Like water, amino acids are amphoteric, and "amphoteric" is just a fancy word for a molecule that can react either as an acid or as a base depending on the situation. And the reason I mentioned water is because water is the most famous example of an amphoteric molecule. Remember that water could either accept a proton or it could donate a proton based on its surroundings, based on the pKa and the pH of the solution. Well, the same thing is true of amino acids because amino acids have an acidic carboxylic acid group and a basic amine group. So, we're going to need to figure out when it's reacting as an acid and when it's reacting as a base. And for that, we're going to be using exact pKa values. We use exact pKa values given to us on a table to figure out at a certain pH, does that amino acid exist in a charged state or is it in a neutral state, positive or negative. Now, I've gone ahead and provided all these pKas for you in your amino acid breakdown sheet. We're not going to look at it right now, but I just want to remind you, remember all those numbers you saw next to the amino acids? This is what we're using them for. We're going to use them to determine the charges of the amino acids at different pHs, okay?
Before we look at an example, I want to remind you about the theory behind this. Remember that in general chemistry, there was an equation that we learned called the Henderson-Hasselbalch equation. Now, we're not going to have to use that equation in this lesson, but there is just this idea from it that you should be aware of, which is that when the pH of a solution is exactly equal to the pKa of an acid, then exactly half of that functional group is going to be ionized, okay? So, what that means is that if your pH and your pKa are exactly equal to the same amount, that means that that group, that functional group will exist as a 50% neutral molecule and as a 50% charged molecule, maybe positive, maybe negative. It depends on the situation. That's neutral and charged. I'm sorry, it's a little bit off the screen. So what that means is that when they're at equilibrium like that, the neutral and charged forms are even with each other. But once you make the pH a little bit lower than the pKa, what's going to happen? Well, now, you're going to get a more protonated form that predominates because now there's going to be more acid around, so you're going to get more charges, more positives. And if the pH is higher than the pKa, then it's going to be more basic in the solution, so the deprotonated form is going to predominate, okay?
So I want to show you guys an example of what this could look like because it turns out that if you look at the pKa table for phenylalanine, the pKa1, or the pKa of the carboxylic acid group, it's not actually 2; it's close to 2, it's 1.83. You actually need to look that up, okay? You can't just say 2. The exact pKa2, or the pKa of the nitrogen, is not 9; it's 9.13, very similar but just a little bit higher. So, what that means is that at physiological pH of 7.4, like we already discussed, it's going to exist as a zwitterion. It's going to exist where the oxygen gives up its protons to the nitrogen so that the oxygen is negative, and so that the nitrogen is positive. And actually, this is going to be true for any pH all the way from 1.83 all the way to 9.13. So basically, for like 8 numbers of pH, it's going to look like this because in all these situations, the zwitterionic form is the most stable because in all these situations, the pH hasn't become a greater number than the pKa or a lesser number than the pKa.
But what starts to happen once you reduce the pH of the solution below 1.83? Well, when you think of it conceptually, if the pH is 1.83, is that acidic or basic? Super acidic, right? So, it means there's protons everywhere. So this negative charge has been holding on. It's been negative for a while. But what happens if I start slamming it with protons? Like I keep making it more and more acidic, right? What if I drop it to 1.83? What if I drop it to 1.5? What if I drop it to 1? What if I drop it to 0.5? You get what I'm saying? You keep throwing protons at that thing; eventually, it's going to get protonated. Eventually, at some point, if I add enough protons to it, I'm going to get a proton there. That's what 1.83 is. Any pH below 1.83, you add the proton. So that means that at pH 1.83 or below, the pH is less than the pKa1. pKa1 is 1.83, and that means that the protonated form predominates.
Now, the same thing with the nitrogen, so what happens if I start raising the pH of the solution all the way to like 10, all the way to 11, right? So this thing is getting super basic. There's hydroxide molecules everywhere. And at some point, one of these H? So, how do I know when the hydroxide is going to take away that H? You look at the pKa2 value. The pKa2 value is 9.13. So that means that any pH that is higher than 9.13, there's going to be so much base around that that nitrogen won't be able to hold on to that proton and eventually, it's going to lose that proton. Now, it doesn't become negatively charged, that's a different idea. But now, it used to be positive and now it goes into the neutral. And the deprotonated form, in this sense, doesn't mean that it always has a negative. And protonated doesn't always mean that it has a positive because actually, my protonated version didn't have a positive either. It just depends on the molecule you're starting with. But deprotonated means that I lost a proton, and protonated meant that I gained a proton. Does that make sense? Cool. So this is the way that we're going to be thinking about predominant forms. We're going to have to compare pKas using the sheet and then we're going to use the simple formula to figure out what form it should be in. So that's it for this video. Let's move on to the next one.