1. Molecular Structure & Bonding Theories Chapter 10

2. Why study molecular shapes? The three-dimensional shapes of molecules strongly influence their physical & chemical properties Case Study: The stability of SF 6 and its use to help thwart terrorism via testing how toxic gases might spread through subway systems

3. How do we determine the shape of a molecule based on its chemical formula? Several different models exist to help predict molecular shape, each with their own advantages & disadvantages…. ▫Valence Shell Electron Pair Repulsion ▫Valence Bond Theory ▫Molecular Orbital Theory

4. Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world. Models can be physical as with this DNA model Models can be mathematical Models can be theoretical or philosophical A few notes about models…..

5. Localized Electron Model Lewis structures are an application of the “Localized Electron Model” L.E.M. says: Electron pairs can be thought of as “belonging” to pairs of atoms when bonding Resonance points out a weakness in the Localized Electron Model. A few notes about models…..

6. Fundamental Properties of Models  A model does not equal reality.  Models are oversimplifications, and are therefore often wrong.  Models become more complicated as they age.  We must understand the underlying assumptions in a model so that we don’t misuse it.

7. 10.1 VSEPR Model Objectives To predict the shapes of molecules from the valence shell electron pair repulsion model To use the steric number to assign the electron- pair arrangement of each central atom To distinguish between electron-pair arrangement and molecular shape

8. 10.1 VSEPR Model The valence shell electron pair repulsion (VSEPR) model predicts the shapes of molecules from their Lewis structures The main premise of the model is that electron pairs about an atom repel each other VSEPR Rule 1 ▫A molecule has a shape that minimizes electrostatic repulsions between valence shell electron pairs ▫Minimum repulsion results when the electron pairs are as far apart as possible

9. 10.1 VSEPR Model The VSEPR model predicts the shape around each central atom in a molecule To simplify the process of determining shape, we define the steric number as the number of lone pairs on the central atom plus the number of atoms bonded to the central atom

10. 10.1 VSEPR Model Steric Number Calculations Ex: H 2 O H O H Note that there are two hydrogens bound to the central oxygen and there are also two sets of lone pair electrons. Therefore, the steric number is four (2 bonded atoms + 2 lone pairs)

11. 10.1 VSEPR Model Steric Number Calculations Ex: NH 3 H N H H Note that there are three hydrogen atoms bound to the central nitrogen and there are is one set of lone pair electrons. Therefore, the steric number is four (3 bonded atoms + 1 lone pair)

12. 10.1 VSEPR Model Steric Number Calculations Ex: XeF 4 Xe F F Note that there are four fluorine atoms bound to the central xenon and there are also two sets of lone pair electrons. Therefore, the steric number is six (4 bonded atoms + 2 lone pairs) F

13. 10.1 VSEPR Model Once calculated, the steric number is used to determine the bonded-angle lone-pair arrangement (BALPA) The bonded-angle lone-pair arrangement (BALPA) is the shape that maximizes the distance between regions of electron density about a central atom The arrangement names represent the geometric solids formed when imaginary lines are drawn to connect the bonded atoms and/or electron pairs ▫Ex: tetrahedral, trigonal pyramidal, octahedral

14. VSEPR – Valence Shell Electron Pair Repulsion X + E Overall Structure Forms 2 LinearAX 2 3 Trigonal PlanarAX 3, AX 2 E 4 TetrahedralAX 4, AX 3 E, AX 2 E 2 5 Trigonal bipyramidal AX 5, AX 4 E, AX 3 E 2, AX 2 E 3 6 OctahedralAX 6, AX 5 E, AX 4 E 2 A = central atom X = atoms bonded to A E = nonbonding electron pairs on A

15. Geometric Arrangements

17. VSEPR: Linear AX 2 CO 2

18. VSEPR: Trigonal Planar AX 3 AX 2 E BF 3 SnCl 2

19. VSEPR: Tetrahedral AX 4 AX 3 E AX 2 E 2 CCl 4 PCl 3 Cl 2 O

20. VSEPR: Trigonal Bi-pyramidal AX 5 AX 4 E AX 3 E 2 AX 2 E 3 PCl 5 SF 4 ClF 3 I3-I3-I3-I3-

21. VSEPR: Octahedral AX 6 AX 5 E AX 4 E 2 SF 6 ICl 4 - BrF 5

22. 10.1 VSEPR Model When central atoms have lone pairs of electrons, the molecular shape is not the same as the BALPA The molecular shape = the actual geometric arrangement of the atoms. Lone pairs influence the molecular shape, but are not part of it Ex: H 2 O Steric Number is 4 BALPA is tetrahedral (Predicted bond angles = 109.5°) Molecular Shape is BENT or V-shaped (Actual bond angles = 104.5°)

