CHM 2045 Molecular Geometry & Chemical Bonding Chapter 10

1. CHM 2045 Molecular Geometry & Chemical Bonding Chapter 10

2. Using Lewis Theory to PredictMolecular Shapes • Lewis theory says that these regions of electron groups should repel each other • This idea can then be extended to predict the shapes of molecules • the position of atoms surrounding a central atom will be determined by where the bonding electron groups are • the positions of the electron groups will be determined by trying to minimize repulsions between them

3. VSEPR Theory • Electron groups around the central atom will be most stable when they are as far apart as possible – we call this valence shell electron pair repulsion theory • The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule

4. Electron Group Geometry • There are five basic arrangements of electrongroups around a central atom • Each of these five basic arrangements results in five different basic electron geometries • For molecules that exhibit resonance, it doesn’t matter which resonance form you use – the electron geometry will be the same

5. Linear Electron Geometry • When there are two electron groups around the central atom, they will occupy positions on opposite sides of the central atom • This results in the electron groups taking a linear geometry • The bond angle is 180°

6. Trigonal Planar Electron Geometry • When there are three electron groups around the central atom, they will occupy positions in the shape of a triangle around the central atom • This results in the electron groups taking a trigonal planar geometry • The bond angle is 120°

7. Tetrahedral Electron Geometry • When there are four electron groups around the central atom, they will occupy positions in the shape of a tetrahedron around the central atom • This results in the electron groups taking a tetrahedral geometry • The bond angle is 109.5°

8. Tetrahedral Geometry

9. Trigonal Bipyramidal Electron Geometry • When there are five electron groups around the central atom, they will occupy positions in the shape of two tetrahedrathat are base-to-base with the central atom in the center of the shared bases. • This results in the electron groups taking a trigonalbipyramidal geometry

10. Trigonal Bipyramidal Geometry

11. Octahedral Electron Geometry • When there are six electron groups around the central atom, they will occupy positions in the shape of two square-base pyramids that are base-to-base with the central atom in the center of the shared bases • This results in the electron groups taking an octahedral geometry • All positions are equivalent • The bond angle is 90°

12. Octahedral Geometry

13. Octahedral Geometry

14. Molecular Geometry • The actual geometry of the molecule may be different from the electron geometry • When the electron groups are attached to atoms of different size, • when the bonding to one atom is different than the bonding to another, • When there are lone pairs since they occupy space on the central atom, but are not “seen” as points on the molecular geometry

15. The Effect of Lone Pairs • Lone pair groups “occupy more space” on the central atom • Relative sizes of repulsive force interactions is Lone Pair – Lone Pair > Lone Pair – Bonding Pair > Bonding Pair – Bonding Pair • This affects the bond angles, making the bonding pair – bonding pair angles smaller than expected

16. Bond Angle Distortion from Lone Pairs

17. Bond Angle Distortion from Lone Pairs

19. Molecular Geometry Example 1 Which molecule or ion has a trigonal planar molecular geometry? a) PCl3 b) AsF3 c) HCN d) HCCH e) CO32–

20. Molecular Geometry Example 2 Which molecule or ion has a trigonal pyramidal molecular geometry? a) CO32– b) SO3 c) BF3 d) C2H4 e) SO32–

21. Molecular Geometry Example 3 What is the bond angle in a linear molecule or ion? a) 90° b) 109° c) 120° d) 180°

22. Molecular Geometry Example 4 The approximate H—C—C bond angle in ethane, C2H6, is a) 60°. b) 90°. c) 109°. d) 120°. e) 180°.

23. Molecular Geometry Example 5 What is the H—O—H bond angle in water? a) 180° b) 120° c) 90° d) 109° e) slightly less than 109°

24. Dipole moment ,the measure of charge separation • A molecule has a dipole moment when there is charge separation • The strength of the dipole moment depends on the difference in the electronegativity of the atoms in the molecule

25. Dipole moment ,the measure of charge separation • Dipole moment can be shown with an arrow with a cross on one end: • An arrow indicates the direction in which the electrons concentrate.

26. Molecule Polarity The O─C bond is polar. The bonding electrons are pulled equally toward both O ends of the molecule. The net result is a nonpolar molecule.

27. Molecule Polarity The H─O bond is polar. Both sets of bonding electrons are pulled toward the O end of the molecule. The net result is a polar molecule.

