0:14
The VSEPR model allows us to predict the 3D shape, or
geometry, of molecules in space.
When we draw molecules on paper,
such as what we did with Lewis structures, those are shown in two dimensions and
they really don't show us what a molecule actually looks like.
And so, VSEPR model allows us to make those predictions.
And VSEPR is an acronym that stands for valent shell electron pair repulsion.
So what we want to think about is this repulsion part is kind of like a magnet.
If I put the two light poles of a magnet together,
try to force them together, I feel that repulsion between them.
Well, the same thing is going on with our electrons.
So, when we put two pairs of electrons, or
two groups of electrons near each other, they're both negatively charged.
They're going to have repelling forces between them.
The other part of the name the valence shell part is because we're only
looking at the valence electrons.
Just as we looked at the Lewis structures,
we were worried about the valence electrons.
It's those that we're concerned about when we look at the VSEPR geometry,
because it's those electrons that are involved in the bonding.
Remember that our core electrons don't actually interact with other atoms,
it's the valence electrons that we have to be concerned about.
1:28
So when we look at VSEPR geometry, we need to first draw our Lewis structure,
because without the Lewis structure we don't understand how those
electron groups are arranged around that central atom.
And so we're going to look at some examples here of how many electron
groups we have around the central atom.
And notice here that we keep talking about groups.
Even though the name VSEPR includes electron pair,
what we really want to look at are the number of groups.
And so when we look at a molecule, such as H2Cl, formaldehyde.
What I'm looking at here is, I have one group here, there's a pair of electrons.
A second group there, and this double bond is actually a single group.
So we say that, that molecule has three electron groups around the central atom.
Our multiple bonds count as a single group.
When I look at something like water, here I have two single bonds, so
that's two groups.
But I also have two nonbonding pairs.
And so here, I have four groups around my central atom.
For our molecule here, we have one, two, three, four.
2:36
Here, for ICl5, I actually have six groups because I have five bonding groups and
one non-bonding group.
So I've got six groups there.
Here for SCN minus, I have two groups.
Because I have a double bond and a double bond.
Each of those double bonds only counts as a single group.
Now notice in many of these molecules some of the terminal atoms
have lone pairs on them.
However, I'm not worried about those electrons or those lone pairs.
When I'm worried about geometry,
I'm only looking at what's going on around the central atom.
Because those are the only electron groups that are going to
actually affect the geometry or the 3D arrangement of those bonds.
When I get to my last one, CH3Cl again, I see I have four groups, okay?
I don't have to worry about these electrons around chlorine,
because they're not going to affect the geometry of carbon.
So I'm only looking at the electron groups around my central atom, okay?
So focusing on the central atom and knowing that if I have a non bonding group
on my central atom, a lone pair, that counts as a single group.
If I have two pairs, of non-bonding electrons that counts as two groups.
And each bonding group counts as one.
Whether it be a single bond, a double bond, or
a triple bond, it's still considered a single group.
4:00
So when we look at our geometry this is an example of some of
the geometries we're going to look at and I want to talk about how we name them.
But the key thing we're worried about here is how these
angles form to get the maximum distance between these electrons groups.
And so these are all examples where we're only dealing with bonding groups.
We're not going to deal with non bonding groups just yet,
we'll get to that shortly ,but here we're looking at bonding groups.
So if I have my central atom, and I have two groups around my central atom,
the way that they can be as far apart from one another as possible.
Is for them to be 180 degrees from one another.
When I look at my next geometry,
where I have three groups, now, I'm looking at 120 degree angle.
Because that's how these groups that can be as far apart from one
another as possible.
When I go beyond having three groups, when I go to four groups,
things change a little bit.
Because here looking at two groups and
three groups, I'm looking at situations where I can do them in a single plane.
Now when I get to the four groups around the central atom, now I have to look at
my bond angles but I have to really go into that sphere shape, okay.
So, I imagine a kind of a sphere encircling this molecule.
