Lewis diagrams for molecules | Chemistry | Khan Academy

Let's draw LS draw structures for certain molecules. It's a lot of fun to do that.

Okay, now the first thing we need to do to draw these structures is to identify the number of valence electrons. Okay, and we've talked about these valence electrons in our previous videos, but this is such a central concept, so just let's quickly recap. Valence electrons represent the number of electrons in the outermost shell.

Okay, and if we bring up our periodic table, then the valence electrons for all the elements belonging to a particular group is given this way. So the first group will have the elements of the first group will have one valence electron. The elements of the second group will have two electrons in its outermost shell, two valence electrons, and so on and so forth.

Um, when it comes to group 18, helium is an exception; it has two valence electrons, but the rest of them will have eight valence electrons—eight electrons in its outermost shell. Okay, and just to get a sneak peek of why this is the case, well, here are some structures of some of the elements. You can see hydrogen has one electron in its outermost shell, lithium has one electron in its outermost shell, sodium, which has a total of 11 electrons, but one electron in its outermost shell—same thing will continue.

Okay, but what about the elements of these groups? What about their valence electrons? Well, we don't have to worry about them because we're drawing structures for molecules that are formed by covalent bonds, meaning bonds formed by sharing of electrons. Now, metals, which I drew in red, okay, they usually do not participate in covalent bonds.

Okay, it's only the non-metals that usually do that, and the non-metals are in green, and you can see all the valence electrons for the non-metals—we already have them, so we only got to worry about them and not worry too much about the metals.

Okay, all right with that recap, let's now draw the Lewis structure for hydrogen. So, the first step is how many valence electrons does hydrogen have? It has one valence electron. So let's write that one valence electron. And now hydrogen would love to have two electrons because it'll make it stable in its outermost shell.

And how do I know that? Because if I look at helium, look, helium has two electrons, and it doesn't bond with anybody because helium is stable. All noble gases are stable—that's how I know hydrogen would love to have two electrons in its outermost shell.

Now, that means it needs one more electron, and here's the thing. If it needs one more electron, it's going to say, "Hey, I’m going to share one electron. If you need one, you share one." Okay? So hydrogen says, "Hey, I'm willing to share my one valence electron." So it’s going to share one electron, and if somebody else is willing to share their one valence electron, it's going to form a nice bond with them.

Okay, that's why it forms a bond with itself. And so now we can draw the structure this way. So here's the hydrogen, and here is the valence electron, and it's ready to share that. Here’s another hydrogen; its valence electron is going to—it's ready to share that. And look, they're going to form a bond.

This is called the covalent bond, and by sharing, each atom has access to two electrons, making them both stable—that's why they stay together. Okay, covalent bond—how do we draw the Lewis structure, the final structure? Well, we just going to represent this bond with a single line, so a single line represents a bond between two electrons.

Okay, this is the Lewis structure for hydrogen.

Okay, let's take another example. Let's take oxygen. If I go over here, oxygen has six valence electrons—that's where we start. Oxygen has six valence electrons, and how many does it need to attain stability? Well, if you look at neon, then you can see that it has eight electrons in its outermost shell, and I know neon is very stable, so that helps me understand that eight electrons in the outermost shell gives you stability.

In fact, that is the case for most other elements. For hydrogen, it's two, but for most other elements, they require eight—other eight electrons in their outermost shell. We call this the octet rule.

All right, so since oxygen requires eight electrons in its outermost shell, it has six, so it needs two more. If you need two, you share two—that's the rule. Okay? So it’s going to share two electrons, and if somebody else is willing to share their two electrons, it's going to form a bond with them.

That's why it forms a bond with itself. Okay, so how do we do that? Well, again, we’re going to draw two. So, we’re going to draw the Lewis structure; we’re going to draw two oxygen atoms.

Now, since it has six valence electrons, here’s how I show the valence electrons: 1, 2, 3, 4, 5, 6. Why do I draw it this way? Because just by drawing it this way, again, you can see just from the dot structure that there's one pair, there's a second pair. Oh, there are no pairs here—oh, there are no pairs over here. These electrons are lonely; they want to get paired, and that's how I know that these two electrons are going to participate in bonding—it’s going to be shared.

