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I. VSEPR = Valence Shell Electron-Pair Repulsion Electron pairs repel each otherPowerPoint Presentation

I. VSEPR = Valence Shell Electron-Pair Repulsion Electron pairs repel each other

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4. Double and triple bonds are treated as only 1 pair of electrons

sp2 Hybridization

- I. VSEPR = Valence Shell Electron-Pair Repulsion
- Electron pairs repel each other
- This applies to bonding pairs and lone pairs alike
- Steps in Applying VSEPR
- Draw the Lewis Structure
- Count atoms and lone pairs and arrange them as far apart as possible
- Determine the name of the geometry based only on where atoms are

- Electron pairs repel each other

- B. Complications to VSEPR
- Lone pairs count for arranging electrons, but not for naming geometry
- Example: NH3 (ammonia)
- Lone pairs are larger than bonding pairs, resulting in adjusted geometries
- Bond angles are “squeezed” to accommodate lone pairs

Trigonal pyramidal

Tetrahedral

linear

trigonalplanar

bent

trigonal pyramidal

tetrahedral

bent

trigonalbipyramidal.

T-shaped

linear

see-saw

octahedral

square pyramidal

square planar

II. Hybridization

- Localized Electron Model
- Molecules = atoms bound together by sharing e- pairs between atomic orbitals
- Lewis structures show the arrangement of electron pairs
- VSEPR Theory predicts the geometry based on e- pair repulsion
- Hybridization describes formation and properties of the orbitals involved in bonding more specifically

- sp3 Hybridization
- Methane, CH4, will be our example
- Bonding only involves valence e-
- a. H = 1s
- b. C = 2s, 2p

- If we use atomic orbitals directly, the predicted shape doesn’t match what we predict by VSEPR and what we observe experimentally
- One C2s--H1s bond would be shorter than others
- Three C2p--H1s bonds would be at right
- angles to each other
- We know CH4 is tetrahedral, symmetric

- Free elements (C) use pure atomic orbitals
- Elements involved in bonding will modify their orbitals to reach the minimum energy configuration

- When C bonds to four other atoms, it hybridized its 2s and 2p atomic orbitals to form 4 new sp3 hybrid orbitals
- The sp3 orbital shape is between 2s/2p; one large lobe dominates
- When you hybridize atomic orbitals, you always get an equal number of new orbitals

- The new sp 2p atomic orbitals to form 4 new sp3 orbitals are degenerate due to the mixing
- The H atoms of methane can only use their 1s orbitals for bonding. The shared electron pair can be found in the overlap area of H1s—Csp3

- NH 2p atomic orbitals to form 4 new sp3 and H2O also use sp3 hybridization, even though lone pairs occupy some of the hybrid orbitals.

- C2H4, ethylene is our example
- Lewis and VSEPR structures tell us what to expect
- H atoms still can only use 1s orbitals
- C atom hybridizes 2s and two 2p orbitals into 3 sp2 hybrid orbitals

- The new sp 2p atomic orbitals to form 4 new sp2 orbitals are degenerate and in the same plane
- One 2p orbital is left unchanged and is perpendicular to that plane
- One C—C bond of ethylene forms by overlap of an sp2 orbital from each of the two sp2 hybridized C atoms. This is a sigma (s) bond because the overlap is along the internuclear axis.
- The second C—C bond forms by overlap of the remaining single 2p orbital on each of the carbon atoms. This is a pi (p) bond because the overlap is perpendicular to the internuclear axis.

- 9. Pictorial views of the orbitals in ethylene 2p atomic orbitals to form 4 new sp

- sp Hybridization 2p atomic orbitals to form 4 new sp
- CO2, carbon dioxide is our example
- Lewis and VSEPR predict linear structure
- C atom uses 2s and one 2p orbital to make two sp hybrid orbitals that are 180 degrees apart
- We get 2 degenerate sp orbitals and two unaltered 2p orbitals

- Oxygen uses sp 2p atomic orbitals to form 4 new sp2 orbitals to overlap and form sigma bonds with C
- Free p orbitals on the O and C atoms form pi bonds to complete bonding

- d-orbitals can also be involved in hybridization 2p atomic orbitals to form 4 new sp
- dsp3 hybridization in PCl5
- d2sp3 hybridization in SF6

F. The Localized Electron Model 2p atomic orbitals to form 4 new sp

1. Draw the Lewis structure(s)

2. Determine the arrangement of electron pairs (VSEPR model).

3. Specify the necessary hybrid orbitals.

- Polar Bonds 2p atomic orbitals to form 4 new sp
- A. Any bond between atoms of different electronegativities is polar
- Electrons concentrate on one side of the bond
- One end of the molecule is (+) and one end is (-)

- B. Complex molecules: vector addition of all bond dipoles

- A. Any bond between atoms of different electronegativities is polar

- Examples: 2p atomic orbitals to form 4 new sp
- Polarity greatly effects molecular properties

Oppositely charged ends of polar molecules

Attract each other like opposite poles of magnets

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