Chapter 8 Covalent Bonding w videos 2017.pptx

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1. COVALENT BONDING 1

2. Chemical Bond A Quick Review…. A bond results from the attraction of nuclei for electrons All atoms are trying to achieve a stable octet IN OTHER WORDS the protons ( +) in one nucleus are attracted to the electrons (-) of another atom This is Electronegativity !! 2 Lewis Diagrams Made Easy (7min)

3. What did the atom of fluorine say to the atom of sodium? You complete me. 3

4. Three Major Types of Bonding Ionic Bonding forms ionic compounds transfer of valence e - Metallic Bonding Covalent Bonding forms molecules sharing of valence e - This is our focus this chapter 4

5. Ion ic Bonding Always formed between metal cations and non-metals anions The oppositely charged ions stick like magnets [METALS ] + [NON-METALS ] - Lost e - Gained e - 5

6. Metal l ic Bonding Always formed between 2 metals (pure metals) Solid gold, silver, lead, etc … 6

7. Covalent Bonding Pairs of e- are shared between 2 non-metal atoms to acquire the electron configuration of a noble gas. molecules 7

8. Covalent Bonding Occurs between nonmetal atoms which need to gain electrons to get a stable octet of electrons or a filled outer shell. nonmetals

9. How to Draw Lewis Structures: Five Easy Steps Drawing molecules (covalent) using Lewis Dot Structures Symbol represents the KERNEL of the atom (nucleus and inner electrons) dots represent valence electrons The ones place of the group number indicates the number of valence electrons on an atom. Draw a valence electron on each side (top, right, bottom, left) before pairing them. 9

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11. Always remember atoms are trying to complete their valence shell! “2 will do but 8 is great!” The number of electrons the atoms needs is the total number of bonds they can make. Ex. … H? O? F? N? Cl? C? one two one three one four 11

12. Draw Lewis Dot Structures You may represent valence electrons from different atoms with the following symbols x , , H or H or H x 12

13. Covalent bonding The atoms form a covalent bond by sharing their valence electrons to get a stable octet of electrons.(filled valence shell of 8 electrons) There are two electrons per bond, each atom donates one electron to the bond. Electron-Dot Diagrams of the atoms are combined to show the covalent bonds Covalently bonded atoms form MOLECULES

14. Methane CH 4 This is the finished Lewis dot structure Every atom has a filled valence shell How did we get here? OR 14

15. General Rules for Drawing Lewis Structures All valence electrons of the atoms in Lewis structures must be shown. Generally each atom needs eight electrons in its valence shell (except Hydrogen needs only two electrons and Boron needs only 6). Multiple bonds (double and triple bonds) can be formed by C, N, O, P, and S. Central atoms have the most unpaired electrons. Terminal atoms have the fewest unpaired electrons. 15

16. When carbon is one of you atoms, it will always be in the center Sometimes you only have two atoms, so there is no central atom Cl 2 HBr H 2 O 2 N 2 HCl We will use a method called ANS (Available, Needed, Shared) to help us draw our Lewis dot structures for molecules 16

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19. Sometimes multiple bonds must be formed to get the numbers of electrons to work out DOUBLE bond atoms that share two e- pairs (4 e-) O O TRIPLE bond atoms that share three e- pairs (6 e-) N N 19

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22. Let’s Practice H 2 Available = 2 valence e- Needed = H: 2, H:2 = 4 e- Shared = 4 - 2= 2 e- shared ÷ 2 = 1 bond DRAW 22

23. Let’s Practice CH 4 Available = 8 Needed = C: 8 H:2 , H:2, H:2, H:2 = 16 Shared = (N-A)16 – 8 = 8 e- shared ÷2 = 4 bonds DRAW 23 Carbon is the central atom because it has the most unpaired (single) valence electrons

