1. Properties of Molecular Substances Low melting & boiling points Low heats of fusion & vaporization High vapor pressure May be soft, as wax May be crystalline, as sugar (weak lattice, based on dipole-dipole or H bonding) Molecules are neutral. NEVER conduct.

2. Describe changes in chemical potential energy that accompany bond formation or bond breaking. Breaking a bond is endothermic. Making a bond is exothermic. Breaking a bond is endothermic. Making a bond is exothermic. Breaking   Making

3. Describe the relationship between stability & potential energy As PE , stability . As PE , stability .

4. Does the above equation represent an endothermic or exothermic process? How do you know? H + H  H 2 + energy Exothermic - energy term is on product side - bond formation releases energy

5. Does the above equation represent an endothermic or exothermic process? How do you know? H 2 + energy  H + H Endothermic - energy term is on reactant side - bond breaking absorbs energy

6. Why do atoms form bonds? To achieve the electron configuration of the nearest noble gas!

7. Electrons are shared Covalent Bonds

8. How do you identify a covalent formula? All Nonmetals in formula

9. Covalent bonds result from … The simultaneous attraction of electrons and two different nuclei.

10. What kinds of formulas do molecular substances have? Molecular, which give exact composition of molecule Empirical, which give lowest whole number ratio of atoms in molecule Sometimes they are the same. Otherwise, molecular is a whole number multiple of empirical H2OH2O C6H6C6H6 CH

11. Structural Formula Shows which elements & how many atoms of each. Shows connectivity or how atoms are linked Shows type of bond – single, double or triple H-O-HO=C=OO-N=O

12. Lewis Diagram Shows type & number of atoms. Shows connectivity Shows type of bonds Shows all nonbonding valence electrons, in addition to the bonding valence electrons Bonding electrons are in between two atoms!

13. Single Bond One electron pair or Two electrons Shared between atoms Represented by 2 dots or 1 dash between atoms

14. Double Bond Two electron pairs or Four electrons Shared between atoms Represented by 4 dots or 2 dashes between atoms

15. Triple Bond Three electron pairs or Six electrons Shared between atoms Represented by 6 dots or 3 dashes between atoms

16. Bond Energy Energy change that occurs when a bond is formed between two atoms. Symbol = D 0

17. Bond Energy As the number of electrons shared between 2 atoms increases, the attractive interactions increase, & the bond energy increases. Triple > Double > Single

18. Bond Length Distance between two bonded nuclei. The more shared electrons between 2 nuclei, the greater the attractive interactions, the shorter the bond length.

19.  Bond Overlap of two orbitals occurs on the line directly connecting the two nuclei. All single bonds are sigma bonds.

20.  Bond Overlap of two orbitals occurs above and below the line that directly connects the two nuclei. Double bonds consist of one  & one  bond. Triple bonds consist of one  & two  bonds.

21. Procedure for writing a Lewis structure for a molecular substance 1)Total up the valence electrons from all the atoms in the molecule. 2)Draw the skeleton, including a single bond between every atom. 3)Compare numbers: 1)# of electrons needed for each atom in the skeleton to have an octet/duet 2)# of electrons available after drawing skeleton 4)Distribute electrons 5)Verify by performing two validity checks

22. Lewis Diagram of H 2 1.Total # of valence electrons = 2 2.Skeleton: H – H 3.Compare: 0 electrons needed & 0 electrons available 4.No electrons to distribute 5.Verify.

23. Lewis Diagram of Cl 2 1.# of valence electrons = 14 2.Skeleton: Cl – Cl 3.Compare: Need 12, Have 12 4.Distribute 5.Verify

24. Bonding vs. Nonbonding electron pairs Bonding electrons are located BETWEEN two atoms. Nonbonding electrons are located on one atom only.

