2. 1 Chapter 8 Covalent bonding

3. 2 I. Octet Rule l What is the Octet Rule? l The octet rule states that atoms lose gain or share electrons in to acquire a full set of 8 valence electrons Create a drawing of Mg and Cl, Al and Cl

4. 3 II. Covalent Bonds A Chemical Bond occurs when electron are shared. Hydrogen and Hydrogen A molecules are formed from the overlap of orbitals and sharing of electrons

5. 4 Covalent bonds a. Nonmetals hold onto their valence electrons. b. They can’t give away electrons to bond. Still want noble gas configuration. c. Get it by sharing valence electrons with each other. d. By sharing both atoms get to count the electrons toward noble gas configuration.

6. 5 Its all in the distance a. Too far no bond b. Too close electrons repel c. Just right and a molecule is born A molecules are formed from the overlap of orbitals and sharing of electrons

7. 6 Covalent clip

8. 7 III. Molecule l A. a covalently bonded compound. –1. Tend to occur between non metals that are close together on the periodic table. – a. diatomic molecules – occur naturally in nature a. this is a more stable arrangement. H H F F Br Br Cl Cl N N

9. 8

10. 9 I. Single Covalent Bond A. A sharing of two valence electrons. 1.Only nonmetals and Hydrogen. 2.Different from an ionic bond because they actually form molecules. 3. Two specific atoms are joined.

11. 10 How does H 2 form? l The nuclei repel ++

12. 11 When Atoms Combine to make Molecules Fig 8-1 Atoms contain both positive and negative charges. When they come Together they arrange themselves so that the attractive forges of opposite Charges is greater than the repulsive forces of like charges

13. 12 How does H 2 form? ++ l The nuclei repel l But they are attracted to electrons l They share the electrons

14. 13 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.

15. 14 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

16. 15 Water l Put the pieces together l The first hydrogen is happy l The oxygen still wants one more H O

17. 16 Water l The second hydrogen attaches l Every atom has full energy levels H O H

18. 17 l use acetate sheets to work on page 244 1-5 to create lewis structures l Use a square of acetate to show each atom (use vis a vis pens only) l Show the overlap of orbitals l Draw the outcome in your notebook-use structural formula (line to represent pair) l Circle the shared pairs

19. 18 Lewis structure l 1. PH 3 l 2. H 2 S l 3. HCl l 4. CCl 4 l 5. SiH 4

20. 19 Lewis structure l Use molecular model kit to build l 1. PH 3 l 2. H 2 S l 3. HCl l 4. CCl 4

21. 20 Covalent Bond Formation l Covalent bond forms by overlap of orbitals. l Two types of bonds  Sigma bond: all single bonds are sigma bonds  Pi bond: in multiple bonds: the first one is sigma, all other bonds are pi. There are  Single bonds  Multiple bonds (double and triple only)

22. 21 Two types of Bonds l Sigma bonds from overlap of orbitals along the axis connecting the nuclei between the atoms Pi bond (  bond): perpendicular overlap of p-orbitals above and below the axis connecting the atoms

23. 22 Sigma bond: s-s Orbital Overlap

24. 23 l Pg 247- # 12 a-e

25. 24 III. Multiple Bonds A. Sometimes atoms share more than one pair of valence electrons. B. A double bond is when atoms share two pair (4) of electrons. C. A triple bond is when atoms share three pair (6) of electrons.

26. 25 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

27. 26 Carbon dioxide l Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short O C

28. 27 Carbon dioxide l Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short O C O

29. 28 Carbon dioxide l The only solution is to share more O C O

30. 29 Carbon dioxide l The only solution is to share more O C O

31. 30 Carbon dioxide l The only solution is to share more O CO

32. 31 Carbon dioxide l The only solution is to share more O CO

33. 32 Carbon dioxide l The only solution is to share more O CO

34. 33 Carbon dioxide l The only solution is to share more O CO

35. 34 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

36. 35 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

37. 36 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

38. 37 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

39. 38 How to draw them l Add up all the valence electrons. l Count up the total number of electrons to make all atoms complete octet rule. 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.

