Exam Questions Theories Of Chemical Bonding 301 Chapter 7

CHAPTER 7

THEORIES OF CHEMICAL BONDING

CHAPTER STUDY OBJECTIVES

1. Use the orbital overlap model to explain the bonding in simple molecules.

SKILLS TO MASTER: Describing the bonding in molecules using the orbital overlap model

KEY CONCEPTS: Bonding orbitals are constructed by combining atomic orbitals from adjacent atoms. Only the valence orbitals are needed to describe bonding.

2. Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

SKILLS TO MASTER: Correlating the steric number with the electron group geometry, hybridization, number of hybrid orbitals, and number of unused p orbitals on an inner atom

KEY CONCEPTS: The number of hybrid orbitals generated by the hybridization process equals the number of valence atomic orbitals participating in hybridization. The steric number of an inner atom can be used to infer the hybrid orbitals it is using. Elements beyond the second row of the periodic table can form bonds to more than four ligands and can be associated with more than an octet of electrons.

3. Describe the σ and π bonding systems in multiple bonds.

SKILLS TO MASTER: Constructing the σ and π bonding framework of a molecule

KEY CONCEPTS: A sigma (σ) bond has high electron density distributed symmetrically along the bond axis. A pi (π) bond has high electron density concentrated above and below the bond axis.

4. Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

SKILLS TO MASTER: Calculating the bond order; drawing MO diagrams for heteronuclear diatomic molecules; predicting the effects of ionization on bond length and energy

KEY CONCEPTS: The total number of molecular orbitals produced by a set of interacting atomic orbitals is equal to the number of interacting atomic orbitals. Delocalized electrons occupy molecular orbitals (MOs), so called because they may span entire molecules. A bonding MO results from positive overlap of atomic orbital wave functions and stabilizes the bond. An antibonding MO results from negative overlap of atomic orbital wave functions and destabilizes the bond. Paramagnetism arises from unpaired electrons.

5. Describe the bonding in three-atom π systems.

SKILLS TO MASTER: Using the composite model of bonding

KEY CONCEPTS: Hybrid orbitals can be used to construct the molecular framework. Molecular orbitals can then be used to construct the π bonding system.

6. Describe the bonding in extended π systems.

KEY CONCEPTS: A delocalized π system is present whenever p orbitals on more than two adjacent atoms are in position for side-by-side overlap. Molecules that have alternating single and double bonds are said to have conjugated π systems. Delocalized π systems result in increased stability and often in coloured compounds.

7. Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

KEY CONCEPTS: Metals conduct heat and electricity because their band gaps are so small.

Semiconductors have intermediate band gaps. A doped semiconductor has almost the same band structure as the pure material, but different electron populations in its bands.

Multiple Choice QUESTIONS

1. Which of the following molecules is NOT well described by the valence bond model?

a) SiH₄

b) H₂S

c) SbH₃

d) H₂Te

e) HF

Difficulty: Medium

Learning Objective: Use the orbital overlap model to explain the bonding in simple molecules.

Section Reference: 7.1 Localized Bonds

2. According to the valence bond model, which of the following is NOT a requirement of bonding?

a) overlap of atomic orbitals

b) singly occupied atomic orbitals

c) no two electrons in a molecule have identical descriptions

d) atoms in molecule contain unpaired valence electrons

e) orbitals must have same azimuthal quantum number to overlap

Difficulty: Easy

Learning Objective: Use the orbital overlap model to explain the bonding in simple molecules.

Section Reference: 7.1 Localized Bonds

Feedback: a)VB theory is based on atomic orbital; b) bonding requires that each atom in the molecule contribute an electron; c) Pauli exclusion principle holds; d) VB theory describes bonding in terms of valence electrons; e) this is incorrect, for example s and p orbitals may overlap to for bond

3. Valence bond theory based unhybrized orbitals fails to adequately describe the structure of methane

a) as carbon has only two unpaired electrons.

b) because p orbitals on carbon cannot overlap with s type orbitals on H to form bonds.

c) because it fails to adequately describe the shape of the methane molecule.

d) because the Aufbau principle is violated.

e) because methane is a square planar molecule.

Difficulty: Easy

Learning Objective: Use the orbital overlap model to explain the bonding in simple molecules.

