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# Formal charge on oxygen

How to calculate the formal charge on oxygen.

## Want to join the conversation?

- We see that the oxygen with the single bond contains 7 electrons in the valence shell but from where does it gets the 7th electron?(3 votes)
- In CH₃CH₂O⁻, the O atom has eight atoms in its valence shell, but seven electrons "belong" to it
**according to the rules for formal charge**.

The extra electron came from an atom that was originally attached to the O atom but left without its half of the covalent bond.

For example, if CH₃CH₂:OH → CH₃CH₂O:⁻ + H⁺, the extra electron comes from the H atom.(7 votes)

- Why do we add electrons as lone pairs? why not a single electron, or 3, or 4, or whatever?(2 votes)
- this arises from filled orbitals of an atom (nonbonding electrons). 3 electrons are just one pair + one electron, 4 are just 2 pairs of electrons. one electron is specific for radicals (3 electrons would also make a radical)(2 votes)

- Is there any case in which the formal charge is +2?(2 votes)
- When
*oxygen*bonds we have found it to either have a formal charge of 0 (2 bonds and 2 lone pairs), +1 (3 bonds and 1 lone pair), and -1 (1 bond and 3 lone pairs).

There are a couple other possibilities which you may run into when studying**free radical reactions**and such.**Answer**

From what I've heard, oxygen will never have a formal charge of 2, at least in naturally occurring bonds.

Hope this helps,

- Convenient Colleague(2 votes)

- Is there a difference in symbolizing an ionic charge vs the oxidation state of an element? Does the plus sign move from in front of to behind?(2 votes)
- To indicate the oxidation state, you draw plus/minus sign first

To show the ionic charge, vice versa.(2 votes)

- In the case of CHO (H-C-O) in the video why is the formal charge on the Oxygen -1 and not +1. I look at the oxygen and see one bond. That in my mind leaves 2 pairs of lone electrons which gives a FC of +1. In the video there magically appears 3 pairs of lone electrons which give the FC of -1. My question is how do I know if it should be 2 pairs of lone electrons or 3 pairs of lone electrons.(2 votes)
- What time in the video is this? There are many molecules in this video so I'm not sure exactly which you're meaning.

One thing to note is that oxygen is always going to follow the octet rule, that is it will always (at this level at least) have 8 electrons around it even if they are not drawn in.

If you see oxygen with 2 bonds you can assume there are also 2 lone pairs

If you see oxygen with 1 bond you can assume there are 3 lone pairs

If you see oxygen with 3 bonds you can assume there is 1 lone pair

And from that information you can instantly know the formal charge on any oxygen too. Note I'm using a different method to calculate formal charge from Jay, I feel this one shows you where the numbers come from better. Remember each bond is 2 electrons, and each lone pair is 2 electrons.

Formal charge = # of valence electrons - # of lone pair electrons - # of bonding electrons/2

2 bonds and 2 lone pairs = 6 - 4 - 4/2 = 0 formal charge

1 bond and 3 lone pairs = 6 - 6 - 2/2 = -1 formal charge

3 bonds and 1 lone pair = 6 - 2 - 6/2 = +1 formal charge(2 votes)

- At4:04oxygen is more electronegative than hydrogen. So why didn't the oxygen get hydrogen electron?(1 vote)
- Remember there’s a difference between formal charge and oxidation state.

For oxidation state the more electronegative element gets both electrons in a bond

For formal charge each atom gets 1 electron from each bond(2 votes)

- at9:20, the molecule has 3 carbons instead of 4, right? because the double bond is between two carbon molecules(1 vote)
- If the FC changes depending on the molecular structure how do can I solve a problem? Do I assume FC is just 0? But then again it can be +1 or -1 or....(1 vote)
- No each time you would count the number of bonds to calculate the FC. FC is really important later in organic chemistry.(1 vote)

- when oxygen has a positive charge, who gets it one electron?(1 vote)
- The rules for counting electrons say that each atom gets half of the bonding electrons between them.

