1. Chapter 9 Molecular Geometry

2. Introduction 1.Lewis Structures help us understand the compositions of molecules & their covalent bonds, but not their overall shapes. 2.The properties of a substance largely depend on the shape & size of its molecules, together with the strength & polarity of its bonds. 3. Examples: Taxol, smell, vision

3. VSEPR Theory Valence-shell electron-pair repulsion 1.The overall shape of a molecule is determined by its bond angles. 2. The VSEPR Model is used to predict Molecular Shapes or Molecular Geometries. 3. The VSEPR Theory assumes that each atom in a molecule will be positioned so that there is minimal repulsion between the valence electrons of that atom.valence electrons

4. Five Basic Shapes 1.Linear 2. Trigonal Planar 3. Tetrahedral 4.Trigonal Bipyramidal 5.Octahedral

5. Chart to Memorize for Basic Shapes Bond Angle(s) #Electron-Pairs Hybridization 180 2 sp 120 3 sp 2 109.5 4 sp 3 90 & 120 5 sp 3 d 90 6 sp 3 d 2

6. Using VSEPR Model to Predict Shapes 1.Draw the Lewis dot structure to determine the total # of electron pairs around the central atom. 2.Multiple bonds (double and triple) count as one. 3.The total number of bonding and nonbonding electron pairs determines the geometry of the electron pairs (one of the 5 basic shapes). 4.Then use bonding electron pairs only to determine the molecular geometry (actual shape of molecule).

7. Special Rules 1.Lone pairs effect geometry more than bonding pairs. 2.NH 3 has one lone pair: reduces angle from 109.5 to 107 3.H 2 O with two lone pairs: reduces angle from 109.5 to 105. 4.Multiple bonds affect geometry more than single bonds 5.H 2 C=O (116 instead of 120) 6.H 2 C=CH 2 (117 instead of 120) 7.Lone pairs occupy the axial (middle) position for trigonal bipyramidal structures.

8. Practice Determine the geometry of the following: BeCl 2 CO 2 BF 3 O 3 SO 2 CH 4 PCl 3 H 2 O PCl 5 SF 4 ClF 3 XeF 2 SF 6 IF 5 XeF 4 Consult the next 3 pages to help you.

12. Linear 180 o BeCl 2 valence e - = 2 +(2 x 7)= 16e - Cl.. Be Cl.. Two Electron Pairs = Linear Molecule

13. linear180 o CO 2 valence e - = 4 +(2 x 6) = 16e - C O.. O C O O valence pairs on C ignore double bondstwo single and double bonds same linearmolecular shape molecular geometry linear

14. trigonal planar120 o SO 2 valence e - = 6+(2 x 6)= 18e - valence pairs on Sthree one lone pair molecular geometry molecular shape bent trigonal S O.. O : S O O : S O O : two bonding pairs < 120 o

15. tetrahedral109.5 o CH 4 valence e - = 4+(4 x 1)= 8e - valence pairs on Cfour C HH H H 109.5 o molecular geometry molecular shape tetrahedral

16. 109.5 o NH 3 valence e - = 5+(3 x 1)= 8e - valence pairs on Nfour N HH H : < 109.5 o molecular geometry molecular shape trigonal pyramid tetrahedral one lone pair three bonding pairs

17. bipyramidal120 o and 180 0 PCl 5 valence e - = 5+(5 x 7)= 40e - valence pairs on Pfive molecular geometry molecular shape bipyramidal P Cl.. Cl.. Cl.. Cl.. Cl.. 90 o 120 o 180 o

18. bipyramidal120 o and 180 0 SF 4 valence e - = 6+(4 x 7)= 34e - valence pairs on Sfive molecular geometry molecular shape seesaw bipyramidal one lone pair four bonding pairs S.. F F F F : < 180 o

19. bipyramidal120 o and 180 0 ClF 3 valence e - = 7+(3 x 7)= 28e - valence pairs on Clfive molecular geometry molecular shape T bipyramidal two lone pair three bonding pairs Cl.. F F F : : 180 o 90 o

20. bipyramidal120 o and 180 0 ICl 2 - valence e - = 7+(2 x 7)+ e - valence pairs on Ifive molecular geometry molecular shape linear bipyramidal three lone pair on I two bonding pairs = 22e - I.. Cl.. : : Cl.. :