Section Reference: 7.1 Localized Bonds

Feedback: a) valence bond theory allows for promotion of one of the s electrons to an empty low-lying 2p orbital allowing for 4 bonds; b) orbital overlap can be between different “kinds” of orbitals; c) correct answer; d) Aufbau principle is not violated in the molecule; e) methane is tetrahedral

4. The concept of hybridization of atomic orbitals was introduced

a) to predict shapes of molecules that cannot be described using traditional atomic orbitals.

b) to explain shapes of molecules that cannot be described using traditional atomic orbitals.

c) to account for the fact that carbon can form 4 bonds.

d) to explain π bonding.

e) to explain delocalized π bonding.

Difficulty: Medium

Learning Objective: Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Section Reference: 7.2 Hybridization of Atomic Orbitals

Feedback: Student should be clear in understanding that hybridization in valence bond theory is not predictive and that this theory has limitations (even in describing bonding and physical properties of a simple model like O₂); one needs to know structure of molecule to suggest hybridization; this is a good example of evolution of scientific theory.

5. What is the hybridization of the central atom in BeCl₂?

a) sp

b) sp²

c) sp³

d) sp⁴

e) sp³d

Difficulty: Easy

Learning Objective: Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Section Reference: 7.2 Hybridization of Atomic Orbitals

6. What is the hybridization of the central atom in ICl₂^-1?

a) sp

b) sp²

c) sp³

d) sp⁴

e) sp³d

Difficulty: Medium

Learning Objective: Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Section Reference: 7.2 Hybridization of Atomic Orbitals

7. What is the hybridization of the central atom in SO₂?

a) sp

b) sp²

c) sp³

d) sp⁴

e) sp³d

Difficulty: Medium

Learning Objective: Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Section Reference: 7.2 Hybridization of Atomic Orbitals

8. Identify the steric number, shape, hybridization about the central atom and molecular polarity of sulphuryl chloride, SO₂Cl₂.

a) 4, square planar, sp³, non-polar

b) 4, tetrahedral, sp³, polar

c) 6, tetrahedral, sp³d², polar

d) 4, square based pyramid, sp³, polar

e) 6, octahedral, sp³d², non-polar

Difficulty: Medium

Learning Objective: Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Section Reference: 7.2 Hybridization of Atomic Orbitals

Feedback: Combines concepts from Chapter 6 which are critical to using the valence bond model.

9. Identify the steric number, shape, hybridization about the central atom in XeF₂Cl₂.

a) 4, tetrahedral, sp³

b) 4, square planar, sp³

c) 6, octahedral, sp³d²

d) 6, tetrahedral, sp³d²

e) 6, square planar, sp³d²

Difficulty: Medium

Learning Objective: Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Section Reference: 7.2 Hybridization of Atomic Orbitals

Feedback: Combines concepts of Chapter 6. a) fails to account for lone pairs; b) correct shape but fails to account for lone pairs in SN and hybridization; c) hybridization is correct but shape fails to account for lone pairs; d) shape is incorrect; e) correct answer

10. What is the hybridization of Xe in the molecule XeF₄?

a) sp³

b) sp²

c) sp³d²

d) sp³d

e) sp⁴

Difficulty: Medium

Learning Objective: Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Section Reference: 7.2 Hybridization of Atomic Orbitals

11. The best description of the CO bond in acetaldehyde, CH₃-(C=O)-H is

a) a sigma bond and two pi bonds.

b) two sigma bonds.

c) two pi bonds.

d) a sigma bond and a pi bond.

e) a delta bond and a pi bond.

