how to calculate the formal charge of o3 electron

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

how to calculate the formal charge of o3 electron

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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O3 Lewis Structure, Molecular Geometry, Hybridization, and MO Diagram

We have all heard of the ozone layer depletion, haven’t we? Due to vast global warming and the rapid increase of temperature on earth, the ozone layer of the stratosphere has a hole in it.

This causes severe climate change and environmental damage. A pale blue gas with a molar mass of 47.99 g/ml, this molecular compound is often termed the activated oxygen.

It can be useful for killing bacterial growth and also releases a pungent smell.

Formed from the dioxide molecule, this molecule having three oxygen atoms is very crucial from the chemistry point of view. If you want to dive into his molecule, let us fasten our seat belts!

Because I am going to make you travel through all the essential concepts and explanations related to bonding within ozone.

Lewis Structure

To be very precise, Lewis Structure is the name given to the structural representation of a molecule. It is the diagrammatic layout for understanding the nitty-gritty of chemical bonding.

A very essential concept of molecular chemistry, the following steps dictate how you can successfully draw Lewis Structure:

The initial step towards forming this structure is to find out the total number of valence electrons .

‘+’ stands for positive charge i.e giving away(loss) of electrons.

‘-’ stands for the gain of electrons, or in other words, negative charge.

While calculating the valence electrons, we need to work with these two signs.

We now need to determine the central atom . How can we do so? We can easily find the solution to this with one simple trick!

First, point out the least electronegative atom. You can check this out by calculating the valence number. The one with the highest valence usually has the least electronegativity.

This atom will consist of higher sites of bonding compared to the others.

In this step, the task is to visualize the position of single bonds present in the molecule as a whole to the central atom.

This is carried out by sketching the skeleton diagram of the respective molecules as per requirements

Do you know that when atoms contain less than eight electrons in their outermost electron shell, they are still in their reactive state?

Hence, they react accordingly and tend to form more stable molecular compounds. So, the octet rule is based on the fact that every atom should have eight electrons in its valence shell.

The fourth step of Lewis Structure formation is based on achieving this. Starting with the electropositive ones, slowly fulfill the octet of the atoms.

Once, octet fulfillment has been done, we now need to find out if bond formation is left. Accordingly, multiple bond formation can be done.

We are now done with all kinds of bond formation, The last step is focused on the formal charge concept.

The bond formation (single and multiple ones) leads us to the final step of the process, i.e. the sixth step. Here, we will focus on calculating the formal charge.

We need to check whether all the atoms inside the given molecule are maintained at their least formal charge.

Below is the formula for formal charge:

Lewis Structure of O3

Here, we will be dealing with ozone, the molecular formula is O3.

The below discussion, therefore, will be based on finding out the Lewis Structure of O3.

Ozone consists of three oxygen atoms. Oxygen belongs to group VI of the periodic table with an atomic no of 8. It thus has 6 valence electrons.

Thus, the total number of valence electrons in ozone= 3*6 = 18

Just like triiodide ion where all the atoms are iodine, here, all the atoms are oxygen. So, we will just consider one of the three to be the central atom and place the other two on lateral sides.

Now, we will draw the skeletal structure of ozone based on step no. 3. While drawing, let us now place the valence shell electrons( total count=18) for octet fulfillment.

As discussed above, we can place six electrons surrounding each oxygen as per the periodic table knowledge. Have a look at the above diagram now.

You can see that the side oxygen atoms have both achieved octet. They both have eight electrons surrounding them. But, the central atom only has six electrons around itself.

So, to fulfill the octet rule, what we need to do is:

We have to shift two electrons from either one lateral oxygen atom and place it beside the central oxygen atom.

The octet formation is now successful.

We now have one double bond and one single bond concerning the central O atom.

Since we could have drawn, either way, we now have resonance structures. After checking the formal charge, the final Lewis Structure or electron dot structure of O3 has been done.

Hybridization of O3

What do we mean by hybridization? Why is this such a common topic in chemical bonding and why do we need to study this?

Well, hybridization is one of the vast and major topics in molecular chemistry.

It refers to the process of intermixing orbitals to form hybrid orbitals. How and why several atoms tend to combine with one another forms the basis of hybridization.

