1. Bonding

2. Video 5.1 Types of Bonds

3. Octet Rule Review Atoms bond with other atoms by sharing or transferring electrons in order to achieve a stable octet (8 valence electrons). Bonding creates stability! *When bonds are formed energy is ___________. *When bonds are broken energy is ___________. released absorbed

4. Ionic Bonds O Transfer of electrons from the cation to the anion (metal to nonmetal). O High melting point and boiling point O Mostly hard crystalline solids O Conduct as liquid (either melted or dissolved) due to mobile ions.

5. Ionic Bonds Sodium Chloride: NaCl (table salt) properties: O Hard O Solid crystals O High melting point, forget boiling! O Liquid phase conducts (electrolytes are salts)

6. Metallic Bonds O Metals only O All metals lose their valence electrons and form a sea of electrons O High melting point and boiling point O Insoluble in water O Always able to conduct heat and electric due to mobile electrons O Malleable O Ductile

7. Metallic Bonds Copper (Cu) properties: O Hard solid O High melting point, forget boiling! O Malleable and ductile O Conductor O Can’t dissolve

8. Metallic Bonds Sea of electrons Copper (I) ions Copper (II) ions

9. Covalent Bonds (Molecular) O Nonmetals only O Share electrons between atoms O Low melting point and boiling point O Never conduct heat or electricity O Soft solid or gas

10. Covalent Bonds Dextrose C 6 H 12 O 6 (Sugar) properties: O Soft O Melts easily in sauce pans for caramel O Doesn’t conduct (nonelectrolyte)

11. What type of bond is created? 1. Ca + O 2. K + Br 3. S + Cl 4. I + S 5. Li + Mg 6. Ba + S M+ NM = Ionic NM + NM = Covalent M + NM = Ionic NM + NM = Covalent M + M = Metallic

12. Video 5.2 Ionic Compounds

13. Review: Find the ionic formula: 1. K + Br 2. Mg + Cl 3. Na + S 4. Ca + S KBr MgCl 2 Na 2 S CaS + - +2- + -2 +2 -2

14. Draw Lewis structures: KBr MgCl 2 Na 2 S CaS

15. Which subatomic particle is involved in bonding? Electrons only!

16. Geometry of ionic crystals Ions Ionic crystal

17. Video 5.3 Covalent Compounds

18. Covalent Lewis Structures Rules: CCl4 1. Add up all valence e - 2. Draw a skeletal structure with bonds between elements. Least frequent element in the middle. 3. Subtract 2e - from total for each bond drawn. 4. Draw in remaining e - to fill each atom’s octet. 5. Evaluate: each atom should have 8 e- only. C: 4 + 4Cl: 7 = 32 valence e - Cl Cl—C—Cl Cl 32-8=24

19. VSEPR “Valence shell electron pair repulsion” is a model for molecules. Lone electron pairs are repelled by one another and should be placed as far apart as possible.

20. Geometry 1. Linear: The molecule is on one plane (flat) such as CO 2 or H 2. 2. Bent: The molecule is bent at angle like H 2 O due to unshared electrons and two bonding pairs on the central atom.

21. Geometry 3. Pyramidal: The molecule has a triangular shape like NH 3 due to a lone pair and three bonding pairs on the central atom. 4. Tetrahedral: The molecule has four bonding pairs and no lone pairs on the central atom like CH 4.

22. Examples: O Draw the following molecules and identify their geometry: 1. PCl 3 2. SiCl 2 H 2 3. Br 2 4. H 2 S pyramidal tetrahedral linear bent

23. Video 5.4 Bond Polarity

24. The earth has two poles; North and South. A magnet also has two poles. Bonds may have two poles. This means one element is charged different than the other. If a bond is polar, the two elements have different electronegativities. The element with a higher electronegativity will be more negative.

25. Bond Polarity

27. Nonpolar Bond

28. Bond Polarity Electronegativity difference Bond type 0-0.4Nonpolar 0.5-1.0Polar 1.1-2.0Very Polar 2.0-4.0Ionic

29. Ionic, polar or nonpolar? 1. C-Br 2. Na-S 3. C-C 4. H-O 5. K-O 6. Be-B 7. As-O 8. N-O 9. C-O 10. F-F 11. S-C 12. N-H P I NP P I I P P P P

30. Covalent Bonding O If 2 atoms or more form a bond with the same electronegativity the bonds are nonpolar and they share e- equally. ( F-F ) O If there is an electronegativity difference between bonded atoms, the bonds are polar and e- are pulled toward the more electronegative atom. (H-F) O If a bond is polar, the molecule will have a slightly negative and slightly positive side, like 2 poles of a magnet.

31. Video 5.5 Molecular Polarity

32. O A polar molecule will be asymmetrical. O A nonpolar molecule will have a symmetrical shape or all nonpolar bonds.

33. Molecular Polarity Which are polar molecules? Show charges. - - - - - - - - - + + + + + + NP P P

34. Molecular Polarity Water is polar, and like dissolves like, so only polar molecules are soluble in water. Polar molecules are also attracted to an electric field.

35. Molecular Polarity O As you can see, normally polar molecules are unaligned. O When a electric source comes by, the molecules quickly align themselves.

36. Video 5.6 IMF

37. O Intramolecular forces is another name for bonds, that keep elements together in compounds. O Intermolecular forces of attraction are weaker than bonds, but are responsible for holding a substance together (multiple molecules in a confined area).

38. IMF O The stronger the IMF, the tighter the structure (solid). The melting and boiling points will be high. O The weaker the IMF, the looser the structure (gas). The melting and boiling points will be low.

39. Dipole-Dipole O Dipole-Dipole attractions are strong forces between polar molecules. It is like static holding the + and – charges together.

40. Hydrogen Bonding A special case: Hydrogen Bonds are the strongest bonds between Hydrogen and very electronegative atoms such as F, O and N. (H bonds are FON!) For example, H 2 O and HF, due to their polarity, they will attract each other.

41. London Dispersion Forces (LDF) The weakest attraction between nonpolar molecules occur because electrons temporarily shift creating a temporary + and – charge. The more electrons the compound has, the stronger the force is.

42. Summary O From weak to strong: O Nonpolar LDF O Polar Dipole Dipole forces O Hydrogen bonds O Covalent Bonds O Ionic Bonds O Metallic Bonds

43. Class Notes

44. Show the individual and bonded Lewis structures: 1. Li and F 2. Mg and O 3. Be and S 4. What did all the cations do? 5. What did all the anions do? 6. Which of the subatomic particles were changed and how were they changed?

46. Type of Bonding? 1. CaCl 2 2. CO 2 3. H 2 O 4. BaSO 4 5. K 2 O 6. NaF 7. Na 2 CO 3 8. CH 4 9. SO 3 10. LiBr 11. MgO 12. NH 4 Cl 13. HCl 14. KI 15. NaOH 16. NO 2 17. AlPO 4 18. FeCl 3 19. P 2 O 5 20. N 2 O 3