Lewis Dot Structures

1. Lewis Dot Structures • Developed by G. N. Lewis to serve as a way to describe bonding in polyatomic systems. • Central idea: the most stable arrangement of electrons is one in which all atoms have a “noble” gas configuration. • Example: NaCl versus Na+Cl- Na: [Ne]3s1 Cl: [Ne]3s23p5 Na+: [Ne] Cl-: [Ne]3s23p6 = [Ar]

2. LDS Mechanics • Atoms are represented by atomic symbols surrounded by valence electrons. Lone Pair (6 x) • Electron pairs between atoms indicate bond formation. Bonding Pair

3. LDS Mechanics (cont.) Three steps for “basic” Lewis structures: Sum the valence electrons for all atoms to determine total number of electrons. Use pairs of electrons to form a bond between each pair of atoms (bonding pairs). Arrange remaining electrons around atoms (lone pairs) to satisfy the “octet rule” (“duet” rule for hydrogen).

4. LDS Mechanics (cont.) An example: Cl2O 20 e- 16 e- left

5. LDS Mechanics (cont.) An example: CH4 8 e- 0 e- left Done!

6. LDS Mechanics (cont.) An example: CO2 16 e- 12 e- left Octet Violation 0 e- left CO double bond more than one pair of electrons is shared

7. Lewis Diagram for CO2

8. LDS Mechanics (cont.) + An example: NO+ + 10 e- 8 e- left +

9. Whiteboarding! • In a group of 3 try the following Lewis Structures!

11. Resonance Structures • We have assumed up to this point that there is one correct Lewis structure. • There are systems for which more than one Lewis structure is possible: • Different atomic linkages: Structural Isomers • Same atomic linkages, different bonding: Resonance

12. Resonance Structures (cont.) The classic example: O3 Both structures are correct!

13. Resonance Structures (cont.) In this example, O3 has two resonance structures: • Conceptually, we think of the bonding being an average of these two structures. • Electrons are delocalized between the oxygens such that on average the bond strength is equivalent to 1.5 O-O bonds.

14. Resonance Structures for CO3-2

15. Formal Charge Formal Charge: Compare the nuclear charge (+Z) to the number of electrons (dividing bonding electron pairs by 2). Difference is known as the “formal charge”. #e- 7 6 7 7 6 7 Z+ 7 6 7 7 7 6 Formal C. 0 0 0 0 +1 -1 • Structure with less F. C. is more correct.

16. Formal Charge Example: CO2 e- 6 4 6 6 4 6 7 4 5 Z+ 6 4 6 6 6 4 6 6 4 FC 0 0 0 0 +2 -2 -1 +2 -1 More Correct

17. Beyond the Octet Rule • There are numerous exceptions to the octet rule. • We willdeal with two classes of violation here: • Sub-octet systems • Valence shell expansion

18. Beyond the Octet Rule (cont.) • Some atoms (Be and B in particular) undergo bonding, but will form stable molecules that do not fulfill the octet rule. • Experiments demonstrate that the B-F bond strength is consistent with single bonds only.

19. Beyond the Octet Rule (cont.) • For third-row elements (“Period 3”), the energetic proximity of the d orbitals allows for the participation of these orbitals in bonding. • When this occurs, more than 8 electrons can surround a third-row element. • Example: ClF3 (a 28 e- system) F obey octet rule Cl has 10e-

21. Summary Remember the following: • C, N, O, and F obey the octet rule • B and Be are often sub-octet • Second row (Period 2) elements never exceed the octet rule • Third Row elements and beyond can use valence shell expansion to exceed the octet rule. • In the end, you have to practice…..a lot!