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how to calculate the formal charge of o3 pure

  • Formal Charge

how to calculate the formal charge of o3 pure

How do you calculate the formal charge of ${{O}_{3}}$ ?

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What are the formal charges in O 3 (ozone)?

Formal charge: the formal charge of an atom in a molecule is the charge which might exist on the atom if all bonding electrons were evenly shared. a formal charge value is equal to an atom's valence electrons deducting the number of electrons given to it. f . c . = [ total no . of valence e – in free state ] – [ total no . of non - bonding pair e – ( lone pair ) ] – 1 2 [ total no . of bonding e – ] structure of ozone: ozone has a dipole moment of 0 . 53 d and thus a polar molecule. the molecule can be described as a resonance hybrid with significant contributing structures, one with a single bond on one side and the other with a double bond. both sides have an overall bond order of 1 . 5 in this arrangement. formal charge in o 3 ( ozone): in an o 3 molecule, the formal charge on the middle oxygen atom( 2 ) is + 1 . f . c = 6 – 2 – 1 2 ( 6 ) f . c = 6 – 5 f . c = 1 in an o 3 molecule, the formal charge on the left oxygen atom( 3 ) is - 1 . f . c = 6 – 6 – 1 2 ( 2 ) f . c = 6 - 7 f . c = - 1 in an o 3 molecule, the formal charge on the right oxygen atom( 1 ) is 0 . f . c = 6 – 4 – 1 2 ( 4 ) f . c = 6 – 6 f . c = 0.

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The formal charge on oxygen which is single bonded in ozone is:

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Home / A Key Skill: How to Calculate Formal Charge

Bonding, Structure, and Resonance

By James Ashenhurst

Last updated: December 13th, 2022 |

How To Calculate Formal Charge

To calculate the formal charge of an atom, we start by:

The formal charge   FC is then calculated by subtracting NBE  and B  from VE .

FC = VE – ( NBE + B ) 

which is equivalent to

FC = VE – NBE – B

The calculation is pretty straightforward if all the information is given to you. However, for brevity’s sake, there are many times when lone pairs and C-H bonds are not explicitly drawn out .

So part of the trick for you will be to calculate the formal charge in situations where you have to take account of implicit  lone pairs and C-H bonds.

In the article below, we’ll address many of these situations. We’ll also warn you of the situations where the calculated formal charge of an atom is not necessarily a good clue as to its reactivity , which is extremely important going forward.

Table of Contents

Quiz Yourself!

(advanced) references and further reading, 1. formal charge.

Formal charge is a book-keeping formalism for assigning a charge to a specific atom.

To obtain the formal charge of an atom, we start by counting the number of valence electrons [ Note 1 ] for the neutral atom , and then subtract from it the number of electrons that it “ owns ” ( i.e. electrons in lone pairs, or singly-occupied orbitals ) and half of the electrons that it shares ( half the number of bonding electrons, which is equivalent to the number of bonds )

The simplest way to write the formula for formal charge   ( FC)  is:

It’s called “ formal ” charge because it assumes that all bonding electrons are shared equally . It doesn’t account for electronegativity differences (i.e. dipoles).

For that reason formal charge isn’t always a good guide to where the electrons actually are in a molecule and can be an unreliable guide to reactivity. We’ll have more to say on that below .

2. Simple Examples For First-Row Elements

When all the lone pairs are drawn out for you, calculating formal charge is fairly straightforward.

Let’s work through the first example in the quiz below.

See if you can fill in the rest for the examples below.

If that went well, you could try filling in the formal charges for all of the examples in this table.

It will take some getting used to formal charge , but after a period of time it will be  assumed that you understand how to calculate formal charge , and that you can recognize structures where atoms will have a formal charge .

Let’s deal with some slightly trickier cases.

3. Formal Charge Calculations When You Aren’t Given All The Details

When we draw a stick figure of a person and don’t draw in their fingers, it doesn’t mean we’re drawing someone who had a bad day working with a table saw . We just assume that you could fill in the fingers if you really needed to, but you’re skipping it just to save time.