Hence, the study is important to know more in detail about a molecule and its properties.

To learn about ozone, therefore, we need to have some knowledge about its hybridization.

Now, how can we find out here?

How many electrons does the central oxygen atom have? 8.

2s orbital has two electrons and the rest six are present in 2px and 2py.

Total no of orbitals= 1s and 2p

Hence, the hybridization of the O3 molecule is sp2.

Molecular Geometry of O3

To find out the molecular geometry of ozone we need to check the VSEPR theory model.

How to proceed with this?

At the very beginning, you have to check the terminal atoms.

Terminal atom no.= 2

Now, find out the no of lone electrons or lone pairs

Lone electrons=2 in the central atom

Lone pair=1

After this, we have to match this with the VSEPR model graph

The ozone molecule is found to be bent trigonal planar shape due to the presence of resonance. Repulsion causes the bond angle to come to about 116 degrees.

The polarity concept is based on the distribution of positive and negative charges inside a molecule surrounding the constituent atoms.

The dipole moment is used to measure or calculate the polarity. This only has a net value when there is a charge difference.

In the case of ozone, the usual dipole moment value varies and there is the presence of partial + and + charges inside the molecule.

The ozone atom at the center will carry the partial+ + charge.

Dipole moments then are responsible for moving the ozone molecule in a downward direction The lone electron pair results in a net dipole in O3 and hence the ozone molecule is considered to be polar in nature.

O3 Molecular Orbital Diagram (MO)

The molecular orbital theory is one of the major revolutionary concepts of chemical bonding.

It uses quantum mechanics to give us a detailed almost explanatory diagram of the bonding nature inside a molecule.

Here is a diagrammatic representation of the MO diagram of ozone.

Ozone is a trigonal planar molecule. Hence, as we take one p orbital from each atom of oxygen(O3), we focus on the 4 electron H3- anion.

According to hybrid orbital approximation, we will consider the 2s, 2py, and 2pz orbitals. Then we will perform semi-empirical Molecular Orbital calculations and involve the concept of group orbitals.

Chemical bonding is indeed one of the vast and yet very important chapters for the entire life. If you are here to learn small details about any molecule, you have come to the right place.

Now that you have gone through the major concepts of the ozone molecule in detail, I hope you have got a basic understanding.

We have discussed Lewis structure and Hybridization. We have mentioned the shape and the polarity of the molecule and also the MO Diagram. So, did you have a good time?

I think so. Always keep learning.

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This article illustrates the o3 lewis structure , How to draw the lewis structure, hybridization, shapes, charges, and some other properties of the o3 lewis structure.

The highly reactive gas ozone (o3) naturally exists in our atmosphere.

How to draw lewis structure for O3?

Draw the o3 lewis structure using different steps:

Step 1: calculate the total no. of Valence electron in the o3 lewis structure:

Step 2: draw the skeleton of the given molecule and find the central atom:, step 3: assign the valence electron:.

Thus total valence electrons are 18 in the o3 lewis structure molecule, assign them in between the three atoms of o3 with the help of dots. Complete the octet of each atom that is around the central atom.

Step 5: Perform the following octet for the central atom:

O3 lewis structure shape.

Due to the distortion in the o3 lewis structure , the O3 molecule shape is frequently said to be bent.

O3 lewis structure formal charges

The formal charge on O: 6-2-½ (6)= 1, thus the formal charge on o3 lewis structure is +1 on the central oxygen atomSimilarly, two adjacent oxygen atoms carry (-½ ) partial negative charge, and central oxygen carries +1 formal charge as shown in the figure below.

O3 lewis structure lone pairs

O3 hybridization.

Hybridization refers to joining two or more atomic levels with the same or different energies to form a new orbital. Ozone has only one central oxygen atom with eight electrons in its outermost shell.

O3 lewis structure resonance

Except for the placement of the electrons, these Lewis structures are identical.

From a stability standpoint, these structures are identical because each has a positive and negative formal charge on two oxygen atoms.

O3 lewis structure octet rule

O3 polar or nonpolar.

The ozone molecule is thought to be polar because the lone electron pair produces a net dipole. The typical dipole moment value changes in an ozone molecule, and the molecule contains partial positive and negative charges.

