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The Nature of the Chemical Bond
Figure 1:
The formation of a single covalent bond in a
hydrogen molecule by the overlap of two 1s orbitals
or individual hydrogen atoms. This represents a
new, lowerenergy state of the two atoms.
ANIMATION OF ELECTRON DENSITY IN THE BOND REGION
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Figure 2:
A hydrogen atom has only one occupied orbital, the
1s orbital. For simplicity, only the one halffilled p
orbital of the fluorine atom is shown. The final
combined, bonding orbital contains a pair of
electrons and the fluorine atom now has a complete
octet.
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Figure 3:
Two overlaps are possible to produce two shared
pairs of electrons forming two covalent bonds. As
before, the oxygen atom completes its octet of
electrons.
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Hybrid Orbitals
Consider the 2nd energy level for a groundstate carbon atom.
It would appear that carbon has one lone pair and only 2
bonding electrons. Why does it form four bonds? Linus Pauling
was the first to suggest “electron promotion”. One s electron
gets promoted to the empty p orbital. He justified this by
suggesting that the energy gained by the molecule when it
bonds would be greater than the energy required for promotion
to a slightly higher energy level. However, experimental
evidence suggests that the electron orbitals are all equivalent in
shape and energy there isn’t one that’s more “s” than the
others. As a result, the four bonds for carbon in molecules such
as methane are explained by hybridization to four identical sp3
atomic orbitals.
These hybridized orbitals only exist when bonding occurs and
are NOT found in an isolated atom.
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Example 1
What are the bonding orbitals and the structure of the BF3
molecule?
Solution
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Example 2
Provide the groundstate and the promoted state configurations
for beryllium and then describe the bonding and structure of a
BeH2 molecule.
Solution
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Double and Triple Covalent Bonds
sigma (σ)bond:a bond created by the end to end overlap of
atomic orbitals (like those below)
Figure 4:
Sigma bonds form with the overlap of
(a) s orbitals (b) p orbitals and (c) hybrid orbitals.
ANIMATION OF SIGMA BOND FORMATION
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pi (π) bond:
a bond created by the sidebyside (or parallel)
overlap of atomic orbitals (usually p orbitals)
Figure 5:
p orbitals form with the sidebyside overlap of
orbitals
ANIMATION OF PI BOND FORMATION
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Double Bonds
The carbon atom is the most common central atom in molecules
with double and triple covalent bonds. It is thought that there is a
partial hybridization of the available orbitals leaving one or two p
orbitals with single unpaired electrons.
Figure 6: For this carbon atom,
the sp2 hybrids are
planar at 120° to
each other and the p
orbital is at right
angles to the plane of
the hybrid orbitals.
For example, after promoting
an electron in carbon’s 2s
orbital to a 2p orbital, we form
only three sp2 hybridized
orbitals but three of these are
hybrids and one is a “normal”
p orbital.
ANIMATION OF ETHYLENE BONDS
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Figure 7:
(a) The sigma bonds for a C2H4
molecule use the sp2 hybrid
orbitals.
(b) The two halffilled p
orbitals of the adjacent
carbon atoms overlap sideways.
(c) The complete
bonding orbitals for a
C2H4 molecule.
Therefore, a double bond consists of a σ bond and a π bond.
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Triple Bonds
For C2H2:
Figure 8: Instead of mixing all four orbitals, valence bond
theory suggests that only two are mixed to form sp hybrid
orbitals and two unhybridized p orbitals for a carbon atom.
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Figure 9: (a)
The sigma bonds for a C2H2 molecule
use the sp hybrid orbitals.
(b) The two pairs of halffilled
p orbitals of the adjacent
carbon atoms overlap sideways.
(c) The complete bonding
orbitals for a C2H2 molecule.
Therefore, triple bonds involve a σ bond and 2 π bonds.
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Sigma Bonds in Ethylene.mov
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