Chemical line drawings are like stick figures. They omit a lot of detail but still assume you know that certain things are there.

One note. If we draw a stick figure, and we do draw the fingers, and took the time to only draw in only 3 , then we can safely assume that the person really does only have 3 fingers . So in  the last two examples on that quiz we had to draw in the hydrogens in order for you to know that it was a carbocation, otherwise you would have to assume that it had a full octet!  

Oxygen and nitrogen (and the halogens) are dealt with slightly differently.

So even when the lone pairs aren’t drawn in, assume that enough are present to make a full octet . And when bonds from these atoms to hydrogen are missing , that means exactly what it seems to be: there really isn’t any hydrogen!

Try these examples:

Now see if you can put these examples together!

(Note that some of these are not stable molecules, but instead represent are resonance forms that you will encounter at various points during the course!)

4. Some Classic Formal Charge Questions

We can use the exact same formal charge formula, above, along with the rules for implicit lone pairs and hydrogens, to figure out the formal charge of atoms in some pretty exotic-looking molecules.

Here are some classic formal charge problems.

Note that although the structures might look weird,  the formal charge formula remains the same. 

The formal charge formula can even be applied to some fairly exotic reactive intermediates we’ll meet later in the semester.

Don’t get spooked out. Just count the electrons and the bonds, and that will lead you to the right answer.

5. Formal Charges and Curved Arrows

We use curved arrows to show the movement of electron pairs in reactions and in resonance structures. ( See post: Curved Arrows For Reactions )

For example, here is a curved arrow that shows the reaction of the hydroxide ion HO(-) with a proton (H+).

The arrow shows movement of two electrons from oxygen to form a new O–H bond .

Curved arrows are also useful for keeping track of changes in formal charge .  Note that the formal charge at the initial tail of the curved arrow (the oxygen) becomes more positive (from -1 to 0) and the formal charge at the final tail (the H+) becomes more negative (from +1 to 0). 

When acid is added to water, we form the hydronium ion , H 3 O + .

Here’s a quiz. See if you can draw the curved arrow going from the hydroxide ion to H 3 O+.

If you did it successfully – congratulations!

But I’m willing to bet that at least a small percentage of you drew the arrow going to the positively charged oxygen .

What’s wrong with that?

There isn’t an empty orbital on oxygen that can accept the lone pair.  If you follow the logic of curved arrows, that would result in a new O–O bond, and 10 electrons on the oxygen, breaking the octet rule.

Hold on a minute, you might say. “ I thought oxygen was positively charged? I f it doesn’t react on oxygen, where is it supposed to react ?”

On the hydrogens! H 3 O+ is Brønsted acid, after all. Right?

This is a great illustration of the reason why it’s called “ formal charge”, and how formal charge not the same as  electrostatic charge (a.ka. “partial charges” or “electron density”).

Formal charge is ultimately a book-keeping formalism, a little bit like assigning the “win” to one of the 5 pitchers in a baseball game. [ Note 3 ] It doesn’t take into account the fact that the electrons in the oxygen-hydrogen bond are unequally shared, with a substantial dipole.

So although we draw a “formal” charge on oxygen, the partial positive charges are all on  hydrogen. Despite bearing a positive formal charge bears a partially negative electrostatic charge.

This is why bases such as HO(-) react at the H, not the oxygen.

Just to reiterate:

6. Halogens

Positive formal charges on halogens fall into two main categories.

We’ll often be found drawing  halonium ions   Cl+ , Br+, and I+ as species with six valence electrons and an empty orbital  ( but never F+ – it’s a ravenous beast )

It’s OK to think of these species as bearing an empty orbital since they are large and relatively polarizable .  They can distribute the positive charge over their relatively large volume.

These species can accept a lone pair of electrons from a Lewis base , resulting in a full octet.

Cl, Br, and I can also bear positive formal charges as a result of being bonded to two atoms.

It’s important to realize in these cases that the halogen bears a  full octet and not an empty orbital. They will therefore not directly accept a pair of electrons from Lewis bases; it’s often the case that the atom adjacent to the halogen accepts the electrons.

