 I'm going to try recording this with a different device and see if this video looks any better than what the last one did. Alright, we're going to try to combine together everything. The Lewis structures, the shapes of the molecules all into a single video here and see if it's a little better. Alright, so in covalent bonding, it's a lot different than ionic. In covalent bonding, when we do Lewis structures and we write formulas, I tell you what elements I want you to use. So I say calcium and chlorine. And you do the Lewis structures, you write everything out, you tell me what the formula is going to be and so on. So you start with the Lewis structures. Calcium is in group two and because that has two valence electrons, we represent those with two dots around the calcium symbol. Now we have chlorine. Chlorine is in group 17. Group 17 has seven valence electrons and we draw its Lewis structure like this to show the seven valence electrons. Our goal is to move electrons from the metal to the non-metal to make everything stable. This is our metal so we can move an electron there and that would make that chlorine nice and stable, but it wouldn't be good for the calcium. The calcium goes from two valence electrons to one and one is a more unstable condition. It's not as good for it as having two. So we have to find another home for that lone electron there. So the way we do that is by drawing in another chlorine atom like so. And that gives us a second place to put that second electron. Now the metal is stable because it lost all the electrons it needs to lose. Both the non-metal atoms are stable because they both have eight electrons. It took one calcium and two chlorines to do it, so the formula would be CaCl2. Again, that's the way we approach it in ionic bonding and a lot of you did well with that so I don't think it's too big of an issue. As we move to covalent bonding, however, we have to recognize that there's not going to be any atom that wants to lose electrons. We're dealing with atoms that have a high electron affinity, high ionization energy, so transferring electrons is just not feasible. They're going to share electrons instead. And the big issue when you do that is there's more than one way to share electrons. You know, when you take, for example, nitrogen and oxygen, there's half a dozen or more different formulas that would be correct that would show the right number of electrons being shared and you'd have seven, eight different formulas and we don't want that. So what we do instead with the covalent bonding is I give you the formula. I say, for example, SiO2 and I want you to write the Lewis structure to show me how it bonds. I want you to show me the stability and then we take it a step further and we draw something called a structural formula that shows the shape and how it's all put together. Again, it all starts with the Lewis structures. So this time we already know how many atoms to draw. So we know I need one silicon and silicon is in group 14, just like carbon is, so it has four valence electrons. Oxygen is in group 16, so it has six valence electrons and there are two of them. Now remember, you can put the single dots wherever you want to. You can put the paired dots wherever you want to. You can twist that Lewis structure, turn that Lewis structure to get the electrons where they need to be. And our electrons should be somewhere close to our center atom. Now, center atom, we didn't have to worry about that with the ionic bonding because there's no molecules being produced, but we do have to worry about that here. You have to decide, do I bond the silicon and that oxygen to this oxygen down here? Do I bond them to the oxygen over here? Do they both bond to the silicon? We have to figure out what the center atom is and it's pretty straightforward. The center atom is the one that has the few, has the most single dots, the fewest pairs, the most single dots. This one has no pairs and it has four single dots. This one has two pairs and two single dots, two pairs, two single dots. The one with the most single dots, the silicon is the center atom and everything just connects to that. So what we do is we just start pairing up singles, circling them this time because we want to say, we're going to share these two electrons. Silicon is going to share that one electron with the oxygen and oxygen is going to share that one electron with the silicon all the way around. Silicon shares that one with the oxygen and oxygen shares that one with the silicon. It's a sharing idea. That's why we don't draw arrows. Again, nothing's going to lose electrons, nothing's going to take them. They're just going to share these electrons all the way around. Now, that's all the further I want to go with the Lewis structures. Because that's all I need to know to draw a structural formula. This is all based on something called Vesper-Velenshell electron pair repulsion theory. And that theory tells us that there's some very predictable shapes that we can get. Based on the number of atoms, whether there's unshared electrons, we can use those two things to figure out what the shape of our molecule