CHAPTER 8 n Molecular Structure Covalent Bonding Theories

CHAPTER 8 n Molecular Structure & Covalent Bonding Theories 1

Chapter Goals A Preview of the Chapter 2. Valence Shell Electron Pair Repulsion (VSEPR) Theory 3. Polar Molecules: The Influence of Molecular Geometry 4. Valence Bond (VB) Theory 1. 2

Chapter Goals Molecular Shapes and Bonding 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. Linear Electronic Geometry: AB 2 Species Trigonal Planar Electronic Geometry: AB 3 Species Tetrahedral Electronic Geometry: AB 4 Species Tetrahedral Electronic Geometry: AB 3 U Species Tetrahedral Electronic Geometry: AB 2 U 2 Species Tetrahedral Electronic Geometry – ABU 3 Species Trigonal Bipyramidal Geometry Octahedral Geometry Compounds Containing Double Bonds Compounds Containing Triple Bonds A Summary of Electronic and Molecular Geometries 3

Stereochemistry is the study of the three dimensional shapes of molecules. n Some questions to examine in this chapter are: n 1. Why are we interested in shapes? 2. What role does molecular shape play in life? 3. How do we determine molecular shapes? 4. How do we predict molecular shapes? 4

Molecular Shapes The shape of a molecule plays an important role in its reactivity. n By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule. n 5

Two Simple Theories of Covalent Bonding n Valence Shell Electron Pair Repulsion Theory – Commonly designated as VSEPR – Principal originator • R. J. Gillespie in the 1950’s n Valence Bond Theory – Involves the use of hybridized atomic orbitals – Principal originator • L. Pauling in the 1930’s & 40’s 6

VSEPR Theory n In order to attain maximum stability, each atom in a molecule or ion arranges the electron pairs in its valence shell in such a way to minimize the repulsion of their regions of high electron density: (a) Lone (unshared or nonbonding) pairs of electrons (b) Single bond (c) Double bond (d) Triple bond 7

VSEPR Theory n These four types of regions of high electron density (where the electron are) want to be as far apart as possible. The electrons repel each other. n There are five basic molecular shapes based on the number of regions of high electron density around the central atom. 8

VSEPR Theory These are the regions of high electron density around the central atom for two through six electron densities around a central atom. 9

Electron-Density Geometries All one must do is count the number of electron density in the Lewis structure. n The geometry will be that which corresponds to that number of electron density. : n : H H H Tetrahedral 10

VSEPR Theory Electronic geometry is determined by the locations of regions of high electron density around the central atom(s). 2. Molecular geometry determined by the arrangement of atoms around the central atom(s). 1. Electron pairs are not used in the molecular geometry determination just the positions of the atoms in the molecule are used. 11

Molecular Geometries electron-density Geometry - tetrahedral The electron-density geometry is often not the shape of the molecule, however. n The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs. 12 n

VSEPR Theory n An example of a molecule that has the same electronic and molecular geometries is methane - CH 4. n Electronic and molecular geometries are tetrahedral. 13

VSEPR Theory n An example of a molecule that has different electronic and molecular geometries is water - H 2 O. n Electronic geometry is tetrahedral. n Molecular geometry is bent or angular. 14

VSEPR Theory n Lone pairs of electrons (unshared pairs) require more volume than shared pairs. – Consequently, there is an ordering of repulsions of electrons around central atom. n Criteria for the ordering of the repulsions: 15

VSEPR Theory Lone pair to lone pair is the strongest repulsion. 2 Lone pair to bonding pair is intermediate repulsion. 3 Bonding pair to bonding pair is weakest repulsion. n Mnemonic for repulsion strengths 1 lp/lp > lp/bp > bp/bp n Lone pair to lone pair repulsion is why bond angles in water are less than 109. 5 o. 16

VSEPR Theory lp/bp 17

Multiple Bonds and Bond Angles Double and triple bonds place greater electron density on one side of the central atom than do single bonds. n Therefore, they also affect bond angles. n bp/bp repulsion 18

Nonbonding Pairs and Bond Angle Nonbonding pairs are physically larger than bonding pairs. n Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule. n 19

Nonbonding Pairs and Bond Angle 20

Polarity In Chapter 7 we discussed bond dipoles. n But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar. n 21

Polar Molecules: The Influence of Molecular Geometry n Molecular geometry affects molecular polarity. – Due to the effect of the bond dipoles and how they either cancel or reinforce each other. A B A linear molecule nonpolar A B A angular molecule polar 22

Polarity By adding the individual bond dipoles, one can determine the overall dipole moment for 23 the molecule.

