 Last thing, we went over a structure of ions or ionic compounds. So if you look at these, right, this is a simple ionic compound, okay? And this is the crystal structure. Well, it's not really the crystal structure, it's the representation of it, right? But it indicates the crystal structure of something that would be like sodium nitrate or something like that. So it's a simple ion plus a polyatomic ion. So a monoatomic ion plus a polyatomic ion. And this is, of course, monoatomic and monoatomic. So this would be something like sodium chloride, like what's the picture that's there? So I guess today we're really going to be talking about bonding, covalent bonding, as opposed to ionic bonding. How we'll be drawing arrows showing electrons moving and what not. Okay, just like we did with the ionic bonding. The one thing I want you to remember, again, is the charges are oxidation states. I might call them oxidation states and charge are the same thing. But with covalent bonds, right, we're going to be worried mostly about the non-metal. So starting with carbon all the way over. So if we look at the periodic table up here, we can think of anything like. So the first thing we're going to need to know about is electronegativity. And what does that mean? Well, electronegativity is the measure of the ability of an atom to attract electrons in a chemical bond. So what does that mean? It means that when you make a, remember, all bonds are made because of electron transfer or electron sharing. What happens is usually unless you have the same type of atom, like a bromine and a bromine bonding together, what will happen is that one of them will have more of a pull on the electrons than the other one does within that bond. Okay, so it'll be like an unequal sharing. So you can imagine two guys playing tug of war. One guy's a really big muscular, huge guy. The other guy's like a little tiny, skinny, whatever guy. So the big guy will have more of a pull on the rope than the little guy will. That's like electronegativity. So the big guy would be the more electronegative atom. Ironically, in this case, they're the smaller atoms. So the smaller atoms have more of a pull than the bigger atoms. And in fact, what you find is electronegativity increases going like this across the periodic table. This is going to be a handout you'll have from now on on every test because it's such an important concept when you're talking about bonding. So I want you to become familiar with this. Of course, don't put it to memory or anything like that because you'll have it on the test. But you're going to be using this quite frequently for the rest of the class. Trying to determine which atoms hold more of the electrons closer to them. So elements with high electronegativity have greater ability to attract electrons than elements with low electronegativity. So this gives rise to what we call bond polarity. Polar things have positive and negative charges. And we'll get more and more into bond polarity in a little bit. So let's look at covalent bonding. So covalent compounds form from the... So you've got to remember all that electronegativity stuff from now on. So everything we talk about you always want to keep in mind electronegativity. So covalent compounds form from the sharing of electrons between one atom and another. They're typically formed when two non-glyols react. The bond results from the electrostatic attraction between the nucleus of one atom and the electrons of the others. And remember, covalent compounds consist of discrete molecules like this, rather than massive 3D crystal structures like this. Everybody can see the difference, hopefully. Some common covalent compounds that you might be familiar with or will be after this slide are water, H2O, methane, CH4, ammonia, NH3, and carbon dioxide, CO2. We're going to be learning how to build those molecules and their bonding codes. So let's look at covalent bonding. The simplest molecule that has a covalent bond in it is H2. So hydrogen gas. Hydrogen is a diatomic gas. Remember, if we look at hydrogen valence electrons, hydrogen only has one valence electron. But it needs another one to completely fill its orbital. It wants to have two electrons in its orbital. So what it can do, it can either gain one electron and be H minus. So let's look at this. Remember, the first energy level, two electrons fill the ship, and then it's got the noble gas configuration of neon. So we've learned that hydrogen can do one of two things already. It can take this electron and give it up, right? And that would effectively make it not have any electrons, right? So it's kind of a pseudo-noble gas configuration. Or it can gain an electron from some other source, like that, right? And then it's got the helium configuration. So both of these are noble gas configurations, if you would. The top one's a pseudo-noble gas configuration. But what typically happens when you've got two hydrogen atoms together, so you've got one hydrogen atom here, that wants to have an electron, and another hydrogen atom here, that wants to have an electron. So what they'll do is instead of one of them giving it up to the other one, what'll happen is they'll share those two electrons in between them. So they each want to have an electron, so if they share it, then they can say, well, we both have two electrons in. And the proper way to describe a covalent bond, so remember with ionic bonds we've described it like, with the electron just going away or going to another atom, what we would describe with a covalent bond. We want to describe the electron going there, the other electron going like that. Remember the one-sided arrows show the movement of one electron. So when we do that, we get... So if you notice they're both