory to answer the following: (a) Write the molecular orbital configuration of each ion (ignore the electrons). (b) Predict the bond order of each ion. (c) Which of these ions is paramagnetic? Which is diamagnetic? (d) Which of these ions do you think has the greater bond length? Explain.

Construct the molecular orbital diagram for CF. Would you expect the bond length of to be longer or shorter than that of CF?

Construct the molecular orbital diagram for CaF. Would you expect the bond length of to be longer or shorter than that of CaF?

Explain why the concept of delocalized molecular orbitals is essential to an understanding of bonding in the benzene molecule,

Explain how it is possible to avoid the concept of resonance by using molecular orbital theory. C6H6.

In which of the following molecules would you expect to find delocalized molecular orbitals: (a) (b) (c) Explain.

In which of the following ions would you expect to find delocalized molecular orbitals: (a) (b) (c) CH3 Explain.

Which of the following factors are especially important in determining whether a substance has metallic properties: (a) atomic number; (b) atomic mass; (c) number of valence electrons; (d) number of vacant atomic orbitals; (e) total number of electronic shells in the atom? Explain.

Based on the ground-state electron configurations of the atoms, how would you expect the melting points and hardnesses of sodium, iron, and zinc to compare? Explain.

How many energy levels are present in the 3s conduction band of a single crystal of sodium weighing 26.8 mg? How many electrons are present in this band?

Magnesium is an excellent electrical conductor even though it has a full 3s subshell with the electron configuration: Use band theory to explain why magnesium conducts electricity.

From this list of terms electrical conductor, insulator, semiconductor choose the one that best characterizes each of the following materials: (a) stainless steel; (b) solid sodium chloride; (c) sulfur; (d) germanium; (e) seawater; (f) solid iodine.

In what type of material is the energy gap between the valence band and the conduction band greatest: metal, semiconductor, or insulator? Explain.

Which of the following substances, when added in trace amounts to silicon, would produce a p-type semiconductor: (a) sulfur, (b) arsenic, (c) lead, (d) boron, (e) gallium arsenide, (f) gallium? Explain.

Which of the following substances, when added in trace amounts to germanium, would produce an n-type semiconductor: (a) sulfur, (b) aluminum, (c) tin, (d) cadmium sulfide, (e) arsenic, (f) gallium arsenide? Explain.

The effect of temperature change on the electrical conductivity of ultrapure silicon is quite different from that on silicon containing a minute trace of arsenic. Why is this so?

Explain why the electrical conductivity of a semiconductor is significantly increased if trace amounts of either donor or acceptor atoms are present, but is unchanged if both are present in equal number.

The energy gap, for silicon is What is the minimum wavelength of light that can promote an electron from the valence band to the conduction band in silicon? In what region of the electromagnetic spectrum is this light?

Explain why the solar cell in Figure 11-40 operates over a broad range of wavelengths rather than at a single wavelength (often the case when quantum effects are involved)?

The Lewis structure of indicates that the nitrogento- nitrogen bond is a triple covalent bond. Other evidence suggests that the bond in this molecule involves the overlap of hybrid orbitals. (a) Draw orbital diagrams for the N atoms to describe bonding in (b) Can this bonding be described by either or hybridization of the N atoms? Can bonding in be described in terms of unhybridized orbitals? Explain.

Show that both the valence-bond method and molecular orbital theory provide an explanation for the existence of the covalent molecule in the gaseous state. Would you predict by the Lewis theory?

A group of spectroscopists believe that they have detected one of the following species: or Assume that the energy-level diagrams of Figure 11-26 apply, and describe bonding in these species. Which of these species would you expect the spectroscopists to have observed?

Lewis theory is satisfactory to explain bonding in the ionic compound but it does not readily explain formation of the ionic compounds potassium superoxide, and potassium peroxide, (a) Show that molecular orbital theory can provide this explanation. (b) Write Lewis structures consistent with the molecular orbital explanation.

The compound potassium sesquoxide has the empirical formula Show that this compound can be described by an appropriate combination of potassium, peroxide, and superoxide ions. Write a Lewis structure for a formula unit of the compound.

