Solubility and intermolecular forces

Now that we know what a solution is, let's think a little bit about what it takes to get a molecule to be soluble into a solution or into a solvent. So let's say I start off with a salt, and I'll do a little side here, because in chemistry, you'll hear the word salt all the time. Let me write it down: salt. And in our everyday language, salt is table salt. It makes food salty, or sodium chloride. And this indeed is both a salt from the Food Channel point of view and from the chemistry point of view, although the chemistry point of view does not care about what it does to season your food. The chemistry point of view, the reason why it's called a salt is because it's a neutral compound that's made with ions. So we all know that this is made when you take sodium. Sodium wants to lose its one electron in its valence shell. Chloride really wants to take it, so it does. Chloride becomes a negative ion and sodium is a positive ion, and they stick to each other really strongly because this guy's positive now, and this guy's negative after he took away his electron. Imagine your house is too small, so you have to give away your dog to someone who has room for the dog, but now you have to hang out at that person's house all the time because they have the dog you love. I don't know if that analogy was at all appropriate. But I think you get the idea. A salt is just any compound that's neutral. The other common ones, potassium chloride, you could do calcium bromide, or I could do a bunch of them, but these are all salts. And what we want to think about is what happens when you try to essentially dissolve these salts in water. So we know what water is doing, liquid water. So let me draw some liquid water. So if that's the oxygen and then you have two hydrogens that are kind of lumping off of it, I'll draw it like that. I'll draw a couple of them. And then, of course, you have another oxygen here. Maybe the hydrogens are in this orientation because the hydrogen ends are attracted through hydrogen bonds-- we've learned this-- to the oxygen ends because this has a slight negative charge here, a slight positive charge here. These are the hydrogen bonds that we've talked so much about. And maybe you have another oxygen here and it's got its hydrogens there and there. You have some hydrogen bonds there. I could do another oxygen here, and you can kind of see the structure that forms, although what I'm drawing, this is actually more of a-- if you were in a solid state, this would be kind of rigid and they would just vibrate in place. In the liquid state, they're all moving around. They're rubbing up against each other, but they're staying very close. Actually, the liquid state for water is actually the most compact state for water. Now, when you're dealing with stuff like this-- these are moving around, maybe this guy's moving that way, that guy's moving that way-- and you want to dissolve something like sodium chloride. Sodium chloride's actually quite a large molecule. If you look at the Periodic Table up here, oxygen is a Period 2 element. Hydrogen is very small. We know when it gets into a hydrogen bond with oxygen, it's really just a proton sitting out there because all the electrons like to hang out with the oxygen, while, say, sodium and chloride, they're considerably larger. I won't go into the exact molecular sizes, but maybe sodium-- let's do sodium-- which actually, just as a review, which is larger. We know that it becomes smaller as you go to the right of the Periodic Table, so sodium is quite a large atom, while chloride is a good bit smaller, but they're both bigger than oxygen and a lot bigger than hydrogen. So let me draw that. So sodium-- I'll do sodium as a positive. It's pretty big. Maybe it looks like this. Sodium is positive and then you have the chloride. The chloride I'll do in purple. They're still pretty big. The chloride, it'll look like this. And what happens when you put it into water, it disassociates. Even though these guys in a solid state, they're jam-packed to each other. When you put it into water, the positive cations are attracted to the negative partial charges on the oxygen side of the water, and the negative anions are attracted to the positive sides of the hydrogen. But in order to get, for example, this sodium ion into the water, it has to fit in there. So, for example, I drew this as a liquid initially, but if this was a solid and you had this structure, it would be extremely difficult. In fact, it would be next to impossible to squeeze these huge sodium ions in place to make it soluble into, say, solid ice. And as even cold water, these bonds are still going to be pretty strong and they're going to be just kind of barely moving past each other because there's not a lot of kinetic energy. So what you need to do is, the warmer the water you have-- I mean, you can fit it into cold water, because at least cold water has some give, but the warmer the better, because you have some kinetic energy, and that essentially gives space. Or it makes room for this sodium ion that's entering in to kind of bump its way into a configuration that's reasonably stable. And a reasonably stable configuration would look something like this. Sodium would look-- and then you'd have a bunch of-- sodium is positive. It would be attracted to the negative end of the water molecules, so the oxygen end. So it looks like that, the oxygen end, and then the hydrogen ends are going to be pointing in the other direction. The hydrogen ends are going to be on the other side. And, of course, the chlorine atom is going to be very attracted to that other side, so the chlorine atom might be right over here. So the chlorine atom might want to hang out right here. In order to get as much of the sodium chloride into your water sample, you want to heat up the water as much as possible. Because what that does is it allows these bonds to not be taken as seriously and these relatively huge atoms to kind of bump their way in. So, in general, if you think about solubility of a solute in water-- or especially if you think of a solid solute, which is sodium chloride-- into a liquid solvent, then the higher the temperature while