Difficulty: Easy

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

12. Carbon dioxide, CO₂, exists as discrete molecular compounds, whereas SiO₂ exists as a network covalent compound. Identify the hybridization and the number of sigma and pi bonds on each C and Si.

a) C(sp³_, 2,2) S(sp³,2,2)

b) C(sp_,2,2) S(sp³,0,4)

c) C(sp_, 2,2) S(sp³,4,0)

d) C(sp_, 2,2) S(sp,2,2)

e) C(sp³_,4,0) S(sp³,4,0)

Difficulty: Medium

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

Feedback: a) CO₂ has SN of 2, therefore it must be sp hybridized; b) S has only sigma bonds; c) correct answer; d) if this were correct both SiO₂ and CO₂ would be discrete linear molecules; e) CO₂ has SN of 2

13. Given that the bond energy of C-O single bond is 360 kJ/mol, a C-O double bond is 750 to 800 kJ/mol and Si-O single bond is 450 kJ/mol, what is the bond energy of a Si-O double bond?

a) 900 kJ/mol

b) significantly greater than 900 kJ/mol

c) significantly less than 900 kJ/mol

d) not possible to predict

e) slightly greater than 900 kJ/mol

Difficulty: Medium

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

Feedback: Discussion of strength of pi bonding in Si vs. C is applied to bond strength.

14. Below is the line structure for the amino acid tryptophan. Determine the type of hybridization and number of each type for the starred atoms.

a) 1 sp, 2 sp², 2 sp³

b) 1 sp, 3 sp², 1 sp³

c) 2 sp², 3 sp³

d) 3 sp², 2 sp³

e) 1 sp², 4 sp³

Difficulty: Hard

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

15. How many sigma and pi bonds are in acetic acid, CH­₃CO₂H?

a) 4 , 3 

b) 5 , 2 

c) 6 , 2 

d) 7 , 1 

e) 7 , 2 

Difficulty: Medium

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

16. Which of the following cations (O₂⁺, C₂⁺, N₂⁺; F₂⁺) would have a higher bond order than the neutral molecule?

a) all of them, removing electrons always increases bond order

b) none of them, removing electrons always decreases bond order

c) O₂⁺, C₂⁺

d) N₂⁺; F₂⁺

e) O₂⁺, F₂⁺

Difficulty: Medium

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

17. What is the bond order in Mg₂?

a) 0

b) ½

c) 1

d) 1.5

e) 2

Difficulty: Easy

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

18. What property of diatomic oxygen, O₂, is explained by MO theory but not by Lewis, valence bond on VSEPR?

a) bond strength

b) paramagnetism

c) molecular shape

d) bond length

e) ability to form multiple bonds

Difficulty: Easy

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

19. Consider HF, a heteronuclear diatomic molecule. Which of the following atomic orbitals will combine to form a σ molecular orbital?

a) H(1s) F(1s)

b) H(2p_x) F(2p_x)

c) H(1s) F(2s)

d) H(1s) F(2p_x)

e) H(1s) F(2p_x, 2p_­y, 2p_z)

Difficulty: Hard

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

Feedback: a) energy of F(1s) is too low for it to interact in a significant way with H(1s); b) energy of H (2p) is too high for it to interact in a significant way with F(2p); c) energy of F(2s) is too low for it to interact in a significant way with H(1s); d) correct answer; e) H(1s) will combine with only one of the 2p orbitals

20. Which of the following molecules/ions (O₃, NO₂^-, NCO^-, N₂O, NO₂) have the same bonding scheme as CO₂?

a) NO₂^-, NO₂

b) O₃, N₂O, NO₂

c) NCO^-, N₂O, NO₂

d) NCO^-, N₂O

e) O₃, NO₂^-, NCO^-, N₂O, NO₂

Difficulty: Easy

Learning Objective: Describe the bonding in three-atom π systems.

Section Reference: 7.5 Three-Centre π Orbitals

Feedback: Student must realize that all structures with the same number of valence electrons as CO₂ will have the same bonding scheme.

21. Which of the following best shows one of the  bonds in H₂C=C=CH₂?

a)

b)

c)

d)

e)

Difficulty: Medium

Learning Objective: Describe the bonding in three-atom π systems.

Section Reference: 7.5 Three-Centre π Orbitals

22. Which of the following does NOT have delocalized  bonding?

a) NO₂

b) SO₂

c) F₂O

d) NO₂⁺

e) O₃

Difficulty: Medium

Learning Objective: Describe the bonding in three-atom π systems.

Section Reference: 7.5 Three-Centre π Orbitals

23. In the allene (H₂C=C=CH₂) molecule,

a) all the atoms lie in one plane.

b) the  bonds are delocalized over all three C atoms.

c) the two H atoms on one C lie above and below the plane defined by the other CH₂ group.

d) the non-bonding orbitals are full.

e) delocalized bonds are formed from overlap of hybridized orbitals.