23. 10.1 VSEPR Model When central atoms have lone pairs of electrons, the molecular shape is not the same as the BALPA Lone pairs of electrons repel other electron pairs more strongly and need more space than bonding pairs VSEPR Rule 2 ▫Forces between electron pairs vary as follows: Lone-pair-lone-pair repulsion > lone-pair-bonding-pair repulsion > bonding pair-bonding-pair repulsion

24. P. 378 P. 378 Steric #, BALPA, and Molecular Shape

25. p. 378 Steric #, BALPA, and Molecular Shape Steric # Bonded # of BALPAShape Bond Ex. Atoms Lone Pairs Angles

26. 10.1 VSEPR Model When more than one central atom is present, the shape around each central atom can be determined separately using the rules already described The VSEPR model does not, however, predict how the geometry of one central atom is oriented with respect to other central atoms in the molecule. This is a weakness of the VSEPR model Ex: ethylene C 2 H 4

27. 10.2 Polarity of Molecules Objectives To predict the polarity of a molecule from bond polarities and molecular shape

28. 10.2 Polarity of Molecules Polar molecules contain an unequal distribution of charge Generally: ▫polar molecules interact with other polar molecules, and ▫nonpolar molecules react with other nonpolar molecules ▫Ex: sugar dissolves in water/oil does not

29. 10.2 Polarity of Molecules The degree of polarity can be measured by a dipole moment ▫A dipole moment equals the magnitude of the separated charges times the distance between them ▫Differences in electronegativity between two atoms are used to predict the polarity of each bond ▫Remember dipoles are typically represented by arrows pointing toward the more electronegative atom. Longer arrows represent more polar bonds H F H Br

30. Polarity of Molecules The bond dipoles in CO 2 cancel because the linear shape orients the equal magnitude bond dipoles in exactly opposite directions.

31. Polarity of Molecules The bond dipoles do not cancel in COSe; they are oriented in the same direction and are of unequal length. They do not cancel in OF 2 because the V-shape of the molecule does not orient them in opposite directions.

32. Polarity of Molecules The bond dipoles in BCl 3 and CCl 4 cancel because of the regular shape and equal magnitude.

33. Polarity of Molecules The bond dipoles in BCl 2 F and CHCl 3 do not cancel because they are not of the same magnitude.

34. 10.2 Polarity of Molecules Based on the given examples: ▫Note that a molecule with polar bonds can be nonpolar if the geometry causes the bond polarities to sum to zero. ▫Note that molecules are nonpolar when there are no lone pairs on the central atom and all of the atoms bonded to the central atom are identical ▫Note that molecules with lone pairs of electrons on a central atom are generally polar, but there are exceptions

35. 10.2 Polarity of Molecules Some problem solving….. Predict which of the following molecules are polar and which are non-polar. 1)HCN 2)PF3 3)XeF4 1) polar 2) polar 3) nonpolar

36. 10.2 Polarity of Molecules More classwork problem solving….. 1)An experiment shows that the molecule PF 2 Cl 3 is nonpolar. What is the molecular shape, and what is the arrangement of the atoms in that shape? 2)If the hypothetical molecule SF 2 Cl 2 Br 2 is nonpolar, what is its shape, and what is the arrangement of the atoms?

37. 10.3 Valence Bond Theory Objectives To identify the orbitals used to form the bonds in any specific molecule To assign the hybrid orbitals used by a central atom for bonds and lone pairs

38. 10.3 Valence Bond Theory Valence Bond Theory describes covalent bonds as being formed by atoms sharing valence electrons in overlapping valence orbitals Ex: The bond in H 2 would be formed by the overlap of an unpaired electron from the 1s orbital of each hydrogen atom

39. 10.3 Valence Bond Theory The bond in HF would be formed by the overlap of the unpaired electron from the 1s orbital of a hydrogen atom and an unpaired electron from a 2p orbital of the fluorine atom

40. 10.3 Valence Bond Theory Unfortunately, this simplified version of valence bond theory does not account for the bonding of atoms in numerous other molecules such as ▫NH 3 (predicts the wrong bond angles: 90° vs the correct 107°), or ▫CH 4 (how do you get four equivalent bonds) Therefore, the hybridization of atomic orbitals was proposed Hybrid orbitals are orbitals obtained by mixing two or more atomic orbitals of the same central atom

41. We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds. Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it. What Proof Exists for Hybridization? Lets look at a molecule of methane, CH 4.

