An important way in which ionic compounds are formed is through so-called precipitation reactions. And that's what I'm going to talk about right now. Then you should first look at this picture on the right. It first contained a solution of colorless silver nitrate. But then they dripped in a solution of sodium chloride, also colorless. Then a white precipitate formed in this way. Why is it formed, and what does it consist of? Bring your notes now and we'll go through that really carefully! We write this way to begin with, that silver ions are reagents for chloride ions. This means that you can use silver ions to find out whether a solution contains chloride ions. This is how it works. Then it's just like in this picture that I showed you first, that we have a test tube with some silver nitrate in it. Notice that it says "aq" here. This means that the silver nitrate is dissolved in water, and that in turn means that we have free silver ions and nitrate ions swimming around here in the solution. Now, to this, we'll add some sodium chloride solution. Notice that this one is also in aqueous solution, NaCl(aq) in other words. That means we also have some sodium ions in aqueous solution and chloride ions in aqueous solution, and they're swimming around here inside this little pipette that I have here. And then we add it. Like this! And then a precipitate forms down here. It consists of solid silver chloride, so we'll print that directly here. The fact that a precipitate of silver chloride forms is because silver chloride is very insoluble in water. The solution thus immediately becomes saturated with silver and chloride ions, and the remainder then precipitates out as a solid precipitate. Now if we were to write a reaction formula for this, we could write it this way, that we have silver ions in aqueous solution plus nitrate ions in aqueous solution plus sodium ions in aqueous solution and chloride ions in aqueous solution, and so this forms solid silver chloride. But the sodium ions and the nitrate ions, they remain in the solution here, nothing happens to them. And ions like these that just remain in the solution, and nothing happens to them, it's as if they're just standing on the sidelines and watching. That's why they are actually called spectator ions. And the spectator ions, we can actually completely exclude them from the reaction formula – I mean, how interesting is it to say that we have nitrate ions to begin with and then it becomes – nitrate ions! It's not particularly exciting, so we can actually get rid of them altogether, both the sodium ions and the nitrate ions, and just write like this: Silver ions in aqueous solution plus chloride ions in aqueous solution becomes solid silver chloride. A precipitate of solid silver chloride is thus formed. We take another example, and then we write like this, that barium ions are reagents for sulfate ions. In the same way as before, we have a test tube here, with barium nitrate in aqueous solution, and then we have barium ions and nitrate ions in aqueous solution like this swimming around in the test tube. So we add some sodium sulfate solution in this case. And the sodium sulfate is also in aqueous solution, so that means we have sodium ions in aqueous solution and sulfate ions in aqueous solution here as well, swimming around here inside the pipette. Now we add this solution, and then this time we also get a white precipitate, but this time the precipitate is barium sulfate. We write down a reaction formula for that too, and write with all the ions that we have in the aqueous solutions, and then we'll see what is formed then. So we have barium ions in aqueous solution, we have nitrate ions in aqueous solution – two nitrate ions even, because we have two nitrate ions for each barium ion (the barium ion is divalent positive). So we have two sodium ions, that's exactly what it says up here, Na₂SO₄, and we have a sulfate ion, also in aqueous solution. What is then formed is this solid precipitate of barium sulfate. But the sodium ions and nitrate ions, they still hang around in the solution, nothing happens to them, so they are spectator ions too. We can also shorten them so that we get a simpler reaction formula just like this: Barium ions in aqueous solution plus sulfate ions in aqueous solution, it becomes solid barium sulfate this way. This leads us to some rules of thumb about soluble salts. They say this, that all salts of sodium and potassium and ammonium ions, they are actually readily soluble. And all nitrates, i.e. salts with nitrate ions, are also easily soluble. It is also the case that most chlorides are readily soluble, with the notable exception of silver chloride and then also lead chloride, PbCl₂. The final rule of thumb is that most sulfates are readily soluble, except for barium sulfate and also lead sulfate, which is very insoluble, and to a certain extent calcium sulfate, which is somewhat semi-insoluble. There are some rules of thumb for poorly soluble salts as well. They are like this, that all carbonates are poorly soluble, all phosphates are poorly soluble and all hydroxides are poorly soluble, except if they are together with these sodium, potassium and ammonium ions that I talked about just now. Then they are easily soluble. So if we have, for example, sodium carbonate or sodium phosphate or sodium hydroxide, they are readily soluble. Or potassium carbonate or ammonium phosphate or whatever you want, they are all easily soluble. But other salts with these negative ions, such as calcium carbonate or calcium phosphate, are poorly soluble. Now we can take some examples of this. Does any precipitate form? Then this is like a typical exam question that could appear. A solution of barium nitrate is mixed with a solution of ammonium chloride. The question now is, does any precipitate form? Yes, just to be safe, we'll draw this out the same way we did before. So we have a solution of barium nitrate here, which means we have barium ions and nitrate ions swimming around and having a good time here in this test tube. And then we add a little ammonium chloride in aqueous solution to this – that's "aq" here. That means we have free ammonium ions and free