Study Guide to Chemical Bonding

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DISCLAIMER: This Guide is not meant to be exhaustive. That is, I have tried to summarize the essential points of the lectures on this topic. The presence of a topic here does not guarantee a related question on an exam, nor are exam topics limited to what appears in this Guide. As with any Chemistry class, you are responsible for ALL of the assigned readings, problems and lecture material. Lectures will often contain information not covered or given less emphasis in the text.

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Text Reading

Chapter 13 (pages 582-641)
Chapter 14 (pages 650-680)

What should you learn from this section of the course?

We have worked our way through three different approaches to the understanding of the formation of chemical bonds: Lewis dot structures, the Valence Shell Electron Pair Repulsion (VSEPR) Model and molecular orbitals (and hybridization). These satisfy our criteria for a "good" model to different degrees. In general, you should be aware of the qualitative and quantitative differences among these three models. Which gives accurate predictions of chemical properties? Which yields the molecular formula? Can we estimate the shapes of molecules easily?

General

• Definitions of chemical bonding properties, such as bond order and bond length
• Correlate bond order with bond energy and/or length
• Differences between ionic and covalent bonds
• Use dipole moments to predict molecular geometry and "ionic character"

Lewis dot structures

• Be able to write structures whether or not they violate the "octet rule"
• Be comfortable with resonance structures
• Be able to calculate formal charges and use them to differentiate resonance structures

VSEPR

• Use steric numbers to predict geometries

Molecular orbital theory

• Be able to construct a diatomic molecular orbital energy diagram
• Use a molecular orbital energy diagram to write electron configurations for molecules
• Relate bond order to electron configuration
• Relate photoelectron spectroscopy to molecular orbitals

Orbital hybridization

• Understand the origin of hybrid orbitals
• Understand the relationship to molecular orbital theory
• Understand the relationship to VSEPR
• Use steric numbers to find the orbital hybridization
• Be able to describe molecular bonding in terms of orbital hybridization

Ionic bonding

• Understand the simple theory of formation of an ionic bond
• Understand the concept of lattice energy
• Understand the relationship of the Madelung constant to crystal structure
• Be able to relate lattice energy to physical properties

Recommended Problems from the Text

Chapter 13 Problems

Bonds & electronegativity: 11, 13, 17. Ionic compounds: 20, 21, 25, 27, 29. Bond energies: 31. Lewis structures: 47, 53, 57, 59, 63, 65. Formal charge: 67. Structure & polarity: 71 ( as refers to 47), 73 (as refers to 59), 75 (as refers to 73), 83, 85. Mixed concepts: 95, 99

Chapter 14 Problems

Hybrid orbitals: 9, 13, 15, 21. Molecular orbitals: 25, 31, 35, 41. Mixed concepts: 49, 55, 59, 61.

Additional Problems

1. The empirical formulae and molecular skeletons of a number of molecules composed of the elements H, C, N and O are given below. In each case enumerate all the resonance structures, which contain the correct number of electrons and for which each atom obeys the octet rule.

• NO₂^- (O N O)

o H₂CNH (H C N)

o C₆H₆

o H₂CCO (H C C O)

o HN₃ (H N N N)

The structures in parentheses are only meant to indicate which atoms are connected.

2. For the molecules considered in Problem 1, compute the formal charges on all the atoms for all the resonance structures that satisfy the octet rule. Evaluate the importance of the contribution that each of the resonance structures makes to the actual structure of each molecule or ion.

3. Predict the actual molecular structures of the following central atom molecules : OCl₂, PH₃, AsF₅, ClF₅, ClO₄^-, SO₃^2-, Cl₂F⁺, AsF₃, BrF₄^-, AgCl₂^-, SiF₆^2-, BrCl₂^-.

4. Predict the observed geometry of each of the following molecules or ions indicating any deviations from idealized bond angles: BeH₄^2-, H₂O, NO₃^-, NF₃, PCl₄⁺, XeF₄, HCN, S₃^2-.

5. Discuss the geometry of the sulfur dichloride (SCl₂) molecule from the viewpoint of the Valence-Shell-Electron-Pair-Repulsion (VSEPR) model. Do the predictions agree with the experimental bond angle of 100.3 degrees?

6. Use the VSEPR model to explain why NO₂⁺ is linear, NO₂ is bent (< ONO = 134 degrees) and NO₂^- is bent (< ONO = 115 degrees).

7. The two carbon-oxygen bonds in acetic acid (CH₃COOH) differ in length, but the two carbon-oxygen bonds in the acetate ion (CH₃COO^-) are equal. Explain using Lewis structures.

8. Write three Lewis dot structures that follow the octet rule for cyanate ion, NCO^-, and isocyanate ion, CNO^-. Use formal charges to identify the most important resonance structures for each ion.

9. Use the VSEPR method to predict the geometry of each of the following molecules: SO₂, PbCl₄ and SbH₃.

10. The dipole moment of gaseous CsF is 7.88D with a bond length of 0.255 nm. Determine the percent ionic character in this molecule.

11. Using electronegativities, arrange the following in order of increasing ionic character of the bonds: HCl, ClF, KCl, CCl₄ and Cl₂.

12. Arrange the following molecules in order of increasing dipole moment, noting any molecules with a dipole moment of zero: CH₃F, CH₃Cl, CH₄, CCl₄ and CH₃Br.

13. From each pair of molecules/ions, select the one with the greater bond energy: B₂ or B₂⁺; O₂⁺ or O₂^-; Be₂ or Be₂⁺; F₂ or F₂⁺; F₂ or F₂^-.

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