In this video, we'll use VSEPR theory to practice identifying the major molecular geometries, including bond angles. The VSEPR model is a simple yet powerful tool to predict the shape of molecules. So we'll look at the valence shell electrons.
These are the electrons on the outside of the atom, and they come in general in pairs. Since electrons are negative, and they're on the outside of the atom, each atom has a cloud around it, and that cloud will repair other atoms when they come close together. For molecules we also have that repulsion between atoms.
So here we have a central atom in purple and we're going to add other atoms to it to make a molecule. So we add one and then when we add another remember that we have a negative cloud of electrons around each atom. So when we add another one they push away. If we try to put them close to each other you can see that they move away.
They spread out to be as far away as they can be. but still bonded. And we could keep adding atoms to this, and you can see that the shape changes. Now that we have an idea about Vespa, let's put it to work and go through each of the major molecular shapes.
We use our chart here to help us organize our ideas and build our memory as we learn. It's often assumed that you'll know that if we have two atoms, it's going to be linear. What you should know though, it really doesn't matter if we have single bonds, double bonds, or triple bonds, it's still going to have a linear molecular geometry. And that goes for all of the structures we'll look at today.
The type of bond, single, double, triple, doesn't influence the molecular geometry. Back to our chart, it's organized by steric number and by the number of lone pair electrons. The steric number is the number of things attached to that central atom. That can be atoms, or it can be lone pair electrons, also called unbonded electron pairs.
We use the steric number along with the number of lone pairs to figure out the molecular geometry of the molecule. When we have a steric number of two with no lone pairs, that means there's two atoms attached to the central atom and there are no lone pairs, we have a linear molecular geometry. CO2 is a good example.
We'll start with the Lewis structure and you can see that central carbon atom There are no lone pairs. We can see there's two oxygen atoms, and we're not concerned about the lone pairs on the oxygen atoms. We only worry about the central atom. So the purple, that's our central atom, and we'll add the two oxygens. Each one is double bonded.
Here's one oxygen, and then when I add the other, they spread out to be as far away from each other as possible, and we have a linear geometry, and the bond angle is 180 degrees. Next we move to trigonal planar where we have three atoms and no lone pairs bonded to a central atom. BF3 has a trigonal planar molecular shape.
In the Lewis structure you can see three atoms and no lone pairs bonded to that central boron atom. Here we have the central atom, the boron, and then we'll add fluorine atoms. There's one, two, they spread out, three, to form the trigonal planar molecular geometry. We have the bond angles of 120 degrees and there are no lone pairs. So this is the trigonal planar molecular geometry.
When we have a steric number of three with two atoms and one lone pair, we have a bent molecular geometry. Lone pairs, sometimes called non-bonding electron pairs, are extremely important and influence the shape of molecules. Let's look at SO2.
We have our two atoms bonded to that central sulfur, and then we have a lone pair. So that makes this a bent molecule. But you can't tell this by looking just at the Lewis structure. We need Vespa theory to help us visualize how these atoms are going to spread out, the two atoms and the lone pair.
we'll push the atoms away as well to figure out the molecular geometry. Let's see what that looks like. We have our central sulfur atom and we'll put two oxygen atoms on that.
They're double bonded and next we'll add the lone pair and watch what happens when we do that. The lone pair, it pushes down the two oxygen atoms and now we have a bent molecular geometry. It's important to note that these lone pairs have their own orbital here, and they do occupy space, and that's what leads them to repel and push the oxygen atoms away.
Before we move on, let's lock this into our minds with some practice. The first thing is we'll draw the Lewis structures. I've done that for you to speed things along. At this point you'll want to try to visualize what these structures might look like in three dimensions as those atoms spread out and push away. away from each other.
Then find the steric number and then count all the atoms and lone pairs on the central atom and then remember or look up the molecular geometry. Pause and give these a try. For BF3, we have the central boron atom and then we have three atoms attached to it with no lone pairs on that central boron. You can imagine the fluorine atoms will spread out to be as far away from each other as possible and they might look something like this. Since we have a steric number of three and no lone pairs, that makes this trigonal planar.