24. “ANS” Method (available , needed, shared) Used to determine the number of bonds in a covalent molecule “A” is the total number of valence electrons from all atoms in the chemical formula “N” is the total number of valence electrons needed for an octet for all of the atoms in the chemical formula “S” is the number of valence electrons that must be shared (N > A). S ÷2 = number of bonds in molecule 24

25. Let’s Practice NH 3 A = total of 8 valence e- N = N:8 H:2, H:2, H:2 = 14 valence e- S = 14-8 = 6 shared e- ÷2 = 3 bonds DRAW 25

26. Let’s Practice CO 2 Available = C:4, O:6, O:6, = 16 Needed = C:8 O:8, O:8 total = 24 Shared = 24-16 = 8 valence e- must be shared = 4 bonds- put 2 valence e- in each bond. Remaining = (A-S) 16 – 8 = 8 not bonding DRAW – carbon is the central atom 26

27. Let’s Practice BCl 3 boron only needs 6 valence electrons, it is an exception. A = = 24 N = B-6, Cl-8, Cl-8 , Cl-8 30 S = 30-24 = 6 ÷ 2 = 3 bonds Remaining = 24 – 6 = 18 e- not bonding DRAW 27 Lewis Diagrams Made Easy (7min)

28. Naming Molecular Compounds (Covalent) Type III Nonmetal + nonmetal

29. T he C ova l ent B ond Sh a r i n g o f e l ectr o ns

30. Pr o p er t i e s of M o lec u l ar or C ova le n t C om p o u n d s Made f rom 2 or mo r e non­ m e t a l s C ons i st of mo l ecu l es not i ons

31. M o le c u la r F o r m u la s Sh ow t h e k in d s a n d nu m b er s of atoms pr e s e n t i n a mo l e c u l e of a c om p o und . Molecular Formula = H 2 O

32. H N H H NH 3 S t ructur a l f o r m u l a M o le c u la r f o r m u l a

33. Mo l e cu l ar For m ulas Examples C O 2 S O 3 N 2 O 5

34. Rules for Naming Mo l e cu l ar c o m po u n ds Th e mo s t “me t a lli c” nonme t al e l e m ent i s w r i t t en f i rst ( t he one that i s f urthe s t l e f t) Th e mo s t no n ­ m e t a l l i c of the t w o nonme t a l s is w r i t t en l ast i n the f or m u l a N O 2 not O 2 N A l l binary mo l e c u l ar co m po u nds end i n - ­ i de

35. Ion i c co m pou nd s u se c h arges to det e r m i ne the chemical f or m u l a The mo l ecu l ar co m pou n d‘s na m e tells you the number of each element in the chemical fo r mu l a. U ses pr e f i x es to te l l y o u the quantity of each element. Y ou ne e d to m e mo r i z e the pre f i x es ! Mo l e cu l ar c o m po u n ds

36. Pr e fi x es 1 mo n o­ 2 d i ­ 3 t r i­ 4 t e t r a­ 5 p e n ta­ 6 h e xa­ 7 h e p ta­ 8 o c ta­ 9 n o n a­ 10 d ec a ­ Me m o r i z e!

37. If the r e i s o n l y o ne o f the f i rst e l e m ent do n ot p ut (prefix) mon o ­ Example: carbon monoxide (not monocarbon monoxide) If the nonmetal starts with a vowel, drop the vowel ending from all prefixes except di and tri m ono x i de not m onoo x i de t e t r o x i de not te t rao x i de More M o le c u la r C o m p o u n d R ul e s

