Once the molecule file is fully loaded the image at right
will become live. At that time the "activate 3-D" icon
Valence Bond Model of the Double Bond
(ethylene as the example)
At right is a 3-D version of the Lewis structure of the ethylene
molecule. You may rotate the molecule yourself by holding the
mouse button down while dragging it around within the image
frame. Dark grey spheres represent C atoms and light grey
(white) spheres represent hydrogen atoms. Single bonds are
represented as single sticks and double bonds as a pair of
sticks. To return to this view at any time click on the button
Each carbon atom is surrounded by three "groups" as specified in the VSEPR
model for molecular shapes. An atom surrounded by three
groups will be sp2
hybridized to get orbitals pointing in the correct directions to
make bonds at 120 degrees from each other. Click the button below to
show these hybrid orbitals on one C.
Single bonds (σ-bonds) are made by overlapping hybrid orbitals pointed
along the bonds. The σ-bond that is part of the double bond in ethylene
is made by the overlapping sp2 hybrids on the neighboring C
atoms. Click on the button below to show the hybrid orbitals on the
other C atom and the overlap forming the σ-bond.
The bonds to the H atoms are made by overlapping the remaining hybrid
orbitals with the H 1s orbitals. Click on the button below to add
the H 1s orbitals and thus show the whole σ-bond network.
π-bonds are made by the side-on overlap of p orbitals on neighboring
atoms. To add the two p orbitals that are left over on the
carbons from the formation of the sp2 hybrid orbitals click
on the button below.
The yellow orbitals are the p orbitals.
To show how the two p orbitals combine to make a π-bond click on the
The red surface indicates how the two yellow p orbitals overlap. To
remove the p orbitals and leave the π-bond click on the button below.
This model shows the double bond as the σ-type overlap of the sp2
hybrid orbitals (in blue) plus the π-type overlap of the p orbitals (in
red). The electron density above and below the bond in the π-bond
prevents the two ends of the molecule from rotating relative to each
other about the C–C double bond; thus maintaining the planar geometry of the
To show just the π-bond click on the button below.
You may look at any of these intermediate views again by clicking on
the appropriate button.
Last Update: Dec. 24, 2013 J. Gutow
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