O3 lewis structure bond angle

A double bond exists between the central oxygen atom and one of the lateral oxygen atoms in the O3 Lewis structure. In the O3 Lewis structure, a double bond exists between the central oxygen atom and one of the lateral oxygen atoms .

O3 lewis structure electron geometry

But the molecular geometry is bent.

O3 valence electrons

To calculate the total number of valence electrons provided by a given element, multiply the number of electrons in the valance shell by the number of atoms in that element.

An oxygen atom has six electrons in its valence shell. There are three oxygen atoms in an O3 molecule. As a result, the ozone molecule has a total of 18 valence electrons.

Conclusion:

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How to calculate formal charge

How to calculate formal charge Examples

ot all atoms within a neutral molecule need be neutral. An atom can have the following charges: positive , negative , or neutral , depending on the electron distribution. This is often useful for understanding or predicting reactivity. Identifying formal charges helps you keep track of the electrons.

The formal charge is the charge on the atom in the molecule. The term “formal” means that this charge is not necessarily on the presented atom because in some cases, it is also prevalent on other atoms present in the molecule. It is actually spread out through the other atoms and is not only on the one atom. Identifying a formal charge involves:

The formal charge on an atom can be calculated using the following mathematical equation.

Lewis structures also show how atoms in the molecule are bonded. They can be drawn as lines (bonds) or dots (electrons). One line corresponds to two electrons . The nonbonding electrons, on the other hand, are the unshared electrons and these are shown as dots. One dot is equal to one nonbonding electron. The valence electrons are the electrons in the outermost shell of the atom.

CH 4 , methane

A number of non-bonding electrons: 0 for both H and C

[ Formal charge ] H = 1 – (1/2) × 2 – 0 = 0 ⇒ This applies to each hydrogen. These hydrogens are all zero.

[ Formal charge ] C = 4 – (1/2) × 8 – 0 = 0

⇒ This molecule is neutral .

CH 3 + , methyl cation

[ Formal charge ] H = 1 – (1/2) × 2 – 0 = 0 ⇒ This applies to each hydrogen. These hydrogens are all zero. [ Formal charge ] C = 4 – (1/2) × 6 – 0 = 4 – 3 – 0 = +1

⇒ This is a cation .

CH 3 – , methyl cation

A number of non-bonding electrons: 0 for H, 2 for C

[ Formal charge ] C = 4 – (1/2) × 6 – 2 = 4 – 3 – 2 = -1

⇒ This is a anion .

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How to Calculate Formal Charge Khan Academy

In order to calculate the formal charge of an atom, you will need to know the number of valence electrons on the atom. The number of valence electrons can be found on the periodic table. Once you have determined the number of valence electrons, you will need to subtract the number of non-bonding electrons and half of the bonding electrons from the total number of valence electrons. This will give you the formal charge on the atom.

Table of Contents

Formal charge | Molecular and ionic compound structure and properties | AP Chemistry | Khan Academy

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How to Calculate Formal Charge Formula

When it comes to calculating formal charge, there is a simple formula that you can follow. This will ensure that you get the correct answer every time. The first thing that you need to do is determine the total number of valence electrons in the molecule. This can be done by looking at the periodic table or using an online tool. Once you have determined the total number of valence electrons, you need to subtract the number of non-bonding electrons from this total. This will give you the number of bonding electrons in the molecule. Finally, you need to subtract the total number of lone pairs from the total number of bonding electrons. This will give you your final answer for formal charge.

