0:14

The VSEPR model allows us to predict the 3D shape, or

Â geometry, of molecules in space.

Â When we draw molecules on paper,

Â such as what we did with Lewis structures, those are shown in two dimensions and

Â they really don't show us what a molecule actually looks like.

Â And so, VSEPR model allows us to make those predictions.

Â And VSEPR is an acronym that stands for valent shell electron pair repulsion.

Â So what we want to think about is this repulsion part is kind of like a magnet.

Â If I put the two light poles of a magnet together,

Â try to force them together, I feel that repulsion between them.

Â Well, the same thing is going on with our electrons.

Â So, when we put two pairs of electrons, or

Â two groups of electrons near each other, they're both negatively charged.

Â They're going to have repelling forces between them.

Â The other part of the name the valence shell part is because we're only

Â looking at the valence electrons.

Â Just as we looked at the Lewis structures,

Â we were worried about the valence electrons.

Â It's those that we're concerned about when we look at the VSEPR geometry,

Â because it's those electrons that are involved in the bonding.

Â Remember that our core electrons don't actually interact with other atoms,

Â it's the valence electrons that we have to be concerned about.

Â 1:28

So when we look at VSEPR geometry, we need to first draw our Lewis structure,

Â because without the Lewis structure we don't understand how those

Â electron groups are arranged around that central atom.

Â And so we're going to look at some examples here of how many electron

Â groups we have around the central atom.

Â And notice here that we keep talking about groups.

Â Even though the name VSEPR includes electron pair,

Â what we really want to look at are the number of groups.

Â And so when we look at a molecule, such as H2Cl, formaldehyde.

Â What I'm looking at here is, I have one group here, there's a pair of electrons.

Â A second group there, and this double bond is actually a single group.

Â So we say that, that molecule has three electron groups around the central atom.

Â Our multiple bonds count as a single group.

Â When I look at something like water, here I have two single bonds, so

Â that's two groups.

Â But I also have two nonbonding pairs.

Â And so here, I have four groups around my central atom.

Â For our molecule here, we have one, two, three, four.

Â 2:36

Here, for ICl5, I actually have six groups because I have five bonding groups and

Â one non-bonding group.

Â So I've got six groups there.

Â Here for SCN minus, I have two groups.

Â Because I have a double bond and a double bond.

Â Each of those double bonds only counts as a single group.

Â Now notice in many of these molecules some of the terminal atoms

Â have lone pairs on them.

Â However, I'm not worried about those electrons or those lone pairs.

Â When I'm worried about geometry,

Â I'm only looking at what's going on around the central atom.

Â Because those are the only electron groups that are going to

Â actually affect the geometry or the 3D arrangement of those bonds.

Â When I get to my last one, CH3Cl again, I see I have four groups, okay?

Â I don't have to worry about these electrons around chlorine,

Â because they're not going to affect the geometry of carbon.

Â So I'm only looking at the electron groups around my central atom, okay?

Â So focusing on the central atom and knowing that if I have a non bonding group

Â on my central atom, a lone pair, that counts as a single group.

Â If I have two pairs, of non-bonding electrons that counts as two groups.

Â And each bonding group counts as one.

Â Whether it be a single bond, a double bond, or

Â a triple bond, it's still considered a single group.

Â 4:00

So when we look at our geometry this is an example of some of

Â the geometries we're going to look at and I want to talk about how we name them.

Â But the key thing we're worried about here is how these

Â angles form to get the maximum distance between these electrons groups.

Â And so these are all examples where we're only dealing with bonding groups.

Â We're not going to deal with non bonding groups just yet,

Â we'll get to that shortly ,but here we're looking at bonding groups.

Â So if I have my central atom, and I have two groups around my central atom,

Â the way that they can be as far apart from one another as possible.

Â Is for them to be 180 degrees from one another.

Â When I look at my next geometry,

Â where I have three groups, now, I'm looking at 120 degree angle.

Â Because that's how these groups that can be as far apart from one

Â another as possible.

Â When I go beyond having three groups, when I go to four groups,

Â things change a little bit.

Â Because here looking at two groups and

Â three groups, I'm looking at situations where I can do them in a single plane.

Â Now when I get to the four groups around the central atom, now I have to look at

Â my bond angles but I have to really go into that sphere shape, okay.

Â So, I imagine a kind of a sphere encircling this molecule.