Thus in the C-O bond of CH₃CH₂—O⁻, C gets one of the bonding electrons and O gets the other.(1 vote)

- at4:41how come oxygen is making 3 bonds and has 2 electrons. what about that one electron that wasnt drawn(1 vote)
- I don't understand your question.

The O atom has eight electrons in its valence shell, but only five of them "belong" to the O atom when we count them to determine formal charge.(1 vote)

## Video transcript

- [Voiceover] We've already seen that formal charge is equal to the number of valence electrons in the free atom, minus the number of valence electrons in the bonded atom, and another way of saying that is the formal charge is equal to the number of valence electrons the atom is supposed to have, minus the number of valence electrons the atom actually has in the drawing, so let's
assign a formal charge to oxygen in this molecule. Remember that each bond is made up of two electrons, so this bond right here is made up of two electrons. This bond over here is
made up of two electrons, and our goal is to find the formal charge on oxygen, so the formal charge on oxygen is equal to the number of valence electrons in the free atom, so the number of valence electrons that oxygen is supposed to have, we know that's six, right, oxygen is supposed to have six valence electrons, minus the number of electrons that oxygen actually has in our drawing. Now remember when you have a bond with two electrons, we give one electron to one atom, and the other electron to the other atom. So from these two electrons, oxygen gets one of those electrons. All right, same thing
for this other bond here, oxygen gets one of those electrons, and the other electron goes to hydrogen. So now we have a total of six electrons around our oxygens, so
I'll highlight those. We have one, two, three, four, five, six, so there are six valence electrons around the oxygen in our drawing, so six minus six is equal to zero, so the formal charge on oxygen is equal to zero. And let me go ahead and
write down this pattern that we've just seen. When oxygen has two bonds, and two lone pairs of electrons, one oxygen has two bonds and two lone pairs of electrons, the formal charge is equal to zero. And sometimes, the lone
pairs are just left off for convenience reasons, right, so you could draw this with your oxygen, and your hydrogen like that, or you could even go like this, and all those are just different ways of representing the same molecule. So leaving off lone pairs of electrons. Let's look at some other examples, where the formal charge on oxygen is equal to zero, and we'll look at the
one on the left first, so the formal charge is equal to zero on this oxygen and we can see we have two bonds here to oxygen, here's one of the bonds to oxygen,
and here's the other bond to oxygen. The lone pairs of electrons have been left off this dot structure, but we know, since the
formal charge is zero, we already have our two bonds here. There should be two pairs of electrons on that oxygen, so you
can put them in there or you could leave them off. I'll go ahead and put them in. So we have a total of eight electrons around our oxygen, so oxygen's following the octet rule here,
so I'll highlight them, two, four, six, and eight. On the right, we have another example, where oxygen has a formal charge of zero, and this oxygen has two bonds to it so here's one of the bonds,
and here's the other bond, and this oxygen would also have to have two lone pairs of electrons on it, and again, I didn't draw them in here. Sometimes you don't draw
them in here for convenience, but I could go ahead and add them, so it's easier to see that that oxygen has a formal charge of zero. And a total of an octet
of electrons around it, so let me highlight those, two, four, six, and eight. So these patterns are important, and for oxygen, two
bonds, and two lone pairs of electrons give us a
formal charge of zero. Let's move on to another formal charge situation for oxygen, let's
find the formal charge on oxygen here, so we start by drawing in the electrons in our bonds. Each bond consists of two electrons, so I draw those in. What is the formal charge
on oxygen this time? So the formal charge is equal to the number of valence electrons oxygen is supposed to have, which is six, minus the number of valence electrons oxygen actually has in our drawing. So we divide up our electrons again, so oxygen gets one from this bond, and one from this bond, and one from this bond, so how many electrons
are around oxygen now? This would be one, two, three, four, and five. So six minus five is equal to plus one, so it's like oxygen has
lost an electron here. So oxygen has a formal charge of plus one. I could re-draw that, let me go ahead and do that over here. So I could re-draw that over here, so we have oxygen with