42. What is the expected orbital notation of carbon in its ground state? (Hint: How many unpaired electrons does this carbon atom have available for bonding?) Can you see a problem with this? Carbon ground state configuration

43. You should conclude that carbon only has TWO electrons available for bonding. That is not enough! How does carbon overcome this problem so that it may form four bonds? Carbon’s Bonding Problem

44. The first thought that chemists had was that carbon promotes one of its 2s electrons… …to the empty 2p orbital. Carbon’s Empty Orbital

45. However, they quickly recognized a problem with such an arrangement… Three of the carbon-hydrogen bonds would involve an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.

46. This would mean that three of the bonds in a methane molecule would be identical, because they would involve electron pairs of equal energy. But what about the fourth bond…?

47. The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron. Such a bond would have slightly less energy than the other bonds in a methane molecule.

48. This bond would be slightly different in character than the other three bonds in methane. This difference would be measurable to a chemist by determining the bond length and bond energy. But is this what they observe?

49. The simple answer is, “No”. Chemists have proposed an explanation – they call it Hybridization. Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy. Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane.

50. Hybridization - The Blending of Orbitals Poodle + +Cocker Spaniel = = = = + +s orbitalp orbital Cockapoo sp orbital

51. In the case of methane, they call the hybridization sp 3, meaning that an s orbital is combined with three p orbitals to create four equal hybrid orbitals. These new orbitals have slightly MORE energy than the 2s orbital… … and slightly LESS energy than the 2p orbitals.

52. Here is another way to look at the sp 3 hybridization and energy profile… sp 3 Hybrid Orbitals

53. While sp 3 is the hybridization observed in methane, there are other types of hybridization that atoms undergo. These include sp hybridization, in which one s orbital combines with a single p orbital. This produces two hybrid orbitals, while leaving two normal p orbitals sp Hybrid Orbitals

54. Another hybrid is the sp 2, which combines two orbitals from a p sublevel with one orbital from an s sublevel. One p orbital remains unchanged. sp 2 Hybrid Orbitals

55. Hybridization Involving “d” Orbitals Beginning with elements in the third row, “d” orbitals may also hybridize dsp 3 = five hybrid orbitals of equal energy d 2 sp 3 = six hybrid orbitals of equal energy

56. Another representation of hybridization…. Fig. 10-12, p. 388

57. Another representation of hybridization…. Fig. 10-17, p. 391

58. Hybridization and Molecular Geometry Forms Overall Structure Hybridization of “A” AX 2 Linearsp AX 3, AX 2 ETrigonal Planarsp 2 AX 4, AX 3 E, AX 2 E 2 Tetrahedralsp 3 AX 5, AX 4 E, AX 3 E 2, AX 2 E 3 Trigonal bipyramidal dsp 3 AX 6, AX 5 E, AX 4 E 2 Octahedrald 2 sp 3 A = central atom X = atoms bonded to A E = nonbonding electron pairs on A

59. 10.3 Valence Bond Theory cont’d Note that the steric number and number of hybrid orbitals are equivalent Ex: steric # = 5 (AX 5, AX 4 E, AX 3 E 2, AX 2 E 3 ) is represented by dsp 3 which is 1 d + 1 s + 3 p = 5 dsp 3 hybrids orbitals Ex: steric # = 3 (AX 3, AX 2 E) is represented by sp 2 which is 1 s + 2 p = 3 sp 2 hybrids orbitals

60. Table 10-1, p. 392

61. 10.4 Multiple Bonds Objectives To define and identify  (sigma) and  (pi) bonds To identify the orbitals used to form  and  bonds To describe the bonding in molecules that contain multiple bonds To describe cis and trans isomers

62. Sigma and Pi Bonds Sigma (  ) bonds exist in the region directly between two bonded atoms. Pi (  ) bonds exist in the region above and below a line drawn between two bonded atoms. Single bond1 sigma bond Double Bond1 sigma, 1 pi bond Triple Bond1 sigma, 2 pi bonds

63. Sigma and Pi Bonds Single Bonds Ethane 1  bond

64. Sigma and Pi Bonds: Double bonds Ethene 1  bond 1  bond

65. Sigma and Pi Bonds Triple Bonds Ethyne 1  bond 1  bond

66. The De-Localized Electron Model Pi bonds (  ) contribute to the delocalized model of electrons in bonding, and help explain resonance Electron density from  bonds can be distributed symmetrically all around the ring, above and below the plane.

