In this module, we're going to look at how to find the formulas of ionic compounds. By the end of this module, you should be able to write the formula for ionic compounds, from a variety of types of ions. So the question is, how do we figure out what the formulas are for ionic compounds? We know that they're formed from a cation or positively charged ion and an anion, or a negatively charged ion, and that there's attraction to each other. But how do we know how many of each we have to have together? Well, ultimately our main concern is that we have species that have electrically neutral formulas. In other words, that we have the same number of electrons lost by the cations, as we have gained by the anions. And those subscripts, tell us how many of each of those ions we need to put together, in order to get that electrically neutral compound. Now, when we look at the compounds, we're going to see lots of different types of ions. We've already talked about cations and anions, we're also going to see that we have what we call monatomic, or monoatomic, because there's one atom in those ions. Add polyatomic ions, so many atom ions. And we also see that the way the main group elements and the way the transition elements behave in ionic compounds, is a little bit different. So when we talk about main group ions, what we're talking about are the first two columns here, and the last six columns here. Those are considered our main group elements, and those are the ions we're most concerned about right now. We will talk about our transition metals, which are here in the middle, but they behave a little bit differently, so we're going to start with our main group which are much more predictable. What I see is if I look at elements in this first column, I see that they each want to lose one electron, because if they lose one electron, they'll look like the nearest noble gas. They'll have their complete octet, and they will be a stable ion. If I look at the second column, say magnesium has 12 electrons, it wants to look like neon, which only has ten electrons, so it needs to lose two electrons. So, where it's going to form a plus 2 charge. This continues over here with boron and aluminum. Gallium and indium not quite as predictable. Most of these elements kind of down in this lower area are not quite as predictable in their behavior. Carbon, we don't really see any ionic compounds forming from carbon. If it did, it would depend on what it was partnered with, whether it form a plus 4 or minus 4. But we randomly see those, we don't have to worry about those. When I start from the other end, notice that my noble gases here, do not form ions because they're very stable, that's why they're called the noble gases, they're inert, they don't react. And if I look at my halogens, my column, here I get minus 1 charges, minus 2, and minus 3 charges here for the nitrogen column. Because these elements are trying to gain electrons, to look like the nearest noble gas. Fluorine wants to gain one electron, so it takes one and becomes a minus 1, oxygen wants to gain two, nitrogen wants to gain three. So let's look at an example of how we would actually use this to actually figure out the formula. For example, if we're looking at the compound between lithium and chlorine, we first want to write our symbols out, so Li for lithium, and Cl for chlorine. If I look at the periodic table, I see that lithium is here in the first column, so I know it's going to have a plus 1 charge, and I don't even have to put the one in, I can just write that as plus, that indicates it's a one, on the other side, I see I have chlorine, which is in the column of halogens. And those elements anions with a negative 1 charge, and again I can just put a negative sign without the one. The ones are implied for both of these. Now I need to figure out what the formula's going to be. And there's a couple of ways we can do this. One is we can look at my charges, and see well I've got plus 1 and minus 1, they balance each other out, so all I need one of each. Sometimes it's a little hard to figure out what those ratios are, so we can use what we call the criss-cross down method, where I take the value of my charge on the anion, and it becomes the subscript, on my cation. And I take the value of the charge on my cation, and it becomes the subscript on my anion. Now in chemical formulas, we never write ones in the formula, because it's always assumed it's one if there's no other number there. So we can actually rewrite this as LiCl. And what I find, is that I have one lithium ion, with a plus 1 charge, plus, one chloride ion, with a minus 1 charge, and when I add those values together, I get zero. Which is exactly what I want, because ionic compounds are electrically neutral. Now, let's look at an example where we don't have the plus and minus 1. Now, we have magnesium, which is Mg, and bromine, which is Br. Now, when I look at magnesium, I see that it's in the second column of our periodic table. And it actually has a 2 plus charge. Bromine, has a minus 1 charge. I can take the same approach as I did before, I can criss-cross down, and what I get is MgBr, whoops Mg1Br2. I don't need to include the one in there, so I can rewrite that as MgBr2. I need to leave the two, because I need to show that it takes two bromines, accepting one electron each, to balance out the two electrons donated by the one magnesium. If I look at my charges, I