1. BONDING AND VSEPR THEORY STRUCTURES OF SOLIDS AND LIQUIDS Intermolecular Attractions

2. CHEMICAL BONDS – OCCUR WHEN ELECTRONS ARE TRANSFERRED OR SHARED BY ELEMENTS SO THAT THEY EACH BECOME MORE STABLE Electron Dot (Lewis) Diagrams Explain Chemical Bonding

3. How many electrons do most atoms want in their outer energy level to be stable? 1. 0 2. 1 3. 2 4. 6 5. 8 6. 18

4. Bonds that form between two nonmetals are which type? 1. Ionic 2. Covalent 3. Metallic

5. Drawing Electron Dot Diagrams Electrons usually stay in pairs when bonded. Bonding pairs – pair of electrons that form the bond - can be represented as a line segment Lone (or unbonded) pairs – pairs of electrons that are not involved in bonds and are shown as dots

6. How many bonding pairs are in the following compound? 1. 1 2. 2 3. 3 4. 4 5. 6 6. 9 7. 18

7. How many lone pairs are in the following compound? 1. 1 2. 2 3. 3 4. 4 5. 6 6. 9 7. 18

8. How many bonding pairs and lone pairs are in the following compound? 1. 6 bonding pairs, 18 lone pairs 2. 12 bonding pairs, 18 lone pairs 3. 12 bonding pairs, 36 lone pairs 4. 6bonding pairs, 6 lone pairs

9. Drawing Electron Dot Formulas for Compounds Exceptions: Hydrogen only needs 2 electrons (1 bond) Boron tends to need only 6 electrons (3 bonds) Single atoms go in the center If more than one single atom, middle atom central atom

10. Draw the electron dot formula. Then state how many bonding and unbonding pairs are present. A) NBr 3 B) Water C) Chlorite ion (ClO 2 - ) D) CF 2 Cl 2

11. Multiple Bonds If there are not enough electrons to form full octets, multiple bonds may need to be formed.

12. Draw the electron dot formula E) O 2 F) CO 2

13. Resonance Structures If there are more than one possibility, resonance structures are drawn. Resonance structures show possible locations of the bonds. In reality the electrons exist as an average of the two structures – splitting time equally between them.

14. Resonance Example Each resonance structure is shown followed by the combination with the double bonds shown with a dotted line as one of the bonds.

15. Draw the electron dot formulas including resonance structures G) SO 2 H) N 2 O

16. What is the name of the property that describes the tendency of an atom to attract electrons when bonded to another atom? 1. Ionization energy 2. Conductivity 3. Electronegativity 4. Metallic Character 5. Bond length

17. Classifying Bond Types Chemical bonds can be classified by how much the bonded electrons are shared or are not shared by the elements involved. Electronegativity: The ability of one atoms in a molecule to attract electrons to itself. Wolfgang Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). Electronegativity increases across a period and down a group.

18. Electronegativities of Elements Electronegativity

20. Bond Classification based on Electronegativity Difference As the difference in electronegativity increases, electrons are less equally shared and become more polar.

21. Type of Bonds Electronegativity Difference Nonpolar covalent Polar covalent Ionic x ≤ 0.4 0.4 < x < 1.8 x ≥ 1.8 Bond Classification based on Electronegativity Difference

23. Classify the bond between the following elements: Cl and Cs 1. Ionic 2. Polar Covalent 3. Nonpolar Covalent

24. Classify the bond between the following elements: C and H 1. Ionic 2. Polar Covalent 3. Nonpolar Covalent

25. Classify the bond between the following elements: N and O 1. Ionic 2. Polar Covalent 3. Nonpolar Covalent

26. INTERMOLECULAR FORCES Interactions between Molecules

27. Polarity of a Compound Like bonds, compounds themselves can also be classified as polar or nonpolar. Polarity is based on:  Difference in electronegativity of atoms within a compound  Symmetry of the compound

28. Nonpolar Compounds - Diatomic molecules are always nonpolar. (ex. F 2 ) - Also, compounds that are totally symmetric may be nonpolar as well. (ex. CCl 4 )

29. Nonpolar Compound – the bonds are polar but the dipoles cancel out since the compound is symmetrical (tetrahedral)

30. Nonpolar Compound – the bonds are polar but the dipoles cancel out since the compound is symmetrical (linear)

31. Polar Compounds Polar compounds have one side of the compound that is more positive and another side that is more negative.

