1 Covalent Bonding: Molecular Geometry Hybridization of Atomic Orbitals Molecular Orbitals

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3 AB 2 20linear Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB 3 30 trigonal planar 10.1 AB 4 40 tetrahedral AB 5 50 trigonal bipyramidal trigonal bipyramidal AB 6 60 octahedral Valence shell electron pair repulsion (VSEPR) model: Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs.

4 linear trigonal planar tetrahedral trigonal bipyramidal octahedral VSEPR

5 linear trigonal planar tetrahedral trigonal bipyramidal octahedral Cl Be BF 3 CH 4 PCl 5 SF 6

6 bonding-pair vs. bonding pair repulsion lone-pair vs. lone pair repulsion lone-pair vs. bonding pair repulsion >> Effects of Lone Pairs O HH N HH H C HH H H

7 Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 30 trigonal planar AB 2 E21 trigonal planar bent

8 Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 E31 AB 4 40 tetrahedral trigonal pyramidal

9 Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 4 40 tetrahedral AB 3 E31tetrahedral trigonal pyramidal AB 2 E 2 22tetrahedral bent H O H

10 Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 5 50 trigonal bipyramidal trigonal bipyramidal AB 4 E41 trigonal bipyramidal distorted tetrahedron

11 Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 5 50 trigonal bipyramidal trigonal bipyramidal AB 4 E41 trigonal bipyramidal distorted tetrahedron AB 3 E 2 32 trigonal bipyramidal T-shaped Cl F F F

12 Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 5 50 trigonal bipyramidal trigonal bipyramidal AB 4 E41 trigonal bipyramidal distorted tetrahedron AB 3 E 2 32 trigonal bipyramidal T-shaped AB 2 E 3 23 trigonal bipyramidal linear I I I

13 Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 6 60 octahedral AB 5 E51 octahedral square pyramidal Br FF FF F

14 Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 6 60 octahedral AB 5 E51 octahedral square pyramidal AB 4 E 2 42 octahedral square planar Xe FF FF

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16 Predicting Molecular Geometry 1.Draw Lewis structure for molecule. 2.Count number of lone pairs on the central atom and number of atoms bonded to the central atom. 3.Use VSEPR to predict the geometry of the molecule. What are the molecular geometries of SO 2 and SF 4 ? SO O AB 2 E bent S F F F F AB 4 E distorted tetrahedron

17 Dipole Moments and Polar Molecules H F electron rich region electron poor region    = Q x r Q is the charge r is the distance between charges 1 D = 3.36 x C m

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19 Which of the following molecules have a dipole moment? H 2 O, CO 2, SO 2, and CH 4 O H H dipole moment polar molecule S O O CO O no dipole moment nonpolar molecule dipole moment polar molecule C H H H H no dipole moment nonpolar molecule

20 Does CH 2 Cl 2 have a dipole moment?

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23 Chemistry In Action: Microwave Ovens

24 Bond Dissociation EnergyBond Length H2H2 F2F kJ/mole kJ/mole 74 pm 142 pm Valence bond theory – bonds are formed by sharing of e - from overlapping atomic orbitals. Overlap Of 2 1s 2 2p How does Lewis theory explain the bonds in H 2 and F 2 ? Sharing of two electrons between the two atoms.

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26 Change in electron density as two hydrogen atoms approach each other.

27 Valence Bond Theory and NH 3 N – 1s 2 2s 2 2p 3 3 H – 1s 1 If the bonds form from overlap of 3 2p orbitals on nitrogen with the 1s orbital on each hydrogen atom, what would the molecular geometry of NH 3 be? If use the 3 2p orbitals predict 90 0 Actual H-N-H bond angle is

28 Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals. 1.Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals. 2.Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process. 3.Covalent bonds are formed by: a.Overlap of hybrid orbitals with atomic orbitals b.Overlap of hybrid orbitals with other hybrid orbitals

