HF H: 1s 1 F: 1s 2 2s 2 2p 5 Overlap between the valence orbital of H (1s) and valence orbital of F...
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Transcript of HF H: 1s 1 F: 1s 2 2s 2 2p 5 Overlap between the valence orbital of H (1s) and valence orbital of F...
HF H: 1s1 F: 1s22s22p5
Overlap between the valence orbital of H (1s) and valence orbital of F (2p) to form a bonds
Note: electron spin is paired in the orbital
By definition: z is the direction along the internuclear axis
N2 N: 1s22s22p3
The two pz orbitals from each N can overlap to form a orbital.
The px and py orbitals are perpendicular to the internuclear axis
bond - overlap of two pz orbitals
bond - overlap of two px orbitals and/or two py orbitals
In a bond, electron density has a nodal plane that contains the bond axis
According to the VB theory
A single bond is a -bond
A double bond is -bond plus a -bond
A triple bond is a -bond and two -bonds.
VB theory: assumes bonds form
when unpaired electrons in valence shell atomic orbitals pair
the atomic orbitals overlap end to end to form -bonds or side by side to form -bonds.
Hybridization of Orbitals
C: Is2 2s2 2p2
VB theory, as described so far, would predict that C can form just two bonds
In CH4, C forms four bonds.
C needs four unpaired electrons so that each can pair with a H atom - need to revise valence-bond theory
“Promote” a 2s electron to a 2p orbital - this requires energy.
But now C has four unpaired electron and since bonding releases energy the cost of promoting is overcome by the lowering of energy on bond formation
CH4
Promoting a 2s electron to 2p allows C to have four unpaired electrons.
All bonds on CH4 are equivalent
“Mix” the 2s and the three 2p orbitals to form four hybrid orbitals all of the same energy and spatial distribution - hybridization.
One s + three p = four sp3 orbitals
C
H
HH
H
Hybrid orbitals are constructed on an atom to reproduce the electron arrangement of the experimentally determined shape of the molecule.
In CH4: each sp3 orbital has one unpaired electron
Each overlaps with a 1s orbital of H to form -bond
The four resulting -bonds point towards the corners of a tetrahedron.
All four -bonds are identical
http://www.whfreeman.com/chemicalprinciples/
Ethane: C2H6
Each C has four sp3 hybrid orbitals, pointing towards the corner of a tetrahedron, each with one electron
Three of these four overlap with three H atoms forming -bonds (C sp3, H 1s).
The C-C bond is formed by an overlap of the remaining sp3 orbital on each C forming a -bond (C sp3, C sp3).
NH3
H: Is1 N: Is2 2s2 2p3
Hybridize the 2s and 2p orbitals in N to form four sp3 hybrid orbitals.
One of the sp3 has two paired electrons - the lone pair on N
The three other sp3 orbitals form s-orbitals with each of the three H 1s orbitals
NH
H
H
H2O
H: Is1 O: Is2 2s2 2p4
Hybridize the 2s and 2p orbitals in O to form four sp3 hybrid orbitals.
O 2p
2s
O sp3
Two of the sp3 have two paired electrons - the two lone pairs
The two other sp3 orbitals overlap with H 1s orbitals
OH H
An s orbital and two p orbitals can hybridize to form three sp2 hybrid orbitals which point to the corners of an equilateral triangle - trigonal planar geometry
Example: BF3
An s and a p orbital can hybridize into two sp orbitals that point in opposite directions - linear geometry
PCl5 P: [Ne] 3s2 3p3 Cl: [Ne] 3s2 3p5
P Cl 3p 3p
3s 3s
Promote a 3s electron to the 3d orbital
P _ _ _ _
sp3d empty 3d
Valence shell expansion - expansion to include d orbitals along with s and p orbitals
One 2, three p, and one d orbital form five sp3d hybrid orbitals, each pointing towards a corner of a trigonal bipyramid
One 2, three p, and two d orbital form six sp3d2 hybrid orbitals, each pointing towards a corner of a octahedron
SF6 S: [Ne] 3s2 3p4 F: [He] 2s2 2p5
S
3p
3s
Include two 3d orbital and hybridize one s, three p and two d
S _ _ _ sp3d2 empty 3d
Multiple Bonds
Ethylene: CH2CH2
Experimental data: all six atoms lie in the same plane and the H-C-H and C-C-H bond angles are 120o.
Trigonal planar geometry indicates that each C is sp2 hybridized
For each C: two of the sp2 orbitals bond with two H 1s orbitals to form -bonds, The third Csp2 bond on each bond with each other to form a C-C -bond
The “pure” 2p orbitals on each C overlap to form a -bond between the two C atoms
The electron density of this -orbital lies above and below the axis of the C-C -bond
http://www.whfreeman.com/chemicalprinciples/
Acetylene: C2H2
Linear molecule; each C is sp hybridized, leaving two pure p orbitals on each C
http://www.whfreeman.com/chemicalprinciples/
Multiple bonds are formed when an atom forms a -bond by using an sp or sp2 hybrid orbital and one or more -bonds by using un-hybridized p orbitals
Formic acid: HCOOH