And now, when I look at these bond angles the way that they get
as far apart from one another is possible is to a 109.5 degrees.
And that angle’s a little bit harder for us to kind of picture in our head.
We're used to 90 and 120 or
180, but 109.5, it doesn't quite kind of stick out in our head very well.
We'll talk a little bit more about the other two with five groups and
six groups when we get to those later.
Now we're going to spend a little bit more time looking at what happens when we
have lone pairs, so we have four electron groups.
But what happens when they're not all bonding groups around that central atom?
5:52
So let's look at three examples here.
We have methane, CH4, ammonia, NH3, and water, H2O.
And so, we also have a pair, two lone pairs there on our oxygen.
So in each of these they have four groups around them.
Here we have four bonding, here we have three bonding,
and one non-bonding.
And in water, we have two bonding and two non-bonding.
So we still have the same basic shape.
We see here our bond angles are going to be 109.5 degrees here
between any pair of bonds there.
And that's the ideal situation because we
have four identical bonds around that carbon atom.
And so they're going to spread out completely evenly nice and symmetrical.
Now, when I replace one of those bonding groups with a non-bonding group.
So here I have a non-bonding group or a lone pair of electrons on my nitrogen.
What I see is that now my angles actually get compressed a little bit.
And I get down to about 107.8.
And what's happening is, is that remember, this electron group that is
between the nitrogen and the hydrogen is kind of involved in this bonding, and
so it's kind of got other things to worry about.
But remember when I had that lone pair, it's not involved in a bond,
it's not interacting with another atom.
And so as a result, this lone pair of electrons has more repelling power.
Then do the bonding electrons.
And so what's happening is, that we can kind of think
of this lone pair is being larger than any of our bonding groups.
And as a result it's going to start squishing and
compressing these bonding groups together.
And so what we get is a bond angle that is a little bit less 109.5.
So this is our ideal, okay.
Four identical groups around the central atom.
8:05
Here we have two lone pairs, so now if we think about their
electron density we can see they are going to be repelling between each lone pair,
because again those are close by each other in the molecule, and
they're also going to be repelling towards these bonding groups, and
as a result we're going to push or compress these two atoms, together.
And get an even smaller bonding.
So 104.5 or approximately, and so the more nonbonding groups we have,
the smaller our angles are going to get relative to the ideal angle.
And it's not so
important that you memorize the specific values that it's 107.8 or 104.45.
What matters is that you realize that the bond angle of
the NH3 is going to be smaller than the bond angle is the CH4.
And the bond angle in water is going to be smaller than the bond angle of
the NH3 because of the affect of those lone pairs of electrons.
9:09
We see something similar when we look at something like formaldehyde how we
saw this structure before.
This is H2CO, and when we look at those bond angles,
here we have two pairs of electrons and still just one electron group.
So we have three electron groups, and so we know our ideal angle.
9:30
There's 120 degrees.
So our ideal angle's 120 degrees.
But because we have two pairs of electrons, or four electrons there,
what we see is that causes more repelling power than a single bond.
Now, not as much as having a lone pair, but
a little bit more than the single bond.
As a result, it's going to push these atoms together a little bit,
and so we actually get compression of this angle, so we end up
with a bond angle of about 115.5 degrees, compared to our ideal which is 120.
And we also see that this bond angle has gotten a little bit larger.
Now in the previous slide we didn't talk about the bond angle with the lone pairs
because there's nothing there to measure the angle with.
We have to have atoms there in order to measure those bond angles, and so
here we see a smaller bond here we see a larger bond.
10:32
So here our bond angle is going to be a little bit greater than 120, because we
have this double bond, and as a result we have more electrons than there and so
we're going to have more repelling power and so it's actually going to
push these two groups together, so we'll have a smaller bond here.
But this bond will be greater than 120 degrees.
We're again not worried about the specific amount just
knowing whether it's greater than or less than that ideal bond angle.
Dr. Allison SoultLecturer
Dr. Kim WoodrumSr. Lecturer