Khan Academy Formal Charge Practice

In chemistry, formal charge is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. Formal charge is important because it allows chemists to predict the predominant structure or ionic character of a molecule or polyatomic ion. It also provides insight into possible electron-sharing arrangements (resonance structures) in molecules and how those resonance structures contribute to the overall stability of the molecule. The concept of formal charge is based on Lewis dot structures, which are two-dimensional representations of molecules and ions that show how valence electrons are arranged around atoms. In a Lewis dot structure, each dot represents one valence electron. The number of dots equals the number of valence electrons for that particular atom. For example, carbon has four valence electrons, so its Lewis dot structure would look like this: C . Formal charge is calculated by subtracting the number of valence electrons on an atom from the total number of electrons surrounding that atom in a Lewis dot structure. The resulting number is the formal charge on that atom. For example, consider this Lewis dot structure for carbon dioxide: O=C=O . In this molecule, each oxygenatom has sixvalenceelectrons (three lone pairs and one bond), whilecarbonhasfourvalenceelectrons(two bondsand two lone pairs). Therefore, we can calculate the formal charges as follows: Formal Charge on Oxygen = 6 – (3 + 1) = 2 Formal Charge on Carbon = 4 – (2 + 2) = 0 Therefore, in this molecule there are two oxygen atoms with a formal chargeof+2andonecarbonatomwithaformalchargeof0 . This particular arrangement results in what is called acarbonateion . Keepinmindthatintheseexamplesweareneglectingthepresenceofanynon-bondingelectrons(also known as “lone pairs”). Now let’s look at another example where we will take into account non-bonding electrons.

How to Calculate Formal Charge on Resonance Structures

When it comes to chemistry, resonance is an important concept to understand. In short, resonance occurs when a molecule has more than one possible structure that can be drawn. The actual structure of the molecule is a hybrid of these different structures. In order to calculate the formal charge on a resonance structure, you’ll need to follow a few steps. First, determine which atoms are bonded and which are not. Then, assign electrons to each atom according to its electronegativity. Finally, use the following equation: Formal Charge = Number of Valence Electrons – (Number of Bonds + Number of Unshared Electrons) Let’s walk through an example using methane, CH₄. In this molecule, all four hydrogen atoms are bonded to the carbon atom. We also know that carbon has 4 valence electrons and hydrogen has 1 valence electron. So our equation would look like this: Formal Charge (carbon) = 4 – (4 + 0) = 0 Formal Charge (hydrogen) = 1 – (1 + 0) = 0 All in all, we can see that the formal charge on each atom in this methane molecule is zero!

How to Calculate Formal Charge of O3

O3, or ozone, is a molecule made up of three oxygen atoms. The formal charge of a molecule is the sum of the charges on all of its atoms. In order to calculate the formal charge of O3, we need to add up the charges on each of its atoms. The first oxygen atom has a charge of -2, the second has a charge of 0, and the third has a charge of +1. When we add these up, we get a formal charge of -1 for O3.

How to Determine Formal Charge from Lewis Structure

In chemistry, formal charge is the overall charge on an atom in a molecule if we were to assume that electrons in all chemical bonds are shared equally between atoms. This is different from the actual charge on an atom, which takes into account the individual electron pairs that each atom has. To determine the formal charge of an atom, first draw out the Lewis structure of the molecule. Then, count up the number of valence electrons on each atom. Formal charge is calculated by subtracting the number of valence electrons from the number of lone pair electrons: Formal Charge = Number of Valence Electrons – Number of Lone Pair Electrons For example, consider carbon dioxide (CO2).

How To Calculate Formal Charge Khan Academy

Credit: www.khanacademy.org

How Do You Calculate the Formal Charge?

In order to calculate the formal charge on an atom, you must first determine the number of valence electrons on the atom. This can be done by looking at the position of the element on the periodic table. Once you have determined the number of valence electrons, you must then subtract the number of non-bonding electrons and half of the bonding electrons from that total. The resulting number is the formal charge on the atom. For example, let’s say we want to calculate the formal charge on carbon. Carbon has 4 valence electrons. If we subtract 2 non-bonding electrons (1s2) and 1/2 of 2 bonding electrons (sp2), we are left with a formal charge of +1 on carbon.

What is Formal Charge And How is It Calculated?

In chemistry, formal charge is the overall charge of an atom in a molecule or polyatomic ion where the valence electrons have been divided up according to the rules: # In Lewis structures, covalent bonds are shared equally between atoms. Formal charges can be calculated by assigning dots for lone pairs (non-bonded electrons) and crosses for unpaired electrons. The number of dots minus crosses equals the formal charge on that atom. # Where there are multiple Lewis structures with different arrangements of lone pairs and unpaired electrons, the one with the lowest formal charges on all atoms is generally preferred. The concept of formal charge is useful in predicting trends in reactivity and bonding properties in molecules and ions. Compounds tend to be more stable when they have a lower overall formal charge. This is because molecules or ions with higher formal charges tend to be more reactive, since they have a greater tendency to lose electrons and form cations or to gain electrons and form anions.

