1. 1 Covalent bonding And hybridization of electrons

2. 2 How does H 2 form? l The nuclei repel ++

3. 3 How does H 2 form? ++ l The nuclei repel l But they are attracted to electrons l They share the electrons

4. 4 Covalent bonds l Nonmetals hold onto their valence electrons. l They can’t give away electrons to bond. l Still want noble gas configuration. l Get it by sharing valence electrons with each other. l By sharing both atoms get to count the electrons toward noble gas configuration.

5. 5 Covalent bonding l Fluorine has seven valence electrons F

6. 6 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven FF

7. 7 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons FF

8. 8 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons FF

9. 9 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons FF

10. 10 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons FF

11. 11 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons FF

12. 12 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals FF

13. 13 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals FF 8 Valence electrons

14. 14 Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals FF 8 Valence electrons

15. 15 Single Covalent Bond l A sharing of two valence electrons. l Only nonmetals and Hydrogen. l Different from an ionic bond because they actually form molecules. l Two specific atoms are joined. l In an ionic solid you can’t tell which atom the electrons moved from or to.

16. 16 How to show how they formed l It’s like a jigsaw puzzle. l I have to tell you what the final formula is. l You put the pieces together to end up with the right formula. l For example- show how water is formed with covalent bonds.

17. 17 Water H O Each hydrogen has 1 valence electron Each hydrogen wants 1 more The oxygen has 6 valence electrons The oxygen wants 2 more They share to make each other happy

18. 18 Water l Put the pieces together l The first hydrogen is happy l The oxygen still wants one more H O

19. 19 Water l The second hydrogen attaches l Every atom has full energy levels H O H

20. 20 Multiple Bonds l Sometimes atoms share more than one pair of valence electrons. l A double bond is when atoms share two pair (4) of electrons. l A triple bond is when atoms share three pair (6) of electrons.

21. 21 Carbon dioxide l CO 2 - Carbon is central atom ( I have to tell you) l Carbon has 4 valence electrons l Wants 4 more l Oxygen has 6 valence electrons l Wants 2 more O C

22. 22 Carbon dioxide l Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short O C

23. 23 Carbon dioxide l Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short O C O

24. 24 Carbon dioxide l The only solution is to share more O C O

25. 25 Carbon dioxide l The only solution is to share more O C O

26. 26 Carbon dioxide l The only solution is to share more O CO

27. 27 Carbon dioxide l The only solution is to share more O CO

28. 28 Carbon dioxide l The only solution is to share more O CO

29. 29 Carbon dioxide l The only solution is to share more O CO

30. 30 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO

31. 31 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

32. 32 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

33. 33 Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

34. 34 How to draw them l Add up all the valence electrons. l Count up the total number of electrons to make all atoms happy. l Subtract. l Divide by 2 l Tells you how many bonds - draw them. l Fill in the rest of the valence electrons to fill atoms up.

35. 35 Examples l NH 3 l N - has 5 valence electrons wants 8 l H - has 1 valence electrons wants 2 l NH 3 has 5+3(1) = 8 l NH 3 wants 8+3(2) = 14 l (14-8)/2= 3 bonds l 4 atoms with 3 bonds N H

36. 36 NHH H Examples l Draw in the bonds l All 8 electrons are accounted for l Everything is full

37. 37 Examples l HCN C is central atom l N - has 5 valence electrons wants 8 l C - has 4 valence electrons wants 8 l H - has 1 valence electrons wants 2 l HCN has 5+4+1 = 10 l HCN wants 8+8+2 = 18 l (18-10)/2= 4 bonds l 3 atoms with 4 bonds -will require multiple bonds - not to H

38. 38 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N NHC

39. 39 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N l Uses 8 electrons - 2 more to add NHC

40. 40 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N l Uses 8 electrons - 2 more to add l Must go on N to fill octet NHC

41. 41 Another way of indicating bonds l Often use a line to indicate a bond l Called a structural formula l Each line is 2 valence electrons HHO = HHO

42. 42 Structural Examples H CN C O H H l C has 8 electrons because each line is 2 electrons l Ditto for N l Ditto for C here l Ditto for O

43. 43 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC

44. 44 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC

45. 45 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC

46. 46 How do we know if l Have to draw the diagram and see what happens. l Often happens with polyatomic ions and acids.

47. 47 Resonance l When more than one dot diagram with the same connections are possible. l NO 2 - l Which one is it? l Does it go back and forth. l It is a mixture of both, like a mule. l NO 3 -

48. 48 VSEPR l Valence Shell Electron Pair Repulsion. l Predicts three dimensional geometry of molecules. l Name tells you the theory. l Valence shell - outside electrons. l Electron Pair repulsion - electron pairs try to get as far away as possible. l Can determine the angles of bonds.

49. 49 VSEPR l Based on the number of pairs of valence electrons both bonded and unbonded. l Unbonded pair are called lone pair. l CH 4 - draw the structural formula l Has 4 + 4(1) = 8 l wants 8 + 4(2) = 16 l (16-8)/2 = 4 bonds

50. 50 VSEPR l Single bonds fill all atoms. l There are 4 pairs of electrons pushing away. l The furthest they can get away is 109.5º. CHH H H

51. 51 4 atoms bonded l Basic shape is tetrahedral. l A pyramid with a triangular base. l Same shape for everything with 4 pairs. C HH H H 109.5º

52. 52 3 bonded - 1 lone pair NHH H N HH H <109.5º l Still basic tetrahedral but you can’t see the electron pair. l Shape is called trigonal pyramidal.

53. 53 2 bonded - 2 lone pair OH H O H H <109.5º l Still basic tetrahedral but you can’t see the 2 lone pair. l Shape is called bent.