Draw a Lewis structure for the urea molecule, and predict its geometric shape with the VSEPR theory. Then revise your assessment of this molecule, given the fact that all the atoms lie in the same plane, and all the bond angles are Propose a hybridization and bonding scheme consistent with these experimental observations.

Methyl nitrate, is used as a rocket propellant. The skeletal structure of the molecule is The N and three O atoms all lie in the same plane, but the group is not in the same plane as the group. The bond angle is and the bond angle is One nitrogen-to-oxygen bond length is 136 pm, and the other two are 126 pm. (a) Draw a sketch of the molecule showing its geometric shape. (b) Label all the bonds in the molecule as or and indicate the probable orbital overlaps involved. (c) Explain why all three nitrogen-to-oxygen bond lengths are not the same.

Fluorine nitrate, is an oxidizing agent used as a rocket propellant. A reference source lists the following data for (The subscript a shows that this O atom is different from the other two.) Bond lengths: NOa OaF = 142 pm = 139 pm; NO = 129 pm; FOaNO2. FONO2, s p, ONO 125. CON 105, CH3 NO3 CH3ONO2. CH3NO3, 120. CO1NH222, K2O3. KO2, K2O2. K2O, NeF-. NeF, NeF+, Na2 Na2 N2 sp 3 sp 2 N2. sp s N2 Bond angles: plane is perpendicular to the plane Use these data to construct a Lewis structure(s), a three-dimensional sketch of the molecule, and a plausible bonding scheme showing hybridization and orbital overlaps.

Draw a Lewis structure(s) for the nitrite ion, Then propose a bonding scheme to describe the and bonding in this ion. What conclusion can you reach about the number and types of molecular orbitals in this ion? Explain.

Think of the reaction shown here as involving the transfer of a fluoride ion from to to form the ions and As a result, the hybridization scheme of each central atom must change. For each reactant molecule and product ion, indicate (a) its geometric structure and (b) the hybridization scheme for its central atom.

In the gaseous state, molecules have two nitrogen-to-oxygen bond distances of 121 pm and one of 140 pm. Draw a plausible Lewis structure(s) to represent this fact, and propose a bonding scheme in the manner of Figure 11-18.

does not exist as a stable molecule, but there is evidence that such a molecule can be formed between electronically excited He atoms. Write a molecular orbital diagram to account for this.

The molecule formamide, has the approximate bond angles The bond length is 138 pm. Two Lewis structures can be written for this molecule, with the true structure being a resonance hybrid of the two. Propose a hybridization and bonding scheme for each structure.

Pyridine, is used in the synthesis of vitamins and drugs. The molecule can be thought of in terms of replacing one CH unit in benzene with a N atom. Draw orbital diagrams to show the orbitals of the C and N atoms involved in the and bonding in pyridine. How many bonding and antibonding molecular orbitals are present? How many delocalized electrons are present?

One of the characteristics of antibonding molecular orbitals is the presence of a nodal plane. Which of the bonding molecular orbitals considered in this chapter have nodal planes? Explain how a molecular orbital can have a nodal plane and still be a bonding molecular orbital.

The ion is linear, but the ion is bent. Describe hybridization schemes for the central Cl atom consistent with this difference in structure.

Ethyl cyanoacetate, a chemical used in the synthesis of dyes and pharmaceuticals, has the mass percent composition: 53.09% C, 6.24% H, 12.39% N, and 28.29% O. In the manner of Figure 11-18, show a bonding scheme for this substance. The scheme should designate orbital overlaps, and bonds, and expected bond angles.

A certain monomer used in the production of polymers has one nitrogen atom and the mass composition 67.90% C, 5.70% H, and 26.40% N. Sketch the probable geometric structure of this molecule, labeling all the expected bond lengths and bond angles.

A solar cell that is 15% efficient in converting solar to electric energy produces an energy flow of when exposed to full sunlight. (a) If the cell has an area of what is the power output of the cell, in watts? (b) If the power calculated in part (a) is produced at 0.45 V, how much current does the cell deliver?