Polarity 24

Polar Molecules: The Influence of Molecular Geometry n Polar Molecules must meet two requirements: 1. One polar bond or one lone pair of electrons on central atom. 2. Neither bonds nor lone pairs can be symmetrically arranged that their polarities cancel. 25

Polarity 26

Valence Bond (VB) Theory Covalent bonds are formed by the overlap of atomic orbitals. n Atomic orbitals on the central atom can mix and exchange their character with other atoms in a molecule. n – Process is called hybridization n Hybrids are common: 1. Pink flowers 2. Mules § Hybrid Orbitals have the same shapes as predicted by VSEPR. 27

Valence Bond (VB) Theory Regions of High Electron Density 2 3 4 5 6 Electronic Geometry Hybridization Linear Trigonal planar Tetrahedral Trigonal bipyramidal Octahedral sp sp 2 sp 3 d 2 28

Molecular Shapes and Bonding n In the next sections we will use the following terminology: A = central atom B = bonding pairs around central atom U = lone pairs around central atom n For example: AB 3 U designates that there are 3 bonding pairs and 1 lone pair around the central atom. 29

Linear Electronic Geometry: AB 2 Species (No Lone Pairs of Electrons on A) n Some examples of molecules with this geometry are: Be. Cl 2, Be. Br 2, Be. I 2, Hg. Cl 2, Cd. Cl 2 n All of these examples are linear, nonpolar molecules. n Important exceptions occur when the two substituents are not the same! Be. Cl. Br or Be. IBr will be linear and polar! 30

Linear Electronic Geometry: AB 2 Species (No Lone Pairs of Electrons on A) Electronic Geometry 31

Linear Electronic Geometry: AB 2 Species (No Lone Pairs of Electrons on A) Polarity 32

Linear Electronic Geometry: AB 2 Species (No Lone Pairs of Electrons on A) Valence Bond Theory (Hybridization) 33

Linear Electronic Geometry: AB 2 Species (No Lone Pairs of Electrons on A) 34

Linear Electronic Geometry: AB 2 Species (No Lone Pairs of Electrons on A) 35

Trigonal Planar Electronic Geometry: AB 3 Species (No Lone Pairs of Electrons on A) n Some examples of molecules with this geometry are: BF 3, BCl 3 n All of these examples are trigonal planar, nonpolar molecules. n Important exceptions occur when the three substituents are not the same! BF 2 Cl or BCI 2 Br will be trigonal planar and polar! 36

Trigonal Planar Electronic Geometry: AB 3 Species (No Lone Pairs of Electrons on A) Dot Formula Electronic Geometry 37

Trigonal Planar Electronic Geometry: AB 3 Species (No Lone Pairs of Electrons on A) Polarity 38

Trigonal Planar Electronic Geometry: AB 3 Species (No Lone Pairs of Electrons on A) Valence Bond Theory (Hybridization) 3 s 3 p Cl [Ne] 39

Trigonal Planar Electronic Geometry: AB 3 Species (No Lone Pairs of Electrons on A) 40

Trigonal Planar Electronic Geometry: AB 3 Species (No Lone Pairs of Electrons on A) 41

Tetrahedral Electronic Geometry: AB 4 Species (No Lone Pairs of Electrons on A) n Some examples of molecules with this geometry are: CH 4, CF 4, CCl 4, Si. H 4, Si. F 4 n All of these examples are tetrahedral, nonpolar molecules. n Important exceptions occur when the four substituents are not the same! CF 3 Cl or CH 2 CI 2 will be tetrahedral and polar! 42

Tetrahedral Electronic Geometry: AB 4 Species (No Lone Pairs of Electrons on A) 43

Tetrahedral Electronic Geometry: AB 4 Species (No Lone Pairs of Electrons on A) 44

Tetrahedral Electronic Geometry: AB 4 Species (No Lone Pairs of Electrons on A) 45

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Tetrahedral Electronic Geometry: AB 3 U Species (One Lone Pair of Electrons on A) Some examples of molecules with this geometry are: NH 3, NF 3, PH 3, PCl 3, As. H 3 n These molecules are our first examples of central atoms with lone pairs of electrons. Thus, the electronic and molecular geometries are different. All three substituents are the same but molecule is polar n NH 3 and NF 3 are trigonal pyramidal, polar molecules. 49 n

Tetrahedral Electronic Geometry: AB 3 U Species (One Lone Pair of Electrons on A) Valence Bond Theory 50

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Tetrahedral Electronic Geometry: AB 3 U Species (One Lone Pair of Electrons on A) 52

