 And this one is on the reaction between acids and metals. This is not a neutralization reaction. Water is not a product of this particular reaction. In fact, this is better described as a displacement reaction where in most cases what you'll find is the metal, a more active metal, will be displacing hydrogen ions from the solution. And so we have some slightly different products. Now one thing I should say at the beginning is that the concentration of the acid and the type of acid can actually change the products of this reaction. We're going to simplify things in this particular video in looking at the reaction of metals with dilute acid solutions, but where the solutions are concentrated you can actually find some of the less reactive metals will react, but they won't always produce the same product. So we might have a look at a few of those examples as extension examples or as examples of interest in class, but for now we just want to focus on these key general types of reactions. So the first thing we want to do is we have with the last two groups of reactions is to look at the general equation. So here's a general equation, an acid and again as I mentioned we're talking about dilute acids plus metal produces salt and hydrogen gas. We'll use the same sort of examples that we've looked at previously to give us an example. So hydrochloric acid, which as we know is a solution reacting with magnesium solid will form hydrogen gas and a salt. And again the salt we can determine by looking at, well firstly the metal itself, the metal effectively is going to push the cation out of the solution. So magnesium is going to be the first part of this salt and the cation is going to come from the acid as it usually does. So this particular salt would be magnesium chloride and using our NAG SAGs we know that group 7 elements form soluble salts with a couple of exceptions. Magnesium is not on the exception list so therefore this will be a soluble salt in solution. And so that brings us of course to the net ionic equations which we've been looking at previously for some of our other examples. So in the net ionic equation if we split up our acids as we have and of course I've done this without balancing the equation at the top you'll notice that when the equation is not balanced you'll end up with problems later on. So there's a small oversight, hopefully you picked that up before I told you about it. But balancing equations is obviously a very, very important part of writing of your chemistry. So we've also got magnesium and as a solid it is not an ionic form. And they're going to produce magnesium ions and chloride ions too and hydrogen gas. So as we've done previously where we have spectator ions we can always cancel those out. In this case we only have one spectator ion which is the chloride ions. When we cancel those out our net ionic equation is going to be 2H+, plus mg solid goes to mg2+, plus h2 gas. And of course these are in aqueous solution so I'll just throw those in as well there. So there's our net ionic equation. The chloride ions, the anion, has been cancelled out because it's common, it remains in the solution from the beginning to the end. But what we've done is we've basically used our more active metal to push out the hydrogen ions from the solution.