Toluene-2,4-diisocyanate is used in the manufacture of polyurethane foam. Its structural formula is shown 40.0 cm2, 1.00 kW>m2 s p here. Describe the hybridization scheme for the atoms marked with an asterisk, and indicate the values of the bond angles marked and

The anion is linear, and the anion is V-shaped, with a angle between the two arms of the V. For the central atoms in these ions, propose hybridization schemes that are consistent with these observations.

Pentadiene, has three isomers, depending on the position of the two double bonds. Determine the shape of these isomers by using VSEPR theory. Describe the bonding in these molecules by using the valence-bond method. Do the shapes agree in the two theories? Use molecular orbital theory to decide which of these molecules has a delocalized system. Sketch the molecular orbital and an energy-level diagram.

A conjugated hydrocarbon has an alternation of double and single bonds. Draw the molecular orbitals of the system of 1,3,5-hexatriene. If the energy required to excite an electron from the HOMO to the LUMO corresponds to a wavelength of 256 nm, do you expect the wavelength for the corresponding excitation in 1,3,5,7-octatetraene to be a longer or shorter wavelength?

Resonance energy is the difference in energy between a real molecule a resonance hybrid and its most important contributing structure. To determine the resonance energy for benzene, we can determine an energy change for benzene and the corresponding change for one of the Kekul structures. The resonance energy is the difference between these two quantities. (a) Use data from Appendix D to determine the enthalpy of hydrogenation of liquid benzene to liquid cyclohexane. (b) Use data from Appendix D to determine the enthalpy of hydrogenation of liquid cyclohexene to liquid cyclohexane. For the enthalpy of formation of liquid cyclohexene, use (c) Assume that the enthalpy of hydrogenation of 1,3,5-cyclohexatriene is three times as great as that of cyclohexene, and calculate the resonance energy of benzene. (d) Another way to assess resonance energy is through bond energies. Use bond energies from Table 10.3 (page 435) to determine the total enthalpy change required to break all the bonds in a Kekul structure of benzene. Next, determine the enthalpy change for the dissociation of into its gaseous atoms by using data from Table 10.3 and Appendix D. Then calculate the resonance energy of benzene.

The 60-cycle alternating electric current (AC) commonly used in households changes direction 120 times per second. That is, in a one-second time period a terminal at an electric outlet is positive 60 times and negative 60 times. In direct electric current (DC), the flow between terminals is in one direction only. A rectifier is a device that converts alternating to direct current. One type of rectifier is the p n junction rectifier. It is commonly incorporated in adapters required to operate electronic devices from ordinary house current. In the operation of this rectifier, a p-type semiconductor and an n-type semiconductor are in contact along a boundary, or junction. Each semiconductor is connected to one of the terminals in an AC electrical outlet. Describe how this rectifier works. That is, show that when the semiconductors are connected to the terminals in an AC outlet, half the time a large flow of charge occurs and half the time essentially no charge flows across the p n junction.

Furan, is a substance derivable from oat hulls, corn cobs, and other cellulosic waste. It is a starting material for the synthesis of other chemicals used as pharmaceuticals and herbicides. The furan molecule is planar and the C and O atoms are bonded into a fivemembered pentagonal ring. The H atoms are attached to the C atoms. The chemical behavior of the molecule suggests that it is a resonance hybrid of several contributing structures. These structures show that the double bond character is associated with the entire ring in the form of a electron cloud. (a) Draw Lewis structures for the several contributing structures to the resonance hybrid mentioned above. (b) Draw orbital diagrams to show the orbitals that are involved in the and bonding in furan. [Hint: You need use only one of the contributing structures, such as the one with no formal charges.] (c) How many electrons are there in the furan molecule? Show that this number of p electrons is the same, regardless of the contributing structure you use for this assessment.