7. Conclusion

If you have reached the end and did all the quizzes, you should be well prepared for all the examples of formal charge you see in the rest of the course.

Related Articles

Note 1. Using “valence electrons” gets you the right answer. But if you think about it, it doesn’t quite make sense. Where do positive charges come from? From the positively charged protons in the nucleus, of course!

So the “valence electrons” part of this equation is more properly thought of as a proxy for valence protons – which is another way of saying the “ effective nuclear charge” ; the charge felt by each valence electron from the nucleus, not counting the filled inner shells.

Note 2. Nitrenes are an exception. Another exception is when we want to draw  bad resonance forms.

Note 3 . In baseball, every game results in a win or a loss for the team . Back in the days of   Old Hoss Radborn , where complete games were the norm, a logical extension of this was to assign the win to the individual pitcher. In today’s era, with multiple relief pitchers, there are rules for determining which pitcher gets credited with the win. It’s very possible for a pitcher to get completely shelled on the mound and yet, through fortuitous circumstance, still be credited for the win.  See post: Maybe They Should Call Them, “Formal Wins” ? 

In the same way, oxygen is given individual credit for the charge of +1 on the hydronium ion , H 3 O+, even though the actual positive electrostatic charge is distributed among the hydrogens.

Note 4. This image from a previous incarnation of this post demonstates some relationships for the geometry of various compounds of first-row elements.

1. Valence, Oxidation Number, and Formal Charge : Three Related but Fundamentally Different Concepts Gerard Parkin Journal of Chemical Education 2006 83 (5), 791 DOI : 10.1021/ed083p791 

2. Lewis structures, formal charge , and oxidation numbers: A more user-friendly approach John E. Packer and Sheila D. Woodgate Journal of Chemical Education   1991   68  (6), 456 DOI : 10.1021/ed068p456

00 General Chemistry Review

01 Bonding, Structure, and Resonance

02 Acid Base Reactions

03 Alkanes and Nomenclature

04 Conformations and Cycloalkanes

05 A Primer On Organic Reactions

06 Free Radical Reactions

07 Stereochemistry and Chirality

08 Substitution Reactions

09 Elimination Reactions

10 Rearrangements

11 SN1/SN2/E1/E2 Decision

12 Alkene Reactions

13 Alkyne Reactions

14 Alcohols, Epoxides and Ethers

15 Organometallics

16 Spectroscopy

17 Dienes and MO Theory

18 Aromaticity

19 Reactions of Aromatic Molecules

20 Aldehydes and Ketones

21 Carboxylic Acid Derivatives

22 Enols and Enolates

24 Carbohydrates

25 Fun and Miscellaneous

Comment section

57 thoughts on “ a key skill: how to calculate formal charge ”.

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

nice, concise explanation

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

Glad you found it helpful.

thank you for excellent explanation

Glad you found it useful Peter!

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.

thank you for collaboration of formal charge

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.

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

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.

Great!i can use this for my exam!thanks!

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.

Sure, try oxygen with no bonds and a full octet of electrons.

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?

If formal charges bear no resemblance to reality, what are their significance?

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.

http://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Aqueous_Solutions/The_hydronium_Ion

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?

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

Glad it was helpful Haley!

Thank you very very more for the simple explanation! Unbelievably easy and saves so much time!!!!!!

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!

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

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.

what does it means if we determine a molecule with zero charge ?

It’s neutral!

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

AM REALLY LOST NOW ON THAT EXAMPLE OF CH3 CARBON # OF VALENCE ELECTRON=4 # OF BONDING=3 # OF UNSHARED=1

SO WHEN I CALCULATE

FORMAL CHARGE=(#OF VALENCE ELEC)+[(1/2#OF BOND)+(#OF UNSHARED)] FORMAL CHARGE=4+[(1/2*3)+1] =1.5

PLZ HELP IF AM MAKING MISTAKE

Should be 1/2 [# of bonding ELECTRONS] + # unshared. This gives you 4 – [3 – 1] = 0 for ch3 radical.

Should be for CH3(+), not the methyl radical •CH3 .