is going to end up being. Now, the simplest molecules are two atom molecules. Whenever there's a two atom molecule, we don't even have to worry about the unshared electron thing. That's why it's so simple. There's only one shape a two atom molecule can be and that is linear. No matter how you orient it, there's only one way to put them together and that's in a straight line. Element A will bond to element B like so. The only thing we really have to worry about is whether this is a single bond, a double bond, or a triple bond. And that's where drawing the Lewis structures really come in handy. I'll show you what I mean. If I have fluorine, F2, I've got two atoms of fluorine so I already know it's going to be a linear molecule. But I've got to figure out if it's single, double, or triple bonded. Fluorine's in group 17 so it has seven valence electrons, three pairs, and a single. And again, it's F2 so I know I have two of them. And again, when I do this Lewis structure stuff, I want to just connect the single dots. When I make that one little loop there, I'm done. I have used up all the single dots that I have. I have a single loop here which indicates a single bond. So again, two atom molecules are going to be linear. When I connect these two fluorines together, I'll do it with just a single dash to represent that single loop, that single bond that's being produced there, that single pair of electrons that's being shared. If I have oxygen on the other hand, O2, it's a different story. Oxygen's in group 16 and has six valence electrons, like so. And again, I want to pair up these singles like that. I've got two loops this time. I've got two connections, two pairs of electrons being shared, two covalent bonds, what we call a double bond. And again, it's two atoms so I know it's linear but this time when I connect them together, I'm going to do it with a double dash, double dash representing that double bond. Two loops, two dashes. Finally, nitrogen and two. Nitrogen has five valence electrons, one pair and three singles, like so. And again, I just pair up the singles. I've got three pairs I have to share between them. I have three connections, three covalent bonds, that's a triple bond. And I show the triple bond with three dashes. So again, two atom molecules are all linear. Every single one of those is a linear molecule. It's just a question of whether it's single bonded, double bonded, or triple bonded. Again, simple, two atom linear idea, just worry about the bonding. When we get to three atoms, it gets more complicated because there's two different shapes they can form. One with no unshared electrons, one that has unshared electrons. If there are none, it'll be linear. You'll take your central atom and you'll put it right in the middle. And you'll connect your other ones on the left and right so that all the atoms make a straight line. The three atom one is going to be bent. A pair of unshared electrons will push those atoms that are on the side downward. So we'll get a bent shape. Let's take a look at how this works. The linear one will look at carbon dioxide, CO2. I've got one carbon and two oxygens. Carbon's got four valence electrons, represented by four single dots. And oxygen has got six. So two single dots. Again, the one with the most single dots has to be your center atom. So both oxygens will connect up to the carbon. Again, we just start looping around dots. So this oxygen and carbon will have a double bond. That oxygen and that carbon will have a double bond. So we have a pair of double bonds in here. Now, three atoms, one, two, three. We know it's either going to be linear or bent. Depends on if there's unshared electrons. So here's how we figure that out. We look only at the middle atom like that. And we ask ourselves, did we circle all the dots? And if you look at the drawing, we have circled all the dots that the carbon has. So there are no electrons that have been unshared. There are no unshared electrons. Every electron is part of a covalent bond. Every electron has been shared. No unshared electrons is linear. So again, we'll put our middle atom in the middle. We'll flank it with the oxygens on the left and right. And again, since there's two loops here, we'll do a double bond. Two loops here, we'll do a double bond. Makes it nice and linear. Now, let's take a look at one that's not. Let's take a look at water. H2O. Two hydrogens, one oxygen. Again, oxygen has six valence electrons. Two pairs, two singles. Hydrogen only has one. And we know that we've got two of them. Again, circle the singles like so. Look at the middle atom. And again, only the middle atom. It's the only one we ever look at, the one that everything else connects to. That's the center atom. You see these electrons here that we did not circle? Those are unshared electrons. Now, whenever you have a three atom molecule with unshared electrons, it's going to be bent. And again, we just push those ones that are on the end down to the side. Here's what's going on. Those unshared electrons will be sitting right there. And they will be pushing these hydrogens down and away from it. So whenever we have these