stuck together, they both have that filled orbital. Not the way that chemists usually like to write covalent bonds. They don't like to stick two dots in between them. So what we've done by convention is given a bond a different symbol, and that's just a line in between the two atoms that the bond has made between them. So anytime you see one of these lines, you know it's two electrons being shared in between those two atoms. So this is typically what hydrogen does. It doesn't give one of its electrons, so it doesn't do this. It gives its electron to there, go here, go here, and then these two combine. You can do the same thing with any two non-metals. Non-metals and non-metals do this when they react with each other. Let's try another non-metal, non-metal interaction. Instead of H here, H and H, let's do H and F. Because F is also a non-metal, right? Fluorines and non-metals, does everybody agree with me? How many valence electrons does fluorine have? Seven valence electrons. So let's draw fluorine. Everybody draw the Lewis structure of fluorine. So let's draw it together. One, remember, one, two, three, four. How many electrons does fluorine want to fill up its valence show? One. One, right? Because it wants eight electrons, right? And how many electrons does hydrogen want to fill up its valence show? Seven. No, no. How many electrons up here? All right. Just one, just one, right? Because the first shell, two electrons and three also eight electrons. So let's try this again. How many electrons does hydrogen want to fill up its valence show? One. And how many does fluorine need? One. One. So what do you think they're going to do? They're going to do the same thing that hydrogen and hydrogen do. They're going to take this electron, put it like that. And this electron, right? Because there's not two on this side. So it's going to go like that. And then you can think of it going like this. Having the two dots. When we've got the two dots in between the two atoms, we do a different convention. We say HF. But we still need to put all the dots, the rest of the dots around fluorine. Notice how many electrons does fluorine now have around it? Eight. Eight. How do we count that? We go one, two, three, four, five, six, seven, eight. Okay? There's two in that bond, remember? How many electrons does hydrogen have around it here? Two. Two. How do we count that? One, two. Okay? Does everybody get that? So there's two kind of different types of electrons now within this structure, right? There's these electrons that are not bonding electrons or non-bonding electrons, we call them. Non-bonding or lone pair, different names for the same thing. And these electrons, that's what we call them. Bonding. Bonding electrons, then bonding. Does everybody understand what we've done here? So the shared, let's just look at this one again. The shared electron pair is called a covalent bond. This is a covalent bond. Remember, bonds are nothing more than electrons. They're kind of the glue that sticks things together. So each hydrogen atom, in this example, now has two electrons around it. Remember, a hydrogen atom with two electrons is now considered having the helium configuration. Here you can see a picture of what it might look like if we were a little guy or gal small enough to be able to watch two hydrogen atoms interact, okay? They might come close. You see their electrons coming closer and closer. The electrons are influenced by the positive charge of the other nucleus. This guy is influenced by this thing, so they kind of want to stick to it. They stick together, form the bond. Now they both have the electronic structure of helium. So now you're going to, you don't need to necessarily memorize these because you can do these covalent bonds now. You can figure out which ones would be diatomic gases. But there's not a very big list here. So what I would recommend doing is just memorizing them. So the diatomic elements are all the halogens, oxygen, and nitrogen. Group 7, yeah, all of them. But you're never going to see that. Nitrogen, nitrogen, and oxygen. Not sulfur. No. Just the ones that I'm pointing out. So that little inverted L and then the hydrogen. Those are the ones that I'm talking about. Just this little inverted L. So, of course, compounds containing covalent bonds are known as covalent compounds. These diatomic elements, they share their electrons completely. So 50% of, if we look at hydrogen, this gas here, those two atoms share their electrons completely. Why is that? Because they're the same strength when they're playing tug-of-war. They both want the electrons more, but they can't pull the rope more because they're the same half. Remember, the same atoms have all the same properties. So let's look at these guys over here. Hf. Hf. Are those the same atoms there? No. No way, right? So what's going to be happening? There's going to be a tug-of-war here and somebody's going to win. Always. Always when you have different atoms, somebody's going to win. So what we say is the more electronegative atom pulls the electrons closer to it. Remember what we said about electronegativity. It increases going this way, right? Increases going up this way and across this way, okay? So, and chlorine is going to be more electronegative or less electronegative than pay, way, way more. In fact, chlorine is the most electronegative atom on the periodic table. That reacts with things, okay? The noble gases don't react, so we don't include them, okay? Chlorine's the most electronegative atom. So what do you think when we're looking at this little chord in between them? More of it is going to be on the chlorine here. Remember, these things are sharing electrons now. So it's not like an ionic bond where the F takes the electrons