Bonding, Structure, and Resonance

By James Ashenhurst

A Key Skill: How to Calculate Formal Charge

Last updated: March 21st, 2021 |

Hey! Welcome to Master Organic Chemistry, just in case you’re a first time visitor. 

In this blog post I explain how to calculate formal charge for molecules.  However, you might find my videos containing 10 solved examples of formal charge problems to be even more useful. Just thought you should know! 

Need to figure out if an atom is negative, positive, or neutral? Here’s the formula for figuring out the “formal charge” of an atom:

Formal charge = [# of valence electrons] – [electrons in lone pairs + 1/2 the number of bonding electrons]

This formula explicitly spells out the relationship between the number of bonding electrons and their relationship to how many are formally “owned” by the atom.

For example, applying this to BH4 (top left corner in the image below) we get:

  • The number of valence electrons for boron is 3.
  • The number of non-bonded electrons is zero.
  • The total number of bonding electrons around boron is 8 (full octet). One half of this is 4.

So formal charge = 3 – (0 + 4)  = 3 – 4  = –1

There is a slightly easier way to do this, however.

Since a chemical bond has two electrons, the  “number of bonding electrons divided by 2” is by definition equal to the number of bonds surrounding the atom. So we can instead use this shortcut formula:

Formal Charge = [# of valence electrons on atom] – [non-bonded electrons + number of bonds].

Applying this again to BH4 (top left corner).

  • The number of valence electrons for boron is 3.
  • The number of non-bonded electrons is zero.
  • The number of bonds around boron is 4.

So formal charge = 3 – (0 + 4)  = 3 – 4  = –1

The formal charge of B in BH4 is negative 1. 

Let’s apply it to :CH3 (one to the right from BH4)

  • The number of valence electrons for carbon is 4
  • The number of non-bonded electrons is two (it has a lone pair)
  • The number of bonds around carbon is 3.

So formal charge = 4 – (2 +3) = 4 – 5  = –1

The formal charge of C in :CH3 is negative 1. 

Same formal charge as BH4!

Let’s do one last example. Let’s do CH3+ (with no lone pairs on carbon). It’s the orange one on the bottom row.

  • The number of valence electrons for carbon is 4
  • The number of non-bonded electrons is zero
  • The number of bonds around carbon is 3.

So formal charge = 4 – (0 +3) = 4 – 3 = +1

You can apply this formula to any atom you care to name.

Here is a chart for some simple molecules along the series B C N O . I hope beryllium and fluorine aren’t too offended that I skipped them, but they’re really not that interesting for the purposes of this table.


Note the interesting pattern in the geometries (highlighted in colour):  BH4(–), CH4, and NH4(+) all have the same geometries, as do CH3(–), NH3, and OH3(+).  Carbocation CH3(+) has the same electronic configuration (and geometry) as neutral borane, BH3. The familiar bent structure of water, H2O, is shared by the amide anion, NH2(–). These shared geometries are one of the interesting consequences of valence shell electron pair repulsion theory (VSEPR – pronounced “vesper“, just like “Favre” is pronounced “Farve”.)

The formal charge formula also works for double and triple bonds:


Here’s a question. Alkanes, alkenes, and alkynes are neutral, since there are four bonds and no unbonded electrons:  4 – [4+0] = 0.  For what other values of [bonds +  nonbonded electrons] will you also get a value of zero, and what might these structures look like? (You’ll meet some of these structures later in the course).

One final question – why do you think this is called “formal charge”?

Think about what the formal charge of BF4 would be. Negative charge on the boron. What’s the most electronegative element here? Fluoride, of course, with an electronegativity of  4.0, with boron clocking in at 2.0. Where do you think that negative charge really resides?

Well, it ain’t on boron. It’s actually spread out through the more electronegative fluoride ions, which become more electron-rich. So although the “formal” address of the negative charge is on boron, the electron density is actually spread out over the fluorides. In other words, in this case the formal charge bears no resemblance to reality.

Another reminder – 10 videos with solved examples of formal charge problems, right here (look at the very top of the page) 


Comment section

51 thoughts on “A Key Skill: How to Calculate Formal Charge

  1. sir
    the sheet posted by u is really very excellent.i m teacher of chemistry in india for pre engineering test.if u send me complete flow chart of chemistry i will great full for u

  2. Very good explanation.I finally understood how to calculate the formal charge,was having some trouble with it.Thanks:)

  3. The answer to the question in the post above is “carbenes” – they have two substitutents, one pair of electrons, and an empty p orbital – so a total of four electrons “to itself”, making it neutral.

  4. Shouldn’t the formal charge of CH3 be -1? I was just wondering because in your example its +1 and in the chart its -1.

    1. In the question.. its mentioned that CH3 without any lone pairs.. which means the valence would be 4 but there will not be any (2electrons) lone pairs left.. Hence it will be (4-)-(0+3)= 1

      1. In CH3 i think FC on C should be -1 as carbon valency is 4 it has already bonded with 3 hydrogen atom one electron is left free on carbon to get bond with or share with one electron H hence, number of non bonded electrons lone pair of electrons is considered as 2. 4-(2+3) = -1.
        In your case if we take 0 than valency of c is not satisfied.

  5. Hey great explanation. I have a question though. Why is the FC commonly +/- 1? Could you give me an example when the FC is not +/- 1? Thanks.

