Chapter 9 Chemical Bonding II: Molecular Geometry and Bonding Theories. Copyright McGraw-Hill

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1 Chapter 9 Chemical Bonding II: Molecular Geometry and Bonding Theories 1

2 9.1 Molecular Geometry Molecular geometry is the threedimensional shape of a molecule. CCl 4 Geometry can be predicted using Lewis dot structures VSEPR model 2

3 Molecules of the type AB x will be considered A is the central atom B atoms surround the central atom x commonly has integer values from 1 to 6 Examples: All AB 1 molecules are linear. 3

4 AB 2 AB 3 4

5 VSEPR Model: Valence-Shell Electron- Pair Repulsion Model Electron pairs move as far apart as possible to minimize repulsions. Electron domain is a lone pair or a bond (the bond may be single, double, or triple). Strategy to predict geometry: Lewis structure Electron-domain geometry Molecular geometry 5

6 Steps to determine Geometry Step #1: Draw the molecule s Lewis structure. Step #2: Count the number of electron domains on the central atom. Step #3: Determine the electron-domain geometry. Examples The electron-domain geometry is based on the number of electron domains around the central atom. H C N: :O: :O N = O: : : : : :F Xe F: : : : : : 6

7 Electron domains and geometry Number of Electron Domains Electron-Domain Geometry 2 Linear 3 Trigonal planar 4 Tetrahedral 7

8 Number of Electron Domains Electron-Domain Geometry 5 Trigonal bipyramidal 6 Octahedral 8

9 Step #4: Determine the molecular geometry. The electron-domain geometry and the number of bonded atoms determine the molecular geometry. Example: Ammonia, NH 3 Step #1 Step #2 : H N H H 4 e domains electron-domain geometry tetrahedral Step #3 molecular geometry = trigonal pyramidal 9

10 Note: The common molecular geometries are all derived from these 5 electron-domain geometries. Linear Bent Trigonal planar Bent Trigonal pyramidal Tetrahedral T-shaped Seesaw Trigonal bipyramidal Linear T-shaped Square planar Square pyramidal Octahedral 10

11 Axial and equatorial positions 90º 120º The 5 electron domains are not all equivalent. ax = axial eq = equatorial eq eq ax ax eq For SF 4, which geometry is correct? Why? or Fewest lone-pair bond-pair interactions at angles of 90 o 11

12 12

13 13

14 Exercise: For each of the following species, i) draw the Lewis structure. ii) determine the number of electron domains on the central atom and its electron-domain geometry. iii) predict the molecular geometry. a) NF 3 b) CO

15 a) NF 3 i) F N F F ii) 4 electron domains on the central atom. Electron-domain geometry: tetrahedral iii) One lone pair on the central atom. Molecular geometry: trigonal pyramidal 15

16 a) CO 3 2 i) O C O 2 O ii) 3 electron domains on the central atom. Electron-domain geometry: trigonal planar iii) No lone pairs on the central atom. Molecular geometry: trigonal planar 16

17 Deviations from ideal bond angles All electron domains repel each other. The repulsion between domains depends on the types of domains involved. single bond double bond triple bond < < < Increasingly repulsion lone pair 17

18 lone-pair - lone-pair repulsion is greater than lone-pair - bonding-pair repulsion is greater than bonding-pair - bonding-pair repulsion 18

19 Geometry with more than one central atom No lone pairs tetrahedral CH 3 OH AX 4 AX 4 H H C O H H 2 lone pairs bent H H C H O H 19

20 9.2 Molecular Geometry and Polarity The HF bond is polar and HF has a dipole moment ( ). + - H F H F Bond dipoles are vectors and therefore are additive. 20

21 Molecules with more than two atoms Remember bond dipoles are additive since they are vectors. H 2 O dipole moment > 0 CO 2 dipole moment = 0 21

22 Example: Dichloroethene, C 2 H 2 Cl 2, exists as three isomers. Cl H C= C Cl H H Cl C = C Cl H H H C = C Cl Cl cis-1,2-dichloroethene trans-1,2-dichloroethene 1,1-dichloroethene polar = 1.90 D bp = 60.3ºC nonpolar = 0 D bp = 47.5ºC polar = 1.34 D bp = 31.7ºC 22

23 9.3 Valence Bond Theory Electrons in molecules occupy atomic orbitals. Covalent bonding results from the overlap of atomic orbitals. 23

24 Representation of singly-occupied and doubly-occupied s and p atomic orbitals. Singly-occupied orbitals appear light; doubly-occupied orbitals appear darker. 24

25 Example: H(1s 1 ) + H(1s 1 ) H 2 25

26 Example: F(1s 2 2s 2 2p 5 ) + F(1s 2 2s 2 2p 5 ) F 2 26

27 Example: H(1s 1 ) + F(1s 2 2s 2 2p 5 ) HF 27

28 9.4 Hybridization of Atomic Orbitals Valence bond theory cannot account for many experimental observations. Beryllium Chloride, BeCl 2 Cl Be Cl VSEPR linear both bonds equivalent #1 No unpaired electrons #2 2 types of overlap with 2s and 2p 28

29 29

30 The atomic orbitals on an atom mix to form hybrid orbitals. Orbital shapes (boundary surfaces) are pictorial representations of wave functions. Wave functions are mathematical functions. Mathematical functions can be combined. Hybridization of s and p orbitals sp hybridization 30