and becomes a minus charge, okay? So notice there's no minus charge here on the F, okay? It's because they're sharing those electrons. Even though the F is stronger than the H, it's still sharing with it. But since electrons are negatively charged and the fluorine has more of the electron density around it, it is slightly negatively charged relative to the hydrogen atom, okay? So we're going to use this symbol here, delta, okay? So Greek will be delta plus and delta minus. These symbols here indicate a partial negative and a partial positive charge, okay? So if we look here, these guys have the same strength, right? So neither one of them can pull the electrons more or less, okay? So there's no partial negative partial positive here. So as we look over here, since we know the electronegativity trend, we can say the fluorine's going to have more of those electrons around it. So it's going to be delta minus. H, of course, is going to be delta plus. Because it's got less of that electronegativity. So what you would find is that when you've got, when I'm talking about hydrogen fluoride, which is what this is, or hydrofluoric acid, whatever you want to call it, when you're talking about this substance, right? I'm not only talking about one molecule of it, okay? When I have it in a bottle, oh, I don't have a bottle, say this was HF. It wouldn't stay in this glass, but say this is HF, right? There's like billions and billions and billions and billions of billions of HF molecules in this thing, right? If this was HF. It's water, so we could talk about a different covalent, polar covalent compound, but this is a more simple one. So just imagine, so when we're talking about HF, we're not really only concerned with the one molecule of it. We're concerned with a bunch of it, of course, because we don't work with just one molecule. Okay, so if we looked at two HF molecules kind of interacting together, what you would find is that when they're charged, when they're together, when they're interacting with each other, they're going to align in this sort of fashion so they can have their delta minuses and delta pluses align with each other. So remember, minus likes plus, so it kind of does something like that. So there's kind of this slight interaction between them. So what you can think of these molecules, I know this is going to sound weird, but you can think of these molecules as little magnets. So remember, like you got a magnet, probably when you were a kid, or maybe still now, you tried to stick magnets together, and one side they would stick together, and if you flipped it over, the other side they would kind of repel or kind of float on top of each other. This is like these little magnets. These magnets are not very strong. This is the strongest of the not very strong magnets. So if we look at these guys, these are super-duper. These are the ones where you take the things like way apart and they'll like smash in the air. You know the magnets that really want to get together. So that's like those guys. These guys are like the refrigerator magnets. You know what I'm saying? They'll fall off if you put a piece of paper underneath them. So that's what you want to think about. This is the strongest one of those refrigerator magnets. Again, I want to emphasize that these non-metal to non-metal, they share their electrons so they can attain the noble gas configuration. Like flooring in this case wants to attain the neon configuration. So let's think about covalent bonding with molecules with more than one bond. So let's look at, instead of looking at the halogens and hydrogen, let's look, for example, at water or oxygen in particular first. Okay, draw everybody if you could, draw the Lewis structure of the oxygen atom. So let's do that all together. How many valence electrons is without looking up there? Six. So how do we do that? One, two, three, four, five, six. How many valence electrons does oxygen? One. Six. Well, eight, right? But six plus two is eight, yeah. So, yeah, it wants two more, I guess, right? Two more, okay? We look at hydrogen. How many valence electrons does hydrogen add to that? One. One, okay? So if I get one hydrogen and one oxygen together, would that fill the valence shell of hydrogen? No. Yes, it will fill the valence shell of hydrogen. Or why? Because hydrogen only wants to have two electrons, right? But will it fill the valence shell of oxygen? No. How many more electrons would you need to fill the valence shell of oxygen? One more. One more, okay. Yeah, so this is why water's molecular formula is H2O. Because it needs that other hydrogen with this other electron to make it bond, right? So let's make the covalent bond now. So if we make that one, right? And then we make this one. It helps you to do this. Remember on the test, I want to see bonds drawn as lines, okay? But if it helps you to do this intermediate step, please do it, okay? Like that. And then, of course, that's going to get a long pair of electrons. Everybody see that? Those lines are just drawn unevenly because I didn't have enough room, okay? So don't get too caught up. Just don't worry about it, you know? If you think about it that way, you're going to get all mixed up. Just do it. Just do it this way, okay? However, the structure and the Lewis structure are totally different, okay? The structure is going to look much different, okay? We're not into structures, so don't try to confuse yourself. So does everybody understand how water is bonded together? Okay, let's try methane. So carbon, how many valence electrons does carbon have? Four. Four, right? Everyone, four. True that. How many more does it need? Four. Four, right? So if I were to combine carbon and hydrogen together, how many hydrogen atoms would I need? Four. Four. Good job, guys. The stuff that comes out of the Bunsen burner? CH4, methane. Why is it molecular formula like that? Because now I'm going to skip that intermediate step. We got that similar thing without looking up there. Okay, so remember how we were saying that every atom has a different electromegativity, and if these atoms were different, then they would have a delta plus and a delta minus associated with them. Okay, so let's look at... If you could go back to your electromegativity table in the slides, I don't want to keep looking back and forth. But you can... There's actually numbers that are associated with each of the individual atoms, okay? So can anybody tell me what... I'm going to erase this whole thing so I don't have to run. Can anybody tell me the electromegativity number of hydrogen? 