2. 1 Molecular Geometry Mr. Bruder

3. 2 Covalent Bonding  A metal and a nonmetal transfer electrons –An ionic bond  Two metals just mix and don’t react –An alloy  What do two nonmetals do? –Neither one will give away an electron –So they share their valence electrons –This is a covalent bond

4. 3 Covalent bonding  Makes molecules –Specific atoms joined by sharing electrons  Two kinds of molecules:  Molecular compound –Sharing by different elements  Diatomic molecules –Two of the same atom –O 2 N 2

5. 4 Diatomic elements  There are 8 elements that always form molecules  H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2, and At 2  Oxygen by itself means O 2  The –ogens and the –ines  1 + 7 pattern on the periodic table

6. 5 1 and 7

7. 6 Molecular compounds  Tend to have low melting and boiling points  Have a molecular formula which shows type and number of atoms in a molecule  Not necessarily the lowest ratio  C 6 H 12 O 6  Formula doesn’t tell you about how atoms are arranged

8. 7 Polar Bonds  When the atoms in a bond are the same, the electrons are shared equally.  This is a nonpolar covalent bond.  When two different atoms are connected, the electrons may not be shared equally.  This is a polar covalent bond.  How do we measure how strong the atoms pull on electrons?

9. 8 Electronegativity  A measure of how strongly the atoms attract electrons in a bond.  The bigger the electronegativity difference the more polar the bond.  Use table 12-3 Pg. 285  0.0 - 0.4 Covalent nonpolar  0.5 - 1.0 Covalent moderately polar  1.0 -2.0 Covalent polar  >2.0 Ionic

10. 9 Covalent Bonding  Electrons are shared by atoms.  These are two extremes.  In between are polar covalent bonds.  The electrons are not shared evenly.  One end is slightly positive, the other negative.  Indicated using small delta 

11. 10 How to show a bond is polar  Isn’t a whole charge just a partial charge   means a partially positive   means a partially negative  The Cl pulls harder on the electrons  The electrons spend more time near the Cl H Cl  

12. 11 H - F ++ --

13. 12 H - F ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ --

14. 13 H - F ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- ++ -- + -

15. 14 Polar Molecules Molecules with ends

16. 15 Polar Molecules  Molecules with a partially positive end and a partially negative end  Requires two things to be true ¬ The molecule must contain polar bonds This can be determined from differences in electronegativity. ­ Symmetry can not cancel out the effects of the polar bonds. Must determine geometry first.

17. 16 Polar Molecules SS ymmetrical shapes are those without lone pair on central atom –T–Tetrahedral –T–Trigonal planar –L–Linear WW ill be nonpolar if all the atoms are the same SS hapes with lone pair on central atom are not symmetrical CC an be polar even with the same atom

18. 17 Is it polar?  HF H2OH2O  NH 3  CCl 4  CO 2  CH 3 Cl

19. 18 Intermolecular Forces What holds molecules to each other

20. 19 Chapter 11: States of Matter and Intermolecular Forces19 Intermolecular Forces Intramolecular forces determine such molecular properties as molecular geometries and dipole moments EOS Intermolecular forces determine the macroscopic physical properties of liquids and solids

21. Intermolecular Forces © 2009, Prentice-Hall, Inc. Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.

22. Intermolecular Forces © 2009, Prentice-Hall, Inc. Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

23. Intermolecular Forces © 2009, Prentice-Hall, Inc. Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces.

24. Intermolecular Forces © 2009, Prentice-Hall, Inc. van der Waals Forces Dipole-dipole interactions Hydrogen bonding London dispersion forces

25. + Intermolecular forces and melting/boiling point

26. Intermolecular Forces © 2009, Prentice-Hall, Inc. Ion-Dipole Interactions Ion-dipole interactions (a fourth type of force), are important in solutions of ions. The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

27. Chapter 11: States of Matter and Intermolecular Forces26 Dispersion Forces Dispersion forces are forces of attraction between an instantaneous dipole and an induced dipole … also called London forces after Fritz London who offered a theoretical explanation in 1928 The polarizability of an atom or molecule is a measure of the ease with which electron charge density is distorted by an external electrical field EOS Dipoles can be induced in molecules

28. London forces Instantaneous dipole:Induced dipole: Eventually electrons are situated so that tiny dipoles form A dipole forms in one atom or molecule, inducing a dipole in the other

29. Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces These forces are present in all molecules, whether they are polar or nonpolar. The tendency of an electron cloud to distort in this way is called polarizability.