29 Bonding in Methane

30 Formation of sp 3 Hybrid Orbitals Fig. 10.7

31 Formation of sp 3 Hybrid Orbitals

32 Formation of a CH 4 Molecule

33 Formation of a NH 3 Molecule

34 NH 3 CH 4 Stylized Drawing of Valence Bond Theory Predict correct bond angle Sigma bond (  ) – electron density between the 2 atoms

35 Formation of sp 2 Hybrid Orbitals

36 Formation of sp 2 Hybrid Orbitals

37 Formation of sp 2 Hybrid Orbitals 2p z orbital is perpendicular to the plane of hybridized orbitals

38 sp 2 Hybridization of a C atom

39 Bonding in Ethylene C 2 H 4 CC H H H H Sigma bond (  ) – electron density between the 2 atoms

40 Pi bond (  ) – electron density above and below plane of nuclei of the bonding atoms Bonding in Ethylene C 2 H 4

41 Bonding in Ethylene C 2 H 4

42 Formation of sp Hybrid Orbitals

43 Formation of sp Hybrid Orbitals

44 Formation of sp Hybrid Orbitals

45 Bonding in acetylene C 2 H 2

46 # of Lone Pairs + # of Bonded Atoms HybridizationExamples sp sp 2 sp 3 sp 3 d sp 3 d 2 BeCl 2 BF 3 CH 4, NH 3, H 2 O PCl 5 SF 6 How do I predict the hybridization of the central atom? Count the number of lone pairs AND the number of atoms bonded to the central atom

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48 Sigma (  ) and Pi Bonds (  ) Single bond 1 sigma bond Double bond 1 sigma bond and 1 pi bond Triple bond 1 sigma bond and 2 pi bonds How many  and  bonds are in the acetic acid (vinegar) molecule CH 3 COOH? C H H CH O OH  bonds = = 7  bonds = 1

49 Molecular orbital theory – bonds are formed from interaction of atomic orbitals to form molecular orbitals. O O No unpaired e - Should be diamagnetic Experiments show O 2 is paramagnetic Drawback of Valence Bond Theory

50 Amplitudes of wave functions added An analogy between light waves and atomic wave functions Amplitudes of wave functions subtracted.

51 Energy levels of bonding and antibonding molecular orbitals in hydrogen (H 2 ) A bonding molecular orbital has lower energy and greater stability than the atomic orbitals from which it was formed. An antibonding molecular orbital has higher energy and lower stability than the atomic orbitals from which it was formed.

52 Energy levels of bonding and antibonding molecular orbitals in boron (B 2 )

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54 Second-Period Homonuclear Diatomic Molecules

55 1.The number of molecular orbitals (MOs) formed is always equal to the number of atomic orbitals combined. 2.The more stable the bonding MO, the less stable the corresponding antibonding MO. 3.The filling of MOs proceeds from low to high energies. 4.Each MO can accommodate up to two electrons. 5.Use Hund’s rule when adding electrons to MOs of the same energy. 6.The number of electrons in the MOs is equal to the sum of all the electrons on the bonding atoms. Molecular Orbital (MO) Configurations

56 bond order = 1 2 Number of electrons in bonding MOs Number of electrons in antibonding MOs ( - ) bond order ½ 1 0½ Bond Order

57 MO theory predicts that O 2 is paramagnetic! MO for 2 nd Period Homonuclear Diatomic Molecules

58 Molecules with Resonance Structures

59 Delocalized molecular orbitals are not confined between two adjacent bonding atoms, but actually extend over three or more atoms. Delocalized π Molecular Orbitals

60 Electron density above and below the plane of the benzene molecule.

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62 Acknowledgment Some images, animation, and material have been taken from the following sources: Chemistry, Zumdahl, Steven S.; Zumdahl, Susan A.; Houghton Mifflin Co., 6th Ed., 2003; supplements for the instructor General Chemistry: The Essential Concepts, Chang, Raymon; McGraw-Hill Co. Inc., 4 th Ed., 2005; supplements for the instructor Principles of General Chemistry, Silberberg, Martin; McGraw-Hill Co. Inc., 1st Ed., 2006; supplements for the instructor