Difficulty: Medium

Learning Objective: Describe the bonding in three-atom π systems.

Section Reference: 7.5 Three-Centre π Orbitals

24. What is needed to allow a  orbital to be delocalized?

a) 4 or more atoms

b) more than 2 “p” orbitals with appropriate geometry

c) 6 electrons

d) an s and a p orbital in a molecular bond

d) carbon–carbon bonds

Difficulty: Medium

Learning Objective: Describe the bonding in three-atom π systems.

Section Reference: 7.5 Three-Centre π Orbitals

25. P–O double bonds are stronger than P–S double bonds because

a) the O atom had greater electronegativity.

b) there is less polar character to the P-S bond.

c) the O atom has a smaller radius than the S atom.

d) the P atom is less electronegative than the O atom.

e) S has accessible d orbitals.

Difficulty: Medium

Learning Objective: Describe the bonding in extended π systems.

Section Reference: 7.6 Extended π Systems

26. Which of the following sulphur species has the greatest delocalization as judged by the number of resonance structures?

a) H₃SO₄⁺

b) H₂SO₄

c) HSO₄^-

d) SO₄^2-

e) SO₃

Difficulty: Hard

Learning Objective: Describe the bonding in extended π systems.

Section Reference: 7.6 Extended π Systems

27. Consider the perchlorate anion, ClO₄^-:

a) the steric number for this anion is 4, the best description includes resonance structures having two double bonds and two single bonds, electrons are delocalized throughout the ion.

b) the steric number for this anion is 4, the best description includes resonance structures having three double bonds and one single bonds, electrons are delocalized throughout the ion.

c) chlorine is not capable of forming multiple bonds.

d) the steric number for this anion is 4, the best description includes resonance structures having three double bonds and one single bonds, electrons are not delocalized since the structure is not conjugated.

e) the steric number for this anion is 7, the best description includes resonance structures having three double bonds and one single bonds, electrons are delocalized throughout the ion.

Difficulty: Medium

Learning Objective: Describe the bonding in extended π systems.

Section Reference: 7.6 Extended π Systems

Feedback: Requires student to draw correct Lewis structures including resonance and assess the resulting pi system.

28. Which of the following compounds would be expected to absorb the longest wavelength light?

a)

b)

c)

d)

e)

Difficulty: Medium

Learning Objective: Describe the bonding in extended π systems.

Section Reference: 7.6 Extended π Systems

29. Which of the following is true of a conjugate  system?

a) The molecule is more reactive.

b) The molecule is more volatile.

c) The molecule is less reactive.

d) The molecule has the same reactivity.

e) The molecule reacts violently.

Difficulty: Easy

Learning Objective: Describe the bonding in extended π systems.

Section Reference: 7.6 Extended π Systems

30. Which of the following compounds will absorb the most visible light, i.e,. be the most deeply coloured?

a)

b)

c)

d)

e)

Difficulty: Medium

Learning Objective: Describe the bonding in extended π systems.

Section Reference: 7.6 Extended π Systems

31. How many delocalized  electrons are in the following molecule?

a) 6

b) 8

c) 10

d) 12

e) 14

Difficulty: Easy

Learning Objective: Describe the bonding in extended π systems.

Section Reference: 7.6 Extended π Systems

32. Which of the following elements would be added to germanium to produce an n-type semiconductor?

a) gallium

b) silicon

c) aluminum

d) arsenic

e) tin

Difficulty: Easy

Learning Objective: Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Section Reference: 7.7 Band Theory of Solids

33. In n-type seminconductor the

a) donor level of the dopant lies close in energy to the conduction band.

b) donor level of the dopant lies close in energy to the valence band.

c) acceptor level of the dopant lies close in energy to the conduction band.

d) the acceptor level of the dopant lies close in energy to the valence band.

Difficulty: Easy

Learning Objective: Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Section Reference: 7.7 Band Theory of Solids

34. What will happen to a semiconductor made of GaP, if some of the P is replaced with As?

a) The band gap remains the same.

b) The band gap grows larger.

c) The band gap becomes smaller.

d) This cannot happen because As is not isoelectronic with P.

e) It will emit light of shorter wavelength.