our bond to hydrogens, and oxygen has a lone pair of electrons on it, and this oxygen has
a plus one formal charge. And so we can come up
with another pattern here. Here oxygen has three bonds, let me highlight those bonds. I'll use red this time, so here's one bond, two bonds, and then three bonds
and then one lone pair of electrons, so the
pattern of three bonds plus one lone pair of electrons for oxygen will give you a formal charge of plus one, and again, it's good to recognize these patterns. You should be able to do the calculation, and then after you do
enough of these problems, you can just look at it and figure out what the formal charge is. Let's look at some more examples where oxygen has a formal charge of plus one, and the lone pairs were left off of this one, again, for
convenience reasons, so we'll start with this example on the left, we can see that the oxygen with a plus one formal charge has three bonds to it. Here's one, here's two,
and then here's three, so three bonds, in order for that oxygen to have a plus one formal charge, it must also have one lone pair of electrons on it, so you
could just leave them off and know that they're there, or you could go ahead and draw them in, and I'll draw them in on that oxygen. Now to our other example over here on the right. This is the oxygen with
a plus one formal charge, so that oxygen must have three bonds and one lone pair. So here are the three bonds, here's one, two, and three, and again, I didn't draw in the lone pair of electrons on the oxygen, but the lone pair is there, so I'll go ahead and put it in like that. So recognize this pattern, three bonds, plus one
lone pair for oxygen, gives us a formal charge of plus one. You could have also figured out how many electrons are necessary. Let's use this example, by, let me go ahead and re-draw it here. So how else could we figure out how many electrons are on that oxygen, if that oxygen has a
plus one formal charge. Well you could say that, all right, the calculation for formal charge would be six minus x
is what we don't know, but we do know the formal
charge is plus one, so let me just put this
in a little box over here. So six minus x is equal to plus one. Obviously, x would have
to be equal to five, meaning that oxygen
would have five electrons around it, let's think about how many electrons we see right now around oxygens. So let me draw in some electrons here and I'll use red, so this bond is two electrons, and then we have a bunch of electrons in here, so how many electrons around oxygen do we have so far? Well, we would have three, just these three, and we need five, which means we need two more, which means we need a
lone pair of electrons on the oxygen. So that's a little bit too complicated, I think, for figuring it out, but you could use that method, or you could just learn
this pattern, right, and eventually, you'll
have to have this pattern down pretty well. Let's look at another example for assigning formal charge to oxygen. So our goal is to find the formal charge on oxygen in this example, and we put in our electrons in this bond. Each bond represents two electrons, so the formal charge on oxygen, is equal to the number
of valence electrons that oxygen is supposed
to have which is six, minus the number of valence electrons that oxygen actually has, and in this example, right, we would take one of these electrons from this bond, and how many electrons is that total around oxygen? This would be one, two, three, four, five, six, seven, so six minus seven
gives us a formal charge of negative one. So I could re-draw that over here, I could say oxygen has three lone pairs of electrons, and a negative one formal charge, so our pattern, this time, our pattern is one bond. Here's the one bond, and then three lone pairs of electrons. So let me write that down. So the pattern is one bond, when oxygen has one bond, and three lone pairs of electrons the formal charge is negative one, just like we saw up here
with the calculation. So we could leave those electrons off if you wanted to save some time, we could just say, oh, this is oxygen with a negative one formal charge, and we should know that there must be three lone pairs of
electrons on that oxygen. Let's look at one more example, where formal charge is negative one. So right here, this
oxygen has a negative one formal charge, and we can see it already has one bond to it. And so the pattern, of course, is one bond plus three
lone pairs of electrons. So we already have the one bond. In order for that oxygen to have a negative one formal charge, we need three lone pairs of electrons. So we could re-draw this, so that is one way to represent that ion, and we could also represent it like this, with putting three lone pairs of electrons on that oxygen with a negative one formal charge. So again, become familiar
with these patterns.