54. 54 3 atoms no lone pair C H H O l The farthest you can the electron pair apart is 120º

55. 55 3 atoms no lone pair C H H O l The farthest you can the electron pair apart is 120º. l Shape is flat and called trigonal planar. C H HO 120º

56. 56 2 atoms no lone pair l With three atoms the farthest they can get apart is 180º. l Shape called linear. C O O 180º

57. 57 Hybrid Orbitals Combines bonding with geometry

58. 58 Hybridization l The mixing of several atomic orbitals to form the same number of hybrid orbitals. l All the hybrid orbitals that form are the same. l sp 3 -1 s and 3 p orbitals mix to form 4 sp 3 orbitals. l sp 2 -1 s and 2 p orbitals mix to form 3 sp 2 orbitals leaving 1 p orbital. l sp -1 s and 1 p orbitals mix to form 4 sp orbitals leaving 2 p orbitals.

59. 59 Hybridization l We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry. l We combine one s orbital and 3 p orbitals. l sp 3 hybridization has tetrahedral geometry.

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62. 62 sp 3 geometry 109.5º l This leads to tetrahedral shape. l Every molecule with a total of 4 atoms and lone pair is sp 3 hybridized. l Gives us trigonal pyramidal and bent shapes also.

63. 63 How we get to hybridization l We know the geometry from experiment. l We know the orbitals of the atom l hybridizing atomic orbitals can explain the geometry. l So if the geometry requires a tetrahedral shape, it is sp 3 hybridized. l This includes bent and trigonal pyramidal molecules because one of the sp 3 lobes holds the lone pair.

64. 64 sp 2 hybridization l C 2 H 4 l double bond acts as one pair l trigonal planar l Have to end up with three blended orbitals l use one s and two p orbitals to make sp 2 orbitals. l leaves one p orbital perpendicular

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67. 67 Where is the P orbital? l Perpendicular The overlap of orbitals makes a sigma bond (  bond)

68. 68 Two types of Bonds l Sigma bonds from overlap of orbitals l between the atoms Pi bond (  bond) above and below atoms l Between adjacent p orbitals. l The two bonds of a double bond

69. 69 CC H H H H

70. 70 sp 2 hybridization l when three things come off atom l trigonal planar l 120º one  bond

71. 71 What about two l when two things come off l one s and one p hybridize l linear

72. 72 sp hybridization l end up with two lobes 180º apart. l p orbitals are at right angles makes room for two  bonds and two sigma bonds. l a triple bond or two double bonds

73. 73 CO 2 C can make two  and two  O can make one  and one  COO

74. 74 N2N2

75. 75 N2N2

76. 76 Polar Bonds l When the atoms in a bond are the same, the electrons are shared equally. l This is a nonpolar covalent bond. l When two different atoms are connected, the atoms may not be shared equally. l This is a polar covalent bond. l How do we measure how strong the atoms pull on electrons?

77. 77 Electronegativity l A measure of how strongly the atoms attract electrons in a bond. l The bigger the electronegativity difference the more polar the bond. l 0.0 - 0.5 Covalent nonpolar l 0.5 - 1.0 Covalent moderately polar l 1.0 -2.0 Covalent polar l >2.0 Ionic l Use table 12-3 Pg. 285

78. 78 How to show a bond is polar l Isn’t a whole charge just a partial charge  means a partially positive  means a partially negative l The Cl pulls harder on the electrons l The electrons spend more time near the Cl H Cl  

79. 79 Polar Molecules Molecules with ends

80. 80 Polar Molecules l Molecules with a positive and a negative end l 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.

81. 81 Is it polar? l HF lH2OlH2O l NH 3 l CCl 4 l CO 2

82. 82 Bond Dissociation Energy l The energy required to break a bond l C - H + 393 kJ C + H l We get the Bond dissociation energy back when the atoms are put back together If we add up the BDE of the reactants and subtract the BDE of the products we can determine the energy of the reaction (  H)

83. 83 Find the energy change for the reaction l CH 4 + 2O 2 CO 2 + 2H 2 O l For the reactants we need to break 4 C-H bonds at 393 kJ/mol and 2 O=O bonds at 495 kJ/mol= 2562 kJ/mol l For the products we form 2 C=O at 736 kJ/mol and 4 O-H bonds at 464 kJ/mol l = 3328 kJ/mol l reactants - products = 2562-3328 = -766kJ

84. 84 Intermolecular Forces What holds molecules to each other

85. 85 Intermolecular Forces l They are what make solid and liquid molecular compounds possible. l The weakest are called van der Waal’s forces - there are two kinds l Dispersion forces l Dipole Interactions –depend on the number of electrons –more electrons stronger forces –Bigger molecules

86. 86 Dipole interactions l Depend on the number of electrons l More electrons stronger forces l Bigger molecules more electrons Fluorine is a gas Bromine is a liquid Iodine is a solid

87. 87 Dipole interactions l Occur when polar molecules are attracted to each other. l Slightly stronger than dispersion forces. l Opposites attract but not completely hooked like in ionic solids.

88. 88 Dipole interactions l Occur when polar molecules are attracted to each other. l Slightly stronger than dispersion forces. l Opposites attract but not completely hooked like in ionic solids. HFHF  HFHF 

89. 89 Dipole Interactions    

90. 90 Hydrogen bonding l Are the attractive force caused by hydrogen bonded to F, O, or N. l F, O, and N are very electronegative so it is a very strong dipole. l The hydrogen partially share with the lone pair in the molecule next to it. l The strongest of the intermolecular forces.

91. 91 Hydrogen Bonding H H O ++ -- ++ H H O ++ -- ++

92. 92 Hydrogen bonding H H O H H O H H O H H O H H O H H O H H O