chloride ions swimming around in here. The question now is: Does any precipitate form when we add this? We're going to write down a reaction formula, or at least the beginning of a reaction formula, and check. We then have ammonium ions in aqueous solution, chloride ions in aqueous solution, we have barium ions in aqueous solution and nitrate ions in aqueous solution. And the question now is, can any precipitate form? We can imagine, for example, that a precipitate of solid ammonium chloride might form if we combine the ammonium ions with the chloride ions. But, now let's check here. We started here with ammonium chloride in aqueous solution, in this pipette. This means that ammonium chloride is a readily soluble salt! Additionally, we can remember from the rules of thumb that ammonium salts are readily soluble, and so are most chlorides, so no – a precipitate of solid ammonium chloride does not form. We can, in good conscience, put a big red cross over it. Another variant that you can imagine is barium nitrate – we have barium ions here and nitrate ions here. Could it be that they pair up and become a poorly soluble salt? No! They can't. Because in the test tube here, we already had barium nitrate in aqueous solution from the beginning. Then we understand that barium ions and nitrate ions form a readily soluble ionic compound. And by the way, all nitrates are also easily soluble, so we'll put a big red cross over that one too. We're looking at another option here, could it be that a precipitate of solid ammonium nitrate is formed? We have the ammonium ion there and the nitrate ion there. But then we think back to those rules of thumb again, and then we might remember that all ammonium salts are easily soluble – and all nitrates are also easily soluble! Double up on easy solubility here, a big fat red cross over that one, doesn't work. What we are left with now is that it could perhaps be barium chloride that is formed, if we combine barium ions here with chloride ions here. But then we also remember from the rules of thumb that most chlorides are readily soluble except silver chloride and lead chloride. This means that barium chloride is also readily soluble. Check it out too! This means that no precipitate forms in this mixture. Funny, huh? Let's take another example. Now I say this, what precipitate is formed? When you have a solution of lead nitrate and mix it with a solution of potassium iodide, a precipitate forms. But which one? How can we find out about this now? Well, a good tip is to draw a small test tube here and pretend to fill it with lead nitrate. And the lead nitrate, it's just like before, that it's in an aqueous solution, so that means that here we have lead ions and nitrate ions swimming around and having a good time here inside the test tube. To this we now add some potassium iodide – in aqueous solution, of course, which means we have potassium ions and we have iodide ions splashing around here inside the little pipette. Now we already know that a precipitate forms when we add the potassium iodide, and that precipitate is actually a beautiful yellow like this. We take and write down the beginning of a reaction formula here. We write out all the ions that we have, and then we ask ourselves, can we combine these into a sparingly soluble salt? Potassium ions in aqueous solution, iodide ions in aqueous solution, lead ions in aqueous solution and nitrate ions in aqueous solution, what can become of them? Could it be that a precipitate of solid potassium iodide forms? Then we see here that we start with potassium iodide in aqueous solution in our pipette here. Potassium iodide must therefore be a readily soluble salt. And besides, all potassium salts are easily soluble, as we said in those rules of thumb, so no – no potassium iodide precipitate forms. Instead, we try combining some lead ions and nitrate ions and think, could it be that it is a poorly soluble salt? But we remember looking here at what we had from the beginning: Lead nitrate in aqueous solution, swimming around here and having a great time, not at all difficult to dissolve. And besides, all nitrates are easily soluble, so it's not a precipitate of solid lead nitrate that forms either. Then we try a new combination and think, could it be that a precipitate of solid potassium nitrate is forming? We're testing that, the potassium ions here and the nitrate ions here, could those be the ones that form the precipitate? But, no, we remember from those rules of thumb that all potassium salts are readily soluble, and all nitrates, they are also readily soluble. So, no solid potassium nitrate is formed either. We are now left with only lead iodide, that is the only option left, and that is what is formed. A precipitate of solid lead iodide forms. And then we can write a simplified reaction formula for what's happening, where we don't include any spectator ions at all, and we write it like this: Lead ions in aqueous solution together with two iodide ions in aqueous solution, they form solid lead iodide, and it's this beautiful yellow precipitate that forms. We will also look at a third reagent here, and that is that limewater is a reagent for carbon dioxide. If you bubble carbon dioxide gas into a solution of limewater, it becomes cloudy in this way, as you can see in the picture. But what is really happening then? Yes, we start by writing like this, that lime water, it is a saturated solution of calcium hydroxide, Ca(OH)₂(aq). If carbon dioxide gas is blown in, the solution becomes cloudy because insoluble calcium carbonate is formed. It happens according to this reaction formula: Ca(OH)₂ in aqueous solution plus CO₂ in gaseous form becomes CaCO₃ in solid form plus water. And so it is this solid calcium carbonate that causes the cloudiness that you saw in the photo just now. If you add an excess of carbon dioxide, the cloudiness will disappear after a short while. This is because the calcium carbonate then begins to convert into easily soluble calcium bicarbonate. It happens according to this reaction formula: CaCO₃ in solid form plus water and carbon dioxide in gaseous form becomes calcium hydrogen carbonate, i.e. Ca(HCO₃)₂ in aqueous solution.