1. Molecular Geometries and Bonding Theories

2. Molecular Shapes The shape of a molecule plays an important role in its reactivity. The shape of a molecule is determined by its bond angles and bond lengths.

3. Molecular Shapes By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.

4. Repulsion applies to both bonding and non- bonding pairs of electrons Double and triple bonds orient themselves together so the behave as a single unit

5. What Determines the Shape of a Molecule? Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

6. Electron Domains We can refer to the electron pairs as electron domains. In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain. This molecule has four electron domains.

7. Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

8. Molecular Geometry The s hape formed by a molecule is created by the bond angles due to repulsions of in the electron arrangement. Bond angles a determined by the interactions of bonded electrons (X) and lone pair electrons (E). Lone pair – non-bonded pair repulsions are greater than bonded pair repulsions VSEPR notation: NH 3 AX 3 E 1 Central atom

9. Repulsion decreases in the following order: Lone pair-lone pair> lone pair-bonding pair> bonding pair-bonding pair This results in molecules with lone pairs having distortions in their structure which reduces the angle between bonding pairs

10. AX 2 AX 3 AX 4 General Formula

12. AX 5 General Formula AX 6

15. AX 2 AX 3 AX 2 E 117 0

16. AX 4 AX 3 E AX 2 E 2 105 0 107 0 Or V-shaped

17. AX 5 AX 4 E AX 3 E 2 AX 2 E 3 Or unsymmetrical tetrahedron 117 0

18. AX 6 AX 5 E AX 4 E 2

19. Electron-Domain Geometries All one must do is count the number of electron domains in the Lewis structure. The geometry will be that which corresponds to that number of electron domains.

20. Rules for Predicting Molecular Geometry 1. Sketch the Lewis structure of the molecule or ion 2. Count the electron pairs and arrange them in the way that minimizes electron-pair repulsion. 3. Determine the position of the atoms from the way the electron pairs are shared. 4. Determine the name of the molecular structure from the position of the atoms. 5. Double or triple bonds are counted as one bonding pair when predicting geometry.

21. Molecular Geometries The electron-domain geometry is often not the shape of the molecule, however. The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

22. Molecular Geometries Within each electron domain, then, there might be more than one molecular geometry.

23. Linear Electron Domain In this domain, there is only one molecular geometry: linear. NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

24. Trigonal Planar Electron Domain There are two molecular geometries:  Trigonal planar, if all the electron domains are bonding  Bent, if one of the domains is a nonbonding pair.

25. Nonbonding Pairs and Bond Angle Nonbonding pairs are physically larger than bonding pairs. Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.

26. Multiple Bonds and Bond Angles Double and triple bonds place greater electron density on one side of the central atom than do single bonds. Therefore, they also affect bond angles.

27. In which of the following molecules is the H  X  H bond angle smallest (X is non-H atom): CH 4, H 2 O, and NH 3 ? 1.CH 4 2.H 2 O 3.NH 3 4.They are all equal

28. 1.CH 4 2.H 2 O 3.NH 3 4.They are all equal Correct Answer: Lone pairs exert greater repulsive forces on adjacent electron domains. Bond angles are compressed with increasing numbers of lone pairs.

30. Yes, resonance equalizes repulsions between bonds, making all angles 120°.

31. Tetrahedral Electron Domain There are three molecular geometries:  Tetrahedral, if all are bonding pairs  Trigonal pyramidal if one is a nonbonding pair  Bent if there are two nonbonding pairs

32. Trigonal Bipyramidal Electron Domain There are two distinct positions in this geometry:  Axial  Equatorial

33. Trigonal Bipyramidal Electron Domain Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry.

34. Trigonal Bipyramidal Electron Domain There are four distinct molecular geometries in this domain:  Trigonal bipyramidal  Seesaw  T-shaped  Linear

35. Octahedral Electron Domain All positions are equivalent in the octahedral domain. There are three molecular geometries:  Octahedral  Square pyramidal  Square planar

36. Using VSEPR, predict the molecular geometry of NH 3. 1.Linear 2.Trigonal planar 3.Trigonal pyramidal 4.Tetrahedral

37. Correct Answer: 1.Linear 2.Trigonal planar 3.Trigonal pyramidal 4.Tetrahedral As indicated in the illustration, the molecular geometry of NH 3 is trigonal pyramidal.

38. Using VSEPR, predict the molecular geometry of SF 4. 1.T-shaped 2.Seesaw-shaped 3.Trigonal bipyramidal 4.Linear

39. Correct Answer: 1.T-shaped 2.Seesaw-shaped 3.Trigonal bipyramidal 4.Linear The molecular geometry of SF 4 is seesaw.

40. Using VSEPR, predict the molecular geometry of O 3. 1.Bent 2.Trigonal planar 3.Trigonal pyramidal 4.Linear

41. Correct Answer: 1.Bent 2.Trigonal planar 3.Trigonal pyramidal 4.Linear The molecular geometry of O 3 is bent.

42. Larger Molecules In larger molecules, it makes more sense to talk about the geometry about a particular atom rather than the geometry of the molecule as a whole.

