Video Transcript
In this video, we’ll learn how to determine the shape of molecules by using VSEPR, or VESPR theory. We’ll learn how to determine a molecule’s shape and the bond angles in the molecule by looking at its structure.
Before we figure out how to determine the shape of molecules, let’s take a moment to review how we create Lewis structures. Let’s do that by creating Lewis structures for these three molecules: CH4 for methane; NH3 or ammonia; and H2O, water.
The first step for creating a Lewis structure is to determine the number of valence electrons. Carbon has four valence electrons, and hydrogen has one. Since there’s one carbon and four hydrogens in methane, that means we have a total of eight valence electrons to work with when we’re creating our Lewis structure.
Next, we’ll want to place our atoms and connect them with single bonds. We’ll place the atom that can form the greatest number of bonds in the center of our structure. In the case of methane, that would be carbon, since hydrogen can only form one bond. Single bonds each contain two electrons. And since we’ve placed four bonds to connect all the atoms in our structure, that means methane’s eight valence electrons are already accounted for in our structure. So this is the correct Lewis structure for methane.
Now let’s make the Lewis structure for ammonia. Again, we’ll start by determining the number of valence electrons. Nitrogen has five valence electrons, and hydrogen has one. So ammonia has a total of eight valence electrons. Next, we’ll place our atoms and connect them with single bonds. To connect all of our atoms, we’ve placed three single bonds, which has used up six of the eight valence electrons. To account for those last two valence electrons, we’ll put a lone pair on the nitrogen atom, which gives us the correct Lewis structure for ammonia.
That leaves us with our final molecule, water. Oxygen has six valence electrons, which again gives us a total of eight valence electrons for a molecule of water. Next, we’ll place our atoms and connect them with single bonds, which uses up four of our valence electrons, leaving us with four still to place in our structure, which we’ll account for by putting two lone pairs on the central oxygen atom.
Now, do all these structures tell us about the shape of these molecules? Maybe the shape of these molecules looks just like the structures we’ve drawn. Methane is shaped like a cross, and water is oriented in a line. Or perhaps these three molecules have the exact same shape. After all, these three molecules each have four electron domains or groups that are attached to the central atom. These electron domains could either be a bonding pair in between atoms or a lone pair attached to the central atom.
Well, we could deduce the shape of these molecules by measuring the bond angles in between atoms. If we were to do that for these three molecules, we would discover that the bond angles in methane are about 109.5 degrees, the bond angles in ammonia are about 107 degrees, and the bond angle in water is about 104.5 degrees. The shape that these molecules must have based on their bond angles is roughly a pyramid with triangular faces, which is called a tetrahedron.
But why do these molecules have these shapes? How can we generalize this information? And why is there this discrepancy between the bond angles of these three molecules? We’ll soon see how the shape of all molecules is related to the repulsions between electrons. As we know electron domains, that is, those bonded groups and lone pairs around the central atom are negatively charged. And things that are negatively charged tend to repel each other. So in order to minimize the repulsions between these electron domains and a molecule, the distance between them will be maximized. In other words, the electron domains will tend to stay as far apart as possible.
So, for example, let’s take a look at a molecule that has two electron domains like carbon dioxide. In order to maximize the distance between these two groups, they will end up on opposite sides of the molecule, or 180 degrees apart.
But what about if we have a molecule with three electron domains like boron trifluoride? Since there’s three groups bonded to the central atom, they can’t be on opposite ends of the molecule. Instead, these groups will be evenly spaced around the central atom 120 degrees apart. And we’ve already seen how molecules with four electron domains are roughly shaped like a tetrahedron. But we mentioned earlier that the bond angles for these three molecules methane, ammonia, and water are very slightly different. So what causes this?
These three molecules all have four electron domains around the central atom. But they each have a different number of lone pairs. Methane has none, ammonia has one, and water has two. So perhaps the number of lone pairs also plays a role in the shape of molecules. Let’s take a look at this cartoon to try to understand the difference that lone pairs and bonding pairs have on the shape of a molecule. A bonding pair is spread out in space between the two nuclei that are participating in the bond. But a lone pair isn’t spread out between two nuclei, so it takes up more space around the central atom.
As a result, lone pairs repel more strongly than the bonding pairs do. This explains the trend in bond angles that we saw earlier between methane, ammonia, and water. The lone pairs in ammonia and water are going to repel more strongly than the other electron domains in these molecules. So we can think of the lone pairs in these molecules as pushing down on the other bonded groups in the molecule, which decreases the bond angles. All of these ideas that we’ve discussed make up valence shell electron pair repulsion theory, which you’ll usually just see referred to by its acronym: VSEPR or “VESPR.”
VESPR theory allows us to determine the shape of any molecule or polyatomic ion where the central atom is not a metal. VESPR theory is based on the idea that electron pairs in the valence shell repel each other. And so molecules tend to have shapes that maximize the distance between these electron pairs. To use VESPR theory to determine the shape that a molecule has, we just need to know the number of lone pairs and the number of bonded groups around the central atom.
To use VESPR theory to determine the shapes that different molecules have, we’ll be using the A X E method. The A here represents the central atom. The X represents the bonded groups. And the E represents the lone pairs. To use this method, we’ll be counting the number of bonded groups and lone pairs that are around the central atom, which is noted by the letters m and n, respectively. If we add m and n together, this will give us the steric number, which tells us the total number of groups that are around the central atom.
