# 9.E: Exercises

These are homework exercises to accompany the Textmap created for "Chemistry: The Central Science" by Brown et al. Complementary General Chemistry question banks can be found for other Textmaps and can be accessed here. In addition to these publicly available questions, access to private problems bank for use in exams and homework is available to faculty only on an individual basis; please contact Delmar Larsen for an account with access permission.

## 9.2: THE VSEPR MODEL

### Conceptual Problems

1. What is the main difference between the VSEPR model and Lewis electron structures?

2. What are the differences between molecular geometry and Lewis electron structures? Can two molecules with the same Lewis electron structures have different molecular geometries? Can two molecules with the same molecular geometry have different Lewis electron structures? In each case, support your answer with an example.

3. How does the VSEPR model deal with the presence of multiple bonds?

4. Three molecules have the following generic formulas: AX2, AX2E, and AX2E2. Predict the molecular geometry of each, and arrange them in order of increasing X–A–X angle.

5. Which has the smaller angles around the central atom—H2S or SiH4? Why? Do the Lewis electron structures of these molecules predict which has the smaller angle?

6. Discuss in your own words why lone pairs of electrons occupy more space than bonding pairs. How does the presence of lone pairs affect molecular geometry?

• When using VSEPR to predict molecular geometry, the importance of repulsions between electron pairs decreases in the following order: LP–LP, LP–BP, BP–BP. Explain this order. Draw structures of real molecules that separately show each of these interactions.

• How do multiple bonds affect molecular geometry? Does a multiple bond take up more or less space around an atom than a single bond? a lone pair?

• Straight-chain alkanes do not have linear structures but are “kinked.” Using n-hexane as an example, explain why this is so. Compare the geometry of 1-hexene to that of n-hexane.

• How is molecular geometry related to the presence or absence of a molecular dipole moment?

• How are molecular geometry and dipole moments related to physical properties such as melting point and boiling point?

• What two features of a molecule’s structure and bonding are required for a molecule to be considered polar? Is COF2 likely to have a significant dipole moment? Explain your answer.

• When a chemist says that a molecule is polar, what does this mean? What are the general physical properties of polar molecules?

• Use the VSPER model and your knowledge of bonding and dipole moments to predict which molecules will be liquids or solids at room temperature and which will be gases. Explain your rationale for each choice. Justify your answers.

1. CH3Cl
2. PCl3
3. CO
4. SF6
5. IF5
6. CH3OCH3
7. CCl3H
8. H3COH
• The idealized molecular geometry of BrF5 is square pyramidal, with one lone pair. What effect does the lone pair have on the actual molecular geometry of BrF5? If LP–BP repulsions were weaker than BP–BP repulsions, what would be the effect on the molecular geometry of BrF5?

• Which has the smallest bond angle around the central atom—H2S, H2Se, or H2Te? the largest? Justify your answers.

• Which of these molecular geometries always results in a molecule with a net dipole moment: linear, bent, trigonal planar, tetrahedral, seesaw, trigonal pyramidal, square pyramidal, and octahedral? For the geometries that do not always produce a net dipole moment, what factor(s) will result in a net dipole moment?

1. To a first approximation, the VSEPR model assumes that multiple bonds and single bonds have the same effect on electron pair geometry and molecular geometry; in other words, VSEPR treats multiple bonds like single bonds. Only when considering fine points of molecular structure does VSEPR recognize that multiple bonds occupy more space around the central atom than single bonds.

2. Physical properties like boiling point and melting point depend upon the existence and magnitude of the dipole moment of a molecule. In general, molecules that have substantial dipole moments are likely to exhibit greater intermolecular interactions, resulting in higher melting points and boiling points.

3. The term “polar” is generally used to mean that a molecule has an asymmetrical structure and contains polar bonds. The resulting dipole moment causes the substance to have a higher boiling or melting point than a nonpolar substance.

### Numerical Problems

1. Give the number of electron groups around the central atom and the molecular geometry for each molecule. Classify the electron groups in each species as bonding pairs or lone pairs.

1. BF3
2. PCl3
3. XeF2
4. AlCl4
5. CH2Cl2
2. Give the number of electron groups around the central atom and the molecular geometry for each species. Classify the electron groups in each species as bonding pairs or lone pairs.

1. ICl3
2. CCl3+
3. H2Te
4. XeF4
5. NH4+
3. Give the number of electron groups around the central atom and the molecular geometry for each molecule. For structures that are not linear, draw three-dimensional representations, clearly showing the positions of the lone pairs of electrons.

