Worked example: Lewis diagram of formaldehyde (CH₂O) | AP Chemistry | Khan Academy

What we're going to do in this video is get a little bit more practice constructing Lewis diagrams.

In particular, we're going to try to construct the Lewis diagram for formaldehyde. Formaldehyde has one carbon, two hydrogens, and an oxygen, CH₂O. So pause this video and have a go at it. Try to construct a valid Lewis structure or Lewis diagram for formaldehyde.

All right, now let's do this together. Now, the first step—and we saw this in a previous video—we want to think about all of the valence electrons for this molecule. So we want to account for valence electrons.

Now, the reason why we want to do that is so that while we're trying to create this structure we are making use of all of the valence electrons. To figure out how many total valence electrons we have, we can look at a periodic table of elements. We can see that carbon is in that second row, in that second period. So its second shell is its outer shell, and in that shell it has one, two, three, four valence electrons.

So carbon has four valence electrons. A neutral free hydrogen atom is going to have one valence electron, but we have two of them here, so it's going to be two times one. And then oxygen, it also is in the second period, and in its second shell it has one, two, three, four, five, six valence electrons.

So the total valence electrons in this molecule are going to be 4 plus 2, which is 6, plus 6, which is equal to 12 valence electrons. Now, the next step is to try to draw a structure. Try to draw single bonds. I'll say single bonds.

And a key question is, what do we think is going to be our central atom? The rule of thumb is the least electronegative atom that is not hydrogen is a good candidate for our central atom. So we can rule out hydrogen. Between carbon and oxygen, we know that oxygen is one of the most electronegative atoms, one of the most electronegative elements on the periodic table of elements.

It's very close to fluorine and so carbon is a good candidate for the central atom. So let's put the carbon right over here. Then let's put these other atoms around it. We could call them terminal atoms. So let's put our oxygen right over there, and then we have two hydrogens: hydrogen there, a hydrogen there.

And let me draw the bonds. So that's a single bond. That accounts for two valence electrons, that accounts for two valence electrons, that accounts for two valence electrons. So I've just used two, four, six valence electrons.

So if I subtract six valence electrons, I am now left with six valence electrons. Six valence electrons. So the next step is allocate the remaining valence electrons, trying to get to the octet rule for atoms that are not hydrogen and then for hydrogen trying to get it to have two valence electrons.

So allocate remaining valence electrons. All right, so let's start with this oxygen. This oxygen already has these two electrons that it's sharing hanging around. So in order to get to the octet rule, it needs six more. So let's give it six electrons: one, two, three, four, five, six.

Well, I've just used up the remaining six valence electrons, so I don't really have any more to play with. Let's see how the other atoms are feeling. So hydrogen here, it's able to share these two electrons that are in this covalent bond, so it's feeling good.

It can kind of pretend that it has a full outer shell because its outer shell is just that one, that first shell that's filled with two electrons. Same thing for this other hydrogen. So at least the terminal atoms, the oxygen and the two hydrogens, are feeling like they have a full outer shell.

But in the fourth step, we're going to look at our central atom. So let's focus on the central atom and do we need more bonds or do we need to do something interesting here? What we see is that carbon, it's able to have two, four, six electrons hanging around it, but it would love to have eight.

Carbon would love to have a full outer shell. So how could we do that? Well, we could add more bonds. Where could the bonds come from? Well, it would come from some lone pair of electrons.

Well, the only lone pairs of electrons are hanging around this oxygen. So what if we were to take say this lone pair of electrons and then construct another covalent bond with that? Then our Lewis diagram will look like this. I will actually redraw it.

So you have your carbon, you have your three original covalent bonds: you had a hydrogen, a hydrogen, and then you had your oxygen right over here. And now we formed a new covalent bond just like this. And then you have these two other lone pairs around the oxygen.

So let me draw that: two, and then another two around the oxygen. And this is looking pretty good because the oxygen still has eight electrons hanging around: four in lone pairs and then four that in this double bond that it is sharing.

The hydrogens still have two electrons hanging around—they're able to share the electrons in each of these covalent bonds. And now the carbon is participating in, you could think of it as, four covalent bonds: two single bonds and one double bond.

So each of those have two electrons associated with it, so it has eight electrons hanging around. So this is looking really good as a legitimate Lewis structure or Lewis diagram for formaldehyde.