 Apparently, before we start, I wanted to make an announcement. Apparently, there was a mishap with the grading on L2. I'm sure that everybody's aware of it by now. I will figuring out what to do. I'll make sure everybody gets the grades that they're supposed to get. Nothing's going to happen with whatever the grade said. If you did it all and the grade said you got a 40 or something like that, it's no big deal. We'll make sure that it's all taken care of, I promise you. I know you guys got a lot to worry about without having to worry about all that. It was my fault, apparently, and I apologize for that. Just to let everybody know, there is a mishap with the grading on L2. I'll figure it out and we'll get it all straightened up. Other than that, quiz two was passed back today. If you didn't get it, like I said, you could come up and get it after class. There's the key up here for quiz two if you were wondering why you got anything wrong. First thing I would do is add up the points that I took off and if it's different than what the grade says, make sure you come and bring it up to me because I have obviously been known to make errors and grade it. That being said, I guess I wanted to finish up chapter three. I know we kind of sped through at the end talking about polyatomic ions. Of course, this wasn't on the test because I thought we went through it rather fast. Let's just go back to this, 3.6-compounds introduction to Bonnie. Monoatomic ions, hopefully by now you're fairly familiar with, these are ions that consist of one atom, like H plus, Li plus, Ma plus, K plus, Mg2 plus, S2 minus. These are all the common main group ions and the transition metal ions that you are familiar with. We contrast these to polyatomic ions and these are ions that are composed of two or more atoms together with an overall positive or negative charge. I'd just like to compare monoatomic ion with polyatomic ion first. A common monoatomic ion that you may be familiar with is Br minus, of course. A polyatomic ion would be something like this. I'm just going to draw the structure of it just to emphasize, even though we don't know structure yet. Just to emphasize that you can see it's more than one atom. It's one, two, three, four, five atoms. All five of those atoms, we say, have a single positive charge. In fact, the way I've drawn it, I'm saying that the positive charge is actually residing most on that nitrogen. In fact, it is. The way you'll first learn how to do this is to take this and put brackets around it and say the whole thing has one positive charge. This is a conglomerate of atoms, if you will. That has a positive charge amongst them. If we look here, it's only a single atom with a charge on it. That's the difference between a polyatomic, poly meaning many. Polyatomic meaning many atoms. It's a polyatomic ion and a monatomic. Some polyatomic ions I'd like you to be familiar with are these pink here. Ammonium, nitrate, sulfate. This table isn't found in your book, unfortunately, but it is in the notes. Go to the notes, find this table 3.3, and look for the ones in pink. Those are the ones I want you to be familiar with. Ammonium, nitrate, sulfate, hydroxide, cyanide, phosphate, carbonate, bicarbonate, and acetate. I don't think any of the other ones. If we use any of the other ones, I'll specifically tell you what their names are. I definitely want you to know if I were to give you one of the other ones, you could identify it as a polyatomic ion. If I gave you MnO4- you could say that's a polyatomic ion instead of a monatomic ion. I wouldn't expect you to remember that it's the permanganade ion. It's got a negative one charge. Here's just another list of polyatomic ions. It, again, takes some things about polyatomic ions. You can think of polyatomic ions as charged molecules. A molecule consists of more than one atom. In this case it would be like a charged compound. A compound consists of more than one type of atom. Polyatomic ions are covalently bonded to one another. What we're going to learn today is that these lines represent a chemical bond between these two atoms. We kind of started talking about this in lab last week about bonds, breaking bonds, making bonds, and that type of stuff. Those lines in between those atoms represent bonds. If we believe that then we can say these polyatomic ion atoms are covalently bonded to one another, but the overall particle has a positive charge. It's an ion. Even though you've got covalent bonds in it, it's an ion. The whole thing will form ionic bonds. If I got these two guys together, what we would actually form is NH4, because it's got a positive one charge and it's got a negative one charge, so those will cancel out. We've got NH4Br. This would be the formula unit of these two combines. This name of this compound is ammonium bromide, and we'll get to that in a second. If you look, this compound here, or this ionic substance, this salt, has both covalent bonds within it and the cation. It has covalent bonds. Between the cation and anion has ionic bonds. It's got both types of bonds in it. Here's some common, we can get to it, common metal polyatomic ion compounds. Calcium hydroxide, OH- is hydroxide. How do we know it's OH-? Because we know calcium is 2+, and if we've got a