Molecular polarity | Chemistry | Khan Academy

Here's a pretty cool video! If you pour oil in water, you find that the oil does not mix with water. You can see that it’s not mixing. Why not? Well, to answer that question, we need to explore something called molecular polarity, and that's what we'll do in this video.

So let's begin, but first let's quickly recap the ideas of covalent bonds and electronegativities. So let's take an example. Consider hydrogen fluoride. Now both hydrogen and fluorine are going to share one electron each because by sharing these electrons, that shell gets completely filled. This is what we call a covalent bond. But here's the question: Are these shared electrons spending equal amounts of time with hydrogen and fluorine? No! That depends upon something called electronegativity.

Electronegativity is the tendency of a bonded atom to pull the shared pair of electrons towards itself. Here’s an electronegativity chart. You can see, for example, hydrogen has an electronegativity of 2.1, whereas fluorine has much higher, which means fluorine is going to pull that shared pair of electrons more towards itself. It has a higher electronegativity. So, these electrons spend more time towards fluorine than hydrogen.

As a result, fluorine ends up getting a partial negative charge, and the lowercase delta over here represents partial. Why do we say partial? Because it's not a complete negative charge. The electrons are not completely transferred, so it's a partial one. Okay? As a result, hydrogen ends up getting a partial positive charge, and this now is what we call a polar bond.

Another way to represent the polar bond is by drawing an arrow. The arrow represents the direction in which the electrons are being pulled, and then we put a plus sign over here representing that this has become partially positive. As a result, this molecule itself is polar, meaning the charges are unevenly distributed. And that's happening why? Because the shared electrons are spending more time over here compared to over here.

But now let's take some more interesting examples. Here’s carbon dioxide; this is the Lewis dot structure. Again, let's look at their electronegativities. Carbon has 2.5, oxygen has 3.5, so it's more electronegative than carbon. So again, the shared pair of electrons will be pulled more towards oxygen, making these bonds polar. But is the molecule polar? Not necessarily. Here’s the reason why.

Now this, remember, is a Lewis dot structure; it does not in general represent a 3D structure. But how do we think about the 3D structure? Well, we have to use our VSEPR. Remember what VSEPR says: you look at the central atom and then think about the number of electron groups around it. So there's one electron group. Remember in VSEPR, the entire double bond is one group. So a double bond or a single bond or a lone pair is considered as a group.

So there is one electron group and two electron groups. If you have two electron groups, they will be as far as possible, and the angle between them would be 180°. So their 3D structure becomes linear. Now let's think about the electronegativity difference. Because oxygen is pulling the shared pair of electrons more towards itself, oxygen ends up getting a partial negative charge, and carbon ends up getting a partial positive charge.

So we can show it using an arrow mark this way: see partial positive, and the electrons are being pulled towards the oxygen. And the same thing we can do over here; the electrons are being pulled towards the oxygen, giving a partial positive charge. But look at what’s happening: you have two polar bonds but exactly in the opposite direction. Think of it this way: the electronegativity difference between carbon and oxygen is one right?

So this polarity has a value of +1, and we can say that this vector represents -1 because it's exactly in the opposite direction. And we have two equal opposite vectors; they cancel out, giving you zero. It's kind of like if you had some toy over here and you pull it in both directions with the exact same force; well, it's not going to accelerate.

Anyways, nothing's going to happen to it—the forces are going to cancel out. The same thing is going to happen over here; the vectors cancel out, and therefore this whole molecule ends up becoming nonpolar. This is a nonpolar molecule. And again, if we compare with hydrogen fluoride, just like before, the bonds over here are polar because the electrons are being pulled more towards the oxygen.

Therefore, they develop a partial negative charge and the carbon develops a partial positive charge. I've drawn two partial positives here because there are two partial negatives. But the big difference is in this molecule. Look at the ends: one end is partially negative, the other end is partially positive. So charges are unevenly distributed, and that's why this molecule itself is polar.

But over here, look at the ends: both ends are equally partially negative, which means the charges are evenly distributed. And that's why, even though the bonds are polar, the molecule itself is nonpolar. Okay, let's now take another example: consider the water molecule. Why don’t you pause the video and think about whether this molecule is polar or not?