As discussed in the Are You Wondering feature on page 457, the hybrid orbitals are algebraic combinations of the s and p orbitals. The required combinations of and orbitals are (a) By combining the appropriate functions given in Table 8.1, construct a polar plot in the manner of Figure 8-26 for each of the above functions in the plane. In a polar plot, the value of is set at a fixed value (for example, 1). Describe the shapes and phases of the different portions of the hybrid orbitals, and compare them with those shown in Figure 11-11. (b) Convince yourself that the combinations employing the or orbital also give similar hybrid orbitals but pointing in different directions. (c) The combinations for the hybrids in the xy plane are By constructing polar plots (in the plane), show that these functions correspond to the hybrids depicted in Figure 11-9.

In Chapter 10, we saw that electronegativity differences determine whether bond dipoles exist in a molecule and that molecular shape determines whether bond dipoles cancel (nonpolar molecules) or combine to produce a resultant dipole moment (polar molecules). Thus, the ozone molecule, has no bond dipoles because all the atoms are alike. Yet, does have a resultant dipole moment: The electrostatic potential map for ozone is shown below. Use the electrostatic potential map to decide the direction of the dipole. Using the ideas of delocalized bonding in molecules, can you rationalize this electrostatic potential map?

Borazine, is often referred to as inorganic benzene because of its similar structure. Like benzene, borazine has a delocalized system. Describe the molecular orbitals of the system. Identify the highest occupied molecular orbital (HOMO) and the lowest unoccupied molecular orbital (LUMO). How many nodes does the LUMO possess?

Which of the following combinations of orbitals give rise to bonding molecular orbitals? For those combinations that do, label the resulting bonding molecular orbital as s or p.

Construct a molecular orbital diagram for HF, and label the molecular orbitals as bonding, antibonding, or nonbonding.

In your own words, define the following terms or symbols: (a) (b) ; (c) bond order; (d) bond.

Briefly describe each of the following ideas: (a) hybridization of atomic orbitals; (b) framework; (c) Kekul structures of benzene, (d) band theory of metallic bonding.

Explain the important distinctions between the terms in each of the following pairs: (a) and bonds; (b) localized and delocalized electrons; (c) bonding and antibonding molecular orbitals; (d) metal and semiconductor.

A molecule in which hybrid orbitals are used by the central atom in forming covalent bonds is (a) (b) (c) (d)

The bond angle in is best described as (a) between and (b) less than in (c) less than in but not less than (d) less than

The hybridization scheme for the central atom includes a d orbital contribution in (a) (b) (c) (d)

Of the following, the species with a bond order of 1 is (a) (b) (c) (d)

The hybridization scheme for Xe in is (a) (b) (c) (d)

Delocalized molecular orbitals are found in (a) (b) (c) (d)

The best electrical conductor of the following materials is (a) Li(s); (b) (c) Ge(s); (d) Si(s).

A substance in which the valence and conduction bands overlap is (a) a semiconductor; (b) a metalloid; (c) a metal; (d) an insulator. Br2(l); CO3 CH 2-. HS 4; -; H2; sp3; sp3d; sp3d2. XeF2 sp; H2 -H2 Li2; He2; . +; NO3 H2Se. -; I3 PCl3; -; H2S, 90; 90. 109 120; H2S; H2Se N2; SO2; He2. PCl5; sp2 s p C6H6; s-bond s p sp 2*p 2;

Explain why the molecular structure of cannot be adequately described through overlaps involving pure s and p orbitals.

Why does the hybridization not account for bonding in the molecule What hybridization scheme does work? Explain.

What is the total number of (a) bonds and (b) bonds in the molecule

Which of the following species are paramagnetic? (a) (b) (c) Which species has the strongest bond?

Use the valence molecular orbital configuration to determine which of the following species is expected to have the lowest ionization energy: (a) (b) (c)

Use the valence molecular orbital configuration to determine which of the following species is expected to have the greatest electron affinity: (a) (b) (c) (d)

Which of these diatomic molecules do you think has the greater bond energy, or Explain.

Construct a concept map that embodies the ideas of valence bond theory.

Construct a concept map that connects the ideas of molecular orbital theory.

Construct a concept map that describes the interconnection between valence-bond theory and molecular orbital theory in the description of resonance structures.