Difficulty: Medium

Learning Objective: Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Section Reference: 7.7 Band Theory of Solids

35. What would you dope a GaP semiconductor with to decrease the wavelength of light emitted?

a) In

b) As

c) N

d) Ge

e) The band gap cannot be changed.

Difficulty: Medium

Learning Objective: Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Section Reference: 7.7 Band Theory of Solids

ESSAY QUESTIONS

36. Draw a picture showing the bonding between a hydrogen atom and a chlorine atom telling what orbitals from each are making the bond.

Difficulty: Easy

Learning Objective: Use the orbital overlap model to explain the bonding in simple molecules.

Section Reference: 7.1 Localized Bonds

37. Describe the bonding between iodine atoms in molecular iodine, I₂. Make sure to include a drawing to symbolize the overlap of orbitals.

Difficulty: Easy

Learning Objective: Use the orbital overlap model to explain the bonding in simple molecules.

Section Reference: 7.1 Localized Bonds

38. NH₃ has bond angles of 107.3˚. Describe the bonding in NH₃ only using unhybridized orbitals. Discuss why this type of bonding model is faulty.

Difficulty: Medium

Learning Objective: Use the orbital overlap model to explain the bonding in simple molecules.

Section Reference: 7.1 Localized Bonds

39. H₂O has a bond angle of 104.5˚. Describe the bonding in H₂O only using unhybridized orbitals. Discuss why this type of bonding model is faulty.

Difficulty: Medium

Learning Objective: Use the orbital overlap model to explain the bonding in simple molecules.

Section Reference: 7.1 Localized Bonds

40. Methylene chloride, CH₂Cl₂, is a common industrial solvent. Sketch an orbital overlap picture of the bonds and describe the bonding present.

Difficulty: Hard

Learning Objective: Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Section Reference: 7.2 Hybridization of Atomic Orbitals

41. Determine the hybridization of the central atoms of H₂CCHCH₃.

Difficulty: Easy

Learning Objective: Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Section Reference: 7.2 Hybridization of Atomic Orbitals

42. Determine the hybridization of the central atoms in H₂NCHCH₂.

Difficulty: Hard

Learning Objective: Assign the correct hybrid orbitals used by each inner atom in a molecule and the molecular geometry that results.

Section Reference: 7.2 Hybridization of Atomic Orbitals

43. Draw the Lewis structure of the sulphite ion, SO₃^-2. Determine what the hybridization of the central atom is.

Difficulty: Medium

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

44. Draw two sketches describing the bonding in propene, C₃H₆. The first should show any sigma bonding framework and the second sketch should illustrate the π bonds that are formed, if any. Write a brief description of the bonding, naming the orbitals involved.

Difficulty: Hard

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

45. For the molecule ethene, C₂H₄, sketch a picture that shows the location of the  bond as oriented with the plane of the hydrogen atoms.

Difficulty: Easy

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

46. Explain why carbon dioxide includes both pi and sigma bonds whereas SiO₂ is formed from a network of sigma bonds.

Difficulty: Medium

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

47. Draw a picture illustrating the atomic orbitals participating in the bonding in formaldehyde, H₂C=O. List the orbitals used in bond formation.

Difficulty: Medium

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

48. The Lewis structure of vinyl chloride is shown below; Draw orbital pictures of (a) the sigma bonding framework and (b) the  bond.

Difficulty: Hard

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

49. What orbitals overlap to make the bond(s) between C and O in acetone, CH₃COCH₃?

Difficulty: Easy

Learning Objective: Describe the σ and π bonding systems in multiple bonds.

Section Reference: 7.3 Multiple Bonds

50. In which chemical species would you expect the strongest bond: C₂⁺; N₂⁺; O₂⁺? Write the electronic molecular orbital configuration and calculate the bond order for that species.

