0:16

So the question is, how do we figure out what the formulas are for ionic compounds?

Â We know that they're formed from a cation or positively charged ion and an anion, or

Â a negatively charged ion, and that there's attraction to each other.

Â But how do we know how many of each we have to have together?

Â Well, ultimately our main concern is that we have species

Â that have electrically neutral formulas.

Â In other words, that we have the same number of electrons lost by the cations,

Â as we have gained by the anions.

Â And those subscripts, tell us how many of each of those ions we

Â need to put together, in order to get that electrically neutral compound.

Â 1:00

Now, when we look at the compounds,

Â we're going to see lots of different types of ions.

Â We've already talked about cations and

Â anions, we're also going to see that we have what we call monatomic, or

Â monoatomic, because there's one atom in those ions.

Â Add polyatomic ions, so many atom ions.

Â And we also see that the way the main group elements and the way

Â the transition elements behave in ionic compounds, is a little bit different.

Â 1:29

So when we talk about main group ions, what we're talking about are the first two

Â columns here, and the last six columns here.

Â Those are considered our main group elements, and

Â those are the ions we're most concerned about right now.

Â We will talk about our transition metals, which are here in the middle, but

Â they behave a little bit differently, so

Â we're going to start with our main group which are much more predictable.

Â What I see is if I look at elements in this first column,

Â 2:00

I see that they each want to lose one electron,

Â because if they lose one electron, they'll look like the nearest noble gas.

Â They'll have their complete octet, and they will be a stable ion.

Â If I look at the second column, say magnesium has 12 electrons,

Â it wants to look like neon, which only has ten electrons, so

Â it needs to lose two electrons.

Â So, where it's going to form a plus 2 charge.

Â This continues over here with boron and aluminum.

Â Gallium and indium not quite as predictable.

Â Most of these elements kind of down in this lower area are not quite as

Â predictable in their behavior.

Â 2:34

Carbon, we don't really see any ionic compounds forming from carbon.

Â If it did,

Â it would depend on what it was partnered with, whether it form a plus 4 or minus 4.

Â But we randomly see those, we don't have to worry about those.

Â When I start from the other end, notice that my noble gases here,

Â do not form ions because they're very stable,

Â that's why they're called the noble gases, they're inert, they don't react.

Â And if I look at my halogens, my column, here I get minus 1 charges,

Â minus 2, and minus 3 charges here for the nitrogen column.

Â Because these elements are trying to gain electrons,

Â to look like the nearest noble gas.

Â Fluorine wants to gain one electron, so it takes one and

Â becomes a minus 1, oxygen wants to gain two, nitrogen wants to gain three.

Â 3:23

So let's look at an example of how we would actually use this to actually figure

Â out the formula.

Â For example, if we're looking at the compound between lithium and chlorine,

Â we first want to write our symbols out, so Li for lithium, and Cl for chlorine.

Â If I look at the periodic table, I see that lithium is here in the first column,

Â so I know it's going to have a plus 1 charge, and I don't even have to

Â put the one in, I can just write that as plus, that indicates it's a one,

Â on the other side, I see I have chlorine, which is in the column of halogens.

Â And those elements anions with a negative 1 charge, and

Â again I can just put a negative sign without the one.

Â The ones are implied for both of these.

Â Now I need to figure out what the formula's going to be.

Â And there's a couple of ways we can do this.

Â One is we can look at my charges, and see well I've got plus 1 and minus 1,

Â they balance each other out, so all I need one of each.

Â Sometimes it's a little hard to figure out what those ratios are, so

Â we can use what we call the criss-cross down method, where I take the value of

Â my charge on the anion, and it becomes the subscript, on my cation.

Â And I take the value of the charge on my cation, and

Â it becomes the subscript on my anion.

Â Now in chemical formulas, we never write ones in the formula,

Â because it's always assumed it's one if there's no other number there.

Â So we can actually rewrite this as LiCl.

Â 4:54

plus, one chloride ion, with a minus 1 charge, and

Â when I add those values together, I get zero.

Â Which is exactly what I want, because ionic compounds are electrically neutral.

Â 5:12

Now, let's look at an example where we don't have the plus and minus 1.

Â Now, we have magnesium, which is Mg, and bromine, which is Br.

Â Now, when I look at magnesium,

Â I see that it's in the second column of our periodic table.

Â And it actually has a 2 plus charge.

Â Bromine, has a minus 1 charge.

Â I can take the same approach as I did before,

Â I can criss-cross down, and what I get is MgBr, whoops Mg1Br2.

Â I don't need to include the one in there, so I can rewrite that as MgBr2.

Â I need to leave the two, because I need to show that it takes two bromines, accepting

Â one electron each, to balance out the two electrons donated by the one magnesium.

Â If I look at my charges, I see I have one magnesium, with a 2 plus charge,

Â plus two bromines, with a 1 minus charge, and when I add these up,

Â I find that they equal zero.

