 OK, so now we're going to do the Lewis structure of a polyatomic ion, OK? So again, normally you would look at this structure and first want to identify, well, what's the central atom and what are the peripheral atoms? So hopefully you would identify carbon as being the central atom in this case because both because it can make more bonds and it's less than electronegative than oxygen, OK? So I want you to draw carbon in your oxygens around and I'm going to draw two and we'll see if we get the same. So now what I want you to do is draw your Lewis structures of all of those elements, OK? So do that. So carbon's got how many electrons? Four. Four. And oxygen? Six. Six. Six. Like that. So I drew them in such a way because I know what the structure is going to be. And also notice up here it says that I've got two extra electrons, OK? Most of the time those extra electrons will go on the most electronegative atoms, OK? So in this case you've got two electrons, so they're not going to go on the same atom, OK? So put them on the two of the most electronegative atoms. So in this case which one is what you put them on, oxygen or carbon? Oxygen. Oxygen. So I put one here and one here, OK? So or I could put one there, whichever one you want, OK? So I'm going to add one there and if I do that then it's got one extra electron. So that means it has a what on it, a negative charge, yeah, very good, like that. So over here, if I put an extra electron, that has a negative charge, right? So this oxygen and this oxygen carry what we call formal charges, OK? The charge is formally on that oxygen, OK? So you've normally seen it like this and not really known where those charges are, but that they might be throughout this charge. Now you know that they're specifically on a particular atom, OK? So now how many bonds can this oxygen make? One. One. This one? One more. And this one? Two. And this guy? Four. So do we have enough bonds for everybody to go around? So let's draw the structure. There's the Lewis structure of the carbonate atom. So hopefully you got something that looked like that. So again, formally, you can see where those two negative charges are. We call them formal charges. So up here it says, yeah, I know, but like, is there another element in there that you talked about? Do you provide a charge? We just added two electrons to this thing. Don't you have to put those over? I don't know a paper just there, but like, it's an ion, it's a polyatomic ion. So polyatomic ions have charges. So it's just like saying a chloride ion, you know? So like a chloride ion, something like this, right, how do you get that negative charge? Well, at some point in time, it got an electron from somewhere, and that's what I'm saying with the polyatomics. At some point in time, they got an electron from somewhere, OK? This chlorine, like if this chloride is in sodium chloride or something like that, didn't necessarily get that electron from that sodium atom, you know what I'm saying? Just sometime down the line got an electron, OK? The thing is, why does it have this structure is probably the better question, because it's more stable. Why? In this case, everything has a full 10, if that makes sense, OK? That's probably a better question, as opposed to where did it come from? Yeah, I was just saying, like normally, you know, it's at this date, but in order for it to be an anionized, that's a good thing to come from somewhere, is that the one? I mean, yeah, I mean, that's the way we've been presenting, you know what I'm saying? But it doesn't necessarily, the transfer happens right before it becomes an ion, you know? That's not what necessarily happens. Any other questions on this? I won't include...