So just by looking at the dot structure, I can see that it’s going to share two electrons. That's another way to just confirm the same thing that we got over here. Okay? So it's going to share these two electrons. This oxygen atom is going to do the same thing: 1, 2, 3, 4, 5, 6.

It's going to share these two, which means these two electrons and these two electrons can form a bond. Now, just to make it a little convenient for me to draw this, I’m going to show the lonely electron here, and I'm going to show the pair over here. You can draw the pair wherever you want.

Okay, it just makes it so much convenient to draw bonds over here. So this means these two electrons are going to form a bond over here, and these two electrons are going to form a bond over here. And there you have it—that means oxygen is going to form a double bond.

And you can count the number of electrons each oxygen has access to. Now, if you look at this oxygen, it has 1, 2, 3, 4, 5, 6, 7, 8—it has eight electrons access to it. And this oxygen: 1, 2, 3, 4, 5, 6, 7, 8—it has also eight electrons access to it.

And so the octet rule is satisfied. And so now, how do we draw the final Lewis structure? Well, here are the oxygen atoms—one and two. So, two bonds, two lines, and then we have to show the lone pairs. These are called lone pairs because they do not participate in bonding, but it’s important to show them.

So here's the lone pairs—you can show the lone pairs wherever you want, okay? But do show them in the Lewis dot structure. And there you have it—that is the dot structure.

All right, your turn. Why don’t you try drawing a Lewis dot structure for N2? Why don’t you pause the video and give it a try?

All right, here we go again. I look at nitrogen over here, and I know that nitrogen has five valence electrons. That’s where you always start: number of valence electrons and octet rule. Remember, most elements follow the octet rule, so it requires eight electrons in its outermost shell for stability.

It has five, so it needs three more. If you need three, you share three. So it shares three electrons. And again, if it can find some other atom or element that can share three electrons, it will happily bond with them. That’s why it forms a bond with itself.

And so how are we going to show that? Same way, so we’re going to draw the five valence electrons over here in the same way we wrote it over here: 1, 2, 3, 4, 5. And again, you can see just by drawing the dot structure, we can see there's one pair over here, but there are three lonely electrons, which would love to participate in bonding.

That’s why it shares three electrons. Just by looking at the dot structure, you can confirm that same thing over here: 1, 2, 3, 4, 5. So these three electrons can bond with these three.

And again, just to make it a little convenient to draw, okay, I’m going to—sorry, I’m going to draw that pair over here, and I’m going to draw that pair over here. All right, and of course, the more practice we get, the better we’ll be at conveniently drawing them.

Okay, but anyways, now that we have these three electrons, they can form a bond with each other. And so bond number one, bond number two, bond number three—so nitrogen forms a triple bond with itself. And again, if you count the number of electrons, this nitrogen has 1, 2, 3, 4, 5, 6, 7, 8.

And for this nitrogen again: 1, 2, 3, 4, 5, 6, 7, 8. Octet rule is satisfied. So what does the Lewis dot structure, the final structure, look like? Well, I have three bonds, so it’s going to be a triple bond. And then do not forget the lone pairs—never, ever forget the lone pairs.

Okay, and this now explains why nitrogen loves to form three bonds: either a triple bond or a double bond with one element and a single bond with another element or three separate bonds. And oxygen, for example, loves to form two bonds—double bonds with another atom or two single bonds.

All right, let’s kick things up a little bit. How about we draw one for water molecules? H2O. Things are going to get interesting now. The idea is the same, though: you start with the number of valence electrons for each one, and then think about how much they share and then try to form a bond.

Now, of course, because you have three elements over here, there could be a little bit of confusion about how they’re all connected. The general rule is the element which is on the leftmost side of the periodic table—again, if here's a periodic table, look at the—that’s on the leftmost side, which is least electronegative, or you can also think of it as most metallic, because all metals are on the left side, right? So the one that is on the leftmost side will usually be the central atom.

All right, what about hydrogen? Hydrogen is the leftmost, but it is an exception. All right? So hydrogen will always be on the side. So, except for hydrogen, any other elements that you have, you have to consider the one that's on the leftmost side. That will be your central atom.

