1. 1 Chapter 9 Covalent Bonding Molecular Compounds

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 So they share the electrons

4. 4 Covalent Bonds l Nonmetals hold onto their valence electrons l They don’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 their 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 a full outer orbital 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 a full outer orbital 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 a full outer orbital FF 8 Valence electrons

15. 15 Single Covalent Bond l A sharing of two valence electrons l Only nonmetals and Hydrogen form covalent bonds 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 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- let’s look at how water H 2 O 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 total) of electrons. l A triple bond is when atoms share three pair (6 total) 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 l movie movie O CO 8 valence electrons

34. 34 How to Draw Them l Add up all the (total have) valence electrons l Count up the total number of electrons (total want) to make all atoms happy l Subtract total have from total want 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 NH 3 Example l NH 3 – 1 N and 3 H’s l One N - has 5 valence electrons and wants 8 total l Each H’s - has 1 valence electron and wants 2 total l NH 3 has 5+3(1) = 8 Total Have l NH 3 wants 8+3(2) = 14 Total Want l (14-8)/2= 3 bonds l 4 atoms with 3 bonds N H H H

36. 36 NH H NH 3 Example l Draw in the 3 bonds l Account for all 8 valence electrons l Now everything has a full octet H

37. 37 HCN Example l With 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 (triple bond) 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 (triple bond) l Uses 8 electrons of 10 - 2 more to add l 2 must go on N to fill octet NHC

41. 41 Another Way of Indicating Bonds – Dots and Bars l Can use a bar (line) to indicate a bond l Called a Lewis Structure or Structural Formula l Each bar ( ) represents a bonded pair of valence electrons H HO = HHO

42. 42 Structural Examples H CN C O H H l C has 8 electrons because each line is 2 electrons l Same for N l The same for C in this case l Same for O with two lines and 4 electrons

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 Coordinate Covalent? 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 valid dot diagrams with the same connections are possible. l Consider the two ways to draw ozone (O 3 ). l Which one is it? l Does it go back and forth? l Is it a hybrid of both, like a mule? l Consider the three ways to draw NO 3 -

48. 48 It is a Mule! The hybrid bonds are shorter than single bonds, but longer than double bonds.

49. 49 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 l Movie Movie

50. 50 VSEPR l Based on the number of pairs of valence electrons both bonded and unbonded. l Unbonded pairs are called lone pairs. 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

51. 51 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

52. 52 4 Atoms Bonded to Central l Basic shape is called Tetrahedral. l Like a pyramid with a triangular base. l Same shape for everything with 4 pairs. C HH H H 109.5º

53. 53 Tetrahedral Shape

54. 54 3 Bonded - 1 Lone Pair NHH H N HH H <107.3º l Still basic tetrahedral but has an unshared electron pair. l Shape is called Trigonal Pyramidal.

55. 55 Trigonal Pyramidal

56. 56 2 Bonded - 2 Lone Pair OH H O H H <104.5º l Still basic tetrahedral but has 2 lone pair of electrons. l Shape is called Bent.

57. 57 Bent Shape Early Microwave Oven

58. 58 Consequences of H 2 O Polarity Microwave oven

59. 59 3 Atoms No Lone Pairs C H H O l If the atoms all are the same, the farthest you can move the shared electrons apart is 120º B F F F

60. 60 3 Atoms No Lone Pairs B F F F l The farthest you can move the electron pair apart is 120º. l Shape is flat and called Trigonal Planar. B F FF 120º

61. 61 Trigonal Planar

62. 62 2 Atoms No Lone Pairs l With three atoms the farthest they can get apart is 180º. l Shape called Linear. C O O 180º

63. 63 Molecular Shapes

64. 64 Two Types of Bonds l Sigma bonds (σ bond) forms from the overlap of orbitals l Sigma is centered between the atoms Pi bond (  bond) above and below atoms l Between adjacent p orbitals. l The two bonds of a double bond are one sigma and one pi bond.

65. 65 CC H H H H

66. 66 Bond Polarity l Electrons are pulled, as in a tug-of-war, between the atoms nuclei l The greater the difference in electronegativity between two bonded atoms, the more polar the bond. l (-)S N (+) Ionic bond l If the difference is great enough, electrons are transferred from the less electronegative atom to the more electronegative one to form a bond that has ionic character.- Ionic bond.

67. 67 Polar Bonds l In equal sharing (such as diatomic molecules), the bond that results is called a nonpolar covalent bond l When two different atoms are connected, the electrons may not be shared equally. l This unequal sharing is a polar covalent bond or simply a polar bond. l How do we measure how strong the atoms pull on electrons?

68. 68 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 Fig 9-15 gives electronegativity values l 0.0 - 0.4 Covalent nonpolar l 0.41 - 1.7 Covalent polar l >1.7 Ionic l Fig 9-16 shows bond character

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70. 70 Polar and Non-Polar l Like Dissolves Likes! l Polar solutions dissolve polar compounds. l Same for Non-Polar

71. 71 Electronegativity Chart

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74. 74 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  

75. 75 Polar Molecules Molecules with Charged Ends

76. 76 Polar Molecules l Polar Molecules have a positive and a negative end l This 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 to check polarity.

77. 77 Is it polar? l HF lH2OlH2O l NH 3 l CCl 4 l CO 2

78. 78 Bond Dissociation Energy l The energy required to break a CH bond is 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)

79. 79 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 = 3328 kJ/mol l reactants - products = 2562-3328 = -766kJ

80. 80 Endothermic and Exothermic l Endothermic reactions occur greater energy needed to break old bonds then is released forming new bonds. l Exothermic reactions occur when more energy is released making new bonds than is needed to break the old bonds.

81. 81 Intermolecular Forces What Holds Molecules to Each Other?

82. 82 Intermolecular Forces l “Inter” means between or among l So these are forces between molecules l These forces of attraction make solid and liquid molecular compounds possible. l Intermolecular forces are also called van der Waal’s forces. We will talk about three types. l With nonpolar molecules we have dispersion forces or London forces. These are temporary and weak forces.

83. 83 Dispersion Forces l Compressing a gas (colliding atoms) causes a temporary shift in the electron cloud and slight polarity. l Also depends on the # of electrons l More electrons mean stronger forces l Bigger molecules more electrons greater attraction Fluorine is a gas Bromine is a liquid Iodine is a solid

84. 84 Dipole Interactions l A stronger type are the dipole-dipole forces between polar molecules 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.

85. 85 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. HCl      

86. 86 Dipole Interactions    

87. 87 Hydrogen Bonding l A special dipole-dipole attraction force is caused by hydrogen being 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. MovieMovie

88. 88 Hydrogen Bonding H H O ++ -- ++ H H O ++ -- ++

89. 89 Hydrogen bonding H H O H H O H H O H H O H H O H H O H H O