1. HCl
2. NF3
3. ICl2+
4. N3
5. H3O+
4. Give the number of electron groups around the central atom and the molecular geometry for each molecule. For structures that are not linear, draw three-dimensional representations, clearly showing the positions of the lone pairs of electrons.

1. SO3
2. NH2
3. NO3
4. I3
5. OF2
5. What is the molecular geometry of ClF3? Draw a three-dimensional representation of its structure and explain the effect of any lone pairs on the idealized geometry.

6. Predict the molecular geometry of each of the following.

1. ICl3
2. AsF5
3. NO2
4. TeCl4
7. Predict whether each molecule has a net dipole moment. Justify your answers and indicate the direction of any bond dipoles.

1. NO
2. HF
3. PCl3
4. CO2
5. SO2
6. SF4
8. Predict whether each molecule has a net dipole moment. Justify your answers and indicate the direction of any bond dipoles.

1. OF2
2. BCl3
3. CH2Cl2
4. TeF4
5. CH3OH
6. XeO4
9. Of the molecules Cl2C=Cl2, IF3, and SF6, which has a net dipole moment? Explain your reasoning.

10. Of the molecules SO3, XeF4, and H2C=Cl2, which has a net dipole moment? Explain your reasoning.

1. trigonal planar (all electron groups are bonding pairs)
2. tetrahedral (one lone pair on P)
3. trigonal bipyramidal (three lone pairs on Xe)
4. tetrahedral (all electron groups on Al are bonding pairs)
5. tetrahedral (all electron groups on C are bonding pairs)
1. four electron groups, linear molecular geometry
2. four electron groups, pyramidal molecular geometry

3. four electron groups, bent molecular geometry
4. two electron groups, linear molecular geometry
5. four electron groups, pyramidal molecular geometry
1. The idealized geometry is T shaped, but the two lone pairs of electrons on Cl will distort the structure, making the F–Cl–F angle less than 180°.

2. Cl2C=CCl2: Although the C–Cl bonds are rather polar, the individual bond dipoles cancel one another in this symmetrical structure, and Cl2C=CCl2 does not have a net dipole moment.

IF3: In this structure, the individual I–F bond dipoles cannot cancel one another, giving IF3 a net dipole moment.

SF6: The S–F bonds are quite polar, but the individual bond dipoles cancel one another in an octahedral structure. Thus, SF6 has no net dipole moment.

## 9.3: MOLECULAR SHAPE AND MOLECULAR POLARITY

### Conceptual Problems

1. Why do ionic compounds such as KI exhibit substantially less than 100% ionic character in the gas phase?
2. Of the compounds LiI and LiF, which would you expect to behave more like a classical ionic compound? Which would have the greater dipole moment in the gas phase? Explain your answers.

### Numerical Problems

1. Predict whether each compound is purely covalent, purely ionic, or polar covalent.

1. RbCl
2. S8
3. TiCl2
4. SbCl3
5. LiI
6. Br2
2. Based on relative electronegativities, classify the bonding in each compound as ionic, covalent, or polar covalent. Indicate the direction of the bond dipole for each polar covalent bond.

1. NO
2. HF
3. MgO
4. AlCl3
5. SiO2
6. the C=O bond in acetone
7. O3
3. Based on relative electronegativities, classify the bonding in each compound as ionic, covalent, or polar covalent. Indicate the direction of the bond dipole for each polar covalent bond.

1. NaBr
2. OF2
3. BCl3
4. the S–S bond in CH3CH2SSCH2CH3
5. the C–Cl bond in CH2Cl2
6. the O–H bond in CH3OH
7. PtCl42−
4. Classify each species as having 0%–40% ionic character, 40%–60% ionic character, or 60%–100% ionic character based on the type of bonding you would expect. Justify your reasoning.

1. CaO
2. S8
3. AlBr3
4. ICl
5. Na2S
6. SiO2
7. LiBr
5. If the bond distance in HCl (dipole moment = 1.109 D) were double the actual value of 127.46 pm, what would be the effect on the charge localized on each atom? What would be the percent negative charge on Cl? At the actual bond distance, how would doubling the charge on each atom affect the dipole moment? Would this represent more ionic or covalent character?

6. Calculate the percent ionic character of HF (dipole moment = 1.826 D) if the H–F bond distance is 92 pm.

7. Calculate the percent ionic character of CO (dipole moment = 0.110 D) if the C–O distance is 113 pm.

8. Calculate the percent ionic character of PbS and PbO in the gas phase, given the following information: for PbS, r = 228.69 pm and µ = 3.59 D; for PbO, r = 192.18 pm and µ = 4.64 D. Would you classify these compounds as having covalent or polar covalent bonds in the solid state?

