 Next one, a little bit trickier, we're going to do the structure of an ion, a covalent ion. So the carbonate ion, although it's an ion and it participates in ionic bonding, within the ion itself the carbon and the oxygen atoms are joined to each other by covalent bonds. So we're going to draw this structure and we follow exactly the same set of procedures that we did for the neutral molecules, there's just one extra thing that you need to know and it comes in the first step. So first we total up our valence electrons as normal. We have one carbon which has four valence electrons and we have three oxygens and each of those oxygens has six valence electrons because it's in group six, which gives us a total of 18. Alright now here's the bit which is slightly different. This is an anion, a negative ion as opposed to a neutral molecule, which means that it's carrying a couple of extra electrons. Now its charge is 2 minus so it's actually carrying two extra electrons. So in addition to the valence electrons from the carbon and the oxygen we need to add two extra electrons that represent the ions charge. Okay from now on we continue as normal. So our total number of valence electrons is four plus 18 which is 22 plus another two which is 24 electrons. Next step is to look at the bonds. Carbon we know likes to form four bonds and oxygen likes to form two. So we're going to try and draw a skeleton for this ion. Because carbon is the one that forms the most bonds I'm going to try putting it in the middle and I only have three other atoms to join it to. So let's start by putting in single bonds. There we go, I've joined up the carbon with the three oxygens. But we know that the carbon likes to form four bonds and at the moment it's only got three. So I'm going to make one of these a double bond. So now my carbon is happy, it's got four bonds. The oxygen at the top is happy, it's got two bonds. The other two oxygens only have one. But as you'll see in a second we can solve this with non-bonding electrons and it's a situation that arises mostly when you're drawing ions. So let's go on and look at the full outer shell. So carbon has four bonds which means it has access to eight electrons so it's fine. The oxygen at the top has access to four bonding electrons because it's got two bonds. So it needs an extra four electrons so I'm going to give it two lone pairs. So it now has four bonding and four non-bonding electrons so it's happy. The other two electrons have only formed one bond each. So they only have access to two electrons. So I can't give them any more bonds. If I give them another bond it will be too many bonds for the carbon and we just can't do that. So what I'm going to do is I'm going to kind of make it up to them by giving them extra non-bonding electrons. So they've each got two. They need another six. So I'm going to give them three lone pairs like that. Now the oxygens are happy. All right, final step is I'm going to check that I've used the right number of electrons. I had 24 to start off with. So I used up two, four, six, eight, ten, 12, 14, 16 non-bonding electrons. And I've got four bonds. So that's eight bonding electrons, which is a total of 24, which is the right number. So this is a correct structure. Now there's one final thing we need to add to this. As we've drawn it, the Lewis structure looks like it's meant to be a neutral molecule. We have to add something that indicates that it's an ion. So what we do is we put the entire structure into square brackets and we put the charge on the outside like that, two minus, to show that it is in fact an ion. So there are actually three things that were different, not one, about drawing this Lewis structure of a polyatomic ion. First, you need to adjust the number of valence electrons to reflect the charge on the ion. If it's a negative ion, you'll have to add electrons. If it's a positive ion, you'll have to take electrons away. Second, often with oxygens in negative ions, you'll have to make do with oxygens that have only one bond, even though oxygens usually prefer to have two. This is okay as long as all of the atoms have a full outer shell and you have the correct number of electrons in your structure. And third, you need to put your final Lewis structure in brackets and indicate the charge on the ion.