Transcript Chapter 10

Chapter 10
VSEPR - Lewis structures do not help us predict the shape
or geometry of molecules; only what atoms and bonds are
involved. To predict shape another system was developed
that works for many molecules; Valence Shell Electron Pair
Repulsion Theory. The name is derived from the idea that
electron pairs around the central atom (the valence
electrons) want to be as far away from each other (because
of electron repulsion).
Rules:
1. Draw Lewis Structure
2. Count the # of electron pairs around central atom
(both bonding & non-bonding)
Bonding – electrons that are part of a bond.
Non-bonding – electrons that are not in a bond (also
called lone pairs)
(also, for this method we count multiple bonds as 1
pair)
3. This # determines the basic shape to begin with.
2 e-pairs -> linear - bond angle = 180
Bond Angle – The angle formed between the central
atom and 2 terminal atoms.
3 e-pairs -> trigonal planar (a triangle with central atom at
center - bond angle = 120
4 e-pairs -> tetrahedral (illustrate) with central atom in the
middle - bond angle = 109.5
Shape is described this way (i.e. pretending to
connect the terminal atoms)
4. Place e-pairs (with atoms where applicable) at each
vertex of the geometric shape.
4. Actual shape of molecule is determined by the position
of atoms not electrons.
Important to understand that the actual shape is
determined only by actual atoms. We “see” only atoms,
not electrons.
An interesting point is that lone pairs take up more space
than bonding pairs. This may result in bonded atoms
being pushed closer to each other (since bonded pairs
take up less space, they can be closer to each other. An
example is H2O. The bond angle should be 109.5,
(based on 4 electron pairs around O, but only 2 atoms).
The actual bond angle is about 105.
Let’s do some examples. Note that 5 and 6 e-pair systems
exist but we won’t worry about them for now. We will only
consider systems with a maximum of 4 electron pairs
around the central atom.
Let’s do example 10.1, on page 408 (a,b,c and e)
We learned earlier about polar and non-polar bonds. When
a molecule is composed of several bonds, the molecule can
also be polar or non-polar. The definitive diagnosis is
whether or not the molecule has a dipole moment. For our
purposes, this simply means that if there is a net positive
end and a net negative end for the molecule, it has a dipole
moment and is polar.
Many factors contribute to whether a molecule as a whole is
polar or non-polar. If the molecule is composed of all nonpolar bonds, than obviously the molecule, as a whole, will
be non-polar. But even when there are polar bonds in the
molecule, it can be non-polar as a whole. This occurs when
the shape of the molecule is perfectly symmetrical, and
therefore the polarity of the individual bonds cancel out.
Some helpful hints will make your job easier:
My rules for non-polar
1. If all bonds are non-polar, molecule is non-polar
2. All hydrocarbons are non-polar
A hydrocarbon is any molecule consisting of
C and H and nothing else. This is very common as you
will see in organic chemistry.
3. If one of the basic shapes, & if every vertex occupied by
an identical atom, then non-polar
4. If all else fails, does it look symmetrical?
The Lewis structure application along with the VSEPR
model works quite well in predicting structures but to
chemists does not really go far enough. They do not
really explain why chemical bonds exist or why CH4 has 4
equal bonds forming a perfect tetrahedron.
There are 2 Quantum Mechanical models that help here.
We will look at one of them, the Valence Bond Theory.
Basically, in the valence bond model, every atom in a
molecule has atomic orbitals just as they would as free
atoms. Bonding occurs when an atomic orbital on one
atom overlaps with an atomic orbital of another atom. As
long as there is significant overlap (technically are the
energies with the overlap lower than the separate atomic
orbitals) then bonding will occur. It occurs at the point of
maximum overlap. See figure 2-4 below
In order to account for molecular geometry it frequently
becomes necessary to define some of the atomic orbitals
around the central atom differently. It has been discovered
through experimental evidence that when an atom, such as C
bonds to other atoms, in order to obtain maximum
overlapping, the C atom combines s and p orbitals to form
new hybrid orbitals. This process is called hybridization.
These depend on how many atoms are bonding to the C
atom:
4 atoms - sp3 – It is named because it involves mixing of
one s orbital with 3 p orbitals The shape of these new
orbitals will be tetrahedral. See Figure below:
The picture above shows how 4 sp3 bonds to 4 H atoms using
the 1s orbitals of the H atoms, forming 4 identical bonds.
3 atoms - sp2 - trigonal planar and will usually involve one
double bond. The second bond of a double bond is called a
pi bond and has less overlap that a sigma (first bond) bond
and is therefore weaker. It results from the parallel overlap
of p orbitals on both atoms:
2 atoms - sp - linear and will usually involve one triple or 2
double bonds.
We won’t worry about hybridization involving d orbitals
In valence bond theory, the first bond between 2 atoms is the
strongest, because of maximum overlap and are called sigma
() bonds. Any additional bonds between the same 2 atoms
are weaker because of reduced overlap and are called pi ()
bonds.