In this video, we’ll learn how the electronegativities of atoms within a molecule can create polar bonds. We’ll learn how to identify polar bonds within a molecule and determine how this affects the polarity of the molecule as a whole. Before we dive into the topic of polar bonding, let’s take a minute to review electronegativity.
Electronegativity describes the tendency of an atom to attract the shared electrons of a bond. Electronegativity generally increases going from left to right across the periodic table as well as up a group. The Pauling electronegativity scale can be used to quantify the electronegativities of atoms; it’s a unitless scale that goes from zero to four, with four being the greatest. Fluorine has the highest electronegativity value at 3.98. Elements like sodium have much smaller electronegativity values. And we can see elements like carbon with its electronegativity value of 2.55 is less electronegative than fluorine but much more electronegative than sodium.
So, what effect does the electronegativity of an atom have on a bond in a molecule? Well, we might picture a bond between two atoms like this cartoon. Here the electrons that make up the bond are being shared between the two atoms. What we have here shows that the electrons are being equally shared between the two atoms. But depending on the electronegativities of these atoms, there might not be an equal sharing like this. For example, if this pink atom on the left was more electronegative than this orange atom on the right, then it would attract the electrons more strongly than the orange atom does, which means the electrons would be pulled towards the pink atom. And the opposite would be true if the orange atom on the right was the one that was more electronegative. In that case, the electrons would be pulled towards the orange atom.
We can think about bonding like a spectrum. On one end of the spectrum, we would have the completely equal sharing of electrons between the atoms in the bond. For example, this is what we would see in a molecule of chlorine. Since both of these atoms are the same, there’s no difference in electronegativity. So, one atom isn’t going to attract the electrons more than the other. And on the other end of the spectrum, we have ionic bonding, like we see on sodium chloride. Here the difference in electronegativity is so large that the electrons are no longer shared between the two atoms. Instead, the more electronegative atom is attracting those electrons so strongly, then that’s a negatively charged ion and the less electronegative atom has become a positively charged ion. For everything else in the middle of the spectrum, there will be an unequal sharing of electrons, like we might see in a molecule like hydrogen chloride.
Here, we have a difference in electronegativity between the two atoms, but it’s not as large as we see in ionic bonding. A bond like this in a molecule where the electrons aren’t evenly shared is called a polar bond, and bonds where the electrons are evenly shared is called nonpolar unless we know bonds on the other end of the spectrum are ionic. We can roughly use electronegativity differences between the atoms to determine whether a bond will be polar, nonpolar, or ionic.
If the difference in electronegativity between atoms is about 0.4 or less, then we can consider that bond nonpolar. If the difference is greater than 1.8, then we would consider that bond ionic. And of course, anything in the middle would be a polar bond. It’s worth noting that because there’s this range in the electronegativity difference for what makes a bond nonpolar, it’s not just molecules like chlorine with the two atoms are the same that are nonpolar. For example, the bond between carbon and hydrogen in a molecule of methane would also be nonpolar. This is because carbon has an electronegativity value of 2.55 and hydrogen has an electronegativity value of 2.2. So, the difference in electronegativity is 0.35, which is less than 0.4.
So, what effect does this unequal sharing of electrons caused by a difference in electronegativity have on the bond and the molecule? Well, the electrons in a polar bond are closer to one side of the molecule than the other. To see the effect that this has, let’s compare the electron clouds of a polar and a nonpolar molecule. Looking at the nonpolar molecule, we can see that the electron cloud is pretty much uniform around the molecule. There is no areas where there is more or less density of negative charge. But in the polar molecule, the electron cloud has been distorted. Since the electrons are closer to one side of the molecule than the other, the side of the molecule where the electrons are is going to be more negative than the other side. In other words, there is a negative and a positive side of the molecule.
Now, we don’t have a fully negative and fully positive charge like we see in ionic bonding since we don’t actually have ions in a polar bond. Instead, we have partial charges in polar molecules which we indicate with the lowercase Greek letter 𝛿. Now, whenever there is a separation of charges over a distance like we see in this polar molecule where we have these partial negative and partial positive charges, it creates what’s called the dipole moment. So, since there is a separation of charges in this molecule, there will be a dipole moment, which we can indicate with an arrow like this. One end of this arrow has a plus sign on the end, which will always be on the positive side of the molecule, and then the arrow will point towards the negative side of the molecule.
So, so far, we’ve been looking at molecules that have two atoms. So, what effect does polar bonding have on molecules that are larger than that? Well, in our molecule with just two atoms, the polar bond makes the entire molecule polar. But when we have more atoms in a molecule, simply identifying whether or not there’s polar bonds in the molecule isn’t not enough to tell us if the entire molecule is polar. This is because dipole moments can actually cancel each other out. We can think of this like how negative one and positive one will cancel each other on a number line. This is because dipoles are what’s called vectors; they are lines that have a direction to them.
But this terminology isn’t important for understanding the concept here. For example, this molecule might be carbon dioxide. Oxygen is more electronegative than carbon, so there would be polar bonds in this molecule. But because both of the atoms on opposite ends of the molecule are the same, they’re going to be pulling on the central carbon equally. In other words, those dipole moments are going to cancel. In other words, even though this molecule has polar bonds because there is a difference in electronegativity between the atoms in the molecule, the overall molecule is nonpolar because the dipole moments will cancel each other out.