see I have one magnesium, with a 2 plus charge, plus two bromines, with a 1 minus charge, and when I add these up, I find that they equal zero. Again, I have an electrically neutral compound. Now you try one, that's a little bit more challenging. What is the formula for the ionic compound formed from aluminum and sulfur? So, our answer is Al2S3. Remember that Al, has a 3 plus charge. Sulfur, has a 2 minus charge. It's in the same column as oxygen. So when I criss-cross down, I end up with a Al2S3. And I can't actually simplify that anymore. I have to leave both of those subscripts in there, because they're not values of one. So, I can go back, and check my charges. I know I have two aluminiums, each with a 3 plus charge. Three sulfides, each with a 2 minus charge, and when I add those values up, I get zero. So, again, I have an electrically neutral compound. Now, when we look at the transition metals, these elements in the middle of the periodic table, what we see is that their behavior is not quite as predictable as our main group elements. And what we find is that many of our transition metals can actually have multiple charges. What I see in this table, is that I have all of the charges possible for my first row of transition elements. I see that scandium has only a plus 3 charge, and zinc only plus 2. And there are also three transition metals: silver, zinc, and scandium, which only have one possible charge. The remaining transition metals have at least two, if not more possible charges. Now, some of these are more common than others. For example, iron is typically seen as two or three, and not six. But it can happen. Likewise, copper is generally one or two. However, we can't know just by looking, and saying we have a compound with copper, and know that, that's going to be copper plus 1 or plus 2. Some additional information has to be given, to help us figure out what the formulas is with these compounds. And so, if we know we're dealing with, say, a copper one ion, and a compound forming with oxygen, then we know, that the formula's going to be CU2O. Because we have two coppers, each with a plus 1 charge, and one oxygen with a 2 minus charge. So we still get our electrically neutral compound. So, the first examples we looked, at all dealt with monatomic ions. Single atom ions. Now we have to deal with polyatomic ions. And these are ions, which act as a group, okay? But we have multiple atoms together. So we have sulfate, which has the formula SO42 minus, nitrate, carbonate, ammonium, and hydroxide. And these are just a very few of the options. One link that has some more information is given here at the bottom, that you can go and look at some more examples of polyatomic ions. There's also a link in the resources for this module, that gives you another list of those ions. And really, the best way to learn them, is to actually use them in compounds. But generally students are going to have to kind of take some time to sit down and memorize these, until they've done enough problems to feel comfortable with just knowing them. So, these always act as a group, and so, when I look at nitrate for example, over here, what I see is that I have a nitrogen and three oxygens, I have a net minus 1 charge. And this can actually form an ionic compound. Even though these are covalent elements involved, because it forms this charged species, this polyatomic ion, it will act as an ion and form an ionic compound, with a cation. We can also form a polyatomic cation. Most of them are anions, we do have one polyatomic cation here, NH4 plus. And when we look at these compounds, they can actually form ionic compounds with another polyatomic ion, or with a monatomic ion. When we're looking at the formulas for compounds with polyatomic ions, we have to be careful in the way that we write them. When I look at something like barium hydroxide, remember that barium is Ba, and has a 2 plus charge. Hydroxide, has a minus 1 charge. When I form my compound, and get the formula, I have Ba, and I don't need the one there, OH. But I can't just put a two beside the H, because that looks like I have two hydrogens, and I don't. I have two hydroxide units. So, when I have a subscript other than one, for a polyatomic ion, I must include the parentheses in the formula. When I look at something like sodium sulfate, so sodium has a plus 1 charge, sulfate has a 2 minus charge. When I criss-cross down, what I see is that I get Na2SO4. So I don't need parentheses in this case, because I only have one sulfate ion in the formula, so no parentheses are necessary, and it would be wrong to use them. I can also have compounds where I have a subscript on the monatomic element, or monatomic ion, as well as on the polyatomic ion. Such as I have here in the aluminum sulfate. So let's, let you try one. What is the formula for magnesium nitrate? And we're given here the formula for nitrate, until you're a little more familiar with those. So, what we get is magnesium nitrate, Mg(NO3)2. Magnesium, has a 2 plus charge. Nitrate, which we see has a minus 1 charge. When I find my scri, subscripts, I see I have magnesium one, so I don't need to write that in. Nitrate, but I need two nitrates. So, I must include the parentheses, and put the two outside the parentheses. In the next module, what we're going to look at is how we name these compounds. We've talked about the formulas, now we need to know how to name them.