you're in the liquid state, the more of the solid you're going to be able to get into the liquid, or you're going to raise solubility. So temperature goes up, solubility goes up. For example, if you were to take some table salt, and you could experiment with this. It doesn't seem too dangerous and not too expensive because salt is reasonably cheap. Keep putting it into a glass, and at some point it'll dissolve. You could shake it a little bit, just to make sure. You could think about what's happening at the molecular level while you shake it and why does that help to shake or stir things? But at some point, you're going to end up with-- if this is your glass of water, the salt will keep going in there, but at some point, you'll have salt crystals at the bottom of your glass. At that point, your water is saturated with salt at the temperature that you're trying to deal with it. Now, right when you start seeing that, if you were to put it in the microwave or if you were to heat it up, you would see that even these guys are able to be absorbed in the water, and that's because the extra kinetic energy from the temperature is making it more likely that these guys are going to be able to bump out of configuration for just long enough for these guys to bump in. And just a little side note, when you take these salts, which are just ionic compounds that are neutral, they're made of ions, but they cancel each other out. When you put them in water, these compounds by themselves aren't normally-- when they're in the solid state, they don't normally conduct electricity. Even though they're charged, they're very closely stuck to each other, so there's not a lot of room for movement of charge. But once you disassociate them in water or dissolve them in water, now, all of a sudden, you have these floating charges in the water, and this does conduct electricity, so it becomes quite a reasonable conductor of electricity. So the general rule of thumb is, if you're dealing with a solid in a liquid solvent, lowering the temperature will decrease the solubility, because it's harder to jam the molecules in there, and increasing the temperature will increase the solubility. But what about a gas? What if you make some soda and you want to dissolve some carbon dioxide into, let's say, water again? So here, the way to think about it when we did it with salts, these are ionic compounds. They had some natural attraction to the different polar ends of the water molecule. But gases, for the most part, do not have strong attractive forces. That's why they're gases, especially at room temperature. They like to be free. A gas, they have a good bit of kinetic energy, but more important, the bonds between them, for example, in ideal gases we talked about it, they just have their London dispersion forces. They have very weak bonds, and that's why at, say, the same temperature and pressure that water would be a liquid, a lot of these gases are gases. They jump away from each other because they don't want to touch each other. Now, when you put this in liquid, and this is at least my intuition, so let's just say this is a bunch of water molecules here. If you were to dissolve-- let's say it's carbon dioxide. You can ignore this stuff up here. If you were to dissolve carbon dioxide in water-- so if you were to dissolve this in water, so those are some carbon dioxide molecules. I'm just drawing the whole molecule as a circle. What do these molecules want to do? It's natural state is a gas and it is a gas at let's say the standard pressure, so it really wants to escape from this water, but it just can't do it that easily because there's water molecules all around it, right? This guy right here, he might want to bump out, but he's surrounded by water molecules. So what would help him bump out? Well, if you raise the average kinetic energy of the system, if you made all of these guys, that these guys were moving faster, and especially if the carbon dioxide molecules themselves had more kinetic energy, then maybe they could break out. And as you have from personal experience with Coke bottles, you could also shake the system, because if you shake the system, it just moves everything around enough that these guys can escape. So when you're dissolving a gas inside of a liquid solvent, when the solute is a gas, it actually has the opposite effect, that rising temperature. So when temperature goes up, solubility goes down because these guys want to escape. They want to be free. They want to be away from other molecules and they want to bounce around in open-- I shouldn't use the word air-- in open space. And so anything that lets the system move around more, they're going to go up. And likewise, if temperature goes down, solubility goes up. The other factor, and it's not as big of a factor when you talk about a solid solute, but when you talk about a liquid solute-- let me just do it again. So those are the carbon dioxide molecules and then you have a bunch of water molecules-- they should all be the same size-- that it's dissolved in. I think you get the idea. Pressure is also a big factor. I already said that these guys, their natural state is to roam free. They want to get out. They want to somehow bounce out of the water. But if you have a really high pressure up here-- just the atmosphere up here has just tons of molecules bouncing really hard down on the surface of our solution-- so if there's just tons of molecules bouncing really hard off the surface, it'll be harder for anything to escape upwards. And that's why, when you have pressure going up, or at least this is the intuition, when pressure goes up, solubility of a gas also goes up. And this is for a gas. So just the interesting thing to remember is that when you think about solubility, solids do the inverse of gas. Temperature is good for solid solubility, right? We said when you put salt or sugar in water, it's good to increase the temperature. You'll be able put more in there. On the other hand, with a gas, it's the opposite. You want colder temperatures to put more gas into the solution, or you want higher pressure to keep it-- at least in the way my mind works-- from escaping out the top. Anyway, hope you found that useful.