32. Polar Compounds

35. BF3 = Polar or Nonpolar 1. Polar 2. Nonpolar

37. CH3F = Polar or Nonpolar? 1. Polar 2. Nonpolar

39. CF 4 = Polar or Nonpolar? 1. Polar 2. Nonpolar

40. Br 2 = Polar or Nonpolar 1. Polar 2. Nonpolar

41. PBr3 = Polar or Nonpolar 1. Polar 2. Nonpolar

42. Intermolecular Forces Intermolecular Forces are forces that exist between two molecules that hold them together. Intermolecular Forces are caused by charge differences and polarity (because positive and negatives attract) The stronger the polarity, the stronger the attraction between molecules.

43. Intermolecular Forces The stronger the polarity, the stronger the attraction between molecules. The strength of the attraction between molecules determines properties such as:  Boiling point  Melting point  Surface tension  Cohesion  Capillary action

44. Types of Intermolecular Forces Three major types of intermolecular forces:  Dipole-Dipole Interactions  Hydrogen Bonds  Dispersion Forces

45. Dipole-Dipole Interaction Occurs in polar molecules. Positive pole of one molecule is attracted to the negative pole of the next molecule.

46. Hydrogen Bonds Occurs in polar molecules when the hydrogen atom is attracted to the more electronegative nitrogen, oxygen, or fluorine atom of another molecule.

48. Dispersion Forces Dispersion forces are the weakest type of intermolecular forces because they exist between nonpolar molecules. Usually, the electrons are shared equally. But because electrons are constantly moving, sometimes a temporary dipole forms when all the electrons are on one side of the molecule. This temporary dipole would cause an attraction with another temporary dipole.

49. Summary of Intermolecular Forces (from strongest to weakest)

50. What kind of intermolecular force would exist in H 2 O? 1. Hydrogen bonding 2. Dipole-Dipole 3. Dispersion

51. What kind of intermolecular force would exist in PCl 3 ? 1. Hydrogen bonding 2. Dipole-Dipole 3. Dispersion

52. What kind of intermolecular force would exist in Br 2 ? 1. Hydrogen bonding 2. Dipole-Dipole 3. Dispersion

53. What kind of intermolecular force would exist in NH 3 ? 1. Hydrogen bonding 2. Dipole-Dipole 3. Dispersion

54. Melting and Boiling Points In order to melt or boil a substance, intermolecular forces must be broken. Therefore, weaker intermolecular forces require less energy and have lower melting and boiling points. Therefore, stronger intermolecular forces require more energy and have higher melting and boiling points.

55. Which type of intermolecular force would have the lowest boiling point and exist as a gas at room temperature? 1. Ionic Bonds 2. Dispersion Forces 3. Hydrogen Bonds 4. Covalent Bonds 5. Dipole-Dipole attractions

56. Which type of intermolecular force would have the highest melting point and always exist as a solid at room temperature? 1. Ionic Bonds 2. Dispersion Forces 3. Hydrogen Bonds 4. Covalent Bonds 5. Dipole-Dipole attractions

57. Which compound would have the lowest boiling point and exist as a gas at room temperature? 1. H2O 2. N 2 3. CuI 2 4. CO 2 5. LiF

58. Which compound would have the highest melting point? 1. NH 3 2. NaBr 3. F 2 4. CO 2 5. NH 3

59. Effects of Molecular Mass on Melting and Boiling Points The higher molecular masses will have higher melting and boiling points because they have more electrons that form temporary dipoles. Acts as the tiebreaker if the forces are the same types – highest mass has the strongest force (highest melting and boiling point).