1. Chapter Six Representing Molecules

2. Section 6.1 The Octet Rule

3. The Octet Rule Recall: atoms want noble gas configurations
Octet Rule: atoms will gain, lose, or share electrons to achieve a noble gas configuration
Typically 8 valence electrons
Atoms will bond with each other to achieve a full octet

4. Lewis Structures
A pair of shared electrons can be represented by either with 2 dots or with a dash
Unshared electrons are called lone pairs F F F

5. Lewis Structures Types of bonds:
Single bonds: bond containing only 2 electrons
Multiple bonds: bond containing more than 2 electrons
Double bond: 4 electrons (or 2 pairs of electrons)
Triple bond: 6 electrons (or 3 pairs of electrons)

6. Bond Strength In a particular pair of elements
Triple bonds are the shortest
Double bonds are in the middle
Single bonds are the longest
Bond energy is the energy required to BREAK bonds between atoms

7. Section 6.2 Electronegativity and Polarity

8. Electronegativity Electronegativity is a periodic trend
Ability of an atom to attract electrons to itself when bonded to another atom
Quantified by the Pauling Scale

9. Electronegativity

10. Electronegativity

11. Electronegativity

12. Categories of bonds
Let’s consider three molecules:
H2, HF, NaF

13. Ionic, Polar Covalent, Nonpolar
Take the difference in electronegativities of two atoms bonded together
If the difference is 0.5 or lower, the bond is nonpolar covalent
If the difference is between 0.5 and 2.0, the bond is polar covalent
If the difference is greater than 2.0, the bond is ionic

14. Determine if the bond is ionic, polar covalent, or nonpolar covalent
The bond in ClF (chlorine and fluorine)
The bond in CsBr
The carbon-carbon double bond in C2H4
In which of the following molecules are the bonds most polar: H2O, BCl3, PCl5

15. Section 6.3 Drawing Lewis Structures

16. Drawing Lewis Dot Diagrams
1) Determine the central atom and place terminal (“outside”) atoms around central atom
Central atom typically is the least electronegative element in compound, the element with only 1 atom, and/or the element written first in compound
2) Count total # of v.e.
3) Bond all terminal atoms to central using single bond
Each bond is 2 electrons; subtract from total # of v.e.

17. Drawing Lewis Dot Diagrams
4) Complete the octets of terminal atoms w/ remaining v.e.
5) If any electrons left over, put on central atom
6) Use multiple bonds to complete octet of any elements where necessary

18. Examples of Lewis Dot Structures
CH4
H2O O2
CO2
CN-

19. Group Quiz #1 Draw the Lewis Dot Structures for the following: CS2 NF3
ClO3-

20. Section 6.4 Lewis Structures & Formal Charge

21. Formal Charge Another way of keeping track of electrons in a molecule
Formal Charge = (# of v.e.) – (# of associated
electrons)
Ex: Ozone (O3)
Now you try: NO3-

22. Using Formal Charge
Formal Charge can help us determine the best Lewis Structure when there are options
Consider the following two skeletal structures for CH2O. Which one is preferred?

23. Formal Charge Rules
Lewis structures where all formal charges are zero is preferred
Small formal charges (0 and +/-1) are preferred to big formal charges (+/-2, +/-3, etc.)
The best arrangements are where the more electronegative atoms have the more negative formal charge

24. Group Quiz #2
Draw the Lewis Structures for the following compounds and determine the formal charge on EACH atom
SO32-
CO32-

25. Section 6.5 Resonance

26. Resonance Structures
Consider the molecule NO3- and its Lewis Structure

27. Section 6.6 Exceptions to the Octet Rule

28. Exceptions to the Octet Rule
Central atom has fewer than 8 v.e. due to electron shortage
Ex: Boron (happy w/ 6); Beryllium (happy w/ 2)
Central atom has fewer than 8 v.e. due to odd # of electrons (known as radicals)
Ex: Nitrogen (NO2)
Central atom has more than 8 v.e.
Ex: Sulfur (SF6) and Xenon (XeF4)
See pp. 201—204 for examples/explanations

29. Formal Charges, Resonance Structures, AND Exceptions!
Consider the polyatomic ion: SO42-. What would be the BEST Lewis dot structure?
What about PO43-?

30. Group Quiz #3
Draw the Lewis Structure for antimony pentafluoride (SbF5)
Draw the Lewis Structure for Borane (BH3)
Draw the Lewis Structure for Nitrogen Disulfide (NS2)