25. Lewis Diagram of O 2 1.# of valence electrons = 12 2.Skeleton: O - O 3.A) Compare: Need 12, Have 10, deficient by 2, add 1 bond B) O  O. Compare: Need 8, have 8. 4.Distribute 5.Verify

26. Lewis Diagram of N 2 1.# of valence electrons = 10 2.Skeleton: N – N 3.A) Compare: Need 12, Have 8, deficient by 4, add 2 bonds B) N  N. Compare: Need 4, have 4. 4.Distribute 5.Verify

27. Lewis Diagram of H 2 O 1.# of valence electrons = 8 2.Skeleton: H – O – H 3.Compare: Need 4, Have 4 4.Distribute 5.Verify

28. Lewis Diagram of NH 3 1.# of valence electrons = 8 2.Skeleton: 3.Compare: Need 2, Have 2 4.Distribute 5.Verify H H – N – H -

29. Lewis Diagram of CH 4 1.# of valence electrons = 8 2.Skeleton: H – C – H 3.Compare: Need 0, Have 0 4.Distribute (nothing) 5.Verify H H - -

30. Lewis Diagram of CCl 4 1.# of valence electrons = 32 2.Skeleton: Cl – C – Cl 3.Compare: Need 24, Have 24 4.Distribute 5.Verify Cl - -

31. Lewis Diagram of C 2 H 4 1.# of valence electrons = 12 2.Skeleton: H – C – C - H 3.A) Compare: Need 4, Have 2. Add 1 bond. 4.Distribute (nothing) 5.Verify H - - H – C  C - H H - - B) Compare: Need 0, Have 0

32. Lewis Diagram of C 2 H 2 1.# of valence electrons = 10 2.Skeleton: H – C – C - H 3.A) Compare: Need 8, Have 4. Add 2 bonds. B) Compare: Need 0, Have 0 4.Distribute (nothing) 5.Verify H – C  C - H

33. Lewis Diagram of CO 2 1.# of valence electrons = 16 2.Skeleton: O – C - O 3.A) Compare: Need 16, Have 12, deficient by 4, add 2 bonds B) O  C  O. Compare: Need 8, have 8. 4.Distribute 5.Verify

34. What are Resonance Structures? Sometimes, more than one valid Lewis structure can be written for a molecule. For CO 2 :

35. Resonance Structures The atoms are in the same location. The electrons are distributed differently.

36. What are the three general ways the octet rule breaks down? 1.Molecules with an odd # of electrons can never satisfy the octet rule for all their atoms. (NO, NO 2, ClO 2 ) 2.Some molecules have an atom with less than an octet. (BF 3, BeH 2 ) 3.Some molecules have an atom with more than an octet. (PCl 5, SF 6 )

37. Molecules with odd # of electrons 1.Consider NO 5 + 6 = 11 valence electrons 2.Skeleton: N – O 3.A) Compare: the N needs 6 & the O needs 6 for a total of 12. Have only 9 available. Add one bond. B) N = O now each atom needs 4 for a total of 8. Have only 7 available. 4.Distribute: :N = O: 5.Verify... The N atom has only 7 valence e -

38. Molecules with odd # of electrons 1.Consider NO 2. 5 + 2(6) = 17 electrons 2.Skeleton: O – N – O 3.A) Compare: each O needs 6 & N needs 4. Need 16 total. Have 13 available. Add 1 bond. B) O = N – O Compare: Need 4 + 2 + 6 or 12 electrons. Have 11 available. 4.Distribute: :O = N – O: 5.Verify....... The N atom has only 7 e -

39. Molecules with odd # of electrons 1.Consider ClO 2 7 + 2(6) = 19 valence e - 2.Skeleton is O – Cl – O 3.Compare: Need 6 + 4 + 6 = 16 e - Have: 19 – 4 = 15. Deficient by 1 e -. Can’t fix. 4.Distribute :O – Cl – O: 5.Verify........... The Cl atom has only 7 valence e - !

40. Molecules with atoms that have less than an octet Occurs in molecules with Be and B. Be likes to have 4 valence electrons in molecules. B likes to have 6 valence electrons in molecules.

41. Molecules with atoms that contain more than an octet Only the central atom can have more than an octet. And only if it belongs in rows 3-7 of the PT. Consider PF 5. 5 + 5(7) = 40 valence e -. Since P is in row 3 it has empty d orbitals available which can be used for bonding. Skeleton: Distribute the remaining 30 e - by placing 6 e - on each of the 5 F atoms. F F F P F FF F The P has > than an octet.

42. Polyatomic Ions A group of covalently bonded atoms that has gained or lost electrons and hence acquired a charge

43. Lewis Diagrams of Polyatomic Ions When calculating the total # of valence electrons, you must adjust for the charge of the ion. –Total up the electrons contributed from each atom. –Add 1 electron for each negative charge. Or –Subtract 1 electron for each positive charge.