28. Predicting Polarity of Molecules 1. Draw the Lewis structure and determine the molecular geometry 2. Determine whether the bonds in the molecule are polar a) if there are no polar bonds, the molecule is nonpolar 3. Determine whether the polar bonds add together to give a net dipole moment

29. Decide whether the following molecules Are polar Trigonal Bent Trigonal Planar 2.5 1. polar bonds, N-O 2. asymmetrical shape 1. polar bonds, all S-O 2. symmetrical shape polar nonpolar

30. Polar or Nonpolar Example 6 Which of the following molecule is polar? a) CF4 b) SO2 c) CS2 d) C2H4 e) C6H6

31. Polar or Nonpolar Example 7 Which of the following compounds is nonpolar? a) XeF2 b) HCl c) SO2 d) H2S e) N2O

32. Dipole Moment Example 8 Which of the following molecules has a dipole moment? a) NF3 b) CCl4 c) SiCl4 d) SF6 e) BF3

33. Dipole Moment Example 9 For which molecule or ion does the nitrogen atom have the positive end of the dipole moment? a) NH4+ b) Ca3N2 c) HCN d) AlN e) NO

34. Valence Bond Theory • Linus Pauling and others applied the principles of quantum mechanics to molecules • They reasoned that bonds between atoms would occur when the orbitals on those atoms interacted to make a bond • The kind of interaction depends on whether the orbitals align along the axis between the nuclei, or outside the axis

35. Valence Bond Theory • Linus Pauling and others applied the principles of quantum mechanics to molecules • They reasoned that bonds between atoms would occur when the orbitals on those atoms interacted to make a bond • The kind of interaction depends on whether the orbitals align along the axis between the nuclei, or outside the axis

36. Valence Bond Theory – Hybridization • One of the issues that arises is that the number of partially filled or empty atomic orbitals did not predict the number of bonds or orientation of bonds • C = 2s22px12py12pz0 would predict two or three bonds that are 90° apart, rather than four bonds that are 109.5° apart

37. Hybridization • To adjust for these inconsistencies, it was postulated that the valence atomic orbitals could hybridize/mixbefore bonding took place • one hybridization of C is to mix all the 2s and 2porbitals to get four orbitals that point at the corners of a tetrahedron

38. Unhybridized C Orbitals Predict the Wrong Bonding & Geometry

39. Hybridization • Some atoms hybridize their orbitals to maximize bonding • Hybridizing is mixing different types of orbitals in the valence shell to make a new set of degenerate orbitals • sp, sp2, sp3, sp3d, sp3d2 • Same type of atom can have different types of hybridization • C = sp, sp2, sp3

40. Hybrid Orbitals • The number of atomic orbitals combined = the number of hybrid orbitals formed • The number and type of standard atomic orbitals combined determines the shape of the hybrid orbitals

41. Orbital Diagram of thesp3 Hybridization of C

42. Methane Formation with sp3 C

43.   sp3 Hybridized AtomsOrbital Diagrams • Place electrons into hybrid and unhybridizedvalence orbitals as if all the orbitals have equal energy • Lone pairs generally occupy hybrid orbitals Unhybridized atom sp3 hybridized atom C        2p 2sp3 2s N        2p 2sp3 2s

44. Bonding with Valence Bond Theory • According to valence bond theory, bonding takes place between atoms when their atomic or hybrid orbitals “overlap” • To interact, the orbitals must either be aligned along the axis between the atoms, or • The orbitals must be parallel to each other and perpendicular to the interatomic axis

45. Types of Bonds • A sigma (s) bondresults when the interacting atomic orbitals point along the axis connecting the two bonding nuclei • A pi (p) bondresults when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei • The interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore sbonds are stronger than p bonds

47. Orbital Diagrams of Bonding • “Overlap” between a hybrid orbital on one atom with a hybrid or nonhybridized orbital on another atom results in a s bond • “Overlap” between unhybridizedp orbitals on bonded atoms results in a p bond

48.   sp2 Hybridized AtomsOrbital Diagrams Unhybridized atom sp2 hybridized atom C 3 s 1 p        2p 2sp2 2p 2s N 2 s 1 p        2p 2sp2 2p 2s Tro: Chemistry: A Molecular Approach, 2/e

49. sp2 hybridization in Ethylene • Two sp2-hybridized orbitals overlap to form a s bond • p orbitals overlap side-to-sideto formation a pi () bond • sp2–sp2s bond and 2p–2pbond result in sharing four electrons and formation of C-C double bond • Electrons in the  bond occupy regions on either side of a line between nuclei

50. sp2 hybridization in Ethylene, C 2H4