## 9.5: HYBRID ORBITALS

### Conceptual Problems

1. Arrange sp, sp3, and sp2 in order of increasing strength of the bond formed to a hydrogen atom. Explain your reasoning.

2. What atomic orbitals are combined to form sp3, sp, sp3d2, and sp3d? What is the maximum number of electron-pair bonds that can be formed using each set of hybrid orbitals?

3. Why is it incorrect to say that an atom with sp2 hybridization will form only three bonds? The carbon atom in the carbonate anion is sp2 hybridized. How many bonds to carbon are present in the carbonate ion? Which orbitals on carbon are used to form each bond?

4. If hybridization did not occur, how many bonds would N, O, C, and B form in a neutral molecule, and what would be the approximate molecular geometry?

5. How are hybridization and molecular geometry related? Which has a stronger correlation—molecular geometry and hybridization or Lewis structures and hybridization?

6. In the valence bond approach to bonding in BeF2, which step(s) require(s) an energy input, and which release(s) energy?

7. The energies of hybrid orbitals are intermediate between the energies of the atomic orbitals from which they are formed. Why?

8. How are lone pairs on the central atom treated using hybrid orbitals?

9. Because nitrogen bonds to only three hydrogen atoms in ammonia, why doesn’t the nitrogen atom use sp2 hybrid orbitals instead of sp3 hybrids?

10. Using arguments based on orbital hybridization, explain why the CCl62− ion does not exist.

11. Species such as NF52− and OF42− are unknown. If 3d atomic orbitals were much lower energy, low enough to be involved in hybrid orbital formation, what effect would this have on the stability of such species? Why? What molecular geometry, electron-pair geometry, and hybridization would be expected for each molecule?

### Numerical Problems

1. Draw an energy-level diagram showing promotion and hybridization to describe the bonding in CH3. How does your diagram compare with that for methane? What is the molecular geometry?

2. Draw an energy-level diagram showing promotion and hybridization to describe the bonding in CH3+. How does your diagram compare with that for methane? What is the molecular geometry?

3. Draw the molecular structure, including any lone pairs on the central atom, state the hybridization of the central atom, and determine the molecular geometry for each molecule.

1. BBr3
2. PCl3
3. NO3
4. Draw the molecular structure, including any lone pairs on the central atom, state the hybridization of the central atom, and determine the molecular geometry for each species.

1. AsBr3
2. CF3+
3. H2O
5. What is the hybridization of the central atom in each of the following?

1. CF4
2. CCl22−
3. IO3
4. SiH4
6. What is the hybridization of the central atom in each of the following?

1. CCl3+
2. CBr2O
3. CO32−
4. IBr2
7. What is the hybridization of the central atom in PF6? Is this ion likely to exist? Why or why not? What would be the shape of the molecule?

8. What is the hybridization of the central atom in SF5? Is this ion likely to exist? Why or why not? What would be the shape of the molecule?

1. The promotion and hybridization process is exactly the same as shown for CH4 in the chapter. The only difference is that the C atom uses the four singly occupied sp3 hybrid orbitals to form electron-pair bonds with only three H atoms, and an electron is added to the fourth hybrid orbital to give a charge of 1–. The electron-pair geometry is tetrahedral, but the molecular geometry is pyramidal, as in NH3.

1. sp2, trigonal planar

2. sp3, pyramidal

3. sp2, trigonal planar

2. The central atoms in CF4, CCl22–, IO3, and SiH4 are all sp3 hybridized.

3. The phosphorus atom in the PF6 ion is sp3d2 hybridized, and the ion is octahedral. The PF6 ion is isoelectronic with SF6 and has essentially the same structure. It should therefore be a stable species.

## 9.6: MULTIPLE BONDS

### Conceptual Problems

1. What information is obtained by using the molecular orbital approach to bonding in O3 that is not obtained using the VSEPR model? Can this information be obtained using a Lewis electron-pair approach?
2. How is resonance explained using the molecular orbital approach?
3. Indicate what information can be obtained by each method:

Lewis Electron Structures VSEPR Model Valence Bond Theory Molecular Orbital Theory
Geometry
Resonance
Orbital Hybridization
Reactivity
Expanded Valences
Bond Order

### Numerical Problems

1. Using both a hybrid atomic orbital and molecular orbital approaches, describe the bonding in $$BCl_3$$ and $$CS_3^{2−}$$.
2. Use both a hybrid atomic orbital and molecular orbital approaches to describe the bonding in $$CO_2$$ and $$N_3^−$$.