In this lecture were going to extend upon the Valence-Shell Electron Propulsion
model that we developed in the previous lecture and were going to apply it to
some specific properties of poly atomic molecules.
You'll recall that we've developed this model now by which we believe that the
geometry of the molecule was determined by separating as far as we can, the
electrons which are in pairs in the valence shell of the electron, of the
central atom. So over here, we have the methane
molecule and we've separated the four pairs of electrons which are in the
carbon-hydrogen bonds, and when we maximally separate them we get this
tetrahedral arrangement back over here. But recall, even though we don't have
four bonds in the water molecule or in the ammonium molecule We get an
arrangement or bond angles which are comparable in all three of these
molecules. What that says is we are not optimally
separating the bonds. We are optimally separating the pairs of
electrons because in all three of the mo the molecules on the screen here, there
are four pairs of valence electrons around the central atom.
Oxygen in this case, nitrogen in this case, and carbon in this case, and as a
result we wind up with a very similar geometries.
So, in our valence shell model it's important to note that bonded pairs,
bonded pairs of electrons and lone pairs of electrons have exactly the same effect
on the arrangement of the electron pairs in terms of determining what the geometry
of the molecule is going to be. Let's test this a little bit further
though by considering two other molecules.
Let's compare our ethane molecule which we have drawn back over here.
Which is C2H6. If we draw the Lewis structure for C2H6,
it's straightforward enough. Each carbon has three hydrogens
surrounding it. Each carbon, therefore has four pairs of
electrons around it. We expect a tetrahedral geometry, which
is exactly what we observe here. If we consider the Lewis structure for
C2H4, which is Ethene. That Lewis structure, recall, has a
double bond between the two carbons. But each carbon still has four pairs of
electrons around the valence shell. Here's the Lewis structure that I was
drawing. C2H4 has a double bond between the two
carbons. As a consequence, what we might want to
do is look and see, do we get a tetrahedral arrangement?
Do we get bond angles that are about 109.5?
And the answer turns out to be no. As we can see here from our data the bond
angle is actually closer to 120 degrees. In fact let's take a look at what the
molecules actually look at, look like. Here is the Ethene molecule and you can
see a couple of interesting things that distinguish it from the Ethane molecule.
One is that these bond angles are in fact, about a 120 degrees, the other is
its a perfectly planar molecule, all six atoms rest in the same plain, that's
quite different than the Ethane molecule were the molecules can only be viewed in
three dimensions, go back and look at Ethene here[SOUND] .
It's apparent that having the double bond changes the geometry here.
Let's also con consider, Ethyne, which is C2H2, the Lewis structure for that you
may recall from our earlier work, has a triple bond between the carbon atom so we
might draw it something like that. And if we look at the molecular geometry
for Ethyne, we discover the bond angle is around 180 degrees.
So, here we got about 120. Here we get about 180.
In fact, we get exactly 180. If we look at the molecular geometry of
the Ethyne molecule, you can see it here, it's a perfectly linear molecule,
straight down the line. And, you can see the bond angles, which
are there. What does that suggest?
So, in each case, we can see, there are four pairs of electrons around each
central carbon, and all three of the molecules, here.
In all three molecules, Each carbons got four pairs electrons around it.
And yet, the geometries of the molecules are very different.
It must be the case, that something has changed, and what's changed, of course,
is the appearance of a double bond or a triple bond.
Apparently, each double bond or triple bond, behaves differently, then a
separated pairs of electrons. But that actually makes a lot a sense,
because, I would not expect to be able to take this double bond and the ethene
molecule and separate these pairs of electrons.
I wouldn't be able to form a double bond if I did.
Both of these pairs must be oriented towards the, two carbons together.