All right, the first thing I want to do is look at its 3D structure, because that affects the molecule's polarity, right? So how do I know the 3D structure? We're going to use our VSEPR. This is our central atom. Count the number of groups around that central atom. Electron groups around the central atom: we have one group, two, three, four.

So when you have four groups, they will have a tetrahedral structure. So the electron geometry would be tetrahedral, where you have these two lone pairs forming the two corners of the tetrahedron. But if you were to look at just the molecular geometry, I can get rid of the lone pairs and look—we get a bent-shaped geometry.

Now let's go back and look at our chart. Oxygen has 3.5 electronegativity; hydrogen has 2.1. So oxygen is more electronegative, so it's going to pull the shared pair of electrons more towards itself. So we can show that now over here—it’s going to pull the shared pair of electrons more towards itself.

Now, what happens to the overall polarity of the molecules? Again, we have two vectors, but they're not exactly opposite to each other. Right? It's kind of like again if you were to look at this—it's kind of like one vector is this way. You pull a toy, one person pulls it this way, another person pulls it this way. Now they don’t cancel out; now the toy will accelerate in this direction.

You can kind of see that, right? Similarly, over here, the dipole vectors do not cancel out. And as a result, we say that there is a net polarity, which means charges are unevenly distributed. You can see that one end is partially negative, the other end of the molecule is partially positive, so the molecule itself is polar.

Why did that happen? That happened because of the bent shape. The bent shape ensured that one end is partially negative, and the other end is partially positive. Now, the polarity can also be seen more visually using something called the ESP maps, which stand for electrostatic potential maps.

And just like before, if you have a strong concentration of negative charge, we use red; strong concentration of positive charge, we use blue, and you have colors in between—green represents neutral. So going back to carbon dioxide, look; the ends are red because oxygen has a very high electron density. It's also pulling the shared pair of electrons towards itself, so there's a strong concentration of negative charge.

And because electrons are being pulled away from the carbon, it develops a partial positive charge. So we have shown it with bluish color. The same thing over here: partial negative charges, partial positive charges. And again, look at the ends of the molecule. Over here, both ends are reddish, which shows symmetry across the entire molecule, meaning the dipoles cancel out, and the molecule is nonpolar.

But look at the ends over here: one end is reddish; the other end is bluish, representing that there is an uneven distribution of charges making the molecule polar. Okay, let's take a couple more examples. Your turn! Why don’t you try to draw their 3D structures, think about the electronegativity differences, and then eventually conclude whether the molecules are polar or nonpolar?

Pause and try. All right, let's see! Again, let's start with carbon tetrachloride. Because there are four things around the central atom, which is the carbon, VSEPR says this is going to be a tetrahedral structure. Chlorine molecules will be at the corners of the tetrahedron, and so this is what it would look like.

So now we can look at the electronegativity difference: carbon is 2.5, chlorine is 3. Chlorine is slightly more than carbon, so it's going to pull the shared pair of electrons towards itself. So again, each bond is polar, right? Carbon is getting a partial positive charge; electrons are being pulled towards chlorine atoms over here.

So the question now is: Is this whole molecule polar? Well, because this is a tetrahedral structure, again, we have symmetry over here. I mean, it’s kind of like you have a toy which is being pulled in these different directions. Okay? If you pull them in these tetrahedral directions, in these four directions this way—the symmetry because of the symmetry, the whole thing cancels out.

Remember in tetrahedron, the angles between these vectors would be exactly the same, and therefore all of that eventually works out and just cancels out because of the tetrahedral structure over here. So again, here's another example where the bonds are polar, but because of the symmetry, the molecule itself is nonpolar.

Again, what would the ESP map of this look like? We'd probably expect a lot of electron densities over here, but because of the tetrahedral structure, it all balances out. So you get something like this—it’s kind of you do have a lot of electron density around the chlorine atoms over here, but again, because you have the tetrahedral structure, it nicely balances out.

So the charges are evenly distributed across the ends, making the whole molecule nonpolar. What about chloromethane? Well again, it will have a tetrahedral structure just like over here; just that three of those will now be hydrogen, one of them will be chlorine.