Difficulty: Medium

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

51. Write the electron configuration and predict the bond order and number of unpaired electrons O₂^-.

Difficulty: Medium

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

52. Write the electron configuration and predict the bond order and number of unpaired electrons B-C.

Difficulty: Medium

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

53. Determine the bond order for F₂. How many electrons are in antibonding orbitals?

Difficulty: Easy

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

54. Determine the bond order for O₂⁺ and tell how many electrons are in antibonding orbitals.

Difficulty: Easy

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

55. Draw orbital pictures of the highest energy molecular orbital occupied in fluorine F₂ and the atomic orbitals that form it.

Difficulty: Medium

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

56. Predict the bond order and number of unpaired electrons in NO+.

Difficulty: Medium

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

57. Predict using MO theory if NO^- will be paramagnetic or diamagnetic.

Difficulty: Easy

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

58. What kind of bonds are obtained by the overlap of the dxy and the p orbitals shown below?

Difficulty: Easy

Learning Objective: Use molecular orbital theory to calculate a bond order, predict magnetic properties of a molecule, and explain trends in bond length and energy.

Section Reference: 7.4 Molecular Orbital Theory: Diatomic Molecules

59. In the text’s problems, you were asked to find the Lewis structure of HNCO. Compare that Lewis structure with that of the isomer, NCOH and suggest why the HNCO isomer may be more stable.

Difficulty: Hard

Learning Objective: Describe the bonding in three-atom π systems.

Section Reference: 7.5 Three-Centre π Orbitals

60. Draw the bonding and antibonding molecular orbitals for NO₂^-.

Difficulty: Hard

Learning Objective: Describe the bonding in three-atom π systems.

Section Reference: 7.5 Three-Centre π Orbitals

61. Explain the difference between an isolated orbital and a delocalized  orbital.

Difficulty: Medium

Learning Objective: Describe the bonding in three-atom π systems.

Section Reference: 7.5 Three-Centre π Orbitals

62. What are non-bonding molecular orbitals of ozone, O₃, and where do they originate?

Difficulty: Medium

Learning Objective: Describe the bonding in three-atom π systems.

Section Reference: 7.5 Three-Centre π Orbitals

63. Determine the bonding for the nitrite ion, NO₂^-1. Determine what hybridization the central atom has and how many delocalized  orbitals there are.

Difficulty: Medium

Learning Objective: Describe the bonding in three-atom π systems.

Section Reference: 7.5 Three-Centre π Orbitals

64. Draw a combination of a 3d orbital on phosphorus and a 2p orbital on O that allows  bonding.

Difficulty: Hard

Learning Objective: Describe the bonding in extended π systems.

Section Reference: 7.6 Extended π Systems

65. Describe the bonding in the polyatomic anion, nitrate (NO₃^-1) by drawing the resonance structures and the bonding  molecular orbital.

Difficulty: Hard

Learning Objective: Describe the bonding in extended π systems.

Section Reference: 7.6 Extended π Systems

66. Draw the band diagram that is appropriate for a ZnS semiconductor that has been doped with copper.

Difficulty: Medium

Learning Objective: Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Section Reference: 7.7 Band Theory of Solids

67. Explain how band theory can explain the electrical conductivity observed in Mg metal given that atomic Mg has valence electrons fully occupying the 2p atomic orbitals.

Difficulty: Hard

Learning Objective: Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Section Reference: 7.7 Band Theory of Solids

68. A new material is colourless, but is not an insulator. What can you deduce about the energy of the bandgap? Give an estimate of the bandgap in kJ.

Difficulty: Easy

Learning Objective: Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Section Reference: 7.7 Band Theory of Solids

69. Draw a representation of a band gap for graphite and germanium and discuss any differences between them.

Difficulty: Medium

Learning Objective: Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Section Reference: 7.7 Band Theory of Solids

70. Semiconductors are used for solar energy conversion in devices called photovoltaic cells. For high efficiency, the cell should be able to utilize as much light as possible. What band gap (in kJ/mole) would be needed to utilize light of 550 nm (about the middle of the visible spectrum) and shorter wavelengths?

Difficulty: Easy

Learning Objective: Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Section Reference: 7.7 Band Theory of Solids

71. Describe what happens to an electron as it travels through an n🡪p-type junction of a semiconductor.

Difficulty: Hard

Learning Objective: Explain such properties as electrical conductivity and the colour of metals, non-metals, and metalloids in terms of band theory.

Section Reference: 7.7 Band Theory of Solids

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