38. N 2 O 5 Mo l e cu l ar c o m po u n ds

39. N 2 O 5 Mo l e cu l ar c o m po u n ds d i

40. N 2 O 5 d i n i t r og e n Mo l e cu l ar c o m po u n ds

41. N 2 O 5 d i n i t r og e n Mo l e cu l ar c o m po u n ds p e n ta

42. N 2 O 5 d i n i t r og e n Mo l e cu l ar c o m po u n ds p e n taox i d e

43. N 2 O 5 d i n i t r og e n Mo l e cu l ar c o m po u n d Naming Practice p e n taox i d e

44. N 2 O 5 d i n i t r og e n p e n tox i d e Mo l e cu l ar c o m po u n ds d i n i tr og en p ent ox i d e

45. Mo l e cu l ar c o m po u n ds S u lf ur t r i o x i de

46. Mo l e cu l ar c o m po u n ds S u lf ur t r i o x i de S

47. Mo l e cu l ar c o m po u n ds S u lf ur t r i o x i de S

48. Mo l e cu l ar c o m po u n ds S u lf ur t r i o x i de S O 3

49. Mo l e cu l ar c o m po u n ds S u lf ur t r i o x i de S O 3 SO 3

50. Mo l e cu l ar c o m po u n ds CC l 4

51. Mo l e cu l ar c o m po u n ds CC l 4 mo n o c a r b on

52. Mo l e cu l ar c o m po u n ds CC l 4 mo n o c a r b on

53. Mo l e cu l ar c o m po u n ds CC l 4 c a r b on

54. Mo l e cu l ar c o m po u n ds CC l 4 t e t r a c a r b on

55. Mo l e cu l ar c o m po u n ds CC l 4 t e t r a c h lo r ide c a r b on

56. Mo l e cu l ar c o m po u n ds C arbon t et r ach l or i de CC l 4 t e t r a c h lo r ide c a r b on

57. Write molecular fo r m u l a s f o r t h e s e diph o s p h o r us pen t ox i de P 2 O 5 t r i sulfur h e x a f l o ur i de S 3 F 6 nitr o g en t r iio dide NI 3

58. H 2 O NH 3 C o mm o n N a m es

59. H 2 O NH 3 W at er A mm o n i a C o mm o n N a m es

60. Bond Types 3 Possible Bond Types: Ionic Non-Polar Covalent Polar Covalent 60

61. Use Electronegativity Values to Determine Bond Types Ionic bonds Electronegativity (EN) difference > 2.0 Polar Covalent bonds EN difference is between .21 and 1.99 Non-Polar Covalent bonds EN difference is < .20 Electrons shared evenly in the bond 61

62. “Ionic Character” “Ionic Character” refers to a bond’s polarity In a polar covalent bond, the closer the EN difference is to 2.0, the more POLAR its character The closer the EN difference is to .20, the more NON-POLAR its character 62

63. Place these molecules in order of increasing bond polarity using the electronegativity values on your periodic table HCl CH 4 CO 2 NH 3 N 2 HF 63 1 EN difference = 0 2 EN difference = 0.3 4 EN difference = 0.9 4 EN difference = .9 3 EN difference = 0.8 5 EN difference = 1.9

64. Polar vs. Nonpolar MOLECULES Sometimes the bonds within a molecule are polar and yet the molecule itself is non-polar 64

65. Nonpolar Molecules Molecule is Equal on all sides Symmetrical shape of molecule (atoms surrounding central atom are the same on all sides) H H H H C Draw Lewis dot first and see if equal on all sides 65

66. Polar Molecules Molecule is Not Equal on all sides Not a symmetrical shape of molecule (atoms surrounding central atom are not the same on all sides) Cl H H H C 66

67. H Cl  -  + Polar Molecule Unequal Sharing of Electrons 67

68. Cl Non-Polar Molecule Cl Equal Sharing of Electrons 68

69. H Cl Polar Molecule Not symmetrical H B 69

70. H H Non-Polar Molecule Symmetrical H B 70

71. H H O Water is a POLAR molecule ANY time there are unshared pairs of electrons on the central atom, the molecule is POLAR 71

72. Making sense of the polar non-polar thing BONDS Non-polar Polar EN difference EN difference 0 - .2 .21 – 1.99 MOLECULES Non-polar Polar Symmetrical Asymmetrical OR Unshared e - s on Central Atom 72