40. 39 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

41. 40 NHH H Examples l Draw in the bonds l All 8 electrons are accounted for l Everything is full

42. 41 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

43. 42 HCN l Put in single bonds l Need 2 more bonds l Must go between C and N NHC

44. 43 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

45. 44 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

46. 45 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

47. 46 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

48. 47 l Build the molecules using the molecular model kit. l HCN l H 2 O l NH 3 l CH 4 l C 2 H 2 l PH 3 l H 2 S l http://www.youtube.com/watch?v=NYFE 5uslaNo http://www.youtube.com/watch?v=NYFE 5uslaNo

49. 48 l http://www.youtube.com/watch?v=NYFE 5uslaNo http://www.youtube.com/watch?v=NYFE 5uslaNo

50. 49 IV. Bond strength A. bond strength is the energy needed to break a covalent bond 1.Bond length and atom size help determine bond strength a. The shorter the bond length The greater the bond strength i. single bond – longest length ii. Double bond – medium length iii triple bond – shortest length WHICH HAS THE GREATEST BOND STRENGTH?

51. 50 B. Energy is needed to create and break a covalent bond. 1.Energy is released when a covalent bond forms 2. Bond dissociation energy- is needed to break a bond a. always a positive number i. It takes 159KJ/mol to break F 2 Would it take more or less to break N 2 - why?

52. 51 V. Chemical reaction energy A. Endothermic reaction – energy is needed 1. More energy is needed to break the bond than is needed to create new bond. AB + energy  A + B NH 4 SCN + Ba(OH) + energy ( freezes to wood because heat is pulled from the reaction

53. 52 B. Exothermic reaction – energy is released 1. more energy is released when bonds form than is needed to break bonds A + B  AB + energy CaCl + baking soda and water energy is released as heat Hw: pg 247 1-12

54. 53

55. 54 I. Naming compounds A. Two types 1. Ionic - metal and non metal or polyatomics. 2. Covalent- we will just learn the rules for 2 non-metals.

56. 55 Covalent compounds l Two words, with prefixes. l Prefixes tell you how many. l mono, di, tri, tetra, penta, hexa, Hepta, octa, nona, deca l First element whole name with the appropriate prefix, except mono. l Second element, -ide ending with appropriate prefix. l Practice

57. 56 Writing Formulas l Two sets of rules, ionic and covalent l To decide which to use, decide what the first word is. l If is a metal or polyatomic use ionic. l If it is a non-metal use covalent.

58. 57 PREFIXES l Mono - one l di - two l Tri- three l Tetra- four l penta- five l hexa- six l Hepta- seven l Octa - eight l nona - nine l deca - ten

59. 58 l CO 2 l CO l CCl 4 lN2O4lN2O4 l XeF 6 lN4O4lN4O4 l P 2 O 10 Naming Covalent Compounds

60. 59 Covalent compounds l The name tells you how to write the formula l Sulfur dioxide l diflourine monoxide l nitrogen trichloride l diphosphorus pentoxide l Work on naming ditto

61. 60

62. 61 I. Acids l Substances that produce H + ions when dissolved in water. l All acids begin with H. l Two types of acids: l Oxyacids l Non Oxyacids- Binary Acids

63. 62 A. Binary Acids l 1. Binary Acids( Hydrogen and one other element –A. hydro + element name + IC + acid –Examples –1. HCl- –2. HF- 3. HBr-

64. 63 B. Oxyacids l 1. oxyacids- (hydrogen + oxyanions) –a. name anion + (ic or ous) + acid i. use of ic or ous depends on the number of oxygen atoms in the oxyanion Examples: HNO 3- nitric acid HNO 2 –Nitrous acid H 3 PO 4- phosphoric acid H 2 PO 3 phosphorous acid HC 2 H 3 O 2 – acetic acid

65. 64 l H 2 SO 4 l sulfuric acid l H 2 SO 3 sulfurous acid

66. 65 Formulas for acids l hydrofluoric acid- HF l carbonic acid- H 2 CO 3 l hydrosulfuric acid l phosphorous acid

67. 66 Hydrates l Some salts trap water crystals when they form crystals. l These are hydrates. l Both the name and the formula needs to indicate how many water molecules are trapped. l In the name we add the word hydrate with a prefix that tells us how many water molecules.