formula like that, 2+, then since we have two of those OH things, then they both must be minuses to cancel it out. This is how you actually will predict if you don't know the charges inherently on these polyatomic ions, strongly urge you to memorize. But there might be on a test something that you forget, or something like that. You can go back to what you know about the main group elements and infer, if you will, what the charge on the polyatomic ion is. Some other ones, strontium nitrate, sodium bicarbonate, sodium hypochloride. These are all common compounds that you find. Here's another picture of some ionic compounds. Most of these have polyatomic ions associated with them. So cobalt nitrate, cobalt 2 nitrate is this red solution. Potassium dichromate, of course, is the orange solution. Cobalt 2 sulfate is this blue solution over here, and potassium permanganate. Okay, so to name these compounds, I know that you guys haven't memorized the polyatomic ions yet, but to name these, you will have to know all of those polyatomic ions. So this one, of course, would be ammonium chloride, barium sulfate, iron 3 nitrate, copper 1 bicarbonate, and calcium hydroxide. So notice I was putting the, for the iron and the copper, I was putting the Roman numeral in the name. You have to do that with transition metals. Okay, let's look at ionic bonding, strength. What you find is there's another periodic trend here. Kind of, you've got large ions that have 1 plus and a 1 minus charge, right? These are not going to be attracted to each other as if you have smaller ions that have a full positive 1 positive or negative 1 charge. And the reason is because if you think about these two, we can say this is a mon... and this is another monatomic ion. Okay, this would be like comparing, I don't know, rubidium and lithium, we'll say. Okay, what you find is that lithium is a very much smaller ion, and rubidium, due to its placement on the periodic table, but they both have the same amount of charge, 1 plus, okay? So rubidium can actually spread that charge out, and what we'll say is delta plus, okay? Delta, so that means partial positive, okay? So this thing can spread its charge out all the way across its surface, yeah? And so can this one. It can do it too. But of course, it's got a smaller surface to do that. So if you were to think about it in a different way, how many little delta positives could you write on the surface of each one of these things? Okay, of course you could write more on this one than you could on this one due to the fact that this is larger, okay? Does everybody get that? Okay, so what happens is or what's the case is that atoms don't like to be charged in general, okay? They do it to get more stable, but even though they're getting more stable, they're still not like completely stable because they've got this positive or negative charge on them, and they don't like it, okay? Which is the reason why they form ionic compounds, okay? So rubidium or lithium, they don't like to have this positive charge, so if they see some chlorine over here, they'll stick to it, so it'll cancel out that positive and negative charge, okay? So what you find is that when you've got two real small ions, of course with the same charge, and compared to two real big ions with the same charge, what'll happen is that these two... So this is negative, negative, negative. These guys will be very much attracted to each other. These guys will be less so because they don't mind having that charge distributed about their surface area as much as these guys do. They still don't like it, so they'll still be attracted to each other. They just won't have as much strength or force of attraction that these guys do. Does that make sense? Because these ones are small, okay? And then, of course, when you increase the charge amounts, so if you don't increase the size of it and then increase the charge, then it makes it even crazier. It's like, okay, so instead of lithium, we'll do beryllium. So beryllium 2 plus really can't stand it because it's got two pluses on it, right? And it's got to distribute amongst the same amount of surface area. So if you find oxygen 2 minus put it in the presence of beryllium 2 plus, they'll really stick together because they're both so small, okay? And then, of course, when you get bigger 2 plus, of course, the attraction isn't as much. And you can see that by looking at the melting points of these different compounds. So lithium fluoride or lithium oxide, the melting point of lithium oxide is much bigger or higher than the melting point of lithium fluoride. Why is that? Because oxide has the 2 minus charge and fluoride only has the 1 minus charge. And then if you look at calcium sulfide compared to potassium sulfide, again, it's the same sort of instance where the... So they're about the same size, calcium and potassium, right? But look at the difference in melting point. It's fairly dramatic. It's due to the fact that calcium has a 2 plus and sulfur has a 2 plus, whereas potassium only has a 1 plus and sulfur has a 2 minus. Okay, so covalent bonding. We talked about covalent bonding. But again, we went through it kind of fast, so I didn't do it on the test. Covalent compounds form from sharing electrons between one another, so