We'll take more examples; it'll make sense. But at least when it comes to water, since hydrogen is all on the side, the oxygen will be in the center. So we can immediately draw the structure. So the oxygen will be in the center, and the two hydrogens will be connected to it.

And we already know their valence electrons and how much they're going to share. So I'm going to write one for oxygen, going to start with that same technique: 1, 2, 3, 4, 5, 6. So it’s going to share these two, and hydrogen just has one. And so you can kind of see these two are going to form a bond with these two.

And again, just for convenience, what I'm going to do is I’m going to take this and put that pair here so that I can nicely bond these two. And there you have it, so this is going to be bond number one, and this will be bond number two. And that’s it—that's the structure, the Lewis dot structure.

So it’s going to be oxygen—hydrogen, hydrogen—one bond here, one bond here. And do not forget the lone pairs.

Okay, again your turn. How about you draw one for NH3? Pause and try.

All right, there are four elements over here, but again, because hydrogen will always be on the side, we know it's the nitrogen that is going to be the central atom, and the three hydrogens will be connected to it.

So nitrogen will be in the center, and the three hydrogens—we can draw the three hydrogens wherever we want. So let me just draw them over here. And again, we'll start by drawing the central valence electrons on the central atom.

So 1, 2, 3, 4, 5—I know it has five valence electrons. So where do I draw them? I'm going to draw the fifth one over here. You can see that right now I can conveniently draw the bonds over here, and one valence electron for hydrogen is way to share that.

So one bond over here, one bond over here, and one bond over here. And this will be the lone pair. And again, you can check: hydrogen has access to two electrons, nitrogen has access to 1, 2, 3, 4, 5, 6, 7, 8.

Okay, so what's the final structure going to be? Well, I have N and I have single bond H, single bond H, single bond H. Am I forgetting something? Yes, the lone pair—do not forget the lone pairs.

All right, final challenge: carbon dioxide (CO2). So carbon is a new element for us now. Why don’t you pause the video, look at the periodic table, and try to draw the Lewis structure for this one?

All right, since carbon is a new one, we need to think about carbon. Carbon has four valence electrons, and it needs four more. And therefore, let me just write it down: four valence electrons, it needs four more. If you need four, you share four.

Okay, and now is the interesting one. I don't have hydrogen; I have carbon and oxygen. How do I know which is the element that comes in the center? Well, again, CO2 might be simpler; you might be able to guess it. But we’re going to look at which one is the leftmost element.

You can see carbon is the leftmost element on the periodic table; oxygen is over here, carbon is over here, so carbon comes in the center. All right? So carbon will be in the center, and oxygen will be attached to it.

So I’m going to draw two oxygen atoms over here, and now again, I can draw the Lewis structure for carbon. How do I do that? I’m going to show the four valence electrons: 1, 2, 3, 4—that's all there is. And look, they're all lonely; they're all four lonely.

And therefore, they're going to share all four of them to form pairs. And the same thing for oxygen, as we've done before: 1, 2, 3, 4, 5, and 6. Where should I draw 5 and 6? Let me see—I mean you can draw it anywhere, right?

So let me draw the five and six here. Okay, same thing over here: 1, 2, 3, 4. I want to draw the fifth and sixth one over here so we can bond this, and we can bond this, and we can kind of bond this, and we can kind of bond this.

It's a little crooked, but it's okay; we get the idea across. So this is one pair, and then we can bond it over here; we can bond it over here. It's fine, it’s fine. I know it looks a little weird, but we get the gist of it: carbon is going to form a double bond with this oxygen atom, and the carbon is also going to form a double bond with this oxygen atom.

And let’s count: carbon has 1, 2, 2, 3, 4, 5, 6, 7, 8. What about oxygen? 1, 2, 3, 4, 5, 6, 7, 8. And the same thing with this one. And so we can now draw a better final structure. So carbon—oxygen, oxygen—single, sorry, two bonds, so double bond here, double bond over here—do not forget the lone pairs.

So there are two lone pairs over here. Let me draw them over here and here, and there are two lone pairs over here. Here, are there any lone pairs on the carbon? No, there are no lone pairs on the carbon. So there you have it—that was fun, right?