44. Lewis Diagram of NH 4 +1 1.# of valence electrons = 9 – 1 = 8 2.Skeleton: 3.Compare: Need 0, Have 0 4.Distribute (nothing) 5.Verify H – N – H HHHH - - +1

45. Lewis Diagram of OH -1 1.# of valence electrons = 6 + 1 + 1 = 8 2.Skeleton: [O – H] -1 3.Compare: Need 6, Have 6 4.Distribute: [:O - H] -1 5.Verify :

46. Determining Molecular Shape from the Lewis Structure 1.Count up the number of electron domains on the central atom. –Single, double, & triple bonds each count as ONE domain. –Lone electron pairs (or even a lone singleton) counts as ONE domain. 2.Count up the number of atoms that are bonded to the central atom. 3.Compare these two numbers to get the shape.

47. Shapes on the Regents Exam # of electron domains on central atom # of atoms bonded TO the central atom ShapeExample 22LinearCO 2 44Tetrahedral, 109  CH 4 43Trigonal Pyramid, 107  NH 3 42Bent, 105  H2OH2O

48. Additional Shapes # of electron domains on central atom # of atoms bonded TO the central atom ShapeExample 33*Trigonal Planar, bond angle = 120  BF3, BH3 55*Trigonal Bipyramid PF5 54See-Saw 53T-Shape 52Linear 66*OctahedralSF6 65Square Pyramid 64Square Planar

49. Tetrahedral Molecule

50. Trigonal Pyramidal Molecule

51. Bent Molecule

52. Shape of CH 4 1.Inspect central atom in Lewis diagram: 2.Central atom has - 4 electron domains - 4 bonded atoms 3.Shape is tetrahedral with 109  bond angles

53. Shape of NH 3 1.Inspect central atom in Lewis diagram: 2.Central atom has - 4 electron domains - 3 bonded atoms 3.Shape is trigonal pyramid with 107  bond angle Why isn’t the bond angle 109  ? Because the lone pair on the N atom spreads out & squeezes the bonding pairs together.

54. Shape of H 2 O 1.Inspect central atom in Lewis Diagram: 2.Central atom has - 4 electron domains - 2 atoms bonded to it 3.Shape is Bent with a 105  angle Why isn’t the bond angle 109  ? Because the two lone pairs on the O spread out more than the 2 bonding pairs and squeeze the bonding pairs together.

55. Shape of BF 3 1.Inspect central atom in Lewis Diagram: 2.Central atom has - 3 electron domains - 3 atoms bonded to it 3.Shape is Trigonal Planar with a 120  angle Recall: B is an exception to the octet rule! or

56. Shape of BeF 2 1.Inspect central atom in Lewis Diagram: 2.Central atom has - 2 electron domains - 2 atoms bonded to it 3.Shape is Linear with a 180  angle Recall: Be is an exception to the octet rule!

57. Shape of SO 2 1.Inspect central atom in Lewis Diagram: 2.Central atom has - 3 electron domains - 2 atoms bonded to it 3.Shape is Bent with a 120  angle

58. Shape of PF 5 1.Inspect central atom in Lewis Diagram: 2.Central atom has - 5 electron domains - 5 atoms bonded to it 3.Shape is Trigonal Bipyramid F F F P F FF F The P is allowed to have an “expanded octet” because it has empty 3d orbitals that can hold valence electrons.

59. PF 5 Two kinds of F atoms in PF 5 : Axial – set of two Equatorial – set of three The axial F atoms have 3 nearest neighbors at 90 . The equatorial F atoms have 2 nearest neighbors at 90  & 2 nearest neighbors at 120 . The equatorial F atoms are less crowded than the axial F atoms!

60. Shape of SF 6 Inspect central atom in Lewis diagram. Central S atom has –6 electron domains –6 atoms bonded to it Shape is octahedral! The S is allowed to have an “expanded octet” because it has empty 3d orbitals that can hold valence electrons.