2.1. 3.5. Which one of these is bigger? Which number is bigger? 3.5. Okay, so the bigger of the numbers has the higher electromegativity. The delta minus going to be on the hydrogen or on the oxygen. Oxygen. Okay, so we're going to put delta minus there. And where's the delta plus going to be? On the H. Okay, so it's going to be on both this H and on this H. So what you can start to think of now, again, like I was saying, these refrigerator magnets sticking together, sticking together. When you've got a difference in electromegativity, like you do here, so if you wanted to think about it, you could say delta, this just means change of, okay? Electromegativity. We could say 3.5. The units for electromegativity are called device. Okay, it's named after some guy, device who's long since dead, okay? But you can just say D, okay? 3.5 D minus 2.1 D. When we do that calculation, we find that we get 1.4 D. Okay, 1.4 device. So that's the difference in electromegativity between those two atoms. Does that make sense? Yes, it does. Subtract one number from another. Does that make sense? Okay, so this is the difference in electromegativity between these two atoms here. Okay? So when we say that we've got a difference, what we know about that bond is that it's polar, okay? Polar just means it's like one of these little magnets, okay? So this is going to have a pop of delta positive and delta negative charge, so we call this a polar covalent bond. A nonpolar covalent bond would be something that is between a bond that's between two atoms that have the same electromegativity, okay? So like if we look at two hydrogens, that's a great example. What's the electromegativity of hydrogen? 2.1 divide. Yeah, the hydrogen? 2.1 divide. So the difference in electromegativity here is going to be equal to zero divide, right? So we say this is a polar covalent bond. Is that what we say? No. No, it's actually a nonpolar covalent bond. What you'll find is that polar covalent bonds, because they like to, because the molecules themselves, okay, let me start over. If a molecule has polar covalent bonds in it, the molecules remember kind of like to stick to each other, okay? So what you find is that when you've got polar covalent bonds in a molecule or a polar molecule, okay? So overall, this molecule is going to be considered polar water here. And overall, this molecule is going to be considered nonpolar. When you have polarity associated with a molecule, it increases the boiling point, increases the melting point. Because those things like to stick together more. And remember, boiling point is like when the liquids kind of break away and fly away, okay? So if we're just holding on to the other one more, right? It's not going to allow it to break away as easily. Let's talk about multiple bonds, okay? So we've already talked about single bonds. Those are bonds in between two atoms that are just two shared electrons, okay? Just like what we have here or what we have in the example in method here. Let's look at a molecule that contains a double bond, okay? So oxygen, right? Oxygen has how many valence electrons? Six. Six, right? One, two, three, four, five. Let's draw another oxygen. Valence electrons is at oxygen. Six. Two, three, four, five. So we have two oxygens with six valence electrons. So how many valence electrons more do each of these oxygens want? Two, right? Because they want to have eight. But one of the oxygens isn't going to give up its electrons to the other oxygens. But what they can do is share four electrons, okay? And we share these two and these two, right? Like this. Like that. What you'll find is that the new structure will be when you have more than one pair of shared electrons in between the same two atoms, you're going to form a multiple bond, okay? So this is a multiple bond. This one in particular is called a double bond of more than two electrons. Remember, you can only share multiples of two, right? It's actually going to be two, four, or six. So you're not going to have quadruple bonds. Again, you're really ready to talk about a structure. I'm not ready to talk about it, okay? So we'll just talk about bonding for right now. Structure is a whole different ball game. Okay? So does everybody see how to take this to make this? If we look here, how many lone pair electrons does this oxygen have? Four. Four lone pairs, right? So four lone pair electrons but two lone pairs, okay? Yeah, so it's got four lone pair electrons but two lone pairs, okay? Yeah, so that was kind of... But anyway, so this has got four lone pair electrons and how many bonding electrons? Four. It has to have four, right? Why does it have to have four and four? Because it has to have eight. It has to have eight. So does it have two bonding electrons? Four. No, it's got four, right? Is that right? Or am I tripping or what's going on here? Everybody's looking at me like this is some crazy person, right? We just talked about this