they don't transfer their electrons like we talked about with ionic compounds. So they typically form when two or more non-metals react. We'll show more about this next statement in the next slide. But the bond results from the electrostatic attraction between the nucleus of one atom and the electrons of the other atom. Covalent compounds exist as discrete molecules that behave as individual particles, as opposed to the kind of extended networks that ionic compounds result in in compounds here. This ionic compound is actually a polyatomic, monoatomic ionic compound. So this would be like calcium nitrate or something like that. So you can see the polyatoms here with the mono atoms. Just a S2 monoatomic, and hopefully see the difference between those. And then hopefully also see the difference between these and this, right? This is just one little particle. Same thing, same thing, same thing, same thing, same thing. Let's show another covalent compound. What's the difference, right? Very small, discrete individual particles relative to whole conglomerate of things. Very small, discrete particles. Okay? So when we name covalent compounds, usually this only applies, I know you're going to start wanting to name very, very long covalent compounds, but this usually only applies to when you're talking about covalent compounds consisting of two different types of atoms, okay? So you can start naming covalent compounds, but don't try to name them if they've got more than two types of atoms because they have a different name. They won't be named like this. But if they do have two, only two different types of atoms in them, what you can find is that you can name them. So the names of the elements are written in the order which they appear in the formula. Usually they're written with the atomic number of the first atom being less than the atomic number of the second atom, but that's not always the case as you can see. It's not always the case. They're all written very strangely, but if you see the formula written like this, you're going to want to name it in the same way that the formula is written. So there's not really a rhyme or reason the way that they are written. You can kind of say this may be the case that the least, but we haven't learned really about electronegativity yet, but you could say the least of electronegativity is named first, but again it's not always the case. But what I would always do is if I get something like this, SO2, what you're going to want to name it is the first atom first and then the second atom second, no matter what. So this compound, SO2, is called sulfur dioxide because it's only got one sulfur atom and normally you want to put a prefix in front of everything except if the first atom only has one unit to it. So in the case of sulfur dioxide or carbon monoxide, since the first atom only has one, you don't put the mono. So the prefix indicates the number of each type of atom. If only one atom of a particular element is present, the prefix mono is usually made from the first element and the stem of the last element is used, is changed to IDE. So you can see nitrogen triiodide, not mono-nitrogen triiodide. Notice this one has four phosphorus in the first atom that's listed, so we call it tetraphosphorus and five oxygens pentawson. These names, these prefixes, you'll have to remember because they'll be used very commonly. So most of them are pretty straightforward. If there's any you have any trouble with remembering, maybe I can help you thinking of something that reminds me. So these are some very common covalent compounds. Notice they are all just like this, kind of just an individual particle, not a bunch, a bunch, a bunch of atoms stuck together. So notice this one only consists of three atoms, this one consists of five atoms, this one four, this one three, instead of a million. So if you want to try to name these compounds, maybe we can by now. What do you think about this one? What is this one called here? Silicone? Dioxide, right? Si is silicon, for those of you who didn't know. What about this one? Dioxygen? Penta-oxide. So notice you didn't say penta-oxide, it's because the A and the O don't go together, so whenever you have O's, like oxygens, cut off the A of any one of your prefixes. Okay? What about this one? Carbon tetrachloride, carbon tetrachloride. I have to fluoride, yeah. What about this one? This one? I don't know what we'll do. Okay, so I guess you guys kind of got it. Try some more. Try to write the names of the formulas of these things. What would be the formula of this thing? N-O. N-O, yeah. What about this one? N2-O4. N2-O4, N2-O4. And then this one? P2-O5. P2-O5. And then you could probably do that last one. Okay. We're going to go over this at the end of the next chapter, so I'm not going to go over it right now. On over covalent nionic compounds, we're actually going to talk about the natures of the particular bonds that come together to form these different peptides of compounds, covalent nionic. And in fact, all the other types of forces there are between particles. The main ones we're going to focus on are covalent bonds and nionic bonds. Okay, so how are atoms held together? So remember I showed you that picture earlier. Here's another demonstration of