 So let's take a simple example, the combustion of hydrogen in oxygen to produce water. Let's first write out the formulae of the substances. We know that hydrogen and oxygen are part of the diatomic seven, so we can write them as H2 and O2, and water is H2O. So we can write the skeleton of our equation like this. Now let's draw some diagrams of our molecules so that it's easy to see the individual atoms. We'll draw a hydrogen molecule and an oxygen molecule. You can see each of them is made of two atoms joined together, bonded together, and they will somehow combine to give water. Okay, let's do some atom accounting. As we've drawn them, we have two hydrogen atoms on the left and also two on the right, here and here. So we're okay for hydrogen. No hydrogen has mysteriously disappeared or been created. But what of the oxygen? On the left we've got two oxygen atoms, but on the right we only have one, so our oxygen atoms are not balanced. There are two ways that you might deal with this. One would be to change the formulae of water. Let's make it H2O2. Then we've got two hydrogens on the left and two on the right, and two oxygens on the left and two on the right. Can we do this? No! We cannot do it. And the reason that we can't do it is that by changing the formulae, you completely change the compound. H2O2 is hydrogen peroxide, a highly reactive bleach. It's really not something you want to be knocking back a glass of. It might help to try reading this out loud if you're having trouble getting it. Okay, so when hydrogen and oxygen react, you don't get hydrogen peroxide. You get water. And water is H2O, and so the H2O has to stay H2O. Never change the formulae of your compounds. So how else can we get another oxygen there on the right? Well, the only other option is to have another water molecule. Now our oxygen is fine. We've now got two on the left and two on the right, but our hydrogen is unbalanced. We've got two on the left and four on the right. So what do we do? Well, we just need a little more hydrogen on the reactant side. And now we're good. We've got four hydrogens and two oxygens on the left and four hydrogens and two oxygens on the right, but rearranged this time into two water molecules. The final thing we need to do is show these adjustments in the reaction equation. And the way that we do this is by using large numbers at the front of each of the formulae to show how many of each one there are. Now let's start with the oxygen. We have only one oxygen molecule, so we actually leave it blank. The tradition or the convention is that if there is no number out the front of the molecule or the atom, that just means that there's one of them. If there's more than one, then we have to indicate it with a number. So for our hydrogen molecules, we have two of those. So we indicate that with a large two. Two hydrogens plus one oxygen gives and we have two water molecules, so we put a large two out the front there as well. Now if you want to check that your balancing is correct and you don't want to be bothered actually drawing out pictures of all the atoms and molecules, then the way to do it is to scan through your reaction equation and do some simple maths. So again we're checking to make sure that the number and type of atoms on the left is the same as on the right. So we'll check the hydrogens first. On the left, we have an H2 molecule, so that's two hydrogens in one molecule, and we have two of those molecules. So two times two is four hydrogens on the left. And if we go over to the right, we find we have in a single water molecule, there are two hydrogens, but we have two water molecules, so again it's two times two which is four hydrogens. So four on the left, four on the right and we're balanced. For oxygens, on the left we have one oxygen molecule and that molecule is made up of two oxygen atoms, so that's two O's. And on the right, each water molecule has a single oxygen atom and there are two water molecules, so that's two oxygens. So our oxygens are also balanced.