68. 67 Hydrates l In the formula you put a dot and then write the number of molecules. Calcium chloride dihydrate = CaCl 2  2  Chromium (III) nitrate hexahydrate = Cr(NO 3 ) 3  6H 2 O

69. 68

70. 69 I. Resonance A. When more than one dot diagram with the same connections are possible. 1. SO 2 l Which one is it? l Does it go back and forth. l It is a mixture of both,

71. 70 20 Oct 97Bonding and structure (2)70 Sulfur Dioxide, SO 2 These equivalent structures are called: RESONANCE STRUCTURES. The proper Lewis structure is a HYBRID of the two. Each atom has OCTET..... BUT there is a +1 and -1 formal charge + —— + Rules 1-3  O—S —O

72. 71 Draw the resonance structure O 3 SO 3 SO 2

73. 72

74. 73 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC

75. 74 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC

76. 75 Coordinate Covalent Bond l When one atom donates both electrons in a covalent bond. l Carbon monoxide l CO OC

77. 76 How do we know l Have to draw the diagram and see what happens. l Often happens with polyatomic ions and acids. l Work on drawing lewis structures from packet

78. 77 BALLOON ACTIVITY l SHOW -LINEAR –TRIGONAL LANAR AND TETRAHEDRAL – BALLOONS l DRAW LEWIS DOT AND DEFINE LONE AND BONDED PAIRS –SHOW BALLOONS l HCL l H2O l NH4 l CH4

79. 78 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.

80. 79 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 (C)+ 4(1) (H)= 8 l wants 8 (c)+ 4(2) (H) = 16 l (16-8)/2 = 4 bonds l Wants-has /2 = how bonds form

81. 80 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

82. 81 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

83. 82 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.

84. 83 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.

85. 84 3 atoms no lone pair C H H O l The farthest you can the electron pair apart is 120º

86. 85 3 atoms no lone pair C H H O l The farthest you can have the electron pair apart is 120º. l Shape is flat and called trigonal planar. C H HO 120º

87. 86 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

88. 87

89. 88 l http://connected.mcgraw- hill.com/media/repository/protected_con tent/COMPOUND/50000118/10/23/mole cular_structure/sco/main.html?stateCod e=NJ http://connected.mcgraw- hill.com/media/repository/protected_con tent/COMPOUND/50000118/10/23/mole cular_structure/sco/main.html?stateCod e=NJ l PAGE 262- ONLINE TEXT ACTIVITY

90. 89 ABxEy Notation How can geometric structure in 3d be predicted given the numbers of valence electrons?

91. 90 A represents central atom B represents surrounding atoms E represents unshared pairs of valence electrons on Central Atom x and y - subscripts that represent number of surrounding atoms (B) and number of unshared pairs of valence electrons (E)on Central Atom

92. 91 20 Oct 97Bonding and structure (2)91 AB x E y notation a good way to distinguish between electron pair and molecular geometries is the AB x E y notation where: A - atom whose local geometry is of interest (typically the CENTRAL ATOM) B x - x atoms bonded to A E y - y lone pair electrons at A NH 3 is AB 3 E system  pyramidal

93. 92 Example: Water H2O Oxygen Central Atom Hydrogen surrounding atoms 2 surrounding atoms 2 pairs of unshared valence electrons Notation for ABxEy would be:

94. 93 Answer: AB 2 E 2 YOUR TURN…

95. 94 CO 2 CH 4 NH 3 PCl 5 SF 6 BF 3 BrF 5 IF 4

96. 95 VSEPR Electron pairs Bond Angles Underlying Shape 2180° Linear 3120° Trigonal Planar 4109.5° Tetrahedral 5 90° & 120° Trigonal Bipyramidal 6 90°Octagonal

97. 96 Actual shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 220linear 330trigonal planar 321bent 440tetrahedral 431trigonal pyramidal 422bent

98. 97 Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 550trigonal bipyrimidal 541See-saw 532T-shaped 523linear

99. 98 Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 660Octahedral 651Square Pyramidal 642Square Planar 633T-shaped 621linear

100. 99 Hybridization: The concept of orbital mixing The simplest compound formed between carbon and hydrogen is methane, CH 4. Therefore, it is proposed that in CH 4 and other carbon compounds, where the carbon is singly bonded to four other atoms, covalent bonding involves the hybridization of 2s and 2p orbitals of the carbon atom. Hybridization of CH 4 sp3 _

101. 100 HYBRIDIZATION process in which the valence electrons of an atom are rearranged to form four new Identical hybrid orbitals C 2 H 2

102. 101

103. 102 LESSON 11 Types of Covalent Bonding Types l Nonpolar l Polar

104. 103 Non polar and 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 electrons may not be shared equally. l This is a polar covalent bond. l How do we measure how strong the atoms pull on electrons?

105. 104 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 - 0.3 Covalent nonpolar l 0.3 - 1.7 Covalent polar l 1.7> Ionic

106. 105 Covalent: Non-polar 0>0.3 Polar: 0.3 < 1.7 Ionic > 1.7

107. 106

108. 107 Electronegativity A. Nonpolar Covalent Bonds 1. equal sharing of electrons 2. difference in electronegativity is (0 -.3 ) 3. examples ( H 2 O 2 N 2 Cl 2 )

109. 108 B. Polar Covalent Bonds 1. unequal sharing of electrons a. electrons spend more time by one of the atoms b. atoms take on a charge (+dipole or - dipole) c. difference in electronegativity is (.3 - 1.7 ) 2. examples (HBr H 2 S CBr 4 )

110. 109 C. Ionic Nature of Covalent Bonds 1. polar covalent bonds have an ionic nature a. when ionic nature is greater than 50% the bond is considered ionic 1) electronegativity difference is greater than 1.7

111. 110 l Work on polar/nonpolar ditto l Use electronegativity period table

112. 111 Polar Molecules Molecules with ends

113. 112 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 cancel out the effects of the polar bonds. Must determine geometry first.

114. 113 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  

115. 114 Covalent: Non-polar 0>0.3 Polar: 0.3 < 1.7 Ionic > 1.7

116. 115 Bond polarity concept map Bond Covalent Ionic

117. 116 Molecular bonds concept map l Create a molecular bond concept map Molecular Polarity l Non- polar molecules polar-molecules

118. 117 Types of polarity molecules l 1. Non polar molecules –Linear- a—a –Trigonal planar –Tetrahedral –Linear a—b—a 2. Polar molecule Triangular pyramidal Bent Linear- a--b

119. 118

120. 119 D. Polar and Nonpolar Molecules 1. polar molecules a. asymmetrical molecules with polar bonds b. examples (HCl NH 3 H 2 S) 2. nonpolar molecules a. molecules with nonpolar bonds 1) examples (H 2 Cl 2 N 2 ) b. symmetrical molecules with polar bonds

121. 120 E. Properties of Covalent Compounds 1. low b.p. and m.p. compared to ionic compounds 2. solid (often soft), liquid or gas at room temperature 3. may form crystalline solids a. Examples 1) solid CO 2 (dry ice) 2) sucrose C 12 H 22 O 11 (table sugar)

122. 121 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  

123. 122 Is it polar? l HF lH2OlH2O l NH 3 l CCl 4 l CO 2