1. Drawing Lewis Structures “ valence dot diagrams” The valence shell holds up to 8 electrons. 0.Determine the number of valence electrons. 1. Write the element’s symbol. 2. Add one electron to each side. 3. Then double up the electrons as necessary. 4. Any single electrons are available for bonding. How many unpaired electrons?

2. Practicing Lewis Structures Draw: Chlorine, Chlorine Ion Sodium, Sodium ion Oxygen, Oxygen ion Neon

3. Draw Lewis dot diagrams for Oxygen and Sodium ( determine valence electrons, distribute dots appropriately)

4. Covalent compounds are also known as: Molecular Compounds CH 6 Covalent Compounds

5. Molecular Compounds Ch 5 Chemical Names and Formulas 5 Consist of two or more non-metals 1. 2.

6. Naming Covalent Compounds (molecular compounds) Ch 5 Chemical Names and Formulas 6 Ionic charge does NOT dictate ratio of atoms. The name of the compound must indicate the number of each element.

7. Compounds of Carbon and Oxygen Ch 5 Chemical Names and Formulas 7 Both are non-metals Combine in multiple ratios “carbon oxide” does not provide enough information to give the correct ratio of “C” and “O”. Carbon and Oxygen combine to form the following molecular compounds: CO CO 2

8. Use of Numeric Prefixes 1 st element: if more than 1 2 nd element: always Ch 5 Chemical Names and Formulas 8 CO CO 2

9. Read “Naming Covalent Compounds” Pgs. 206-207 Then complete the following chart: FormulaIonic or Covalent?Name CO CO 2 N2ON2O KCl PCl 5 MgCl 2 P4O6P4O6

10. Number Prefixes 1-10 Ch 5 Chemical Names and Formulas 10 Mono Di Tri Tetra Penta Hexa Hepta Octa Nona Deca BF 3 N2ON2O

11. Covalent Bonding The more civilized way to form a relationship

12. Comparing Ionic to Covalent Ionic Electrons are “stolen” or move to another atom to complete octets A net charge is created. The ions are attracted to opposite charge forming an electrically neutral “salt” Usually a metal and non- metal like Sodium and Chlorine…NaCl Covalent Atomic Orbitals [valence] overlap and hybridize to form molecular orbitals. Electrons are “shared” to complete octets. 2 non-metals like Carbon and Oxygen….CO 2

13. Shared Electrons… Sharing occurs to form a stable gas configuration (full octet) It takes two electrons to form a bond These bonding electrons are called a “shared pair” The pair counts toward a full octet for each atom. Draw out the valence dot diagrams for Fluorine and Chlorine

14. How many electrons do each need for a full octet? 1e- needed Why not share a pair??

15. The line represents a bond, consisting of 2 electrons. This shared pair counts for both Cl and F Shared pairs vs. unshared pairs.

16. Showing overlapping P-orbitals… similar to what F and Cl would do.

17. Bonding Tendencies FamilyNumber of valence electrons Electrons needed to make an octet Number of bonds formed Halogens 711 Oxygen Nitrogen Carbon Hydrogen 1 For H and He, an “octet” is only 2 electrons

18. Drawing “Lewis Structures” Valence Dot Diagrams 1. Count the total number of valence electrons. 2. Connect each atom using single bonds. The first atom is usually central to the structure Halogens are usually terminal (end atoms) Hydrogen is always terminal. 3. Add lone pairs to each atom in order to get a full octet. CCl 4 H 2 O NH 3

19. Multiple bonds: double, triple If there are not enough electrons to make every atom single bonded, you will need to use double or triple bonds. For every 2 electrons you are “short”, you will need one more bond. H 2 CO

20. Double and Triple bonds Double bonds use 2 shared pairs Triple bonds use 3 shared pairs Carbon Dioxide CO 2 Nitrogen (gas) N 2

21. Draw Lewis structures for: Oxygen gas NI 3

22. Lewis structures…resonance Resonance helps explain the true structure when 2 or more equally valid structures can be drawn for a molecule. Resonance structures have identical arrangements of atoms, they differ in distribution of electrons. Draw 0 3 (ozone)

23. What type of bonds does ozone have? The bond length shows that there the ozone bonds are neither the single nor the double bond length. The length is in-between single and double This tells us that it is a hybrid, “1 ½ bond” Drawing resonance structures is the chemists way to show this. Type of bondLength in picometers O-O oxygen-oxygen single bond 148 O=O oxygen-oxygen double bond 121 Oxygen-Oxygen bond in OZONE 128

24. Resonance explained

25. Draw Lewis Structures for: Carbonate ion, CO 3 2- 1. Count up available electrons. 2. Identify the central atom, then single bond everything. 3. Assign lone pairs so all atoms have 8 electrons 4. Compare amount used to amount available. 5. If you used 2 many, retry with multiple bonds. 6. Determine if resonance is necessary.

26. Covalent Bonds…shared electrons Non-polar covalent, (Pure covalent)

27. Electronegativity values

28. Polar Covalent

29. 2 different atoms are covalently bonded. The bond is a Polar Bond The more electronegative atom pulls the shared electrons closer to it’s nucleus. H and Cl H and C F and Cl Find the electronegativity differences in the following pairs of atoms. Check pg 194 Figure 6

30. Water Find the Electronegativity difference for H and O Draw Water Label the partial – and partial + charge The polar molecules cause special properties Dipole interactions: A molecule that is polar is said to be a polar molecule… DIPOLE Hydrogen Bonding: Occurs w/ water. The partial positive H is attracted to the lone pairs of oxygen. Use : δ + (delta)

31. Molecular Shapes Ideal Geometries of molecules with a Central Atom. 1. Determine the number of bonds to the central atom. 2. Determine the number of lone pairs around the central atom. Put this information into the “AXE” formula to help categorize the molecule. A = Represents the Central Atom X # = Bonded atoms to central atom E # = Lone pairs around central atom

32. X + E = 4 AXE formulaShapeExample AX 4 Methane CH 4 AX 3 EAmmonia NH 3 AX 4 Water H 2 O

33. VSEPR Valence Shell Electron Pair Repulsion Theory Bonding Angles for a tetrahedral are ______° Example: Methane, CH 4 [AX 4 ] Draw Ammonia, NH 3 AXE: Bond angles Methane: Ammonia:

34. Draw Water (lewis dot) What effect do unshared pairs have?

35. X + E < 4 AXE formulaShapeExample AX 2 CO 2 AX 2 EGeF 2 AX 3 BeF 3 Molecules with 2 atoms are linear

36. Warm UP What are the bond angles for the following molecules: Methane Ammonia Water Carbon Dioxide

37. Exceptions to the octet rule: Odd # of electrons If the total number of valence electrons is odd you end up with a free radical. This unpaired electron is extremely reactive. Examples: NO NO 2