What is Formal Charge Ap Chem?

In chemistry, the formal charge of an atom is the hypothetical charge that an atom would have if we completely ionized all of its bonds and assigned electrons to atoms according to the octet rule. The formal charge on an atom can be calculated by subtracting the number of valence electrons on the free atom from the total number of electrons around the nucleus. For example, carbon has 4 valence electrons. If we add 2 for each C-H bond and 1 for each double bond, we get a total of 8 electrons around the nucleus. This means that the formal charge on carbon is 4 – 8 = -4. Formal charges are important in predicting the structure and reactivity of molecules. They can help us understand why some molecules are more stable than others and why some reactions occur more readily than others.

Why Do We Calculate the Formal Charge?

When we are trying to determine the most likely structure for a molecule, it is important to take into account the formal charges on each atom. The formal charge of an atom is the difference between the number of valence electrons on an uncharged atom and the number of electrons assigned to that atom in the Lewis structure. By taking into account the formal charges on each atom, we can see which Lewis structure is more stable. In general, a molecule will be more stable if the atoms have a lower formal charge. There are several reasons why we calculate the formal charge: 1) To help us determine the most likely Lewis structure for a molecule 2) To see which Lewis structure is more stable

In order to calculate the formal charge on an atom, you’ll need to know the number of valence electrons on that atom. The valence electrons are the outermost electrons on an atom, and they’re what determine how atoms interact with each other. Once you have the number of valence electrons, you can subtract the number of bonds that atom has from that number. The resulting number is the formal charge on that atom.

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Formal Charge Calculator

Enter the number of valence electrons, number of lone and bonded pairs. The calculator will readily determine the formal charge.

No. of Valence Electrons:

No. of Lone Pair Electrons:

No. of Bound Electrons:

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An online formal charge calculator is exclusively designed to calculate formal charge of an atom. Moreover, it should be kept in mind that we must have a sound knowledge of the formal charge chemistry before you start using this free calculator. So, without getting late, let’s dive in!

What Is A Formal Charge?

A formal charge is defined as:

“The individual charge of each atom present in a molecule”

Determining the formal charge is of great significance. This is because it shows how reactive a molecule is and how it will behave while creating bonds with either atoms.

Lewis Structure:

This particular structure displays the bonding electron pairs of atoms in a molecule. Also, the presence of lone pairs in the molecules is also shown using this pattern.

In Lewis Structure:

No matter how complex the dot structure of a molecule is, the free online formal charge calculator assists you in calculating charge in a fragment of seconds.

The Lewis structures of carbon dioxide is below:

Lone Pair Electrons:

“A specific extra pair of electrons that does not take part in the bonding scheme is known as the lone pair”.

In practice, lone pairs of electrons are not usually shown. However, it is mandatory to remember them as well.

Bonding Pair Electrons:

“The electron pairs that do take part in the bond formation are known as the bonding pair of electrons”.

For example:

The bond pair of electrons is Nitrogen Trichloride molecule are represented as “-” and lone pair as “:”

Formal Charge Formula:

You can calculate the formal charge of any atom with the help of the equation below:

$$ FC = V – \left(LP + 0.5BE\right) $$

FC = Formal Charge on Atom V = Number of Valence Electrons LP = Lone Pair Electrons BE = Number of Bonded Electrons

How To Calculate Formal Charge?

Here we will be solving a couple of examples to make your concept even more stronger than before. Let’s start!

Example # 01:

We are seeking to calculate the formal charge of sulphur in sulphur dioxide molecule. How to find formal charge from Lewis structure?

The lewis structure of the sulphur dioxide containing all the bonding and lone pair electrons is as follows:

Valence shell electrons of sulphur = 6 Lone pair electrons of sulphur = 2 Bonded electrons of sulphur = 8

Using formal charge equation:

$$ FC = V – \left(LP + 0.5BE\right) $$ $$ FC = 6 – \left(2 + 0.5*8\right) $$ $$ FC = 6 – \left(2 + 4\right) $$ $$ FC = 6 – 6 $$ $$ FC = 0 $$

Example # 02:

How to determine formal charge on fluorine?

