So this is what happens for the cases of the acetate and the sulphate in the previous examples, they stay together, but the cation dissociates from the anion. When an ionic compound dissolves in water, the cation and the anion are separated from each other by these layers of water that get between them. This is called dissociation. Not all ionic compounds are highly soluble in water, and we'll talk about how we can measure the solubility and express it with a number in one of the later lectures this week. In addition, some groups of primarily nonmetals like to stick together and those are those groups of polyatomic ions that we learned when we were learning how to name ionic compounds. In this case the compound that dissolved was silver nitrate, and when it broke apart the nitrates, the NO3 one minus groups of polyatomic ions stay together. But the silver separated from the nitrate, didn't it? So the cation always separates from the anion but sometimes these polyatomic cations or polyatomic anions stay together. This is, this type of sticking together of groups was what was happening with the acetate and the sulfate in the previous examples. But in all of these examples we've seen the metal ions completely separated and end up being by themselves, completely dissociated from anything else. Let's try one more of these just to make sure you're getting comfortable with the concept. On some paper, let's write out the dissolution equation for sodium phosphate in water. How many moles of ions are produced for each mole of sodium phosphate that dissolves? So I want you to think about dissolving one mole of sodium phosphate. Write the dissolution equation and answer with how many moles of total number of ions are present in that solution. Here is how we would do that problem. Sodium phosphate solid has the equation Na3PO4. We'll dissolve that in water. In this case, we make three sodium cations because there was a three after the metal. We want to make sure that's balanced. And only one phosphate polyatomic anion, which has a charge of minus three. So each sodium phosphate that dissolves, each unit makes four ions. So one mole of sodium phosphate dissolving would make four moles of ions. So far what we've been doing is counting the number of ions that are forming relative to the solid that we've dissolved. But we also need to know how much is dissolved in a given solution. Is the salt water have just a tiny little bit of salt, or is a lot of salt dissolved, for example? The solute which is present in the lesser amount is dissolved in the solvent. In all of the past examples water was the solvent, but there are other liquids that could be the solvent. You could use gasoline as the solvent for example. That would be a completely different type of solution than these aqueous solutions, but you could still use it as a solvent. The solvent is always the item that is present in the greater amount. The relative amounts of species in a solution are expressed in units of concentration. All concentrations are ratios of the solute to either the solvent or the solution. But there are many subtle differences in the ways the concentration could be expressed and some of these are shown here. We can talk about the relative masses of the solute and the solution, or the relative number per million or per billion for less concentrated solutions. So ppm and ppb, which are used a lot in environmental sampling for example, are for, are very useful for less concentrated solutions. Where there's not as much solute in the solution. We could also look at the amount of solute present relative to the volume of the solute and the volume of the solution. One that gets used a lot in chemistry is molarity, and that's the one I'll focus on the most. But there's others, and I'm not even showing all the different ways you can express concentration here. You might say things like mole fraction, and that is just what is the relative number of moles of solute to solution. There's also molality which gets used a lot in engineering and that is the moles of solute per kilogram of solvent in that case. So for all of these the solute is in the numerator, and there's a fraction where either the solvent or the solution is in the denominator. Let's concentrate on molarity. That's the one that I use the most in chemistry. Molarity is defined as the moles of solute per liter of solution. So I'm using N here as the symbol for the number moles, and V as the symbol for the volume. For example, I could write something like this. In square brackets I could write a chloride ion. Now if you're trying to show molar concentrations you have to use square brackets like that. You can't use parentheses and you can't use little squirrely brackets. Those are not correct for showing molarity or molar concentration, another thing we call molarity. So you have to use these very clearly drawn square brackets if you're trying to show that something's got a specific molar concentration. And here this expression that I've written can be read as, the chloride ion concentration is one molar. Sometimes the solution is more dilute. This happens a lot, for example, in biological testing or in environmental testing. And in that case it might be more