1. Bonding

2. Chemical Bonding Types 1)Ionic 2)Covalent Polar Nonpolar 3)Metallic

3. Ionic Bonding [109-114] Occurs between metal and nonmetal Large electronegativity difference –Can you predict ionic vs covalent from position on periodic table? Involves transfer of electrons Metal loses e -, becomes cation (+ ion) Nonmetal gains e -, becomes anion (- ion)

4. Ionic Bonding [109-114] Ions held together by electrostatic attraction Ions form crystal lattice

5. Ionic Bonding [109-114, 150-152] Physical properties of ionic compounds –Hard, brittle solids –High melting and boiling points –Conduct electricity in molten and aqueous states Why not conductive in solid state? –More soluble in water than other solvents

6. Ionic Bonding [109-114, 150-152] Exercises 1-3 page 112 Exercises 4-6 page 114 Examiner’s hint Page 113

7. Covalent Bonding [114] Occurs between nonmetals –Can you predict whether bond will be covalent depending on position in periodic table? Small electronegativity difference Involves sharing of electrons Covalent bond defined attraction between pair of electrons and positively charged nuclei

8. Covalent Bonding [150-152] Physical properties of covalent compounds –Solid, liquid, or gas at room temperature –Low melting and boiling points –Non-Conductive –Solubility depends on polarity

9. Covalent Bonding [115] When two atoms share two electrons (1 pair), a single chemical bond is formed –Designated by a single line A double bond is the sharing of four electrons (2 pair) –Designated by double line A triple bond is the sharing of six electrons (3 pair) –Designated by triple line

10. Dative Bond [117] A dative bond is a covalent bond where both shared electrons are donated by only one atom –aka coordinate bond

11. Covalent Bonding [118] Bond length –single>double>triple Bond strength –triple>double>single The structure of covalent compounds is shown in a Lewis Diagram –aka Lewis Dot Diagram, Electron Dot Diagram

12. Lewis Dot Diagram [115-117] Most elements require an octet; can be combination of unshared or shared pairs (bonds) Remember H only needs 2 electrons, Be* only 4 electrons, B* only 6 electrons Halogens mostly form one bond only Oxygen mostly forms two bonds only Nitrogen mostly forms three bonds only

13. Lewis Dot Diagram [115-117] Rules 1)Sum the valence electrons 2)Bond each atom to the central atom with single bond Central atom is usually written first, or most electronegative, or can form most bonds 3)Give all atoms an octet 4)Count the electrons represented a)Too many, must use multiple bonds b)Too few, add extra to central atom (this is HL)

14. Lewis Dot Diagram [115-117] Exercise 7 page 117 Shapes of O 2, CO 2, H 2 O, and CO should be committed to memory.

15. Lewis Dot Diagram [115-117] Special case of delocalized electrons

16. Shapes of Molecules/Ions [120-122] Shapes of covalent compounds are determined by the repulsion between the valence electron pairs Valence Shell Electron Pair Repulsion –VSEPR Valence electrons will situate themselves to be the furthest away from all other valence electrons, avoiding steric hindrance

17. Shapes of Molecules/Ions [120-122] Electron domains Bonding domains Non bonding domains shapeangle 440Tetrahedral109.5 431Trigonal pyramidal 107 422Bent104.5 330Trigonal planar 120 321Bent<120 220linear180

18. Shapes of Molecules/Ions [120-122] Exercises 10 -13 page 123

19. Bond Polarity [119-120] A covalent bond is either –Non polar – equal sharing of electrons H-H, Cl-Cl, O-O, etc –Polar – unequal sharing of electrons H-F, Cl-O, H-S-H Due to differences in electronegativity Produces slight electrical charges Indicated by arrow or  + and  - Ends of polar bonds are called dipoles

20. Bond Polarity [119-120] Examples –Arrow points towards most electronegative atom –  + placed on least electronegative atom –  - placed on most electronegative atom H – F H – O – H Cl – Cl