Thank u very much my exam is today and i wouldn’t pass without this information

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?

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!

But when I used this formula it works. Thus #valence electrons_#lone pair__#1/2.bond pairs

Thanks for the easy approach.

This was so helpful n the best explanation about the topic…

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

That analysis would be accurate for the methyl radical. However it fails for the methyl carbocation.

That example referred to the carbocation. For the methyl radical, the formal charge is indeed zero.

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.

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.

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.

YOU ARE THE BEST. I GOT THE HIGHEST MARK IN MY FIRST QUIZ, AND I KNOW THAT THROUGH THIS I WILL GET THE BEST IN MIDTERM AND FINAL. I want you guys to go on youtube and follow the steps. THANK YOU VERY MUCH.

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.

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.

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.

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.

Great teaching , can I know where did u studied ??

Nice simple explanation

Thank you so much sir. Finally i understood how to calculate the formal charge

I think for Quiz ID: 2310, the formal charge for the carbon in the fourth molecule should be +1 instead of -1.

Fixed. Thanks for the spot!

Your explanations and examples were clear and easy to understand. I appreciate the detailed step-by-step instructions, which made it easy to follow along and understand the concept. Thank you for taking the time to create this helpful resource

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how to calculate the formal charge of o3 pure

Formal Charge

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Nie wieder prokastinieren mit unseren Lernerinnerungen.

When molecules are formed, they don't just combine in any way they can. Structures are optimized to be as stable as possible. One way molecules do this is by keeping the electrons within the molecule as symmetrical as possible, i.e. we want as neutral of a molecule as possible.

So how do we know whether an atom is neutral inside a molecule? We use a concept called formal charge. In this article, we will be learning all about formal charge: what it is, how to calculate it, and why it's important.

Formal Charge Explanation and its Properties

Let's start by defining formal charge.

Formal charge (FC) is the charge assigned to an atom in a molecule when we assume that electrons in all bonds are shared equally between atoms.

Formal charge ignores the concept of electronegativity . Which is the tendency for an atom/molecule to attract and share electrons unequally. For example, fluorine is very electronegative, while hydrogen is less so, so the electrons in the H-F bond will tend towards fluorine.

Here are some things to remember about formal charge:

1) Every atom can be assigned a formal charge

2) If there are multiples of the same element, they can have different formal charges

3) The formal charge is dependent on:

-The number of bonds

-The number of paired and unpaired electrons

4) Formal charges are assigned based on Lewis structures (2D structure)

When we draw a Lewis structure, we want every element to have 8 total valence electrons.

Valence electrons are the electrons that exist in the highest energy level/shell. They are the electrons that participate in bonding. Atoms want 8 total valence electrons (except H and He, which want 2), because then they would have a filled shell and neutral/low energy.

We can move around bonds and lone pairs, as long as we make sure every element has their valence shell filled. So here's the problem, how do we know how many bonds/lone pairs we should have? That's where formal charge comes in.

Formal charge helps us determine the ideal Lewis structure of a molecule. Ideally, we want every atom to have a FC of 0. This is because having a neutral charge is lower in energy, so it is the most stable state. Here's an example. Let's say you want to draw the Lewis structure for carbon dioxide, so you draw the two possible structures as shown below:

Formal charge CO2 lewis structures StudySmarter

Before we discuss the formal charge, let's do a brief refresher on Lewis structures .

The lines drawn between elements represent a bond, which contains two electrons each. In example 1, you'll see the C=O bond is a double bond, meaning it contains 4 electrons.

The "dots" near our atoms represent lone pairs.

Like I mentioned earlier, we can change the number of bonds and lone pairs so that each element has a full octet. The way we determine this number is first by finding the formal charge.

Now back to our example. Let's look at the formal charge: 1) Carbon: 0 Oxygen: 0 2) Carbon: 0 Oxygen (single): -1 Oxygen (triple): +1

Even though both have a net FC of 0, the first structure is the best option since it minimizes FC for each atom.In the "Calculating Formal Charge" section, we will go over how I got these formal charges together.