67. Multiple Bonds Another view of electron delocalization due to pi bonding…..

68. 10.4 Multiple Bonds Isomers = compounds with the same molecular formula but different structures When double bonds are formed between two carbon atoms, the potential for cis-trans isomerism arises due to lack of rotation around double bonds ▫cis-trans isomerism is observed only if each carbon atom has two different substituents attached

70. 10.4 Multiple Bonds cistrans

71. 10.5 Molecular Orbitals: Homonuclear Diatomic Molecules Objectives To write the molecular orbital diagrams for homonuclear diatomic molecules and ions of the first and second period elements To determine the bond order and the number of unpaired electrons in diatomic species from the molecular orbital diagrams

72. 10.5 Molecular Orbitals: Homonuclear Diatomic Molecules Remember what was said about the limitations of some models……. Valence Bond Theory does not explain the behavior of O 2 ▫O 2 is paramagnetic yet valence bond theory does not predict any unpaired electrons O=O ▫Molecular Orbital Theory can explain this behavior :: ::

73. Fig. 10-34, p. 402 NitrogenOxygen Behavior of liquid nitrogen and oxygen in a magnetic field

74. 10.5 Molecular Orbitals: Homonuclear Diatomic Molecules Molecular Orbital Theory is a model that combines atomic orbitals to form new orbitals that are shared over the entire molecule rather than between two atoms A molecular orbital is a wave function of an electron in a molecule ▫The number of molecular orbitals formed must equal the number of atomic orbitals used to make them

75. 10.5 Molecular Orbitals: Homonuclear Diatomic Molecules These wave functions can be added to generate bonding orbitals (concentrates electron density between atoms) or they can be subtracted to generate anti-bonding orbitals (reduces the electron density between atoms) Let’s use H 2 and He 2 as examples

76. Bonding & Antibonding Orbitals in H 2

77. Fig. 10-36, p. 403

78. 10.5 Molecular Orbital Theory cont’d Bonding molecular orbitals are more stable and antibonding molecular orbitals are less stable than the atomic orbital that are combined.

79. 10.5 Molecular Orbitals: Homonuclear Diatomic Molecules Molecular orbital theory allows us to calculate bond order (measure of stability) and the # of unpaired electrons (paramagnetism) Molecular orbital theory defines bond order as follows: Bond order = ½ [number of electrons in bonding orbitals – number of electrons in anti-bonding orbitals]

80. 10.5 Molecular Orbitals: Homonuclear Diatomic Molecules Let’s analyze O 2 via MOT…. See p. 407 in 3 rd edition for a summary of 2 nd period homonuclear diatomics

81. 10.5 Molecular Orbitals: Homonuclear Diatomic Molecules Notice the two unpaired electrons that contribute to oxygen’s paramagnetism Note that the bond order calculated for diatomic oxygen is two (½*[8-4]), identical to what the Lewis structure indicates

82. 10.5 Molecular Orbitals: Homonuclear Diatomic Molecules

83. 10.6 Heteronuclear Diatomic Molecules and Delocalized Molecular Orbitals Objectives To construct the molecular orbital diagram for heteronuclear diatomic molecules To describe the formation of delocalized molecular orbitals

84. 10.6 Heteronuclear Diatomic Molecules and Delocalized Molecular Orbitals Heteronuclear diatomic molecules contain one atom of each of two different elements The molecular orbital diagrams are close to those for homonuclear diatomic molecules when the valence energies for each atom are fairly close in energy. Let’s try a diagram for NO

85. Molecular Orbital Diagram: NO

86. The electron configuration of NO is (  2s ) 2 (  * 2s ) 2 (  2p ) 4 (  2p ) 2 (   2p ) 1. Bond order for NO is 2.5. NO has one unpaired electron. The two Lewis structures for this molecule, shown below, predict the unpaired electron but not the bond order. Molecular Orbitals: NO

87. 10.6 Heteronuclear Diatomic Molecules and Delocalized Molecular Orbitals Let’s try diagrams for: HHe NO - CN -

88. 10.6 Heteronuclear Diatomic Molecules and Delocalized Molecular Orbitals Sample problems cont’d: What is the electron configuration, the bond order and the number of unpaired electrons in the molecule BC? Answer: The electron configuration is (s 2s ) 2 (s* 2s ) 2 (p 2p ) 3, the bond order is 1.5, and there is one unpaired electron.

89. Note that the molecular orbital diagram is not “one size fits all”. The structure of the diagram is dependent on the energies associated with the atoms that are bonding

90. 10.6 Heteronuclear Diatomic Molecules and Delocalized Molecular Orbitals Delocalized molecular orbitals are: ▫orbitals in which an electron in a molecule is spread over more than two atoms ▫ often useful in describing the behavior of resonance forms ▫Ex: benzene, ozone (O 3 )