1. Chemical Bonding Chp 6 pg 165

2. I. Chemical Bonding A. Intro 1. Chem bond – electrical attraction b/w nuclei and valence electrons of different atoms 2. Naturally occurring elements are chemically bonded (typically) 3. Valence electrons rearrange to be more stable 4. Type of arrangement = type of bond

3. 5. Types of bonds a. Ionic bond 1) Ions bond 2) 1 element takes/gives electrons to another 3) + charge bonds to – charge b. Covalent bond 1) Share electrons 2) Can share 1-3 PAIRS of electrons

4. 3) Polar – uneven distribution of charge 4) Dipole – charge on side of an atom 5)Nonpolar covalent – evenly distributed electrical charge (fig 6-3 pg 163) 6) Polar covalent – Uneven electrical charge c. Metallic bonding

5. B. Covalent Bonding 1. Formation of Covalent Bond a. Elements bond to lower potential energy b. Nucleus and electrons attract – lower pot. energy c. Get closer until repulsive forces (ele. – ele, and proton-proton) equal attractive forces

7. 2. Characteristics of Covalent Bonds a. Electrons move freely in either orbital b. Energy released when bond forms

8. c. Bond energy – energy required (added) to break a bond between atoms d. Bond energy = energy released when bond formed e. Bond energy is in KJ/mol (Fig 2.5 pg 172) Example: It takes 436KJ of energy to break H-H bond in one mole of hydrogen molecules

9. f. Electron cloud has overlap – more electrons found there g. Octet Rule 1) Gain or lose electrons to get 8 in their outer shell (H exception) 2) F 2, I 2, Cl 2, Br 2

10. 3) Number of electrons per cloud a) 1 – 2 e- b) 2 – 8 e- c) 3 – 18 e- d) 4 – 32 e- e) 5 – 50 e- f) 6 – 72 e-

11. 4) Some exceptions h. Lewis dot structures 1) Only shows valence electrons 2) Dots around element symbol

12. 3) Lone pair – electrons that aren’t shared 4) Shared electrons – Represented by a line b/w elements a) Single bond – 1 pair of shared electrons b) Double bond – 2 pairs of shared electrons c) Triple bond – 3 pairs of shared electrons

13. d) Double bond and triple bonds are stronger e) Double bonds shorter than single bonds f) Some might have more than one structure

15. C. Ionic Bonding 1. Bonding of ions 2. Formation of Ionic Compounds a. Normally combine opposite sides of periodic table b. Still complete the octet rule c. Represented by Lewis dot structures d. Attracted by opposite charges

17. 3. Characteristic of ionic bonding (pg 181) a. Form crystal lattice b. Provides many bonds c. High melting point

18. 4. Comparing Ionic and Molecular Compounds a. Ionic compounds 1) Higher melting and boiling points 2) Hard and brittle 3) Can dissolve 4) Not conductors unless dissolved 5) Bonds are stronger

20. b. Molecular compounds 1) Many gases at room temp 2) Low melting points 5. Metallic Bonding a. Bonding of different metals b. Valence electrons can roam through entire atom c. Forms a “sea” of electrons

21. Metallic Bonding

22. d. Properties 1) Conduct heat and electricity 2) Luster – electrons becoming excited 3) Malleable and ductile – electrons can be rearranged 4) Bond strength varies

23.  Bonding  https://www.youtube.com/watch?v=QXT4OVM4vXI https://www.youtube.com/watch?v=QXT4OVM4vXI  Polar and Nonpolar molecules  https://www.youtube.com/watch?v=PVL24HAesnc https://www.youtube.com/watch?v=PVL24HAesnc  Lewis Dot Structures  https://www.youtube.com/watch?v=a8LF7JEb0IAhttps://www.youtube.com/watch?v=a8LF7JEb0IA

24. D. Molecular Geometry (pg 183) 1. Geometry determines polarity 2. VSEPR Theory (valence-shell electron- pair repulsion) a. Repulsion b/w electron pairs keeps atoms as far apart as possible b. Linear (Table 6-5 pg 186)

25. c. Trigonal-planar (120 degrees) d. Tetrahedral (109.5 degrees)

26. e. Unshared electrons 1) Lone pairs take up more space 2) Angular shape – water

27. f. Double and triple bonds treat the same as single bonds 3. Intermolecular Forces a. Boiling can determine intermolecular forces b. Attraction b/w molecules c. When something boils energy is added separates from other molecules d. Higher boiling points = stronger forces e. Weaker than bonds combining elements

28. f. Polarity and dipole-dipole forces 1) Polar molecules have strongest force 2) Dipole – opposite charges separated by short distance 3) Dipole-dipole forces – attraction between molecules due to opposite charge (fig 6-25 pg 191)

29. g. Hydrogen Bonding 1) Type of dipole-dipole force 2) H bonds to F, O, N due to electronegativity 3) No electrons are shared

30. Adhesion cohesion

31. h. London Dispersion Forces 1) Brief attraction between molecules due to moving e- (static electricity) 2) Only happen in noble gases and nonpolar molecules