Formal Charge Formula and Equation

Now that we know what a formal charge is, let's learn how to calculate it. Here is the general formula:

$$FC=(\text{number of valence electrons})-(\text{number of lone pair electrons})-(\text{number of bonds})$$

We can look at the Lewis structure to determine the number of bonds/lone pair electrons, however, to calculate the number of valence electrons, we need to look at the periodic table.

Formal Charge The periodic table

Fig.2-The periodic table

For non-transition metals, you count from left to right, skipping over the transition metals. For example, fluorine is 7 across, so it has 7 valence electrons. The main exception to this is helium (He), which has 2 valence electrons.For transition metals, you also count from left to right. For example, vanadium (V), is 5 across, so it has 5 valence electrons

Calculating Formal Charge

Let's use our example from before to learn how we got those formal charges:

Given the diagram below, what are the formal charges for each possible Lewis structure?

Formal charge calculating CO2 formal charge StudySmarter

Let's start with the first structure:

Counting from left to right, carbon is in the 4th column in the periodic table. This means it has 4 valence electrons. Carbon is double-bonded to each oxygen, so it has 4 bonds in total. This means:

$$FC=(4)-(0)-(4)=0$$

Now for oxygen. Oxygen is in the 6th column, so it has 6 valence electrons. It is double-bonded to carbon, so it has two total bonds. It also has 2 lone pairs (4 electrons in total).

$$FC=(6)-(4)-(2)=0$$

Now for the second structure:

For single-bond oxygen:

$$FC=(6)-(6)-(1)=-1$$

For triple-bond oxygen:

$$FC=(6)-(2)-(3)=+1$$

Let's try another problem:

Given the structures below, which is the most likely structure?

Formal Charge 2O and 1 N compound structure StudySmarter

The first thing you'll probably notice is that this molecule has a charge (-1). This means that the formal charge should add up to -1.

For nitrogen: Nitrogen is in the 5th column, so it has 5 valence electrons.

$$FC=5-4-2=-1$$

For center oxygen:

$$FC=6-2-3=1$$

For right oxygen:

$$FC=6-6-1=-1$$

For nitrogen:

$$FC=5-2-3=0$$

For left oxygen:

$$FC=6-4-2=0$$

$$6-6-1=-1$$

The correct structure is the second option, since it minimizes the formal charge while keeping the net charge on the molecule, -1.

Adding Formal Charges to Resonance Forms

Sometimes when we draw Lewis structures, we may encounter resonance structures.

When two or more Lewis structures with the same arrangement of atoms and number of electrons can be written, these are called resonance structures/forms . In reality, the actual structure is an average of the different possible Lewis structures.

Molecules with the same atoms can have different orientations with different charges, but they are not resonance structures. For example: CO 2 and CO 2 - are similar, but because they have a different number of electrons, they aren't resonance structures of each other

When resonance structures have different formal charges, we can use said FC to determine the "best" structure. When we looked at CO 2 (Figure 3), we were looking at its different resonance forms, which had different formal charges. The "correct" structure is an average of the three possible forms (the third form is just the triple bond being on the opposite oxygen, so it is essentially the same as the second).

When we look at resonance structures with the same formal charge, none of the options are the "best". As an example, here are the three resonance forms of CO 3 2-

Formal Charge Carbonate resonance StudySmarter

Since the bonding is basically the same, so is the formal charge. The "true" form of carbonate is an average of the three forms, where there is a 1 1/3 bond between each oxygen and carbon.

Importance of Formal Charge

Formal charge is important for several reasons. As we discussed earlier, it is helpful for determining the best Lewis structure for both resonance and non-resonance forms.

Another reason why it is important is reactivity. By calculating the formal charge, we can determine where (if any) charges are within the molecule. This helps us understand/predict the kind of reactivity the molecule will have. For example, the right oxygen in the (correct) NO 2 structure (see Figure 4) has a -1 charge, so it can either attract positively charged atoms/molecules and/or donate electrons. Without knowing where the charge is, we can't fully understand a molecule's reactivity.

We often write the formal charge of an atom underneath it, so we can see how it will react!