43. Larger Molecules This approach makes sense, especially because larger molecules tend to react at a particular site in the molecule.

44. Polarity But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.

45. Polarity By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule.

46. Polar bonds result from unequal sharing of electrons Polar bonds result from unequal sharing of electrons This is a result of differences in electronegativities that cause electron pairs to be unequally shared This is a result of differences in electronegativities that cause electron pairs to be unequally shared The greater pulling power causes the electron pairs thus the bond to be asymmetrical with respect to electron distribution. Called a polar bond. The greater pulling power causes the electron pairs thus the bond to be asymmetrical with respect to electron distribution. Called a polar bond.

47. Polarity Molecules with polar bonds are not always polar Polarity of a bond depends on the charge separation between unequally shared electrons in a covalent bond Polarity of a molecule depends on: Polar bonds it contains The shape of the moleculeThe orientation of the polar bonds with respect to each other  The shape of the molecule

48. A polar molecule is a molecule with regions of partial negative (  –) and partial positive (  +) charge. delta The term dipole is used

49. HCl  –  +

50. The Dipole A dipole arises when two electrical charges of equal magnitude but opposite sign are separated by distance.

51. The dipole Symmetricalasymmetrical net dipole

52. If bonds are of equal polarity and are arranged symmetrically with respect to each other, their charge separation will oppose each other and cancel each other out and the molecule will be non-polar If bonds contain bonds of different polarity or its bonds are not symmetrically arranged the molecule will be polar

53. The sum of these vectors will give us the dipole for the molecule The sum of these vectors will give us the dipole for the molecule For a polyatomic molecule we treat the dipoles as 3D vectors

54. Polarity

55. CO 2 : Polar or nonpolar? 1.Polar 2.Nonpolar

56. Correct Answer: The bond dipoles in CO 2 cancel one another such that the molecule will exhibit no net overall dipole. Thus, it is a nonpolar molecule. 1.Polar 2.Nonpolar

58. Yes

59. SO 2 : Polar or nonpolar? 1.Polar 2.Nonpolar

60. Correct Answer: 1.Polar 2.Nonpolar Clearly from the bond dipoles the molecule SO 2 will exhibit a net overall dipole, thus being a polar molecule.

61. Overlap and Bonding We think of covalent bonds forming through the sharing of electrons by adjacent atoms. In such an approach this can only occur when orbitals on the two atoms overlap.

62. Overlap and Bonding Increased overlap brings the electrons and nuclei closer together while simultaneously decreasing electron- electron repulsion. However, if atoms get too close, the internuclear repulsion greatly raises the energy.

65. Yes Cl 2 forms from the overlap of two 3p orbitals while F 2 forms from the overlap of two 2p orbitals which are smaller than 3p orbitals.

66. Hybrid Orbitals But it’s hard to imagine tetrahedral, trigonal bipyramidal, and other geometries arising from the atomic orbitals we recognize.

67. Hybrid Orbitals Consider beryllium:  In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.

68. Hybrid Orbitals But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds.

69. Hybrid Orbitals Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.  These sp hybrid orbitals have two lobes like a p orbital.  One of the lobes is larger and more rounded as is the s orbital.

70. Hybrid Orbitals These two degenerate orbitals would align themselves 180  from each other. This is consistent with the observed geometry of beryllium compounds: linear.

71. Hybrid Orbitals With hybrid orbitals the orbital diagram for beryllium would look like this. The sp orbitals are higher in energy than the 1s orbital but lower than the 2p.

72. Hybrid Orbitals Using a similar model for boron leads to…

73. Hybrid Orbitals …three degenerate sp 2 orbitals.

74. Hybrid Orbitals With carbon we get…

75. Hybrid Orbitals …four degenerate sp 3 orbitals.

76. Molecular Geometries and Bonding

77. Molecular Geometries and Bonding The unhybridized p-orbital is perpendicular to the plane of the sp 2 orbitals.

78. Hybrid Orbitals For geometries involving expanded octets on the central atom, we must use d orbitals in our hybrids.

79. Hybrid Orbitals This leads to five degenerate sp 3 d orbitals… …or six degenerate sp 3 d 2 orbitals.

81. Hybrid Orbitals Once you know the electron-domain geometry, you know the hybridization state of the atom.

82. All you need to know about orbital hybridization Valence shell atomic orbitals (including inner valence d orbitals) on two atoms overlap to produce covalent bonds. Atomic orbitals hybridize to produce orbitals of equal energy with spatial orientations that can be predicted by VSEPR theory. --based on ELECTRON DOMAIN Geometry Linear tri plan tetra tri bypyr octa 2 = sp 3 = sp 2 4 = sp 3 5 = sp 3 d 6 = sp 3 d 2

83. Molecular Geometries and Bonding Which kind of hybridization is about the central atom in PCl 5 ? 1.sp 2 2.sp 3 3.sp 4 4.sp 3 d

84. Molecular Geometries and Bonding Correct Answer: 1.sp 2 2.sp 3 3.sp 4 4.sp 3 d The hybridization for PCl 5 is sp 3 d, as shown to the right.