To use this method, all we’ll need to do is look at a structure and count the number of bonded groups and lone pairs, and then we can determine the steric number. Then there will be a shape and bond angles that are associated with these numbers.
Let’s start off with carbon dioxide, which is a molecule that we’ve already looked at. Carbon dioxide has two groups bonded to the central atom and no lone pairs. This means it has a steric number, two. Molecules with a steric number of two with no lone pairs have a shape that’s linear with bond angles of 180 degrees.
Our next molecule is BF3, boron trifluoride. Boron trifluoride has three groups bonded to the central atom and no lone pairs. That means the steric number for this molecule is three. The shape of molecules with a steric number of three and no lone pairs is called trigonal planar. The bond angles for these molecules is 120 degrees.
Our next molecule is SO2, sulfur dioxide. This molecule has two groups bonded to the central atom and one lone pair. This makes its steric number three. Molecules with a steric number of three and one lone pair have a shape that’s called nonlinear or bent. The bond angles in these molecules will be slightly less than 120 degrees due to the presence of the lone pair.
Our next molecule is CH4 or methane. Methane has four groups bonded to the central atom, and none of them are lone pairs. So the steric number here is four. The shape of molecules like methane that have a steric number of four with no lone pairs is called tetrahedral. The bond angles here will be 109.5 degrees.
Next up is ammonia, NH3. This molecule has three groups bonded to the central atom and one lone pair, so the steric number here is four. Molecules like ammonia that have a steric number of four with one lone pair have a shape that’s called trigonal pyramidal. The bond angles here are 107 degrees.
Next is water or H2O. Water has two groups bonded to the central atom and two lone pairs. So once again we have a steric number of four. The shape of molecules that have a steric number of four with two lone pairs is called bent or nonlinear. And the bond angles for these molecules is 104.5 degrees.
Next, let’s look at this molecule, PCl5 or phosphorus pentachloride. This molecule has five groups bonded to the central atom and no lone pairs, which means that this molecule has a steric number of five. The shape of molecules that are like phosphorus pentachloride with a steric number of five and no lone pairs have a shape that’s called trigonal bipyramidal. We can visualize this shape as two triangular pyramids that are stacked on top of each other.
Bond angles for these molecules are a little bit more complicated because there are two positions that atoms can be in. The top and bottom atoms are in axial positions. These atoms run in a straight line through the molecule like the axis of the Earth. Atoms in the axial positions are 180 degrees from each other and 90 degrees from the other atoms. The remaining three atoms are in the equatorial positions. Atoms in the equatorial positions are evenly spaced around the middle of the molecule like the equator of the Earth. Atoms in the equatorial positions are all 120 degrees apart from each other.
The final molecule we’ll look at is SF6, sulfur hexafluoride. Sulfur hexafluoride has six groups bonded to the central atom and no lone pairs. This makes its steric number six. Molecules like sulfur hexafluoride that have a steric number of six with no lone pairs have a shape that’s called octahedral. We can visualize this shape as two pyramids that have square bases stacked on top of each other.
Just like with the last molecule we looked at, this molecule will also have axial and equatorial positions for the atoms. Again, the axial atoms are going to be 90 degrees from the other atoms in the equatorial positions. But now, since there’s four atoms in the equatorial positions, these atoms will also be 90 degrees apart. This is the last shape that we’ll look at in this video. There are other shapes that molecules can have, but they are less common than the ones that we’ve looked at here.
It’s also worth noting that all of these angles we’ve been discussing aren’t exact. These angles reflect the maximum distance that these electron domains could theoretically be apart from each other. In reality, these bond angles will vary slightly depending on the identity of the atoms that are bonded to the central atom.
As we mentioned earlier, VESPR theory can also be used to determine the shape of polyatomic ions. So let’s determine the shape of the phosphate ion using VESPR theory. The phosphate ion has four oxygens bonded to the central phosphorus atom and no lone pairs. This means that the steric number here is four. Since the phosphate ion has a steric number of four with no lone pairs, we can describe its shape as tetrahedral, And we know that the bond angles would be 109.5 degrees.
We can also use VESPR theory to help us determine the bond angles in larger molecules. So let’s try that with these two molecules, ethene on the left and acetic or ethanoic acid on the right. First, let’s figure out the bond angles around this carbon on the left in the ethene molecule. It has three groups bonded to it and no lone pairs. That means the steric number here is three. Since the steric number here is three and there are no lone pairs, we know that the bond angles around this carbon are all about 120 degrees.
Next let’s take a look at acetic acid. Let’s figure out the bond angles around this oxygen at the end of the molecule. This oxygen has two groups bonded to it and two lone pairs. That makes the steric number four, which means the bond angles here are 104.5 degrees, even though the structure wasn’t drawn that way.
Now let’s wrap up this video with the key points. Valence shell electron pair repulsion or VESPR theory describes the shapes of molecules and polyatomic ions. Electron domains or the groups around the central atom repel each other, so the distance between them is maximized in a molecule. The bond angles in molecules that have lone pairs are smaller because lone pairs have greater repulsions than bonding groups. We can use the A X E method to determine the shape of molecules using VESPR theory. All we need to do is count the number of bonded groups and lone pairs around the central atom denoted by the letters m and n. These two numbers add together to give us the steric number.