61. Shape of all diatomics? Linear 2 points make a line!

62. Shape of NH 3, NF 3, PH 3, etc? Trigonal Pyramid

63. Shape of H 2 O, H 2 S, H 2 Se, etc.? Bent, with bond angle of 105 

64. Shape of CH 4, CCl 4, etc.? Tetrahedral

65. Polar Bond Bond has poles – the ends are different! Bond has a permanent partial separation of charge

66. Polar Bond Electronegativity Difference is 0.5 to 1.7

67. Nonpolar Bond No poles. Ends are the same. Symmetric electron cloud.

68. Nonpolar Bond Electronegativity difference = 0 to 0.5 NO separation of charge in bond.

69. Ionic Bond Electronegativity difference  1.7

70. Ionic Bond Full Separation of Charge At least +1 and -1.

71. Ionic Formula has a metal and a nonmetal

72. Molecular Polarity Depends on how the atoms are arranged in the molecule. Some molecules which contain polar bonds are nonpolar overall. Common. Some molecules which contain nonpolar bonds are polar. Less common.

73. WATER, H 2 O Polar. Water is bent. The O end is a bit negative & the H end is a bit positive.

74. Bonding Results from attractions between nucleus on 1 atom & electrons on another atom.

75. Making a bond … releases energy.

76. Breaking a bond … Absorbs energy.

77. Covalent Compounds are (almost all) Molecular compounds.

78. N2N2 Triple Bond

79. O2O2 Double Bond

80. F 2, Cl 2, Br 2, I 2 Single Bond

81. When bonds are made, energy is … released.

82. When bonds are broken, energy is … absorbed.

83. As the energy of a system , the stability generally … increases.

84. Exothermic System releases energy. Its energy level goes down.

85. Properties of Molecular Substances 1)Soft 2)Low melting point & low boiling point 3)Does not conduct electricity in any phase 4)Does not dissolve in water 5)React slowly

86. Nonpolar Bonds Electrons are shared equally between the two atoms

87. Molecular Polarity Molecule must contain polar bonds and they must be arranged asymmetrically.

88. Molecular Polarity Depends on the shape (Bent & Pyramidal are polar Linear & tetrahedral, polarity depends on composition.)

89. Nonpolar Molecules 1)Noble gas atoms (kickballs) 2) 7 Diatomic Elements (footballs) 3)C X H Y or pure hydrocarbons. 4)Larger molecules that have high symmetry

90. Polar Molecules If it’s not one of the 4 easy categories of nonpolar molecules, it is a polar molecule!

91. Nonpolar Molecules Weak Intermolecular Forces (Dispersion or Van der Waals) Low boiling points & melting points Tend to be gases

92. Polar Molecules Intermolecular Forces are Dipole- dipole forces Stick together better than nonpolar molecules Tend to have higher melting points, boiling points, H f, & H v than nonpolar substances.

93. Coordinate Covalent Bond Covalent bond where both electrons in the bond are donated by 1 atom.

94. Coordinate Covalent Bond H H:N:H H H + H +..  + No electrons to contribute. 2 electrons to contribute. 4 identical N-H bonds!

95. Coordinate Covalent Bond Both electrons in the bond are donated by the same atom.

96. Compound that exhibits both covalent & ionic bonding Compound that contains a polyatomic ion.

97. Lewis Diagram for NaOH [:O - H] -1 : [Na] +1 Metal cation Polyatomic anion

98. Lewis Diagram for NH 4 Cl Polyatomic cation Nonmetal anion H – N – H HHHH - - +1 :Cl: -1 :

99. Lewis Diagram for NH 4 OH Polyatomic cation Polyatomic anion H – N – H HHHH - - +1 [:O - H] -1 :

100. How to calculate the polarity of a bond Subtract the electronegativities of the 2 atoms.

101. To determine molecular polarity Use the symmetry & conmposition of the molecule to help you. Don’t get a number for it. Just polar or nonpolar overall.

102. Low symmetry shapes: POLAR Bent molecules Pyramids, either trigonal or square base See-Saw T-shape

103. High symmetry shapes Linear Trigonal PLANAR Tetrahedrons Trigonal BIpyramids Octahedrons Square Planar

104. For a high symmetry shape to be NONPOLAR All the ends or corners have to match. The corners or ends have to be the same element.

105. Example of Symmetry & Composition CX 4 is nonpolar. Corners match! CXY 3 and CX 2 Y 2 are polar. Corners don’t match!

106. Four network covalent substances are … C dia, C graph, SiO 2, & SiC

107. Network Covalent Substances form … Crystal Lattices!

108. Properties of network covalent substances High melting & boiling points Hard Brittle Nonconductors

109. Network Covalent substances have … Strong directional covalent bonds