Now let's look at... now we have to look at carbon and hydrogen as well. So carbon is 2.5; hydrogen is 2.1. Now if the electronegativity difference is, I mean, there is a difference, so carbon will pull the electrons closer to itself compared to hydrogen. But if the electronegativity difference is less than 0.5, we usually say it's very tiny, and we neglect it. So it’s 0.4 over here, right?

So we’ll neglect it. The bonds are polar, but because the electronegativity difference is pretty small, we can ignore it, and we can neglect it, and we can consider the bonds to be nonpolar. But the electronegativity difference between carbon and chlorine is exactly 0.5, so it’s in the border case, and so we will consider this bond to be polar.

So this bond, like over here, is polar; the electrons are being pulled more towards chlorine. But what about the whole molecule? Is it polar or not? Well, let’s see. Now we only have one vector over here. These don't contribute to polarity at all—not much; there's only one over here. And therefore the whole thing now ends up becoming polar.

Even though it's a tetrahedral structure, unlike what we see over here, we don't have vectors to cancel out with this one. There's only one vector, so the molecule is polar. And so now, if you were to look at the ESP map, because there’s a lot of electron density over here, we would expect this to have a lot of negativity over here.

And because electrons are being pulled away from it, you'll expect, you know, this to be a lot of positive—so red on the top, blue on the bottom, red is negative. So you would get something like this, and you can now clearly see the charges are unevenly distributed across its ends, and so the molecule is polar.

Now, what if the molecules are even bigger? Then rather than thinking about the polarity of the whole molecule, we like to think in terms of regions of polarity. For example, consider this molecule over here. You’ll see that this molecule is now slightly bigger than the ones that we considered so far.

So let's look at the different regions. Let's start with this one. In fact, let me just draw the 3D structure. You can see the 3D structure around carbon is tetrahedral because there are four things around it. Around nitrogen, well, there are four things around it, but because one of them is a lone pair, the electron geometry would be tetrahedral.

But if you consider the molecular geometry, you get rid of the lone pair, and then you get trigonal pyramidal. It doesn't look like that, but you have to think of it as trigonal pyramidal. Okay? Now let’s try to guess what the ESP map looks like for this.

Okay, well, around nitrogen, you can see there’s a high electron density, and also nitrogen is more electronegative compared to hydrogen and even compared to carbon. So it's going to pull the electrons close towards itself. So you can expect a lot of negativity around this one.

So you expect negative as red in our case. So this will be pretty red. And as a result, hydrogen would be pretty blue because electrons are, you know, it's partial positive charge, so this would be blue. So you expect a lot of polarity over here.

What about carbon and hydrogen? Well, we've already seen carbon and hydrogen have very similar electronegativities, and therefore there’s not going to be much polarity over here. And therefore this would expect to be pretty neutral—green. So there'll be charge difference here and not much over here.

So this is what the ESP map looks like. Kind of makes sense, right? You can clearly see this region is polar and this region is nonpolar. Now because this molecule is still pretty small, we could say the whole molecule itself is still polar. But what if you had a really big molecule, like say stearic acid?

We’re not going to try and analyze it, okay? But just let’s look at the ESP map. Now it no longer makes sense to talk about the polarity of the whole molecule; instead, we say that this region of the molecule is nonpolar, and this region of the molecule is polar.

And we are now ready to answer the original question of why oil doesn’t mix with water. Since water molecules are polar, the partial negative end of one molecule will attract the partial positive end of the other molecule, making them stick—kind of like magnets. We call this force of attraction intermolecular forces because it is in between two molecules.

By the way, do not confuse this with intramolecular forces, which is the forces within a molecule—it’s what keeps the molecule itself together. For example, over here, the covalent bonds are causing the intramolecular forces. So "intra" means forces within the molecule that keeps the molecule together; “inter” means forces between two or more molecules.

Now, nonpolar molecules also have intermolecular forces, but they’re much weaker, and guess what? Oil molecules are nonpolar. Because oil molecules are nonpolar, water molecules will just end up sticking to other water molecules and not to oil molecules. And that's why oil molecules cannot come in between.

As a result, oil does not mix with water. In general, if you pour any nonpolar liquid into water, it will not mix with water. Similarly, nonpolar stuff will not even dissolve in water. If you want to mix with water, or if you want to dissolve something in water, their molecules need to be polar.