The dot structure of fluorine is as below:

Valence shell electrons of sulphur = 7 Lone pair electrons of sulphur = 6 Bonded electrons of sulphur = 2

Calculating formal charge:

$$ FC = V – \left(LP + 0.5BE\right) $$ $$ FC = 7 – \left(6 + 0.5*2\right) $$ $$ FC = 7 – \left(6 + 1\right) $$ $$ FC = 7 – \left(7\right) $$ $$ FC = 7 – 7 $$ $$ FC = 0 $$

Except for the formula, our free online formal charge calculator takes a couple of seconds to generate accurate results.

Example # 03:

How to find formal charge on oxygen in $CH_{3}O^-$ ion?

The lewis structure of $CH_{3}O^-$ is as follows:

Carrying out formal charge calculation:

The number of bonding electrons oxygen in $CH_{3}O^-$ = 2 The number of valence shell electrons in $CH_{3}O^-$ = 6 Lone pair electrons = 6

Using formal charge formula:

$$ FC = V – \left(LP + 0.5BE\right) $$ $$ FC = 6 – \left(6 + 0.5*2\right) $$ $$ FC = 6 – 7 $$ $$ FC = -1 $$

How Formal Charge Calculator Works?

Make a use of this free calculator to calculate formal charge on any atom contained within the molecule. Let’s find how!

The free formal charge calculator calculates:

The exact formal charge for the respective atom in the molecule

What is the formal charge rule?

According to the formal charge rule, the number of the electrons shared among the atoms of the molecule must be in equilibrium. Moreover, this knowledge also provides an edge in determining the true Lewis structures.

What is a negative formal charge?

When an atom donates more than a few electrons and there is an octet shell vacant in them, then there exists a possibility of gaining more electrons to fill the octet shell. This gain of electrons results in the negative formal charge on an atom.

Are formal charges real charges?

No, the formal charges are not the real charges. They are just used to suppose the electrons transmission in between atoms for lewis structures.

What is the best formal charge?

In practice, the formal charge of “0” is very best as it shows the atoms have their shells completely filled while bonding with each other in a molecule. Also, it is also known as the lowest formal charge.

Does formal charge affect polarity?

No, the formal charge can never have any effect on the polarity. The formal charges are always assumed for the perfect covalent bonds which means there is no chance of electrons to be shared unequally. This cancels out any chance of polarity within the molecule.

Does formal charge determine the dipole moment?

Yes, of course! The formal charge do determine the dipole moment

Conclusion:

In chemical analysis, the formal charge is very crucial in predicting and establishing a reaction mechanism. As well, it helps you in keeping a proper track of the electron sharing among atoms. That is why using an online formal charge calculator help you in perfect formal charge calculations to avoid any hurdle during any chemical reaction.

References:

From the source of Wikipedia: Chemical bond , Strong chemical bonds, Intermolecular bonding,

From the source of Khan Academy: Types of Bonds, Ionic Bonds , Covalent Bonds, Polar covalent bonds, Comparison Of Covalent And Ionic Compounds

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What is the Charge of O3 (Ozone)? And Why? if(typeof ez_ad_units!='undefined'){ez_ad_units.push([[468,60],'knordslearning_com-box-3','ezslot_5',132,'0','0'])};__ez_fad_position('div-gpt-ad-knordslearning_com-box-3-0');

So let’s calculate the formal charge of O3 (Ozone).

Calculating the formal charge of O3 using lewis structure

Formal charge = Valence electrons – Nonbonding electrons – (Bonding electrons)/2

Formal charge on left Oxygen = Valence electrons – Nonbonding electrons – (Bonding electrons)/2 = 6 – 4 – (4/2) = 0

So the formal charge on left oxygen atom is 0.

So according to the formula of formal charge, you will get;

Formal charge on right Oxygen = Valence electrons – Nonbonding electrons – (Bonding electrons)/2 = 6 – 6 – (2/2) = 1-

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

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Home / A Key Skill: How to Calculate Formal Charge

Bonding, Structure, and Resonance

By James Ashenhurst

A Key Skill: How to Calculate Formal Charge

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.