Formal Charge - Key takeaways

Frequently Asked Questions about Formal Charge

--> how to calculate formal charge.

The formula for formal charge is:

FC=number of valence electrons-number of lone pair electrons-number of bonds

--> What is formal charge?

Formal charge (FC)  is the charge assigned to an atom is a molecule when we assume that electrons in all bonds are shared equally between atoms.

--> How to find formal charge from Lewis structure?

Using the Lewis structure, we can determine the number of bonds and lone pair electrons. Subtracting that from the number of valence electrons, we get the formal charge.

--> How to assign formal charges?

We use the formula for formal charge to calculate the formal charge for each individual atom. We then write that charge beneath each atom. The net charge is written on the top right of the molecule.

--> How to draw Lewis structures with formal charges ?

Formal charges tell us which Lewis structure is the ideal structure. Whichever Lewis structure has its formal charges closest  to zero is the correct structure.

Final Formal Charge Quiz

Why do electrons in bonds get divided by two when counting towards formal charge?

Show answer

Electrons in bonds are shared, so it's assumed on average each atom only has one at a time

Show question

If hydrogen has 1 bond, what is its formal charge?

0, hydrogen almost always only forms one bond and thus almost always has a formal charge of 0 in a molecule

How many resonance forms does CO2 have?

How do you find the number of electrons in a valence shell of an atom?

Looking at a periodic table to find the element and seeing what group it's in

If a carbon atom has 2 bonds and 2 electrons in a molecule, what is its formal charge?

Are molecules with the same structure but with a different amount of electrons resonance structures?

No, a different number of electrons means it's a totally different molecule

Why are formal charges typically assigned after looking at a molecule in a Lewis structure?

Using a Lewis structure allows us to find the number of bonds and lone pairs an atom will have

True or false: If two or more oxygen atoms are in a molecule, they must have the same formal charge

False! No two or more of the same element must have the same formal charge in a molecule

True or false: Formal charge can be used to determine the most commonly occurring form of a molecule?

True! This is why assigning formal charge in Lewis structures is so important

If a chlorine atom has 4 bonds and 2 lone pairs, what is its formal charge?

What is formal charge?

Formal charge (FC)  is the charge assigned to an atom is a molecule when we assume that electrons in all bonds are shared equally between atoms. 

Which of the following is NOT true about formal charge?

If there are multiples of the same element, they will have the same formal charge

What are valence electrons?

Valence electrons  are the electrons that exist in the highest energy level/shell. They are the electrons that participate in bonding .

Which of the following are reasons why we assign a formal charge? Select all that apply

To find the best Lewis structure

What are lone pairs?

Lone pairs  are a set of  valence electrons  that do not  participate in bonding . Because of this, they are also called non-bonding electrons. 

What is the formula for formal charge?

How do we use the Lewis structure(s) to determine formal charge?

Lewis structures tell us the number of bonds and lone pair electrons

What are resonance structures?

When two or more Lewis structures with the same arrangement of atoms and number of electrons can be written, these are called the  resonance structures/forms. 

True or False: For resonance structures, the "true" structure is an average of all the resonance forms.

True or False: Formal charges must  add up to 0

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What are the formal charges in #"O"_3# (ozone)?

how to calculate the formal charge of o3 pure

Consider the resonance structures for #"O"_3# .

Oxygen has #6# valence electrons. Look at the top left oxygen atom. It has two lone pairs ( #4# electrons) and a double bond ( #2# electrons).

Even though a double bond contains #4# electrons total and is counted as such when seeing that oxygen's octet is filled, #2# electrons belong to each oxygen and they are shared among the two.

Let's examine the top resonance structure:

Left: #6# valence #-6# assigned #=color(blue)(0# formal charge Center #6# valence #-5# assigned #=color(blue)(1# formal charge Right: #6# valence #-7# assigned #=color(blue)(-1)# formal charge

Notice that even though the atoms have varying formal charges, the overall charge of #"O"_3# is the sum of the formal charges in the molecule: #0+1+(-1)=0# .

Ions' formal charge sums are #!=0# .

