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Ch. 13 Bonding

Ch. 13 Bonding. Resonance Structures. O O O. Example 5. O 3 ozone O: 6 x 3 = 18 two completely equal arrangements the real structure is an average of these two where each bond is sharing 3 electrons instead of 4 or 2. O O O. O O O. O O O. Resonance Structures.

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Ch. 13 Bonding

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  1. Ch. 13 Bonding Resonance Structures

  2. O O O Example 5 • O3 • ozone • O: 6 x 3 = 18 • two completely equal arrangements • the real structure is an average of these two • where each bond is sharing 3 electrons instead of 4 or 2 O O O

  3. O O O O O O Resonance Structures • resonance – bonding between atoms that cannot be represented in one Lewis structure • show all possible structures with double-ended arrow in between to show that electrons are delocalized

  4. Example 6 • NO31- • N: 5 x 1 = 5 • O: 6 x 3 = 18 • total = 23 + 1 = 24

  5. Resonance Structures • experiments show that all the bonds are actually equal • actual structure is an average of the resonance structures • only the placement of the electrons can be different- not arrangement of atoms

  6. Formal Charge • Helps you to compare various Lewis Structures and choose the best or most likely structure • FC = (# valence e-) – (#e- assigned to it) Assigned: • all of the lone pairs • plus half of the bonding • Valence: • from periodic table

  7. Formal Charge • Not REAL: but provides less extreme charges than oxidation numbers • sum of the FC on molecule must equal overall charge on molecule • goal is to get atoms to FC of 0 • any negative FC must be on most electronegative atom

  8. Example • SO42- sulfate ion • 32 electrons -1 -1 -1 0 +2 0 -1 0 -1 -1

  9. Example • N2O • dinitrogen monoxide • 16 electrons O N N +1 +1 -2 N O N -1 +2 -1 N O N -2 +2 0 N O N 0 +2 -2

  10. Example • NO2 • nitrogen dioxide • 17 electrons

  11. Example • CO2 carbon dioxide • 16 electrons O C O 0 0 0 O C O 1+ 0 1-

  12. S C N -1 0 0 S C N +1 0 -2 S C N 0 0 -1 Example • SCN- thiocyanate ion • 16 electrons

  13. Example • ClO4- perchlorate ion • 32 electrons

  14. V alence S hell E lectron P air R epulsion repulsion between pairs of electrons around an atom cause them to be as far apart as possible used to predict the geometry of molecules VSEPR Theory

  15. Molecular Shapes • diatomic molecules will always be linear • all other molecules can have different shapes based on the steric number of central atom • steric number (number of total electron pairs) includes: • bonding pairs • lone pairs

  16. Parent Geometry

  17. Tips • Draw Lewis Structure • Find parent geometry • Picture without lone pairs to get molecular shape • lone pairs take up more space than bonding pairs • treat all bonds same

  18. (b) is more stable because the lone pairs are further apart

  19. Steric Number 2 • no lone pairs: linear • CO2

  20. no lone pairs: trigonal planar CH2O 1 lone pair: MS: bent SO2 Steric Number 3

  21. 1 lone pair: NH3 trigonal pyramidal no lone pairs: CH4 tetrahedral 2 lone pairs: H2O bent Steric Number 4

  22. Bent Trigonal Pyramidal

  23. Bond Angles

  24. no lone pairs: trigonal bipyramidal PCl5 1 lone pair: seesaw SF4 Steric Number 5

  25. 2 lone pairs: T-shaped ClF3 3 lone pair: Linear I3- Steric Number 5

  26. 2 lone pairs: square planar XeF4 1 lone pair: square pyramidal SbCl52- no lone pairs: octahedral SF6 Steric Number 6 Cl Cl Sb Cl Cl Cl

  27. Where should the lone pairs go?

  28. Practice • Quiz 1 • Quiz 2 • Quiz 3

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