21. octahedral90 o BrF 5 valence e - = 7+(5 x 7) valence pairs on Brsix molecular geometry molecular shape square pyramidal octahedral = 42e - one lone pair five bonding pairs Br F.. F F F F :

22. octahedral90 o XeF 4 valence e - = 8+(4 x 7) valence pairs on Xesix molecular geometry molecular shape square planar octahedral = 36e - two lone pair four bonding pairs Xe F.. F F F : :

24. Cl.. Be Cl.. S O O S O O S O O S OO : C O O C O O B : : F :: : F :: : F : : : C HH H H N HH H :

25. O HH : : P Cl.. Cl.. Cl.. Cl.. Cl.. S F F F F : Cl.. F F F : : I Cl.. : : Cl.. : F S F F F F F

26. Br F.. F F F F : Xe F.. F F F : :

28. Valence Bond Theory 1.Valence Bond Theory explains why molecules have the shapes they do based on a concept called hybridization. 2. Hybridization - A mixture of two or more atomic orbitals.

29. http://www.youtube.com/watch?v=PrNbhuB9W44&feature=related 10 Minutes Valence Bond Theory & Hybridization

31. 1.In 1931, Linus Pauling, proposed that the outermost orbitals of an atom could be combined to form hybrid atomic orbitals. atomic orbitals – Sigma bond (  ) - The end-to-end overlapping of two orbitals. – Pi bond (  ) - The side-to-side overlapping of two p orbitals. Single bonds are made up of one sigma bond. Double bonds are made up of one sigma bond and one pi bond. Triple bonds are made up of one sigma bond and two pi bonds.

32. Sigma Bonds The electron density is concentrated on the internuclear line.

33. More Sigma Bonds

34. Pi Bonds The electron density is concentrated above & below the internuclear line.

35. Molecular Orbital Theory MO Theory is used to predict: 1) whether or not a molecule exists 2) its bond type (single, double, or triple) 3) its bond strength (triple are strongest) and 4) some of its properties (paramagnetic or diamagnetic).

36. MO Diagrams & Magnetic Properties Paramagnetism – Molecules with one or more UNPAIRED electrons in their MO diagram are attracted into a magnetic field. Diamagnetism – Molecules with PAIRED electrons in their MO diagram are weakly repelled from a magnetic field

37. Molecular Orbitals 1.Molecular orbitals, like atomic orbitals, have a definite shape and hold up to 2 electrons of opposite spin. 2.When 2 atomic orbitals combine & overlap, they form 2 molecular orbitals. 3.One of these molecular orbitals in a BONDING orbital and the other is an ANTIBONDING orbital.

38. Bonding molecular orbital - Electrons in this orbital spend most of their time in the region directly between the two nuclei. Antibonding molecular orbital - Electrons placed in this orbital spend most of their time away from the region between the two nuclei. Shown below are a sigma (  ) bonding molecular orbital and a sigma antibonding (  *), molecular orbital. * = Antibond

39. (a)Shows sigma MOs (b) & (c) Shows pi MOs

40. Molecular Orbital Diagram Small 2s-2p Interaction Use for Oxygen, Fluorine, & Neon (Reverse order for  2p and  2p for all other others)

41. Bond Order Bond order Type of Bond 0 Molecule doesn’t exist 1 Single bond 2 Double bond 3 Triple bond

42. Molecular Orbital Diagram Use with the Bond Order Formula to determine if a molecule exists, its type of bond, & whether it is paramagnetic or diamagnetic. (Small 2s-2p Interaction) (Large 2s-2p Interaction) Use for Oxygen, Fluorine, & Neon Use for Boron, Carbon, Nitrogen

43. Molecular Orbital Diagrams - Steps to determine bond type 1) Sum valence electrons for both atoms 2) Place arrows (represent electrons) in center of the diagram 3) Determine Bond Order = (# bonding e- - # antibonding e-)/2 4) ID Bond Type (1 = single, 2 = double, 3 = triple) 5) ID as paramagnetic (unpaired e- = attract magnet) or diamagnetic (all paired e- = repel magnet) Use for Oxygen, Fluorine, & Neon Use for Boron, Carbon, Nitrogen

44. http://www1.teachertube.com/viewVideo.ph p?video_id=57124 http://www1.teachertube.com/viewVideo.ph p?video_id=57124 A Little MO Theory 14 minutes