For N2, which doesn't even show up in our table, we only have two atoms so it's going to be linear. For, oh no, also called NO2-, we have two atoms attached to the central nitrogen there, and then we have a lone pair. So we have a steric number of three, and you can imagine that those two atoms on the outside, the oxygens, are going to be pushed down.
by that lone pair on the nitrogen. That'll give us a bent molecular geometry. Okay, so we'll look at three more important molecular geometries and then do some practice. When we have four atoms bonded to that central atom and no lone pairs, we have a tetrahedral molecular geometry. CH4 methane is often given as an example of the tetrahedral structure.
We have four hydrogen atoms attached to that central carbon and no lone pairs. So here's our central carbon atom. We add one, two, three, and finally the fourth hydrogen atom, and they've spread out into this tetrahedral geometry. We see the bond angle of one of nine point five.
If we look at this, they're all as far away from each other as they can possibly be while still being bonded to that central carbon. For the trigonal pyramidal, we have a steric number of four. That's three atoms and one lone pair attached to that central atom.
NH3, ammonia, is a good example of our trigonal-pyramidal molecular geometry. We have that central nitrogen there with three atoms attached and then the lone pair. And you can imagine that those hydrogens will spread out, but then that lone pair, that's going to push them down to give it the trigonal-pyramidal molecular geometry.
That would look something like this. Three atoms, which all spread out, and then we add our lone pair. which pushes them down to give us this trigonal-pyramidal molecular geometry.
So we've looked at where we have the three atoms and the one lone pair, giving us the trigonal-pyramidal molecular geometry. Next up, we need to look at what happens when we have two lone pairs and two atoms. With two lone pairs and two atoms, asterisk number 4, we have a bent molecular geometry.
A good example is water. In the Lewis structure, you can see you have two atoms and two lone pairs. That makes this a bent molecular geometry. We add our two hydrogen atoms, they spread out, and then we'll add our two lone pairs, one, two, and this is the bent molecular geometry for water. You might note that the bond angle is one If you look the bond angle up for water, it'd be a little bit different.
So this is a general model. It doesn't give us the specific bond angle for specific molecules, just a general idea of what the bond angle will be for a bent molecular geometry. Time for some more practice.
That's the cornerstone of learning. I've given you these Lewis structures and you need to find the molecular geometries. Try to envision how the atoms and lone pairs are spreading out before you go to the table to check the molecular geometry.
So pause and try to figure out the shape of these molecules. For the ones that you had problems with, go back earlier in the video and look at that particular molecular geometry and figure out where the confusion lies. Two more molecular geometries that we want to be aware of are the trigonal, bipyramidal, and the octahedral.
So when we have five atoms attached to that central atom, one, two, three, there's our trigonal. 4, 5, now it's trigonal bipyramidal. This is the shape that we have. So we have a steric number of 5 with 5 atoms attached in no lone pairs.
If we add another atom to this, we end up with what's called octahedron. This is the octahedral shape and you can see all of the atoms they're spreading out to be as far away from each other as they can possibly get in accordance with the VSEPR theory we've been talking about. Note that we could have our five atoms attached to the central atom there for a trigonal bipyramidal, but we could replace one of the atoms with a lone pair and that would give us a different geometry.
That's a little bit beyond the scope of what we want to do here, but do understand there are other ways to do this. There are other geometries for steric number 5 and 6 when you have lone pairs. In this video we've looked at the major molecular geometries using the VSEPR model. You should have an intuitive sense now of how lone pairs and atoms interact with each other and repel, spreading out to give a molecule its shape.
Using the Lewis structure, looking at the steric number, and then just really thinking about how these atoms spread out, you should be able to figure out these major molecular geometries. This is Dr. B with the Vespa model, and thanks for watching.