85. Molecular Geometries and Bonding

86. Molecular Geometries and Bonding Yes, they’d be equivalent and we’d expect the bond angles to be 90°.

87. Valence Bond Theory Hybridization is a major player in this approach to bonding. There are two ways orbitals can overlap to form bonds between atoms.

88. Sigma (  ) Bonds Sigma bonds are characterized by  Head-to-head overlap.  Cylindrical symmetry of electron density about the internuclear axis.

89. Pi (  ) Bonds Pi bonds are characterized by  Side-to-side overlap.  Electron density above and below the internuclear axis.

90. Single Bonds Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy lowering.

91. Multiple Bonds In a multiple bond one of the bonds is a  bond and the rest are  bonds.

92. Multiple Bonds In a molecule like formaldehyde (shown at left) an sp 2 orbital on carbon overlaps in  fashion with the corresponding orbital on the oxygen. The unhybridized p orbitals overlap in  fashion.

93. Multiple Bonds In triple bonds, as in acetylene, two sp orbitals form a  bond between the carbons, and two pairs of p orbitals overlap in  fashion to form the two  bonds.

94. Molecular Geometries and Bonding

95. Molecular Geometries and Bonding The molecule is not linear but is planar. The 2 - 4 electron domains that come off two atoms in a double bond are always planar with the double bond

96. Molecular Geometries and Bonding In the molecule HCN, how many sigma bonds (  ) and how many pi bonds (  ) are there? 1.One sigma, one pi 2.Two sigma, one pi 3.Two sigma, two pi 4.None of the above

97. Molecular Geometries and Bonding Correct Answer: 1.One sigma, one pi 2.Two sigma, one pi 3.Two sigma, two pi 4.None of the above Total = Two sigma + two pi One sigma Two pi

98. Molecular Geometries and Bonding In the molecule C 2 H 4, how many sigma bonds (  ) and how many pi bonds (  ) are there? 1.Five sigma, one pi 2.Four sigma, two pi 3.Six sigma, zero pi 4.None of the above

99. Molecular Geometries and Bonding Correct Answer: 1.Five sigma, one pi 2.Four sigma, two pi 3.Six sigma, zero pi 4.None of the above The sigma and pi bonds for C 2 H 4 are indicated in the illustration.

100. Delocalized Electrons: Resonance When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion.

101. Delocalized Electrons: Resonance In reality, each of the four atoms in the nitrate ion has a p orbital. The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen.

102. Delocalized Electrons: Resonance This means the  electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion.

103. Resonance The organic molecule benzene has six  bonds and a p orbital on each carbon atom.

104. Resonance In reality the  electrons in benzene are not localized, but delocalized. The even distribution of the  electrons in benzene makes the molecule unusually stable.

105. Molecular Orbital (MO) Theory Though valence bond theory effectively conveys most observed properties of ions and molecules, there are some concepts better represented by molecular orbitals.

106. Molecular Orbital (MO) Theory In MO theory, we invoke the wave nature of electrons. If waves interact constructively, the resulting orbital is lower in energy: a bonding molecular orbital.

107. Molecular Orbital (MO) Theory If waves interact destructively, the resulting orbital is higher in energy: an antibonding molecular orbital.

108. MO Theory In H 2 the two electrons go into the bonding molecular orbital. The bond order is one half the difference between the number of bonding and antibonding electrons.

109. MO Theory For hydrogen, with two electrons in the bonding MO and none in the antibonding MO, the bond order is 1212 (2 - 0) = 1

110. MO Theory In the case of He 2, the bond order would be 1212 (2 - 2) = 0 Therefore, He 2 does not exist.

111. MO Theory For atoms with both s and p orbitals, there are two types of interactions:  The s and the p orbitals that face each other overlap in  fashion.  The other two sets of p orbitals overlap in  fashion.

112. MO Theory The resulting MO diagram looks like this. There are both  and  bonding molecular orbitals and  * and  * antibonding molecular orbitals.

113. MO Theory The smaller p-block elements in the second period have a sizeable interaction between the s and p orbitals. This flips the order of the s and p molecular orbitals in these elements.

114. Second-Row MO Diagrams