73. 5 Shapes of Molecules you must know! (memorize) 73

74. Copy this slide VSEPR – Valence Shell Electron Pair Repulsion Theory Covalent molecules assume geometry that minimizes repulsion among valence electrons. Shape of a molecule can be predicted from its Lewis Structure which shows all valence electrons 74

75. 1. Linear (straight line) Ball and stick model Molecule geometry X A X Unshared Pairs = 0 OR A X OR 75

76. 2. Bent Ball and stick model Lewis Diagram A X X .. Unshared Pairs = 1 or 2 76

77. 3. Trigonal Planar (flat) Ball and stick model Molecule geometry X A X X Unshared Pairs = 0 77

78. 4.Trigonal Pyramidal Ball and stick model Molecule geometry Shared Pairs = 3 Unshared Pairs = 1 78

79. 5.Tetrahedral Ball and stick model Molecule geometry Unshared Pairs = 0 79

80. I can describe the 3 inter molecular forces of covalent compounds and explain the effects of each force. 80

81. Like Dissolves Like (memorize this!) Polar substances WILL only dissolve in polar solvents Non-polar substances WILL only dissolve in nonpolar solvents   Non-polar substances WILL NOT dissolve in polar solvents and vice versa Example: acetone and styrofaom 81 Video – Styrofoam and acetone

82. Attractions within or inside molecules, also known as bonds. Ionic Covalent metallic Intra molecular attractions 82 Roads within a state

83. Attractions between molecules Hydrogen “bonding” Strong attraction between special polar molecules (F, O, N, P) Dipole-Dipole Result of polar covalent Bonds Induced Dipole (Dispersion Forces) Result of non-polar covalent bonds Inter molecular attractions 83 Interstate highways connect one state to another

84. More on inter molecular forces Hydrogen “Bonding” STRONG attraction between highly polar molecules HF, H 2 O, NH 3 Occurs between H of one molecule and N, O, F of another molecule Hydrogen “bond”  -  +  +  -  +  +  +  +  - 84 Hydrogen bonding 1 min

85. Why does Hydrogen “bonding” occur? Nitrogen, Oxygen and Fluorine are small atoms with strong nuclear charges powerful atoms N, O & F have very high electronegativities, these atoms strongly attract shared electrons in a bond Create very POLAR molecules 85

86. Dipole-Dipole Interactions WEAK inter molecular force Bonds have high EN differences forming polar covalent molecules, but not as high as those that result in hydrogen bonding. .21<EN<1.99 Partial negative and partial positive charges slightly attracted to each other. Result of polar covalent bonds Only occur between polar covalent molecules 86

87. Dipole-Dipole Interactions 87

88. Induced Dipole Attractions VERY WEAK inter molecular force Result of nonpolar covalent bonds Temporary partial poles are caused by a nearby polar covalent molecule disturbing the arrangement of electrons in the nonpolar molecule Occur briefly between NONPOLAR & POLAR molecules 88 Induced dipole video 30 sec

89. BOND STRENGTH IONIC COVALENT Hydrogen Dipole-Dipole Induced Dipole   intra molecular inter molecular Strongest Weakest 89

90. Intermolecular Forces affect chemical properties For example, strong intermolecular forces cause high boiling point Water has a high boiling point compared to many other liquids due to the hydrogen bonding between water molecules. 90

91. Predict which substance has the highest boiling point. HF NH 3 CO 2 91 19.5  ˚ C (67.1 F ; 292.6 K) −33.3 ˚ C   (−28.0  ˚ F ; 239.8 K) −56.6 ˚ C   (−69.6  ˚ F ; 216.6 K)

92. Why HF has the highest boiling point HF NH 3 CO 2 WHY? The H-F bond has the highest electronegativity difference SO HF is a highly polar molecule resulting in the Hydrogen bonding. T herefore HF needs the most energy to overcome the intermolecular forces and boil. 92

93. The End 93