So now let's go back to our examples, and let's calculate the formal charges for

each atom in each of these structures.

So I'm going to look at my carbon first,

in carbon I have 4 valance electrons in the isolated atom.

I don't have any non bonding electrons assigned to carbon in this structure.

And for half of my bonding, I have 8 bonding electrons.

And I end up with a formal charge of 0.

For oxygen, I know that I have 6 electrons in the isolated atom,

minus the 4 that are at non-bonding electrons assigned to oxygen.

And minus half of the 4 electrons in the bonds.

And again, I get a formal charge of 0.

For chlorine, I only need to do this calculation once,

because the chlorines are exactly the same.

They're both single bonded,

they both have the same number of non-bonding electrons assigned to them.

So I know that chlorine has 7 valence electrons minus 6 electrons

non-bonding assigned to it in this particular molecule, minus half

of the bonding electrons, and I also get 4 more charges of 0 for my chlorine atoms.

So, now I have a set of formal charges of 0, 0, 0.

And that looks pretty good.

I can't get any closer to 0 then that.

But for comparison we want to look at the formal charges of our other structure.

So for carbon, I see that it's going to look exactly the same.

I still have 4 electrons in the isolated atom.

I have no non-bonding electrons, and

I still have 8 bonding electrons around that.

So, I end up with a value of 0.

Now I look at oxygen, which has 6 valence.

Here it has 6 non-bonding electrons, minus half of the bonding,

so I have 6 minus 6 minus 1, gives me minus 1.

Now, I have to look at my chlorine atom separately,

because they have different bonding and different structural features.

So I'm going to look at the chlorine that has the single bond first, and

it's actually going to look just like.

The chlorines in the first structure we looked at,

because it has the exact same structural features there.

When I look at the chlorine with a double bond, I still start with 7 electrons.

I subtract off the 4 that are non-bonding in the chlorine, minus half.

Of the 4 that are bonding, and what I get is the formal charge of plus 1.

So looking at this I have values of 0 minus 1, 0 and plus 1.

Together they really aren't that bad of a set of formal charges, and

to determine the best structure we can't just look at one set of formal charges.

We have to compare it to the alternatives.

Here we really only had two possible structures.

But for some molecules, we'll actually have more structures that

are not equivalent, that we have to decide among.

Now, if it were to be that this were the better of the two structures,

I would also have a resonance structure.

For this particular Lewis structure,

because I can put the double bond between this carbon and chlorine.

Because all my values are zero, I've got the values as close to zero as possible.

I see that this is going to be the better Lewis structure, and

this is equivalent to what I'll actually see in nature.