1. Starter  Which elements naturally exist as diatomic molecules? Remember, the 7 + 1 rule Remember, the 7 + 1 rule  How many valence electrons do each of the halogens have?

2. Octet Rule  representative elements can “fill” their outer energy level by sharing electrons in covalent bonds  Octet Rule- a compound tends to form so that each atom has an octet (8) of electrons in its highest energy level by gaining, losing or sharing electrons  Duet Rule- applies to H and He

3. Octet Rule  Less than 8: Boron: 6 in outer energy level Boron: 6 in outer energy level  More than 8: anything in 3 rd period or heavier anything in 3 rd period or heavier because may use the empty d orbital because may use the empty d orbital ex: S, P, I ex: S, P, I

4. Electron Dot Diagrams  a way to show electron configuration  identifies the number and pairing of valence electrons to show how bonding will occur 1. write the noble gas notation 2. identify the number of valence 3. identify how many are paired and how many are alone 4. do not go by Figure 6-10

5. Example  Nitrogen 1s 2 2s 2 2p 3 1s 2 2s 2 2p 3 5 valence 5 valence 2 are paired 2 are paired 3 are alone 3 are alone  Sulfur 1s 2 2s 2 2p 6 3s 2 3p 4 6 valence 4 paired (2 pairs) 2 are alone N

6. Lewis Structures  like dot diagrams but for entire molecules  atomic symbols represent nucleus and core electrons and dots or dashes represent valence electrons unshared electrons: (lone pairs) pair of electrons not involved in bonding written around only one symbol unshared electrons: (lone pairs) pair of electrons not involved in bonding written around only one symbol bonding electrons: written in between 2 atoms as a dash bonding electrons: written in between 2 atoms as a dash

7. Types of Bonds  single- sharing of one pair of electrons weakest, longest weakest, longest  double- sharing of 2 pairs of electrons stronger and shorter stronger and shorter  triple- sharing of 3 pairs of electrons strongest and shortest strongest and shortest  multiple bonds include double and triple bonds

8. Drawing Lewis Structures 1. find the number of valence electrons in each atom and add them up 2. draw the atoms next to each other in the way they will bond 3. add one bonding pair between each connected atoms 4. add the rest of the electrons until all have 8 (consider exceptions to octet rule)

9. H H C Cl H Example 1  CH 3 Cl CH 3 Cl CH 3 Cl  methyl chloride  C: 4 x 1 = 4  H: 1 x 3 = 3  Cl: 7 x 1 = 7  total = 14 electrons  carbon is central H H C Cl H duet octet

10. Example 2  NH 3 NH 3 NH 3  ammonia  N: 5 x 1 = 5  H: 1 x 3 = 3  total = 8  N is central H N H H

11. Example 3 N2N2N2N2  nitrogen gas  N: 5 x 2 = 10  10 electrons N N

12. Example 4  CH 2 O  formaldehyde  C: 4 x 1 = 4  H: 1 x 2 = 2  O: 1 x 6 = 6  total = 12  C is central H C H O

13. Example 5 O3O3O3O3  ozone  O: 6 x 3 = 18  two completely equalarrangements  the real structure is an average of these two  where each bond is sharing 3 electrons instead of 4 or 2 O O O

14. Resonance Structures  resonance – bonding between atoms that cannot be represented in on Lewis structure  show all possible structures with double-ended arrow in between to show that electrons are delocalized O O O

15. Example 6  NO 3 1- NO 3 1- NO 3 1-  N: 5 x 1 = 5  O: 6 x 3 = 18  total = 23 + 1 = 24

16. Covalent Network Bonding  a different type of covalent bonding  not specific molecules  lots of nonmetal atoms covalently bonded together in a network in all directions  example: diamond diamond diamond silicon dioxide silicon dioxide silicon dioxide silicon dioxide graphite graphite graphite