21. Molecular Polarity [123] Polarity of a molecule depends both on the bond polarity and shape of molecule –Polar molecules contain polar bonds that are nonsymmetrical, they do not cancel out. Exercises 8,9 page 120 Exercise 13 page 123

22. Covalent Crystalline Solid [124] Some covalent compounds form crystals or giant covalent networks –Examples: diamond, SiO 2 –Physical properties: Hard High melting point Non conductor Not soluble in water

23. Allotropes of carbon [124] Allotropes are different forms of an element that exist in the same physical state –Oxygen exists as O 2 or O 3 Draw Lewis dot –Carbon exists as diamond, graphite, or fullerene-60

24. Structure of diamond [124] Giant covalent network Tetrahedral bonding No delocalized electrons non conductor

25. Structure of graphite [124] Trigonal planar bonding Delocalized electrons Conductor

26. Structure of fullerene-60 [124] Trigonal planarish bonding Some delocalization Semi-conductor

27. Metallic Bonding [149] Metal atoms are tightly packed together in a lattice, like oranges in a box The valence electrons delocalize from each atom to surround all the nuclei in the metallic lattice The resulting positive nuclei (ions) are attracted to the electrons and that is what holds metals together Mobile “sea of electrons”

28. Metallic Bonding [149]

29. Metals are conductive because the sea of electrons carries the charge Metals are malleable and ductile because the layers of ions can slide past each other easily Strength of metallic bond depends on how many valence electrons are in the sea and how far the valence electrons are from the nuclei

30. Metallic bonding questions 1. Why are metals conductive in the solid state, but not ionic compounds? 2. Discuss the strength of metallic bond between Na and Mg and Na and K

31. Conductivity [152] Ionic compounds metals Conductive in which states of matter Aqueous or molten NOT solid Solid or molten Species that carry charge ionselectron There must be freely moving ions or electrons for conductivity

33. Intermolecular Forces (IMF) [144-148] Forces between molecules or atoms Affects boiling and melting points

34. Intermolecular Forces (IMF) [144-148] Three types (strongest to weakest) –Hydrogen bonds –Dipole-dipole –Van der Waals or London dispersion Relative strengths giant covalent > Ionic > covalent > IMF

35. Hydrogen Bonds [146-147] Only between H and N, O, or F of another molecule

36. Dipole-Dipole [145-146] Between polar molecules Attraction between  + and  - of neighboring molecules Larger electronegativity differences lead to stronger dipole – dipole interactions

37. Van der Waals [144-145] Actually are present in all molecules Important force between non polar molecules Temporary induced dipoles Larger the molar mass means more electrons in induced dipole, stronger the Van der Waals Important forces between halogens, noble gases

38. Van der Waals [144-145]

39. Physical properties and types of bonding and IMF [148,150-152] Melting and boiling points –Stronger the bonds or IMF that hold substance together, the higher the melting and boiling point will be Volatility –Stronger the bonds or IMF that hold substance together, the lower the volatility will be Conductivity –Ionic compounds conduct electricity in molten and aqueous states, ions are free to move and carry charge

40. Physical properties and types of bonding and IMF [148,150-152] Solubility –“Likes dissolve likes” –Ionic solutes dissolve in ionic and polar solvents –Polar covalent solutes dissolve in ionic and polar solvents –Nonpolar covalent solutes dissolve in nonpolar solvents

41. Solubility [151] Solvation = solute particles being surrounded by solvent particles Hydration = solute particles being surrounded by water molecules

42. Solubility [151] Polar compounds can form IMF with water and will dissolve

43. melting point trends [79-80, 150] melting points trends alkali metals halogens period 3 noble gases Hydrogen halides

44. melting point trends [79-80, 150] Alkali metals 1)what type of bond/IMF? 2)what happens to strength as you move down group?

45. melting point trends [79-80, 150] Halogens and noble gases 1)what type of bond/IMF? 2)what happens to strength as you move down group?

46. melting point trends [79-80, 150] Period 3 1)what type of bond/IMF? 2)what happens to strength as you move across period?

47. melting point trends [79-80, 150]

48. Hydrogen halides