Â Again, I have an electrically neutral compound.

Â 6:45

So when I criss-cross down, I end up with a Al2S3.

Â And I can't actually simplify that anymore.

Â I have to leave both of those subscripts in there,

Â because they're not values of one.

Â 6:58

So, I can go back, and check my charges.

Â I know I have two aluminiums, each with a 3 plus charge.

Â Three sulfides, each with a 2 minus charge, and

Â when I add those values up, I get zero.

Â 7:23

Now, when we look at the transition metals,

Â these elements in the middle of the periodic table, what we see is

Â that their behavior is not quite as predictable as our main group elements.

Â And what we find is that many of our transition metals can

Â actually have multiple charges.

Â 7:41

What I see in this table, is that I have all of the charges possible for

Â my first row of transition elements.

Â I see that scandium has only a plus 3 charge, and zinc only plus 2.

Â And there are also three transition metals: silver, zinc, and

Â scandium, which only have one possible charge.

Â The remaining transition metals have at least two, if not more possible charges.

Â Now, some of these are more common than others.

Â For example, iron is typically seen as two or three, and not six.

Â But it can happen.

Â Likewise, copper is generally one or two.

Â However, we can't know just by looking, and saying we have a compound with copper,

Â and know that, that's going to be copper plus 1 or plus 2.

Â Some additional information has to be given,

Â to help us figure out what the formulas is with these compounds.

Â And so, if we know we're dealing with, say, a copper one ion, and

Â a compound forming with oxygen, then we know, that the formula's going to be CU2O.

Â Because we have two coppers, each with a plus 1 charge,

Â and one oxygen with a 2 minus charge.

Â So we still get our electrically neutral compound.

Â 8:58

So, the first examples we looked, at all dealt with monatomic ions.

Â Single atom ions.

Â Now we have to deal with polyatomic ions.

Â And these are ions, which act as a group, okay?

Â But we have multiple atoms together.

Â So we have sulfate, which has the formula SO42 minus, nitrate, carbonate,

Â ammonium, and hydroxide.

Â And these are just a very few of the options.

Â One link that has some more information is given here at the bottom,

Â that you can go and look at some more examples of polyatomic ions.

Â There's also a link in the resources for

Â this module, that gives you another list of those ions.

Â And really, the best way to learn them, is to actually use them in compounds.

Â But generally students are going to have to kind of take some time to sit down and

Â memorize these,

Â until they've done enough problems to feel comfortable with just knowing them.

Â So, these always act as a group, and so, when I look at nitrate for

Â example, over here, what I see is that I have a nitrogen and

Â three oxygens, I have a net minus 1 charge.

Â And this can actually form an ionic compound.

Â Even though these are covalent elements involved,

Â because it forms this charged species, this polyatomic ion,

Â it will act as an ion and form an ionic compound, with a cation.

Â We can also form a polyatomic cation.

Â Most of them are anions, we do have one polyatomic cation here, NH4 plus.

Â And when we look at these compounds, they can actually form ionic compounds with

Â another polyatomic ion, or with a monatomic ion.

Â 10:34

When we're looking at the formulas for

Â compounds with polyatomic ions, we have to be careful in the way that we write them.

Â When I look at something like barium hydroxide, remember that barium is Ba,

Â and has a 2 plus charge.

Â Hydroxide, has a minus 1 charge.

Â When I form my compound, and get the formula, I have Ba,

Â and I don't need the one there, OH.

Â 11:00

But I can't just put a two beside the H,

Â because that looks like I have two hydrogens, and I don't.

Â I have two hydroxide units.

Â So, when I have a subscript other than one, for

Â a polyatomic ion, I must include the parentheses in the formula.

Â When I look at something like sodium sulfate, so

Â sodium has a plus 1 charge, sulfate has a 2 minus charge.

Â When I criss-cross down, what I see is that I get Na2SO4.

Â So I don't need parentheses in this case,

Â because I only have one sulfate ion in the formula, so

Â no parentheses are necessary, and it would be wrong to use them.

Â I can also have compounds where I have a subscript on the monatomic element, or

Â monatomic ion, as well as on the polyatomic ion.

Â Such as I have here in the aluminum sulfate.

Â 11:53

So let's, let you try one.

Â What is the formula for magnesium nitrate?

Â And we're given here the formula for

Â nitrate, until you're a little more familiar with those.

Â 12:09

So, what we get is magnesium nitrate, Mg(NO3)2.

Â Magnesium, has a 2 plus charge.

Â Nitrate, which we see has a minus 1 charge.

Â When I find my scri, subscripts, I see I have magnesium one,

Â so I don't need to write that in.

Â Nitrate, but I need two nitrates.

Â So, I must include the parentheses, and put the two outside the parentheses.

Â