atoms being held together. So if we think of this as being a hydrogen atom and this being a carbon atom and this being some other type of atom, I don't know. What we can do is stick this carbon atom to this hydrogen atom by putting this little gray thing in between them. So I guess I... So this carbon atom and this hydrogen atom, right? We can stick this little gray thing here in between them. We call this thing a bond, okay? And that kind of sticks these two atoms together. Like that, okay? So you can kind of think of bonds like glue for atoms, okay? So they'll stick the two atoms together. So the more official definition is the force of attraction between any two atoms in a compound. So this attractive force overcomes the repulsion of the positively charged nuclei. So of course both of these nuclei here are positively charged, right? And when they come to really close together, positive and positive, they repel each other, right? But the bond can overcome that repulsion. Yeah, and interaction... And these bonds are just interactions of valence electrons between the two different two elements that are composed in the bond, okay? So we need to remember that valence electrons describe the chemical reactivity of any atom, okay? So it's only the valence electrons. That's why we had all that work on learning of valence electrons and noble gas configurations and all of that business. So it's going to come into play. Okay, so remember the noble gas configuration. If we look at an element, it's conventional or expanded electron. Configuration is for lithium, for example, it would be 1s2 to s1. And then the abbreviated form would be just showing the noble gas plus the valence electrons. Okay, so when you look at this, this really does emphasize how many valence electrons lithium has. But this, again, is kind of a cumbersome way to think about it when you're trying to stick atoms together, okay? So, but we do want to emphasize that it's the noble gas configuration and the valence electrons that really pushes all of this reactivity upon these different elements, okay? So instead of representing it like electron configurations like we've been known to do in the past, for the past two weeks we've been doing, let's go backwards and just look, since we know valence electrons now, we can just look at the periodic table and instead of writing all that electron configuration business, we'll just write dots that represent the valence electrons. And I kind of alluded to this last Wednesday. But what you'll find is that you can represent these atoms plus their valence electrons just through the atomic symbol and dots. So remember, we said Li, we could write as 1s2 to s1, or we could write it as helium to s1 or we could represent it much easier, more concisely, as lithium plus 1 valence electron, okay? So that dot represents its 1 valence electron. And what you'll see is if you look at this kind of makeshift periodic table here, we can see as you go up in group, right, you get one more valence electron, one more dot added to you until you get to the neon or argon configuration, that's the noble gas configuration, they're stable because they have 8 electrons around them. So this representation of these dots and the atomic symbol is known as the Lewis structure named after some guy named Lewis, okay, long since dead guy. As with only valence, since only valence electrons participate in bonding, it makes much easier to work with the octet rule, of course, especially if you're looking at, you know, I don't know, selenium or something like that, you don't want to have all these numbers and letters written out, you can just put the atom, put some dots around it and say you're finished, okay? So remember the number of dots directly corresponds to the number of valence gels that each atom has. So how do we represent these? You put the atomic symbol, you place one dot around the symbol until there are four dots, so this is the way a lot of people like to do it. So for lithium, so let's do boron, for example, we would start here, put one, two, like that, so we're going around the atom, carbon, one, two, then nitrogen would be one, two, three, four, and then pop, like that, okay? So it would have what we call a pair of electrons and three unpaired electrons, okay? So this is the way you want to start drawing your Lewis structures, just one dot, one dot, one dot, one dot, until you get to four, then you put up the next dot with the first one, okay? Remember, two electrons, it's okay to have two electrons, okay? Yeah, and each unpaired dot, so like in nitrogen here, it's got one, two, three unpaired dots, each unpaired dot is available to form a chemical bond, okay? So if we look here, lithium's only got one unpaired dot, and in fact, lithium only participates in ionic bonds, and this is really going to apply mostly to covalent bonds, so you're going to want to really concentrate on the non-metal up here, okay? You can still represent what's going on with ionic bonds, and we're going to a lot, but Lewis structures really help you with covalent a lot more. Again, notice the Lewis dot structure of all of the group one elements here are all the same. It would be the same for all the group two elements, same for all the group three elements, so on