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As per Lewis structure, can oxygen molecules in ozone be interconnected with one another through covalent single bonds?

Ozone has three oxygen molecules. Since all three are the same atom, there is no difference in electronegativity. As per theory, formation of a polar covalent bond is not possible but there are two resonance equivalent structures with one single and one covalent bond.

Image courtesy: ChemLibretext

I am adding one more detail which I found regarding this, called 'formal charge':

The formal charge of an atom in a molecule is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms. Another way of saying this is that formal charge results when we take the number of valence electrons of a neutral atom, subtract the nonbonding electrons, and then subtract the number of bonds connected to that atom in the Lewis structure.

Thus, we calculate formal charge as follows: formal charge = # valence shell electrons (free atom) − lone pair electrons −1/2*(  bonding electrons)

I calculated the formal charge for O1, O2, O3 and all three resulted in +1. Sum of the formal charges didn't give net charge of zero. I guess, that could be one of the explanations. If anyone can confirm, it will be helpful. I will meanwhile also search. Thanks for the help!

Manikandan Chandrasekaran's user avatar

Because reality, not theory, dominates chemistry

Ozone, by the way, has three oxygen atoms not molecules.

But while many theoretical structures are possible and look OK in different bonding theories, that isn't how chemists work out the real structure. Most bonding theories are too weak to make accurate structural predictions for many "difficult" structures.

When it was realised that benzene, for example, was a ring some proposed it was essentially the same as cyclohexatriene, with alternating double and single bonds. But observations of the crystal structure showed that, despite simple bonding theories ideas, the bond lengths were all equal. Better theories were developed to explain the actual structure.

Ozone, in some theories, could be a 3 membered ring or, indeed, a variety of other structures. But it isn't. Microwave spectroscopy show it to be a bent molecule as does the fact is has a dipole moment. That's what is is, theory be damned. The central angle is about 116° and the central oxygen has a net positive charge withe the terminal oxygens having net negative charges but bond lengths between those of a single and double bond.

There are theories that can account for this, but the predictions of simpler theories are useless. What we observe the structure to be is far more important than theory.

matt_black's user avatar

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COMMENTS

  1. Calculating Formal Charges for O3 (Ozone)

    In order to calculate the formal charges for O3 we'll use the equationFormal charge = [# of valence electrons] - [nonbonding val electrons]

  2. How do you calculate the formal charge of ${{O}_{3}}$ ?

    - The formal charge of the oxygen 1 is as follows. ... - The formal charge on oxygen atom 1 is zero. ... - The formal charge on oxygen atom 2 is '1'

  3. Calculate the formal charge on the central oxygen atom in O3.

    The correct option is B 1. Formal charge (FC) = V−L−B2. Where, V = Total number of valence electrons in the atom. L = Total number of non

  4. What are the formal charges in O3ozone?

    A formal charge value is equal to an atom's valence electrons deducting the number ... How do you calculate the formal charge of O3 class 11 chemistry CBSE.

  5. How To Calculate Formal Charge

    A Key Skill: How to Calculate Formal Charge · To obtain the formal charge of an atom, we start by counting the number of valence electrons [Note

  6. Formal Charge: Definition & Calculation

    1) Every atom can be assigned a formal charge · 2) If there are multiples of the same element, they can have different formal charges · 3) The formal charge is

  7. What are the formal charges in "O"_3 (ozone)?

    Explanation: ; 6 valence electrons. Look at the top left oxygen atom. It has two lone pairs ( ; 4 electrons) and a double bond ( ; 2 electrons).

  8. What is formal charge? How can you consider the ozone molecule?

    Formal charge (FC) is the charge in any given Lewis dot structure that results from the assumption that all bonds are 100% covalent (equal sharing of the

  9. As per Lewis structure, can oxygen molecules in ozone be

    I calculated the formal charge for O1, O2, O3 and all three resulted in +1. Sum of the formal charges didn't give net charge of zero.

  10. General Chemistry I, formal charge review and practice

    The formal charge on oxygen is calculated as follows. Oxygen has six valence electrons (GN = 6), two unshared electrons in one lone pair (UE = 2)