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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.

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.

Pingback: Come posso calcolare l’addebito formale? – GufoSaggio

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.

Pingback: ¿Cómo puedo calcular el cargo formal? – ElbuhoSabio

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.

Pingback: Como posso calcular a cobrança formal? – CorujaSabia

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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Organic chemistry

Unit 2: lesson 1.

Comparing formal charges to oxidation states

Formal charge on carbon

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Video transcript

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Chemical Bonding and Molecular Geometry

Formal Charges and Resonance

OpenStaxCollege

[latexpage]

Learning Objectives

By the end of this section, you will be able to:

In the previous section, we discussed how to write Lewis structures for molecules and polyatomic ions. As we have seen, however, in some cases, there is seemingly more than one valid structure for a molecule. We can use the concept of formal charges to help us predict the most appropriate Lewis structure when more than one is reasonable.

Calculating 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:

We can double-check formal charge calculations by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion.

We must remember that the formal charge calculated for an atom is not the actual charge of the atom in the molecule. Formal charge is only a useful bookkeeping procedure; it does not indicate the presence of actual charges.

Calculating Formal Charge from Lewis Structures Assign formal charges to each atom in the interhalogen ion ${\text{ICl}}_{4}{}^{\text{−}}.$

A Lewis structure is shown. An iodine atom with two lone pairs of electrons is single bonded to four chlorine atoms, each of which has three lone pairs of electrons. Brackets surround the structure and there is a superscripted negative sign.

We assign lone pairs of electrons to their atoms . Each Cl atom now has seven electrons assigned to it, and the I atom has eight.

I: 7 – 8 = –1

Cl: 7 – 7 = 0

Check Your Learning Calculate the formal charge for each atom in the carbon monoxide molecule:

A Lewis structure is shown. A carbon atom with one lone pair of electrons is triple bonded to an oxygen with one lone pair of electrons.

Calculating Formal Charge from Lewis Structures Assign formal charges to each atom in the interhalogen molecule BrCl 3 .

A Lewis structure is shown. A bromine atom with two lone pairs of electrons is single bonded to three chlorine atoms, each of which has three lone pairs of electrons.

Assign the lone pairs to their atom. Now each Cl atom has seven electrons and the Br atom has seven electrons.

Br: 7 – 7 = 0

Check Your Learning Determine the formal charge for each atom in NCl 3 .

N: 0; all three Cl atoms: 0

A Lewis structure is shown. A nitrogen atom with one lone pair of electrons is single bonded to three chlorine atoms, each of which has three lone pairs of electrons.

Using Formal Charge to Predict Molecular Structure

The arrangement of atoms in a molecule or ion is called its molecular structure . In many cases, following the steps for writing Lewis structures may lead to more than one possible molecular structure—different multiple bond and lone-pair electron placements or different arrangements of atoms, for instance. A few guidelines involving formal charge can be helpful in deciding which of the possible structures is most likely for a particular molecule or ion:

To see how these guidelines apply, let us consider some possible structures for carbon dioxide, CO 2 . We know from our previous discussion that the less electronegative atom typically occupies the central position, but formal charges allow us to understand why this occurs. We can draw three possibilities for the structure: carbon in the center and double bonds, carbon in the center with a single and triple bond, and oxygen in the center with double bonds:

Three Lewis structures are shown. The left and right structures show a carbon atom double bonded to two oxygen atoms, each of which has two lone pairs of electrons. The center structure shows a carbon atom that is triple bonded to an oxygen atom with one lone pair of electrons and single bonded to an oxygen atom with three lone pairs of electrons. The third structure shows an oxygen atom double bonded to another oxygen atom with to lone pairs of electrons. The first oxygen atom is also double bonded to a carbon atom with two lone pairs of electrons.

Comparing the three formal charges, we can definitively identify the structure on the left as preferable because it has only formal charges of zero (Guideline 1).