30. 29 Dispersion force HH HH HH HH ++ -- HH HH ++ -- ++ 

31. Intermolecular Forces © 2009, Prentice-Hall, Inc. Factors Affecting London Forces The strength of dispersion forces tends to increase with increased molecular weight. Larger atoms have larger electron clouds which are easier to polarize.

32. Chapter 11: States of Matter and Intermolecular Forces31 Dispersion Forces The greater the polarizability of molecules, the stronger the intermolecular forces between them EOS

33. 32 Dispersion Force  Depends only on the number of electrons in the molecule  Bigger molecules more electrons  More electrons stronger forces F 2 is a gas Br 2 is a liquid I 2 is a solid

34. Chapter 11: States of Matter and Intermolecular Forces33 Dipole–Dipole Forces Dipole–dipole forces arise when permanent dipoles align themselves with the positive end of one dipole directed toward the negative ends of neighboring dipoles EOS A permanent dipole in one molecule can induce a dipole in a neighboring molecule, giving rise to a dipole– induced dipole force

35. 34 Dipole interactions  Occur when polar molecules are attracted to each other.  Slightly stronger than dispersion forces.  Opposites attract but not completely hooked like in ionic solids.

36. Intermolecular Forces © 2009, Prentice-Hall, Inc. Dipole-Dipole Interactions Molecules that have permanent dipoles are attracted to each other. –The positive end of one is attracted to the negative end of the other and vice- versa. –These forces are only important when the molecules are close to each other.

37. 36 Dipole interactions  Occur when polar molecules are attracted to each other.  Slightly stronger than dispersion forces.  Opposites attract but not completely hooked like in ionic solids. HFHF  HFHF 

38. 37 + - + - + - + - + - + - + - + - + - + -

39. Intermolecular Forces © 2009, Prentice-Hall, Inc. Dipole-Dipole Interactions The more polar the molecule, the higher is its boiling point.

40. 39 Hydrogen bonding  Are the attractive force caused by hydrogen bonded to F, O, or N.  F, O, and N are very electronegative so it is a very strong dipole.  They are small, so molecules can get close together  The hydrogen partially share with the lone pair in the molecule next to it.  The strongest of the intermolecular forces.

41. Chapter 11: States of Matter and Intermolecular Forces40 Hydrogen Bonds A hydrogen bond is an intermolecular force in which a hydrogen atom covalently bonded to a nonmetal atom in one molecule is simultaneously attracted to a nonmetal atom of a neighboring molecule EOS The strongest hydrogen bonds are formed if the nonmetal atoms are small and highly electronegative – e.g., N, O, F

42. Intermolecular Forces © 2009, Prentice-Hall, Inc. Hydrogen Bonding The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. We call these interactions hydrogen bonds.

43. 42 Hydrogen Bonding H H O ++ -- ++ H H O ++ -- ++

44. 43 Hydrogen bonding H H O H H O H H O H H O H H O H H O H H O

45. Intermolecular Forces © 2009, Prentice-Hall, Inc. Summarizing Intermolecular Forces

46. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

47. 46 VSEPR  Valence Shell Electron Pair Repulsion.  Predicts three dimensional geometry of molecules.  Name tells you the theory.  Valence shell - outside electrons.  Electron Pair repulsion - electron pairs try to get as far away as possible.  Can determine the angles of bonds.  And the shape of molecules

48. 47 VSEPR  Molecules take a shape that puts electron pairs as far away from each other as possible.  Have to draw the Lewis structure to determine electron pairs.  bonding  nonbonding lone pair  Lone pair take more space.  Multiple bonds count as one pair.

49. 48 VSEPR  The number of pairs determines –bond angles –underlying structure  The number of atoms determines –actual shape

50. 49 VSEPR Electron pairs Bond Angles Underlying Shape 2180° Linear 3120° Trigonal Planar 4109.5° Tetrahedral 5 90° & 120° Trigonal Bipyramidal 6 90°Octagonal

51. 50 Actual shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 220linear 330trigonal planar 321bent 440tetrahedral 431trigonal pyramidal 422bent

52. 51 Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 550trigonal bipyrimidal 541See-saw 532T-shaped 523linear

53. 52 Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 660Octahedral 651Square Pyramidal 642Square Planar 633T-shaped 621linear

54. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. What Determines the Shape of a Molecule? Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

55. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Electron Domains We can refer to the electron pairs as electron domains. In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain. The central atom in this molecule, A, has four electron domains.

56. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Electron-Domain Geometries These are the electron-domain geometries for two through six electron domains around a central atom.

57. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Electron-Domain Geometries All one must do is count the number of electron domains in the Lewis structure. The geometry will be that which corresponds to the number of electron domains.

58. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Molecular Geometries The electron-domain geometry is often not the shape of the molecule, however. The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

59. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Molecular Geometries Within each electron domain, then, there might be more than one molecular geometry.

60. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Linear Electron Domain In the linear domain, there is only one molecular geometry: linear. NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

61. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Trigonal Planar Electron Domain There are two molecular geometries: –Trigonal planar, if all the electron domains are bonding, –Bent, if one of the domains is a nonbonding pair.

62. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Nonbonding Pairs and Bond Angle Nonbonding pairs are physically larger than bonding pairs. Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.

63. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Multiple Bonds and Bond Angles Double and triple bonds place greater electron density on one side of the central atom than do single bonds. Therefore, they also affect bond angles.

64. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Tetrahedral Electron Domain There are three molecular geometries: –Tetrahedral, if all are bonding pairs, –Trigonal pyramidal if one is a nonbonding pair, –Bent if there are two nonbonding pairs.

65. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Trigonal Bipyramidal Electron Domain There are two distinct positions in this geometry: –Axial –Equatorial

66. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Trigonal Bipyramidal Electron Domain Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry.

67. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Trigonal Bipyramidal Electron Domain There are four distinct molecular geometries in this domain: –Trigonal bipyramidal –Seesaw –T-shaped –Linear

68. Molecular Geometries and Bonding © 2009, Prentice-Hall, Inc. Octahedral Electron Domain All positions are equivalent in the octahedral domain. There are three molecular geometries: –Octahedral –Square pyramidal –Square planar

69. 68 Molecular Orbitals  The overlap of atomic orbitals from separate atoms makes molecular orbitals  Each molecular orbital has room for two electrons  Two types of MO –Sigma ( σ ) between atoms –Pi ( π ) above and below atoms

70. 69 Sigma bonding orbitals  From s orbitals on separate atoms ++ s orbital +++ Sigma bonding molecular orbital

71. 70 Sigma bonding orbitals  From p orbitals on separate atoms p orbital Sigma bonding molecular orbital  

72. 71 Pi bonding orbitals  P orbitals on separate atoms      Pi bonding molecular orbital

73. 72 Sigma and pi bonds  All single bonds are sigma bonds  A double bond is one sigma and one pi bond  A triple bond is one sigma and two pi bonds.

74. 73 Hybrid Orbitals Combines bonding with geometry

75. 74 Hybridization  The mixing of several atomic orbitals to form the same number of hybrid orbitals.  All the hybrid orbitals that form are the same.  sp 3 -1 s and 3 p orbitals mix to form 4 sp 3 orbitals.  sp 2 -1 s and 2 p orbitals mix to form 3 sp 2 orbitals leaving 1 p orbital.  sp -1 s and 1 p orbitals mix to form 2 sp orbitals leaving 2 p orbitals.

76. 75 Hybridization  109.5º with s and p  Need 4 orbitals.  We combine one s orbital and 3 p orbitals.  Make sp 3 hybrid  sp 3 hybridization has tetrahedral geometry.

77. 76

78. 77

79. 78 sp 3 geometry 109.5º  This leads to tetrahedral shape.  Every molecule with a total of 4 atoms and lone pair is sp 3 hybridized.  Gives us trigonal pyramidal and bent shapes also.

80. 79 How we get to hybridization  We know the geometry from experiment.  We know the orbitals of the atom  hybridizing atomic orbitals can explain the geometry.  So if the geometry requires a 109.5º bond angle, it is sp 3 hybridized.

81. 80 sp 2 hybridization C2H4C2H4  double bond counts as one pair  Two trigonal planar sections  Have to end up with three blended orbitals  use one s and two p orbitals to make sp 2 orbitals.  leaves one p orbital perpendicular CC H HH H

82. 81

83. 82

84. 83 Where is the P orbital?  Perpendicular  The overlap of orbitals makes a sigma bond (  bond)

85. 84 Two types of Bonds  Sigma bonds (  from overlap of orbitals  between the atoms  Pi bond (  bond) between p orbitals.  above and below atoms  All single bonds are  bonds  Double bond is 1   and 1  bond  Triple bond is 1   and 2  bonds

86. 85 CC H H H H

87. 86 sp 2 hybridization  when three things come off atom  trigonal planar  120º  one  bond

88. 87 What about two  when two things come off  one s and one p hybridize  linear

89. 88 sp hybridization  end up with two lobes 180º apart.  p orbitals are at right angles  makes room for two  bonds and two sigma bonds.  a triple bond or two double bonds

90. 89 CO 2  C can make two  and two   O can make one  and one  COO

91. 90 N2N2

92. 91 N2N2