Â And now, when I look at these bond angles the way that they get

Â as far apart from one another is possible is to a 109.5 degrees.

Â And that angleâ€™s a little bit harder for us to kind of picture in our head.

Â We're used to 90 and 120 or

Â 180, but 109.5, it doesn't quite kind of stick out in our head very well.

Â We'll talk a little bit more about the other two with five groups and

Â six groups when we get to those later.

Â Now we're going to spend a little bit more time looking at what happens when we

Â have lone pairs, so we have four electron groups.

Â But what happens when they're not all bonding groups around that central atom?

Â 5:52

So let's look at three examples here.

Â We have methane, CH4, ammonia, NH3, and water, H2O.

Â And so, we also have a pair, two lone pairs there on our oxygen.

Â So in each of these they have four groups around them.

Â Here we have four bonding, here we have three bonding,

Â and one non-bonding.

Â And in water, we have two bonding and two non-bonding.

Â So we still have the same basic shape.

Â We see here our bond angles are going to be 109.5 degrees here

Â between any pair of bonds there.

Â And that's the ideal situation because we

Â have four identical bonds around that carbon atom.

Â And so they're going to spread out completely evenly nice and symmetrical.

Â Now, when I replace one of those bonding groups with a non-bonding group.

Â So here I have a non-bonding group or a lone pair of electrons on my nitrogen.

Â What I see is that now my angles actually get compressed a little bit.

Â And I get down to about 107.8.

Â And what's happening is, is that remember, this electron group that is

Â between the nitrogen and the hydrogen is kind of involved in this bonding, and

Â so it's kind of got other things to worry about.

Â But remember when I had that lone pair, it's not involved in a bond,

Â it's not interacting with another atom.

Â And so as a result, this lone pair of electrons has more repelling power.

Â Then do the bonding electrons.

Â And so what's happening is, that we can kind of think

Â of this lone pair is being larger than any of our bonding groups.

Â And as a result it's going to start squishing and

Â compressing these bonding groups together.

Â And so what we get is a bond angle that is a little bit less 109.5.

Â So this is our ideal, okay.

Â Four identical groups around the central atom.

Â 8:05

Here we have two lone pairs, so now if we think about their

Â electron density we can see they are going to be repelling between each lone pair,

Â because again those are close by each other in the molecule, and

Â they're also going to be repelling towards these bonding groups, and

Â as a result we're going to push or compress these two atoms, together.

Â And get an even smaller bonding.

Â So 104.5 or approximately, and so the more nonbonding groups we have,

Â the smaller our angles are going to get relative to the ideal angle.

Â And it's not so

Â important that you memorize the specific values that it's 107.8 or 104.45.

Â What matters is that you realize that the bond angle of

Â the NH3 is going to be smaller than the bond angle is the CH4.

Â And the bond angle in water is going to be smaller than the bond angle of

Â the NH3 because of the affect of those lone pairs of electrons.

Â 9:09

We see something similar when we look at something like formaldehyde how we

Â saw this structure before.

Â This is H2CO, and when we look at those bond angles,

Â here we have two pairs of electrons and still just one electron group.

Â So we have three electron groups, and so we know our ideal angle.

Â 9:30

There's 120 degrees.

Â So our ideal angle's 120 degrees.

Â But because we have two pairs of electrons, or four electrons there,

Â what we see is that causes more repelling power than a single bond.

Â Now, not as much as having a lone pair, but

Â a little bit more than the single bond.

Â As a result, it's going to push these atoms together a little bit,

Â and so we actually get compression of this angle, so we end up

Â with a bond angle of about 115.5 degrees, compared to our ideal which is 120.

Â And we also see that this bond angle has gotten a little bit larger.

Â Now in the previous slide we didn't talk about the bond angle with the lone pairs

Â because there's nothing there to measure the angle with.

Â We have to have atoms there in order to measure those bond angles, and so

Â here we see a smaller bond here we see a larger bond.

Â 10:32

So here our bond angle is going to be a little bit greater than 120, because we

Â have this double bond, and as a result we have more electrons than there and so

Â we're going to have more repelling power and so it's actually going to

Â push these two groups together, so we'll have a smaller bond here.

Â But this bond will be greater than 120 degrees.

Â We're again not worried about the specific amount just

Â knowing whether it's greater than or less than that ideal bond angle.

Â