and so forth. And here's a good representation of that. So notice transition metals. Again, they're very strange, so we only really want to concentrate on the main group stuff. And again, all of the metals below the staircase and in the group one, group two, those will all be ionic bonders, and everything else up here will be covalent bonds. So remember an ionic bond is a transfer of one or more electrons from one atom to another, okay? So let's form an ionic bond. So one thing, let's draw, so we've drawn the Lewis structure of lithium here. Let's draw the Lewis structure of chlorine. So it would be Cl plus how many dots? One, two, three, four, five, six, seven dots, right? So how do we draw that? We go one, two, three, four, and then now we've got one on each side, so we're going to jump back to there. Five, six, seven. Notice we've got three paired electrons. These paired electrons, I'm going to start calling them lone pair electrons, okay? Because that's what I always call them. They're called lone pair electrons. So anything that's paired up like this. So lone pair electrons are not used in bonding. We've got one unpaired here and one unpaired here, right? What's going to happen? Those two are going to be used in bonding. So what actually happens with a ionic compound which lithium and chlorine will make is that the metal will lose its electron, right? So this one's one electron. This doesn't want to have it. So what will happen? It will transfer its electron there. So notice my arrow only has, and this is different than what they do in the book, okay? But my arrow only has one side to it, so it doesn't look like that, okay? A two-sided arrow, this really means to transfer two electrons, okay? A one-sided arrow means to transfer one electron, okay? So I want you guys to get familiar with this. Even though in the book it really states that two-sided arrows are okay to use for two electrons, but it will become really confusing if you guys try to do any more chemistry than this class, okay? So I want you to use one-sided arrows to show the transfer of one electron. Two-sided arrows will be the transfer of two electrons. In fact, in this class, you'll mostly be using the one-sided arrows, okay? So don't get fixated on them too much. So notice what happens here. We transfer, so we got electron transfer here, okay? From one atom to another. So when we do that, so we're kind of forming a rudimentary, if you will, chemical equation. If you remember, chemical equations from what we talked about last, well, Monday, if you were in my lab, but in any of the labs you probably talked about chemical equations last week. This is a rudimentary one, and again, we're not really anti-chemical equations yet, but what's going to happen here is when that lithium transfers its electrons there, it's going to become positively charged, of course, and now we're going to have a chlorine with a full octet. And since it's gained that electron, it becomes negatively charged. So now we have two ions. Then what happens? Well, these two ions, now they're one of them likes to have a charge, remember we talked about, so they want to stick together. That's called an ionic bond. So does that make sense? So there's kind of these three steps to it. The transfer of electrons, forming of the ions, and then sticking together of the two oppositely charged ions. That's how you make an ionic bond. If you go back to the practice test that I gave you, number five on the practice test is essentially doing this thing with, I don't remember what it was, oxygen and iodine I think, but actually in that case we're talking about covalent bonds. So this is different, so I do have a couple more minutes there. This is different than forming a covalent bond, so notice this is the transfer of electrons. A covalent bond is due to the sharing of electrons. So let's draw, hopefully we get it. Okay, so let's just go over again what I talked about in ionic bonds, and once we finish the ionic bonds then you guys can go. So ion formation takes place by electron transfer. You can see the transfer of the electrons so the sodium plus the electron, the electron then goes to chlorine, forming the full octet here, so both of them gain the noble gas configuration and then they stick together. Again here's another representation of it, but notice you don't have just two atoms sticking together. You've got this whole conglomerate here. Okay, you see that? So ionic bonds form between simple ions when representative elements lose valence electrons, the electrons are gained by other representative non-MLF atoms. Both atoms are changed into ions with noble gas configurations and the resulting ions are attracted to each other sticking together to form this conglomerate. And then you can see the arrangement and the crystal if you want. And then you can see, remember, ionic charges, please. So we'll stop there. We'll stop there for right now. So remember all of that stuff. It's kind of the review from last Wednesday and next time we'll start with the electron negativity and do some in-bonding. That'll probably take up most of it. Do not slime the sheets. If you need to slime the sheets, they're up here. If you didn't get your quiz to, come up and get it. There's a problem with quiz two.