Electronic Geometry 53

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Molecular Geometry 55

Tetrahedral Electronic Geometry: AB 3 U Species (One Lone Pair of Electrons on A) Polarity 56

Tetrahedral Electronic Geometry: AB 2 U 2 Species (Two Lone Pairs of Electrons on A) n Some examples of molecules with this geometry are: H 2 O, OF 2, OCl 2, H 2 S n These molecules are our first examples of central atoms with two lone pairs of electrons. Thus, the electronic and molecular geometries are different. Both substituents are the same but molecule is polar n Molecules are angular, bent, or V-shaped and polar. 57

Tetrahedral Electronic Geometry: AB 2 U 2 Species (Two Lone Pairs of Electrons on A) Valence Bond Theory (Hybridization) O [He] 2 s 2 p ¯ ¯ Þ four sp 3 hybrids ¯ ¯ 58

Tetrahedral Electronic Geometry: ABU 3 Species (Three Lone Pairs of Electrons on A) n Some examples of molecules with this geometry are: HF, HCl, HBr, HI, FCl, IBr n These molecules are examples of central atoms with three lone pairs of electrons. Again, the electronic and molecular geometries are different. n Molecules are linear and polar when the two atoms are different. Cl 2, Br 2, I 2 are nonpolar 59

Tetrahedral Electronic Geometry: ABU 3 Species (Three Lone Pairs of Electrons on A) Dot Formula Electronic Geometry tetrahedral Molecular Geometry Polarity HF is a polar molecule. linear 60

Tetrahedral Electronic Geometry: ABU 3 Species (Three Lone Pairs of Electrons on A) Valence Bond Theory (Hybridization) 2 s 2 p four sp 3 hybrids F [He] Þ tetrahedral 61

Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 § Some examples of molecules with this geometry are: PF 5, As. F 5, PCl 5, etc. n These molecules are examples of central atoms with five bonding pairs of electrons. The electronic and molecular geometries are the same. n Molecules are trigonal bipyramidal and nonpolar when all five substituents are the same. If the five substituents are not the same polar molecules 62 can result, As. F 4 Cl is an example.

Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 Valence Bond Theory 63

Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 64

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Molecular Geometry Trigonal Bipyramidal 66

Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 n If lone pairs are incorporated into the trigonal bipyramidal structure, there are three possible new shapes. 1. One lone pair - Seesaw shape 2. Two lone pairs - T-shape 3. Three lone pairs – linear n The lone pairs occupy equatorial positions because they are 120 o from two bonding pairs and 90 o from the other two bonding pairs. – Results in decreased repulsions compared to lone pair in 67 axial position.

Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 n AB 4 U molecules have: 1. trigonal bipyramid electronic geometry 2. seesaw shaped molecular geometry 3. and are polar n One example of an AB 4 U molecule is SF 4 n Hybridization of S atom is sp 3 d. 68

Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 Lewis Dot Molecular Geometry seesaw Electronic Geometry 69

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Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 n AB 3 U 2 molecules have: 1. trigonal bipyramid electronic geometry 2. T-shaped molecular geometry 3. and are polar n One example of an AB 3 U 2 molecule is IF 3 n Hybridization of I atom is sp 3 d. 71

Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 Molecular Geometry 72

Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 n AB 2 U 3 molecules have: 1. trigonal bipyramid electronic geometry 2. linear molecular geometry 3. and are nonpolar n One example of an AB 3 U 2 molecule is Xe. F 2 n Hybridization of Xe atom is sp 3 d. 73

Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 Molecular Geometry 74

Trigonal Bipyramidal Electronic Geometry: AB 5, AB 4 U, AB 3 U 2, and AB 2 U 3 75

Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 n Some examples of molecules with this geometry are: SF 6, Se. F 6, SCl 6, etc. These molecules are examples of central atoms with six bonding pairs of electrons. n Molecules are octahedral and nonpolar when all six substituents are the same. n If the six substituents are not the same polar molecules can result, SF 5 Cl is an example. 76

Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 Nonpolar 77

Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 78

Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 79

Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 n If lone pairs are incorporated into the octahedral structure, there are two possible new shapes. 1. One lone pair - square pyramidal 2. Two lone pairs - square planar n The lone pairs occupy axial positions because they are 90 o from four bonding pairs. – Results in decreased repulsions compared to lone pairs in equatorial positions. 80

Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 n AB 5 U molecules have: 1. octahedral electronic geometry 2. Square pyramidal molecular geometry 3. and are polar. n One example of an AB 4 U molecule is IF 5 n Hybridization of I atom is sp 3 d 2. 81

Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 Molecular Geometry 82

Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 n AB 4 U 2 molecules have: 1. octahedral electronic geometry 2. square planar molecular geometry 3. and are nonpolar. n One example of an AB 4 U 2 molecule is Xe. F 4 n Hybridization of Xe atom is sp 3 d 2. 83

Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 Molecular Geometry Polarity nonpolar 84

Octahedral Electronic Geometry: AB 6, AB 5 U, and AB 4 U 2 85

Sigma ( ) Bonds n Sigma bonds are characterized by – Head-to-head overlap. – Cylindrical symmetry of electron density about the internuclear axis. 86

Pi ( ) Bonds n Pi bonds are characterized by – Side-to-side overlap. – Electron density above and below the internuclear axis. 87

Single Bonds Single bonds are always bonds, because overlap is greater, resulting in a stronger bond and more energy lowering. 88

Multiple Bonds In a multiple bond one of the bonds is a bond and the rest are bonds. 89

Compounds Containing Double Bonds Ethene or ethylene, C 2 H 4, is the simplest organic compound containing a double bond. Lewis dot formula N = 2(8) + 4(2) = 24 A = 2(4) + 4(1) = 12 S = 12 n Compound must have a double bond to obey octet rule. n 90

Compounds Containing Double Bonds n VSEPR Theory suggests that the C atoms are at center of trigonal planes. H H C C H H Valence Bond Theory (Hybridization) C atom has four electrons. Three electrons from each C atom are in sp 2 hybrids. One electron in each C atom remains in an unhybridized p orbital 91

Compounds Containing Double Bonds n An sp 2 hybridized C atom has this shape. Remember there will be one electron in each of the three lobes. Top view of an sp 2 hybrid n The single 2 p orbital is perpendicular to the trigonal planar sp 2 lobes. The fourth electron is in the p orbital. Side view of sp 2 hybrid with p orbital included. 92

Compounds Containing Double Bonds n n Two sp 2 hybridized C atoms plus p orbitals in proper orientation to form C=C double bond. The head-on overlap of the sp 2 hybrids is designated 93 as a bond.

Compounds Containing Double Bonds n The other portion of the double bond, resulting from the side-on overlap of the p orbitals, is designated as a bond. 94

Compounds Containing Double Bonds n Thus a C=C bond looks like this and is made of two parts, one and one bond. 95

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. n The unhybridized p orbitals overlap in fashion. n 96

Compounds Containing Triple Bonds Ethyne or acetylene, C 2 H 2, is the simplest triple bond containing organic compound. Lewis Dot Formula n N = 2(8) + 2(2) = 20 A = 2(4) + 2(1) =10 S = 10 n Compound must have a triple bond to obey octet rule. 97

Compounds Containing Triple Bonds VSEPR Theory suggests regions of high electron density are 180 o apart. H n C C H Valence Bond Theory (Hybridization) Carbon has 4 electrons. Two of the electrons are in sp hybrids. Two electrons remain in unhybridized p orbitals. 98

Compounds Containing Triple Bonds n An sp hybridized C atom has this shape. Remember there will be one electron in each of the two lobes. n The two 2 p orbital are perpendicular to the sp lobes. The third and fourth electrons are in the p orbitals. n The head-on overlap of the sp 2 hybrids is designated as a s bond. 99

Compounds Containing Triple Bonds A bond results from the head-on overlap of two sp hybrid orbitals. 100

Compounds Containing Triple Bonds n The unhybridized p orbitals form two p bonds. n Note that a triple bond consists of one and two bonds. n The final result is a bond that looks like this. 101

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. 102

Larger Molecules This approach makes sense, especially because larger molecules tend to react at a particular site in the molecule. 103

Here is the structure for most students’ friend: CAFFEINE 1 - Assign hybridization on C, N, and O. Beware I did not put the lone pairs of electrons into the chemical drawing. 2 - How many sigma bonds are present? 3 - How many pi bonds are present? 4 - How many lone pairs of electrons are present? (You have to look for them) 104

Here is the structure Theobromine, one of the components of TEA 1 - Assign hybridization on C, N, and O. Beware I did not put the lone pairs of electrons into the chemical drawing. 2 - How many sigma bonds are present? 3 - How many pi bonds are present? 4 - How many lone pairs of electrons are present? (You have to look for them) 105

End of Chapter 8 This is a difficult chapter. n Essential to your understanding of chemistry! n 106

Homework Assignment One-line Web Learning (OWL): Chapter 8 Exercises and Tutors – Optional 107