31 The two sp orbitals point in opposite directions inline with one another. 31

32 Each Be sp orbital overlaps a Cl 3p orbital to yield BeCl 2. 2Cl + Be BeCl 2 All bond angles 180 o. 32

33 sp 2 hybridization Example: Boron trifluoride, BF 3... F. B.. F. Ḟ VSEPR trigonal planar all bonds equivalent #1 only 1 unpaired electron #2 2 types of overlap with 2s and 2p #3 overlap with p-orbitals = 90 o 33

34 The three sp 2 orbitals point to the corners of an equilateral triangle. All bond angles 120 o. 34

35 Each B sp 2 orbital overlaps a F 2p orbital to yield BF 3. 3F + B BF 3 2s 2 2p 5 All bond angles 120 o. 35

36 sp 3 hybridization Example: Methane, CH 4 H H C H VSEPR H tetrahedral all bonds equivalent #1 only 2 unpaired electrons #2 2 types of overlap with 2s and 2p #3 overlap with p-orbitals = 90 36

37 37

38 The sp 3 hybrid orbitals point to the corners of a tetrahedron. All bond angles o. 38

39 Each C 2sp 3 orbital overlaps a H 1s orbital to yield CH 4. C 4H + CH 4 1s 1 All bond angles o. 39

40 Hybridization of s, p and d orbitals Expanded octets Bond angles 120 o and 90 o. 40

41 Lewis structure Number of electron domains Type of hybridization 41

42 9.5 Hybridization in Molecules Containing Multiple Bonds A sigma ( ) bond is a bond in which the shared electron density is concentrated directly along the internuclear axis between the two nuclei involved in bonding. end-to-end overlap 42

43 Double and triple bonds consist of both sigma ( ) and pi ( ) bonds. A pi ( ) bond forms from the sideways overlap of p orbitals resulting in regions of electron density that are concentrated above and below the plane of the molecule. sideways overlap 43

44 Example: Ethylene, C 2 H 4 C = C H H H H number of = 3 electron domains hybridization = sp 2 Double bond = 1 bond + 1 bond 44

45 Example: Acetylene, C 2 H 2 H C C H number of = 2 electron domains hybridization = sp Triple bond = 1 bond + 2 bonds 45

46 Exercise: How many pi bonds and sigma bonds are in each of the following molecules? Describe the hybridization of each C atom. Cl H sp 3 H sp 2 Cl H sp 2 C Cl C C H 3 C sp 2 sp 3 C C sp 2 H H H H (a) (b) (c) C C H sp sp (a) 4 sigma bonds (b) 5 sigma bonds, 1 pi bond (c) 10 sigma bonds, 3 pi bonds 46

47 9.6 Molecular Orbital Theory Molecular Orbital Theory (MO) Atomic orbitals combine to form new molecular orbitals which are spread out over the entire molecule. Electrons are in orbitals that belong to the molecule as a whole. Molecular orbitals (wave functions) result from adding and/or subtracting atomic orbitals (wave functions). 47

48 Molecular orbital theory (MO) Atomic orbitals combine to form new, molecular orbitals that are spread out over the entire molecule. Electrons are now in orbitals that belong to the molecule as a whole. 48

49 Pictorial representation. H + H H 2 1s 1 1s 1 1s 2 difference antibonding orbital * s 1 bonding orbital 1s 1s H atomic orbitals sum 1s H 2 molecular orbitals 49

50 Types of molecular orbitals Bonding molecular orbital High electron density between the nuclei Lower energy and more stable than the atomic orbitals that were added 50

51 Antibonding molecular orbital Low electron density between the nuclei node Higher energy and less stable than the atomic orbitals that were subtracted 51

52 Bond order bond order = 1/2 (2 0) = 1 single bond bond order = 1/2 (2 2) = 0 no bond 52

53 Examples Using Antibonding Orbitals bond order = 1/2 (2 0) = 1 single bond bond order = 1/2 (2 2) = 0 no bond 53

54 Pi ( ) molecular orbitals Wave functions representing p orbitals combine in two different ways yielding either orbitals or orbitals. End-to-end combination yields sigma ( ) orbitals sum antibonding orbital * p ± - + 2p z Atomic orbitals 2p z difference bonding orbital Molecular orbitals 2p 54

55 Sideways combination yields pi ( ) orbitals + + ± - - 2p y 2p y difference sum * p 2 2p + - ± + - 2p x 2p x Atomic orbitals difference sum * p 2 2p Molecular orbitals

56 Energy order of the 2p and 2p orbitals changes across the period. B 2, C 2, N 2 O 2, F 2, Ne 2 56

57 Molecular orbital diagram of nitrogen, N 2 57

58 Molecular orbital diagram of oxygen, O 2 58

59 Magnetism Diamagnetic substance A substance whose electrons are all paired. Weakly repelled by magnetic fields. Paramagnetic substance A substance with one or more unpaired electrons. Attracted by magnetic fields. 59

60 60

61 9.7 Bonding Theories and Descriptions of Molecules with Delocalized Bonding In localized bonds the and bonding electrons are associated with only two atoms. Resonance requires delocalized bonds when applying valence bond theory. 61

62 In delocalized bonds the bonding electrons are associated with more than two atoms. Examples: NO 3, CO 3 2, C 6 H 6, O 3 62

63 Example: benzene, C 6 H 6 Each C atom has 3 electron domains Hybridization: sp 2 63

64 Key Points Molecular geometry VSEPR model Molecular geometry and polarity Valence bond theory Hybridization of atomic orbitals s and p s, p, and d Hybridization involving multiple bonds 64

65 Key Points Molecular orbital theory Bonding and antibonding orbitals Sigma ( ) molecular orbitals Pi ( ) molecular orbitals MO diagrams 65

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