Let’s take a look at another molecule, boron trifluoride or BF3. Fluorine is more electronegative than boron. So, the fluorines would all have a partial negative charge, while the boron would be partially positively charged. This means all three of these bonds will be polar with a dipole pointing towards the fluorine. But is this molecule polar? Well, according to VSEPR or VSEPR theory, all of the bond angles in this molecule are identical; they’re 120 degrees. This means that we’re going to see something just like we saw with carbon dioxide. Since all three of these bonds are between boron and fluorine, they’re identical in strength and intensity.
In order for a molecule to be polar, there needs to be a region of greater electron density in the molecule. But these three bonds are evenly spaced and they’re pulling on the electrons with equal strength. So, the electron density is going to be the same in all directions. That means that just like we saw with the molecule of carbon dioxide, the bond dipoles here will cancel, which means the molecule is nonpolar although this is a little bit harder to see than it was with carbon dioxide. So, another way we might think about this is that if we rotated the molecule, it would look the same. So, the bond dipoles must cancel out.
Now, let’s take a look at this molecule, carbon tetrachloride. Chlorine is more electronegative than carbon. So, the chlorines will all have a partial negative charge, while the carbon has a partial positive charge. This, of course, means that all of these bonds will be polar. But is this molecule polar? Again, here all of these angles are identical. These bonds are all 109.5 degrees apart from each other. So, just like before, all of these bond dipoles are exactly the same and they’re going to cancel each other out. So, this molecule is nonpolar.
So, I’m sure that you’re seeing a pattern here. All symmetrical molecules, even if they have polar bonds, are going to be nonpolar. For comparison, let’s take a look at a couple more examples of molecules and see if they’re polar or not, starting with this one, CH3Cl. As we mentioned earlier, there’s generally not a large enough difference in electronegativity between carbon and hydrogen to consider that those bonds are polar. But carbon is less electronegative than chlorine, which means there would be a partial negative charge on the chlorine and a partial positive charge on the carbon. This means that there would be a dipole going from the carbon to the chlorine. We can clearly see that this molecule isn’t symmetrical like the ones we were looking at earlier. It has a dipole that isn’t canceled out from another dipole, so this molecule is polar.
Let’s take a look at this molecule next, ammonia or NH3. Ammonia is more electronegative than hydrogen. So, there would be a partial negative charge on the ammonia and partial positive charges on the hydrogen. This means that there would be dipoles going from the hydrogens to the nitrogen. Now, this molecule might look symmetrical, but we can’t forget about this lone pair. We can imagine rotating this molecule so that the lone pair is in a different position. And we can see that this molecule doesn’t look the same as before. This means that the dipoles won’t be able to cancel each other out. There will be a net dipole towards the nitrogen still, meaning that the molecule is polar.
Let’s take a look at one more example, water or H2O. Oxygen is more electronegative than hydrogen. So, the hydrogens would both have a partial positive charge, while the oxygen has a partial negative charge, meaning that each bond will be polar with a dipole pointing towards the oxygen. But just like we saw in ammonia, these dipoles aren’t going to cancel each other out because of the presence of the lone pairs in water. Rather there’s going to be a net attraction of the electrons towards the oxygen. This molecule will be polar.
So now we’ve learned how to identify a polar bond and what effect this has on molecules. So, let’s look at some problems before we conclude this video.
Which of the following bonds is considered polar? A Br-Br single bond, a CH single bond, a CC single bond, an OH single bond, or an OO double bond.
Bonds are polar due to differences in electronegativity between atoms. This is because electronegativity describes a tendency of an atom to attract shared electrons in a bond towards itself. This means that if one atom is more electronegative than the other, the electrons will be attracted towards the more electronegative atom. When this happens, the bond is considered to be polar. So, from this explanation, we know that there needs to be a difference in electronegativity between the atoms, which means that any bond that has between two atoms that are the same isn’t going to be polar because there’s no electronegativity difference between the atoms.
So, we can safely rule out the Br-Br single bond, the carbon-carbon single bond, and the oxygen-oxygen double bond. This leaves us with the carbon-hydrogen single bond and the oxygen-hydrogen single bond. Now, it might seem that perhaps both of these bonds should be polar since there will be some difference in electronegativity between carbon and hydrogen and oxygen and hydrogen.
If we look at the Pauling scale for electronegativities, a common scale used to compare the electronegativities of atoms ranking them from zero to four with four being the highest, we can see that while there’s a pretty large difference in electronegativity between the oxygen and the hydrogen, there’s not that large of a difference between the carbon and hydrogen. Because the difference in electronegativities here are so small, the carbon-hydrogen bond is generally not considered to be polar, but the oxygen-hydrogen bond most certainly is. So, of our list, the bond that is considered to be polar is the oxygen-hydrogen bond.
Now, let’s sum up this video with the key points. Polar bonds are caused by differences in electronegativity between atoms. This will cause the more electronegative atom to have a partial negative charge, which we can indicate using the lowercase Greek letter 𝛿. And the less electronegative atom will have a partial positive charge. This separation of charge across the bond causes there to be a dipole, which we can indicate using this arrow that points towards the more electronegative atom. If we want to determine whether an entire molecule is polar or nonpolar, it’s not enough to determine if it just has polar bonds. Molecules that are symmetrical will be nonpolar as well because the bond dipoles will cancel each other.