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Chapter 1 – Electronic Structure and Covalent Bonding
Section 1.6 – 1.14
Why do atoms form covalent bonds? The answer is STABILITY
Consider the Hydrogen molecule, H2.
As the two orbitals start to overlap to form the covalent bond, energy is released (and
stability increases) because the electron in each atom is attracted to its own nucleus and
to the positive charge of nucleus of other atom
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Bond Strength is given by the standard bond dissociation energy (ΔHo or BDE). It is
defined as the energy required breaking a covalent bond in a hemolytic cleavage.
ΔHo = 104 kcal/mol
The greater the electron density in the region of orbital overlap, the stronger the bond
Two special names for covalent bonds of Organic molecules
Sigma (σ) bonds
Pi (π) bonds
Created when “head on” overlap occurs of
orbital
Created when “side on” overlaps occurs of
orbital (p orbitals)
Pi bonds are usually weaker than sigma bon. From the perspective of quantum mechanics,
this bond's weakness is explained by significantly less overlap between the component porbitals due to their parallel orientation. This is contrasted by sigma bond which form
bonding orbitals directly between the nucleus of the bonding atoms, resulting in greater
overlap and a strong sigma bond.
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The Hybridization Model for Atoms in Molecules
If the four hydrogen atoms in a methane molecule (CH4) were bound to the three 2p orbitals
and the 2s orbital of the carbon atom, what would be the angle about H-C-H?
The following molecules provide examples of all three basic shapes found in organic
chemistry. In these drawings, a simple line indicates a bond in the plan of the paper, a
wedged line indicates a bond coming out in front of the page and a dashed line indicated a
bond projecting behind the page
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Hybridization describes the mixing of atomic orbitals to form special orbital for bonding.
In organic chemistry, our orbital mixtures will be simple combinations of valence
electrons in the 2s and 2p orbital on a single carbon atom. We will mix these orbitals
three ways to generate the three common shape of organic chemistry: linear (2s + 2p),
trigonal planar (2s+2p+2p) and tetrahedral (2s+2p+2p+2p)
The sp Hybridization
- Mixing 2s and one 2p atomic orbitals
- Two sp hybridized orbitals equal in size, energy and shape
- one σ bond ( single bond)
- Two π bond (triple bond)
- Linear
Consider ethyne (C2H2)
3-D representation of ethyne (C2H2)
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The sp2 Hybridization
-
Mixing 2s and two 2p atomic orbitals
Three sp2 hybridized orbitals equal in size, energy and shape
Responsible for σ bond ( single bond)
One π bond (double bond)
Trigonal planar shape
Consider ethane (C2H4)
3-D representation of ethane (C2H4)
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sp3 Hybridization
- Mixing 2s and all 2p atomic orbital
- Four sp3 hybridized orbitals equal in size, energy and shape
- Responsible for sigma bond ( single bond)
- Tetrahedral shape
Consider methane (CH4)
3-D representation of methane (CH4)
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Bonding to O and N
Like Carbon, O and N can participate in single bond and multiple bonds compose of σ and
π
*Note: the lone pair or non bonding e- pair occupies space just as bonded atom
Hybridization for radical carbon and carbon with charges
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E.g Indicate the hybridization at carbon and oxygen, the angle at H-C-O and C-O-H.
Then draw a bonding picture for the formation of sigma in CH3OH
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E.g Assign hybridization for all carbon atoms and identify the angle at C-C-C. Then draw
bonding pictures for the σ and π framework.
E.g Predict the hybridization, geometry, and bond angle for the carbon and oxygen atoms
in acetonitrile (
). Draw a bonding picture for the formation of pi bond(s)
E.g Predict the hybridization, geometry, and bond angle for the carbon and oxygen atoms
in acetaldehyde (CH3CHO)
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