  6. There are meny compounds which bears various structure among these which one is more stable or less energetic is it possible to predicu from the formal charge calculation?

    1. I hope the post doesn’t get interpreted as “formal charges have no significance”. If it does I will have to change some of the wording.

      What I mean to get across is that formal charges assigned to atoms do not *always* accurately depict electron density on that atom, and one has to be careful.

      In other words, formal charge and electron density are two different things and they do not always overlap.

      Formal charge is a book-keeping device, where we count electrons and assign a full charge to one or more of the atoms on a molecule or ion.
      Electron density, on the other hand, is a measurement of where the electrons actually are (or aren’t) on a species, and those charges can be fractional or partial charges.

      First of all, the charge itself is very real. The ions NH4+ , HO-, H3O+ and so on actually do bear a single charge. The thing to remember is that from a charge density perspective, that charge might be distributed over multiple atoms.
      Take an ion like H3O+, for example. H3O *does* bear a charge of +1,

      However, if one thinks about where the electrons are in H3O+, one realizes that oxygen is more electronegative than hydrogen, and is actually “taking’ electrons from each hydrogen. If you look at an electron density map of H3O+ , one will see that the positive charge is distributed on the three hydrogens, and the oxygen actually bears a slight negative charge. There’s a nice map here.

      When we calculate formal charge for H3O+, we assign a charge of +1 to oxygen. This is for book keeping reasons. As a book-keeping device, it would be a royal pain to deal with fractions of charges like this. So that’s why we calculate formal charge and use it.

      Sometimes it does accurately depict electron density. For example, in the hydroxide ion, HO- , the negative charge is almost all on the oxygen.

      If you have a firm grasp of electronegativity then it becomes less confusing.

      Does that help?

  7. Thank you!!! this was awesome, I’m a junior in chemistry and this finally answered all my questions about formal charge :)

  8. This works! I would take your class with organic chemistry if you are a professor. I am taking chemistry 2 now. Organic is next.
    Thank you so much!

  9. you said that non bonded electrons in carbon is 2, but how ?
    because i see it as only 1 because out of the 4 valence electrons in carbon, three are paired with hydrogen so it’s only 1 left

    1. If the charge is -1, there must be an “extra” electron on carbon – this is why there’s a lone pair. If there was only one electron, it would be neutral.

  10. I am beryllium and i got offended!!!!!!……..LOL Just kidding…….BTW, I found this article very useful.Thanks!!!!!!!!!!

    # OF BONDING=3


    FORMAL CHARGE=4+[(1/2*3)+1]


  12. Thanks for the easy approach.
    I have a problem in finding the FC on each O atom in ozone. Can you help me with that ASAP?

    1. The FC on central atom would be +1 because [6-(2+3)]
      FC on O atom with coordinate bond would be: -1 because [6-(6+1)].
      FC on O atom with double bond is: 0 because [6-(4+2)].

      Hope I solved your question!

  13. This method is wrong
    For CH3 , the valence eloctron is 4 , no : of bonds is 3 and no of non bonded electrons is 1
    Then by this equation

    F.C= 4-(1+3) = 0 but here it is given as +1

  14. This really helped for neutral covalent molecules. However, I’m having trouble applying this technique for molecules with an overall charge other than 0. For instance, in (ClO2)- , the formal charge of Cl should be 1. However, with your equation the charge should be 0. With the conventional equation, the charge is indeed 1.

    I’d appreciate it if you replied sooner rather than later, as I do have a chemistry midterm on Friday. I’m quite confused with formal charges :)

    Thanks for the study guide.

  15. I remember learning that in the cyanide ion, the carbon is nucleophilic because the formal negative charge is on carbon, not nitrogen, despite nitrogen being more electronegative. So I think a different explanation could me more accurate, but I’m not sure how to properly address it. I better keep reading.

    1. In cyanide ion, there are two lone pairs – one on carbon, one on nitrogen. The lone pair on carbon is more nucleophilic because it is less tightly held (the atom is less electronegative than nitrogen).
      On all the examples I show that are negatively charged (eg BH4(-) ) there isn’t a lone pair to complicate questions of nucleophilicity.

  16. It was a very great explanation! Now I have a good concept about how to find formula charge.
    And also i am just a grade nine student so i want to say thank you for this.

  17. That was the best i have seen but i have a problem with the formula,i think the side where the shared pair electrons came was suppose to be negative but then yours was positive,so am finfding it difficult to understand because the slides we were given by our lecturer shows that it was subtracted not added. i would love it when u explain it to me.

  18. Hi I am extremely confused. The two formulas for calculating FC that you provided are not the same and don’t produce the same results when I tried them out.

    Formal charge = [# of valence electrons] – [electrons in lone pairs + 1/2 the number of bonding electrons]


    Formal Charge = [# of valence electrons on atom] – [non-bonded electrons + number of bonds].

    They do not produce the same result…
    If I have the formula BH4, and use the first formula provided to find FC of B, I would get:

    (3) – (0 + 2) = +1

    Using the second formula provided:

    (3) – (0+4) = -1

    Aren’t these formulas supposed to produce the same results? I am quite confused and I don’t know if I missed something.

    1. Ah. I should have been more clear. The number of bonding electrons in BH4 equals 8, since each bond has two electrons and there are 4 B-H bonds. Half of this number equals 4. This should give you the same answer.
      I have updated the post to make this more explicit.

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