Chemical Bonding II­15.notebook
February 08, 2016
Drawing Lewis Dot Structures for Molecular Compounds
Predicting Shape, Bond Angle, and Polarity Follow these rules for drawing Lewis structures for molecules.
1.
count the total number of valence electrons
2.
draw the structure keeping in mind the following:
a)
b)
c)
d)
e)
f)
g)
h)
attach atoms bonded to the central atom with dashes (single) bonds. After drawing all the bonds and filling all octets, count the number of valence electrons in your structure. If they equal the number of valence electrons calculated, the structure is correct.
If your structure is short electrons, add them to the central atom (in pairs)
If your structure has too many valence electrons, remove two pairs to make a double bond.
C, N, O, and F as central atoms must have octets (8 electrons or four pairs)
H has a duet – one single bond only
The halogens as ligands (a ligand is any atom bonded to a central atom) can only form single bonds (and only one single bond) and must have octets
Be and B have less than an octet (i.e. they are electron deficient)
P and S may exceed the octet (they can use their d orbitals for additional bonding)
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Chemical Bonding II­15.notebook
February 08, 2016
Effective pairs: single bond, double bond, triple bond or lone pair all count as one effective pair (EP)
VSEPR – Valence Shell Electron Pair Repulsion Theory
This is a fancy way of saying that the bonding orbitals around a central atom will arrange themselves to minimize electron repulsions. This in turn will determine the shape of a molecule. Shape and Polarity
Note these two key ideas:
•
in molecules having a symmetric shape, the individual bond polarities will cancel and the molecule will overall be nonpolar (provided that all atoms bonded to the central atom are identical)
•
molecules having an asymmetric shape will be polar (in an asymmetric shape, the individual bond polarities cannot cancel each other out)
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Chemical Bonding II­15.notebook
Number of EP Around central
Atom
2
February 08, 2016
bonding
pairs
nonbonding VSEPR pairs
predicted shape
bond sym/asym P/NP
angle
2
0
180o sym
linear
beryllium chloride:
carbon dioxide
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Chemical Bonding II­15.notebook
Number of EP Around central
Atom
3
February 08, 2016
bonding
pairs
nonbonding VSEPR pairs
predicted shape
bond sym/asym
angle
3
0
120o sym
bonding
pairs
nonbonding VSEPR pairs
predicted shape
4
0
trigonal planar
boron trichloride
Number of EP Around central
Atom
4
bond sym/asym
angle
tetrahedron 109.5o sym
Carbon tetrachloride
CHCl3
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Chemical Bonding II­15.notebook
Number of EP Around central
Atom
4
February 08, 2016
bonding
pairs
nonbonding VSEPR pairs
predicted shape
bond sym/asym
angle
3
1
<109.5o
bonding
pairs
nonbonding VSEPR predicted pairs
shape
bond sym/asym
angle
2
2
<109.5o
trigonal pyramidal
asym
ammonia
Number of EP Around central
Atom
4
bent asym
Water
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Chemical Bonding II­15.notebook
February 08, 2016
Try these additional molecules: predict shape, bond angle, and overall polarity of molecule
Phosphorous trichloride:
Boron trihydride:
Hydrogen sulfide:
Carbon disulfide:
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Chemical Bonding II­15.notebook
February 08, 2016
Now let’s consider the Lewis structure for sulfur dioxide:
Notice we have one single and one double covalent bond between the S and O atoms. Based on this fact we would predict that the S=O double bond would be shorter in length compared to the S­O single bond. Two shared electron pairs will bring the nuclei closer to each other compared to sharing a single pair of electrons. Measurements have shown, however, that the bond length between the S and O atoms in SO2 is identical! How do chemists account for this discrepancy? We propose a model (theory) to account for this. The model is called resonance and it proposes that a pair of electrons “flips” back and forth between the S and O atoms.
Draw the Lewis structure for sulfur trioxide. Include all resonance forms.
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Chemical Bonding II­15.notebook
February 08, 2016
A Note on hybridization, sigma and pi bonds
Our model of bonding in molecules calls for individual atomic orbitals in atoms (like s and p orbitals) to combine or hybridize to provide bonding between atoms in a molecule. That is, when an atom goes to bond with another atom it hybridizes its atomic orbitals to produce a new set of hybrid bonding orbitals that are net lower in energy. Let’s take a look at a carbon atom as an example:
2p
E
2s
`
ground state
(two bonding orbitals)
hybridized state
(four bonding orbitals)
The four hybrid orbitals are made from one s and three p orbitals and are designetted as sp3 orbitals. (An atom that hybridized one s and two p orbitals would have three hybrid orbitals designated sp2 orbitals with one unhybridized p orbital). The overlap of hybridized orbitals result in sigma bonds while the overlap of unhybridized p
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Chemical Bonding II­15.notebook
February 08, 2016
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Chemical Bonding II­15.notebook
February 08, 2016
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