electrons on the middle atom that we did not circle, it's going to change its shape. In this case, it's going to bend it. All right. Same kind of thing happens with your four atom molecules. There's going to be two types for that. There's four atoms, no unshared. And there's four atoms with unshared. No unshared is going to be trigonal planar, means flat triangle. You put your middle atom in, and then you connect the other ones around so that they represent the corners of the triangle with your center atom sitting right in the middle of the triangle. That's trigonal planar. If you do have unshared electrons, it's pure middle. For that, you draw your center atom, and then everything gets pushed down by, you guessed it, some unshared electrons. So let's take a look at how those work. A good example of something that makes a trigonal planar is boron. Boron only has three valence electrons, and it is an exception to the octet rule. So it does not need eight to be stable. Bcl3 boron trichloride, one boron, three chlorines. Boron has three valence electrons because it's in group 13. There are three chlorines, and each chlorine has seven valence electrons. Draw those in. And remember, orient them so that they are easily accessible. Get the single dots close to the boron so that your loops are nice and neat. And start pairing up all those singles. Everybody needs a partner. There you go. Single bonds all around, so we know we're going to draw nothing but single bonds in our final structural formula. We look only at the center atom. Doesn't matter what's going on out here, it's only what's going on in here. And again, we circled all of boron's dots. So there are no unshared electrons, and that's what makes this trigonal planar. You put the center atom in the center, and then you attach the chlorines around so the chlorines would make up the points of the triangle with the boron sitting smack dab in the middle. Now let's take a look at the different one. Let's take a look at one that will end up being pure middle. Let's do the most common one, ammonia NH3. One nitrogen. Nitrogen has five valence electrons, one pair and three singles. Again, maximum number of singles. Nitrogen only has one valence electron, and again you can orient that so that your circles are easy to draw. Nice and neat. And again, when we look at our covalent bonding and we look at our middle atom and we look at its electrons, we have some electrons we did not circle. Those are unshared electrons. And whenever you have unshared electrons and a full atom molecule that's going to be bent, we draw in the nitrogen and we just push everything downward. These are like the legs on a tripod. This one's coming out of the paper towards us. These two are going back into the paper away from us, so it makes a little tripod, like the legs of a tripod. The final one, the five atom molecule. Again, we don't have to worry about unshared electrons in the five atom molecule, so it makes it a little simpler. That is tetrahedral. How about if I move that down where you can see it? Five atom molecule is tetrahedral. And to draw tetrahedron, we put our atom in the middle, and we go ahead and draw that thing, the trigonal planar. But instead of having unshared electrons on the top, we have another atom, like so. Let's do one example of that. Let's do methane, CH4. We have one carbon with its four valence electrons. And we have four hydrogens, like so. That's how it looks in the bonding. And again, we do this stuff so we can understand octetrable. Carbon has two, four, six, eight electrons around it. That makes it stable. Hydrogen's got two. That makes it stable. It allows us to model the bonding. And so much of our standard is about modeling, being able to do a visual depiction of it to help us understand it. And that is just a model of covalent bonding. Is it exactly the same as covalent bonding? No. But it helps us understand it a little better how it all works. Carbons are center atom, so we draw it in first. And again, do that little tetrahedral shape, which is a pyramidal with the hydrogen on top, instead of the unpaired electrons. Now the next thing we have to do is polarity. And to do polarity, you have to be able to draw the structures. Plain and simple. And you'll notice I didn't put any other valence electrons on my structures, like on this one here. I didn't put the extra dots here around the chlorine. We know where they go. Every chlorine has six dots around it. So we could put them in there if we wanted them. But honestly, we don't need them at this point. They would just get in the way. We've got to put in some plus and minus signs. And if we fill up our picture with all kinds of dots, and I go to start throwing plus and minus signs in there, you're going to have a hard time keeping track of where they all are. That's why I don't put the valence electrons on any of my structural formulas. I hope that was clearer than the other two videos were there. It is a bit longer, but it's both videos in one. Everything that I went through today as I reviewed. So hopefully that'll help. Polarity's next. And if you can do this, you can do polarity.