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Lecture 23: Lewis Dot Structures

Lecture 23: Lewis Dot Structures. Reading: Zumdahl 13.9-13.12 Outline Lewis Dot Structure Basics Resonance Formal Charge. Localized Bond Models. Consider our energy diagram for H 2 bonding:. Localized Model Limitations.

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Lecture 23: Lewis Dot Structures

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  1. Lecture 23: Lewis Dot Structures • Reading: Zumdahl 13.9-13.12 • Outline • Lewis Dot Structure Basics • Resonance • Formal Charge

  2. Localized Bond Models • Consider our energy diagram for H2 bonding:

  3. Localized Model Limitations • It is important to keep in mind that the models we are discussing are just that…..models. • We are operating under the assumption that when forming bonds, atoms “share” electrons using atomic orbitals. • Electrons involved in bonding: “bonding pairs”. Electrons not involved in bonding: “lone pairs”.

  4. Lewis Dot Structures (cont.) • 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]

  5. 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

  6. 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).

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

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

  9. LDS Mechanics (cont.) • An example: CO2 16 e- 12 e- left Octet Violation 0 e- left CO double bond

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

  11. LDS Mechanics (cont.) NO3- 24 e-

  12. 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

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

  14. 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.

  15. Resonance Structures (cont.) • NO3- is a classic example of resonance:

  16. Structural Isomers • What if different sets of atomic linkages can be used to construct correct LDSs: • Both are correct, but which is “more” correct?

  17. 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.

  18. 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

  19. FC -1 +1 -1 -2 +1 0 0 +1 -2 FC 0 0 0 -1 0 +1 +1 0 -1 Triatomic Bonding Patterns Compare CO2 with N3-…both 16 e- systems:

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