1. Warm Up – Name or Write the following compounds
1.Cr2(SO3)3
2.LiCN
3.Copper(II) chlorate
4.Iron(III) hydroxide
5.CS4

2. Writing Lewis Dot Structures
In writing Lewis structures for covalent compounds, you must remember the following:
1. Determine the total number of valence electrons
2. The least electronegative element goes in the center
H is NEVER a center atom
Carbon is usually in the center

3. Writing Lewis Structures
3. Distribute 8 electrons around central atom
4. Place remaining atoms outside of central atom
5. Distribute any remaining electrons around outer atoms

4. Writing Lewis Structures
6. Use double or triple bonds to obey Octet Rule as needed
7. Check to make sure that every atom obeys the Octet Rule

5. Write Lewis Dot Structures for the following:
1. H2O
2. CO2
3. NH3
4. CBr4
5. SO2

6. Practice
Draw the Lewis Structures for the Following:
1) GeF4
2) COS
3) SeO3
4) HOF

7. Practice
Write Lewis structures for the following molecules:
a. HCN d. PH3
b. CHCl e. SiO2
c. SeF2 f. CF4

8. Lewis Structures for Polyatomic Ions
When writing dot structures for polyatomic ions, you must remember to add or subtract the amount of electrons represented by the charge
When writing polyatomic ions, you must include the structure inside brackets, [ ], with the charge outside the bracket

9. Write Lewis Structures for the following polyatomic ions:
1. (SO4) 2-
2. (OH)-
3. (CN)-
4. (NH4)+

10. Warmup Practice Naming or writing Compounds
Mn(ClO2)7
PbS
NH4HSO4
Cd(NO3)2
Lead (IV) perchlorate
Beryllium Acetate
Copper (II) Bisulfate

11. Extended Octet Rule Examples: PCl5 or XeF4
“Expanded/Extended octet” refers to the Lewis structures where the central atom ends up with more than an octet
Examples: PCl5 or XeF4

12. Limitation of Extended Octet Rule
1) The central atom with an expanded octet MUST have an atomic number larger than 10 (beyond neon).
2) Extra electrons should be first placed on the outside atoms. After the outside atoms have fulfilled the Octet Rule, and there are still extra electrons, start with placing them as lone pairs on the central atom.

13. Practice
1) BrF3 2) SF4 3) PCl5 4) XeF4 5) SbF5

14. VSEPR Theory
VSEPR Theory- Valence Shell Electron Pair Repulsion Theory
VSEPR Theory can be used to predict a molecule’s geometry
Electrons repel each other
Molecules adjust their shapes to account for this repulsion

15. VSEPR Theory Rules
The adjusted shape allows for the minimum amount of repulsion between valence shell electrons
Basically, a molecule will arrange itself in such a way that the valence electrons are as far apart from each other as possible

16. Using VSEPR Theory
Unshared pairs of electrons play an important role in the molecule’s geometry
Built off of Domains
Single bonds
Double Bonds
Triple Bonds
Lone Pairs of Electrons

17. Basic Geometric Shapes for Molecules
Linear
Trigonal Planar
Tetrahedral
Trigonal Pyramidal
Trigonal bipyramidal
Octahedral
* The ones in yellow you must memorize

18. Linear

19. Trigonal Planar

20. Tetrahedral

21. Trigonal Pyramidal

22. Bent

23. Trigonal bypyramidal

24. Octahedral

25. Using the VSEPR Theory, Predict the Shape of the following Molecules
1. H2O
2. BF3
3. H2S
4. PH3

26. Practice VSEPR Theory: Draw the Lewis structures for the following compounds and predict the molecular geometry of the compound
1. SCl PI3
2. OCl NH2Cl

27. Draw the Lewis Structures and Predict the Molecular Geometry
a. BeCl2
b. BCl3
c. (SeO4)2-
d. SiCl3Br
f. ONCl

28. Practice
Draw the following Lewis structures for the following and predict the molecular geometry
HCN
SO2
XeF2
PCl5

29. Predict the following molecular geometry
PBr3
N2H2
C2H4
OCl2
HC2F

30. Warm up Practice Drawing Lewis Structures and Predicting Molecular Geometry
1) SCl2
2) BeF2
3) HOF
4) COS

31. Origin of Polar Bonds
Recall, electronegativity refers to the tendency of an atom to attract electrons
Electronegativity really measures the strength that an atom uses to attract electrons
The element with the higher electronegative value will attract electrons more strongly which means that it will pull the bonds closer

32. Origins of Polar Bonds
The greater the difference between each atom’s electronegativity, the more likely it is that the bond is polar

33. Origins of Polar Bonds
The closer the atoms are together, the closer the electronegative values are to each other
Elements that have closer electronegative values form NONPOLAR BONDS
The further the atoms are apart, the more likely they are to form POLAR BONDS

34. Polar Bonds
Polar bonds result when shared electrons are not shared equally
Polar bonds occur only with covalent compounds (means they have an ∆EN< 1.7 but > 0.5)
This causes each atom to develop a temporary charge

35. NonPolar Bond
Non-Polar bonds result when shared electrons are shared equally
Non-Polar bonds occur only with covalent compounds with an ∆EN= 0 to .3

36. Using Electronegative Values to Predict Bond Types
Predict if the following bonds will be: Polar Covalent, Nonpolar Covalent, or Ionic based on Electronegative Values
a. H and O b. Cl and Br
c. K and S d. F and F

37. Polar Molecules
It is possible for a molecule to have polar bonds, but be nonpolar overall
When polar bonds are in opposite directions, they “cancel” each other out
Polar molecules are usually not symmetrical
Example:
CCl4 has polar bonds, but is a NONPOLAR MOLECULE

38. Predict if the following molecules are polar or nonpolar
1. CF4
2. CO2
3. H2O
4. (SO4)2-
1. NonPolar
2. Nonpolar
3. Polar
4. Nonpolar