for the last 20 minutes, right? So each one of those lines represents how many electrons? Two electrons. So since there's two lines in between those two atoms, how many electrons does that represent? Four. Four, okay? And remember, these little dots here, those represent each one electron, right? Right. Okay, so if we look at this oxygen atom, how many electrons does it have around it? Eight. Eight. It's got to have eight. It's got to have eight, right? Why does it have to have eight? Because it wants to fill its shell, right? Just like we talked about all last week, you know, on Monday, you know, the week before that, you know, all this stuff, okay? Okay, let's try this again. How many electrons does that oxygen atom have around it? Eight. Eight. Why does it have to have eight? Because it wants to fill its shell, right? It wants to fill its shell, okay? Okay, so let's try hydrogen. I'll let you try hydrogen on your own. Just set it up for you. So nitrogen, remember, is one of the diatomic gases that you're going to have to know, okay? So, if it's a diatomic gas, it only means it's got two nitrogens, right? It's just two diatomic gases, okay? So, nitrogen, how many valence electrons does nitrogen have around it? Five. Two. Three. Four. Five. It does not matter. You can start with whatever you want, okay? You can start with whatever you want. It does not matter. Let's start with the bottom of the electron, just so we can say, at the top, because I'm very tall and the board is down here usually compared to me, you know what I'm saying? So that's why I start with the top. Maybe somebody who is very small might start with the bottom, you know what I'm saying? But that's just it. And since I'm right-handed, a lot of times I'll start with the right side, you know, but if I'm left-handed, you know, I think my wife goes the other way around. But it does not matter. It does not matter. Do not get caught up in these, like, details. You're going to make yourselves go crazy. Okay? It's just dot, dot, dot, dot, dot. Five dots around a letter, okay? Let's do the other nitrogen atom. Start from up here this time. Five. Okay, so this guy's got his long pair on the bottom. This one's got his long pair on the top. It does not matter, okay? Okay, so from those two atoms, I'd like you to draw a nitrogen molecule. Okay, in your own time. If you've already got it, congratulations. What I can tell you is that I'll have a triple bond in between those two nitrogen atoms. You can come and ask me after class if you're concerned about your structure, okay? At the end of the day, when you make N2, though, how many electrons will each of those nitrogen atoms have around them? Eight. Eight. Why will they have eight again? They want to fill their shells. Why do they want to fill their shells? To be complete. To be complete, like, what? Noble gas. Very stable, right? Noble gases are very stable, okay? If that's what they want to do. And in fact, in a different sort of type of stability, what you'll find is that when you've got a single bond between two atoms, this is very easy to break, those single bonds, okay? If you've got a double bond, it's harder to break that. Okay, it takes more energy. Triple bond, very, very hard to break. Okay, so nitrogen, if you know about the atmosphere, anything about the atmosphere, nitrogen is 78% of the atmosphere. The reason why is because it's very, very stable, because it's got this triple bond. So nothing reacts with it except plants, right? Plants have this molecule in them called rubisco that can take nitrogen and twist it in such a way to break that stable triple bond, okay? But that's the reason why you've got a lot of nitrogen in the atmosphere, for those of you who are. So, already essentially gone over this, so we're going to go over this quite fast. Lewis structure guidelines, so these molecules that we've been building today, these are also called Lewis structures, okay? But they're Lewis structures of molecules, not atoms. So you're going to use the molecular formula to write the skeletal structure of the compound. Put the least electronegative atom in the central position. So if we've got carbon dioxide, oxygen's more electronegative than carbon, so carbon's going to be in the middle between the two oxygens. So you can draw the skeletal structure, that's just drawing an O, then a line, then a C, then a line, then enough, because you know they all have to be bonded together, so they have to be sharing at least two electrons. Find the electronegativity here, put that one in the middle, okay? Determine the number of valence electrons, and you just go through this one here. I don't want to keep doing this. Okay, polyatomics. Here's some good representations of the things that we were doing. Notice they show double-sided arrows, which again, only in this class you'll see, but from now on I want you to write in my single-sided arrows if there are only one electron, okay? And there's some more examples if you need help. But there's nitrogen right there, okay? Nitrogen 2. See, we talked about the bond energy. Triple bond is greater in energy than the double bond. Guys and gals has greater in energy than the single bond. Ready to get out. The distance separating the two nuclei, when you get stronger and stronger, when you go from a single to a double to a triple, the bond length decreases, okay? So single bonds are very long. Double bonds are not as long. Triple bonds are very close to each other. I'd like you to try to draw the Lewis structure of these compounds, one, two, three, and six. Four and five we're going to wait for next time. We're going to talk about drawing Lewis structures of polyatomic ions next time, and then we'll hit this for 30 minutes. Okay? Thanks a lot, guys.