Hydrogen bonding | Intermolecular forces and properties | AP Chemistry | Khan Academy

Let's talk about hydrogen bonds. Depicted here, I have three different types of molecules. On the left, I have ammonia; each ammonia molecule has one nitrogen bonded to three hydrogens. In the middle, I have something you're probably very familiar with. In fact, you're made up of it, which is water. Each oxygen is bonded to two hydrogens. Then here on the right, I have hydrogen fluoride; each fluorine is bonded to one hydrogen.

Now, why are these types of molecules interesting, and what does that have to do with hydrogen bonds? The simple answer is, in each of these cases, you have hydrogen bonded to a much more electronegative atom. Even though these are covalent bonds, they're going to be polar covalent bonds. You are going to have a bond dipole moment that goes from the hydrogen to the more electronegative atom. The more electronegative atom is going to hog the electrons. The electrons are going to spend more time around that, so that end of the molecule is going to have a partial negative charge. The ends with the hydrogen, those are going to have partial positive charges.

Another way to think about it is if you added these dipole moments, you would have a net dipole for the entire molecule. That would look something like that. So, we are dealing with polar molecules, and the polarity comes from both the asymmetry and the very electronegative atom bonded to hydrogen—oxygen, a very electronegative atom bonded to hydrogen. So this end of the molecule is partially negative. This end of the molecule, or these ends of the molecule, are partially positive. For hydrogen fluoride, this end is partially positive, and this end is partially negative.

So, what do you think could happen when these molecules interact with each other? The nitrogen end right over here of this ammonia could be attracted to one of these hydrogens that has a partially positive charge right over there. Or this hydrogen, the partial positive charge, might be attracted to that nitrogen that has a partial negative charge. This attraction between the partially positive hydrogen end and the partially negative end of another molecule—those are hydrogen bonds. They are an intermolecular force that will be additive to the total intermolecular force from, say, things like London dispersion forces. This makes you have a higher boiling point than you would have if you just thought about London dispersion forces.

To make that clear, you can look at this chart. You can see all of these molecules are formed between period two elements and hydrogen. In fact, all of these molecules have similar molar masses: methane, ammonia, hydrogen fluoride, and water. If we were just thinking about London dispersion forces, London dispersion forces are proportional to the polarizability of a molecule, which is proportional to the electron cloud size, which is proportional to the molar mass. Generally speaking, as you go from molecules formed with period two elements to period three elements to period four elements to period five elements, you do see that as the molar mass of those molecules increase, there is that general upward trend of the boiling point. That's due to the London dispersion forces.

But for any given period, you do see the separation. In particular, you see a lot of separation for the molecules formed with oxygen, fluorine, and nitrogen. These molecules, despite having similar molar masses, have very different boiling points. So there must be some other type of intermolecular forces at play, above and beyond London dispersion forces. The simple answer is yes. What you have at play are the hydrogen bonds.

Now, some of you might be wondering, "Well, look at these molecules formed with period three elements and hydrogen or period four elements in hydrogen." They also don't have the same boiling point, even though you would expect similar London dispersion forces because they have similar molar masses. The separation that you see here in boiling points, this too would be due to other things other than London dispersion forces. In particular, dipole-dipole forces would be at play.

But what you can see is the spread is much higher for these molecules formed with nitrogen and hydrogen, fluorine and hydrogen, and oxygen and hydrogen. That's because hydrogen bonds can be viewed as the strongest form of dipole-dipole forces. Hydrogen bonds are a special case of dipole-dipole forces. When we're talking about hydrogen bonds, we're usually talking about a specific bond dipole: the bond between hydrogen and a more electronegative atom like nitrogen, oxygen, and fluorine.

So we're specifically talking about that part of the molecule, that hydrogen part that has a partially positive charge being attracted to the partially negative end of another molecule. It's really about a bond dipole with hydrogen bonds versus a total molecular dipole when we talk about dipole-dipole interactions in general. You can imagine it doesn't even just have to be hydrogen bonds between a like molecule; you could have hydrogen bonds between an ammonia molecule and a water molecule or between a water molecule and a hydrogen fluoride molecule.

And I mentioned that these are really important in biology. This right over here is a close-up of DNA. You can see that the base pairs in DNA, you can imagine the rungs of the ladder, those are formed by hydrogen bonds between base pairs. Those hydrogen bonds are strong enough to keep that double helix together, but then they're not so strong that they can't be pulled apart when it's time to replicate or transcribe the DNA.

Hydrogen bonds are also a big deal in proteins. You learn in biology class that proteins are made up of chains of amino acids, and the function is heavily influenced by the shape of that protein. That shape is influenced by hydrogen bonds that might form between the amino acids that make up the protein. So hydrogen bonds are everywhere. There are many hydrogen bonds in your body right now, mainly not just because of the DNA, mainly because you're mostly water. So life as we know it would not exist without hydrogen bonds.