convenient to express the concentration in millimolar, micromolar, or nanomolar for example, and in those cases, the concentration is more dilute, so we would need to use these prefixes, which mean ten to the minus three, ten to the minus six, and ten to the minus nine. We can convert those to regular molarity using these conversion factors over here. For example, what if I had a concentration of formaldehyde. This is formaldehyde CH2O. That's in water at a concentration of 1.64 times ten to the minus four molar. So that's its concentration. I could say that is 164 micromolar concentration. And then I don't have to use the scientific notation. I don't have to write quite so much. It gets a little simpler to express that in micromolar. So that's why you might see it expressed that way. Now that we've determined all sorts of ways we can express a concentration of our solution, we need to think about the differences in solubility between different compounds. For example, sodium chloride dissolves really readily in water. We can dissolve a lot of salt in water. But there are other chemicals, such as sand, that don't dissolve very well in water. If we took a container of water and dumped in a teaspoon of sand, it wouldn't look like any of it dissolves at all. Whereas if we took a container of water and we dumped in a teaspoon of salt, as long as we had enough water, we could stir it up and the salt will dissolve. So there is different amounts of solute that can dissolve in solvent. There is limits to the amount of a solute that will dissolve in a given type of solvent and that depends on the chemical composition of the solute, the chemical composition of the solvent, the temperature of the reaction mixture, and even the pressure of the reaction mixture. For the most part we're going to assume that we're at room temperature and normal atmospheric pressure at sea level. So not everything dissolves in water to the same extent. We used the example of putting sand in water and seeing that it doesn't dissolve. At least to the naked eye it doesn't look like it dissolves at all. So we use different words to express how much a solute has dissolved in the solvent. If there's a limit to how much goes in solution, we might want to use different molar concentrations as a guide and these are guides I got from my good friend Dennis Wernst's book. He says if we can make a solution of solute that is greater than 0.1 molar in concentration then we can fairly say, with fair certainty, that that compound is soluble. We would describe it as being soluble. So we can say that sodium chloride is soluble in water. Some species, particularly ionic compounds, dissolve but have a limit to their solubility. And this another, this is an arbitrary range that was set in a particular book, but in that book they said if the molar concentration that will dissolve, the maximum amount that will dissolve before solid chunks start to form on the bottom of the container, is between 0.01 molar, and 0.1 molar, then we say that that compound is moderately soluble in that solvent. And if the amount of solute that were dissolved, before it just continues to fall to the bottom of the container as a solid, is less than 0.01 molar in concentration, then we say that particular compound is insoluble in that particular solvent. Now we might say something is insoluble and there is still a very, very tiny bit that does dissolve. But as far as we can tell with our naked eye most of it is not dissolving and if something is not dissolving well in the solvent we say it's insoluble. This brings up another word that we need to define and that word is saturated. Saturated is when we have added enough solute that we have, that we have reached the limit of how much will dissolve in the solvent. So, if we added any more of the solute, let's say we're putting sodium chloride into water, we add and stir it, it dissolves, we add some more and stir, it dissolves. We keep adding sodium chloride and stirring, and it keeps dissolving, and eventually the next little bit that we add can no longer dissolve. There's already so much sodium chloride in the water that the next little bit just falls to the bottom as a solid. It doesn't look like it dissolves. This happens, this is easy to observe with sugar. I have lots of friends, being in the South, who like to drink sweetened iced tea. And, often they put enough sugar in the iced tea that the sugar does not all completely dissolve and it's, there's some sugar slurry at the bottom, sitting at the bottom. Sometimes people observe this in lemonade as well. So something's saturated when it's exactly at that point where the next little, tiny bit that you add will no longer dissolve. Everything you've added up til now has dissolved, and then the next little bit will not dissolve and we say that that solution is saturated. One way you can make a completely saturated solution is to put in more than will dissolve. Stir it readily and then run it through a filter so that the solid is trapped and the liquid that contains the solute passes through, and that gives you a saturated solution. Thanks for watching this video on introduction to solutions. In the next video we'll talk about electrolytes, so check that out.