As another example, the thiocyanate ion, an ion formed from a carbon atom, a nitrogen atom, and a sulfur atom, could have three different molecular structures: CNS – , NCS – , or CSN – . The formal charges present in each of these molecular structures can help us pick the most likely arrangement of atoms. Possible Lewis structures and the formal charges for each of the three possible structures for the thiocyanate ion are shown here:

Two rows of structures and numbers are shown. The top row is labeled, “Structure” and depicts three Lewis structures and the bottom row is labeled, “Formal charge.” The left structure shows a carbon atom double bonded to a nitrogen atom with two lone electron pairs on one side and double bonded to a sulfur atom with two lone electron pairs on the other. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative one, zero, and zero. The middle structure shows a carbon atom with two lone pairs of electrons double bonded to a nitrogen atom that is double bonded to a sulfur atom with two lone electron pairs. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative two, positive one, and zero. The right structure shows a carbon atom with two lone electron pairs double bonded to a sulfur atom that is double bonded to a nitrogen atom with two lone electron pairs. The structure is surrounded by brackets and has a superscripted negative sign. Below this structure are the numbers negative two, positive two, and one.

Note that the sum of the formal charges in each case is equal to the charge of the ion (–1). However, the first arrangement of atoms is preferred because it has the lowest number of atoms with nonzero formal charges (Guideline 2). Also, it places the least electronegative atom in the center, and the negative charge on the more electronegative element (Guideline 4).

Using Formal Charge to Determine Molecular Structure Nitrous oxide, N 2 O, commonly known as laughing gas, is used as an anesthetic in minor surgeries, such as the routine extraction of wisdom teeth. Which is the likely structure for nitrous oxide?

Two Lewis structures are shown with the word “or” in between them. The left structure depicts a nitrogen atom with two lone pairs of electrons double bonded to a nitrogen that is double bonded to an oxygen with two lone pairs of electrons. The right structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom that is double bonded to a nitrogen atom with two lone pairs of electrons.

Solution Determining formal charge yields the following:

The structure with a terminal oxygen atom best satisfies the criteria for the most stable distribution of formal charge:

A Lewis structure is shown. A nitrogen atom with two lone pairs of electrons is double bonded to a nitrogen atom that is double bonded to an oxygen atom with two lone pairs of electrons.

The number of atoms with formal charges are minimized (Guideline 2), and there is no formal charge larger than one (Guideline 2). This is again consistent with the preference for having the less electronegative atom in the central position.

Check Your Learning Which is the most likely molecular structure for the nitrite $\left({\text{NO}}_{2}{}^{\text{−}}\right)$ ion?

Two Lewis structures are shown with the word “or” written between them. The left structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure and there is a superscripted negative sign. The right structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen with three lone pairs of electrons. Brackets surround this structure and there is a superscripted negative sign.

You may have noticed that the nitrite anion in [link] can have two possible structures with the atoms in the same positions. The electrons involved in the N–O double bond, however, are in different positions:

Two Lewis structures are shown. The left structure shows an oxygen atom with three lone pairs of electrons single bonded to a nitrogen atom with one lone pair of electrons that is double bonded to an oxygen with two lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign. The right structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign.

If nitrite ions do indeed contain a single and a double bond, we would expect for the two bond lengths to be different. A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms. Experiments show, however, that both N–O bonds in ${\text{NO}}_{2}{}^{\text{−}}$ have the same strength and length, and are identical in all other properties.

It is not possible to write a single Lewis structure for ${\text{NO}}_{2}{}^{\text{−}}$ in which nitrogen has an octet and both bonds are equivalent. Instead, we use the concept of resonance : if two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an average of that shown by the various Lewis structures. The actual distribution of electrons in each of the nitrogen-oxygen bonds in ${\text{NO}}_{2}{}^{\text{−}}$ is the average of a double bond and a single bond. We call the individual Lewis structures resonance forms . The actual electronic structure of the molecule (the average of the resonance forms) is called a resonance hybrid of the individual resonance forms. A double-headed arrow between Lewis structures indicates that they are resonance forms. Thus, the electronic structure of the ${\text{NO}}_{2}{}^{\text{−}}$ ion is shown as:

Two Lewis structures are shown with a double headed arrow drawn between them. The left structure shows an oxygen atom with two lone pairs of electrons double bonded to a nitrogen atom with one lone pair of electrons that is single bonded to an oxygen atom with three lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign. The right structure shows an oxygen atom with three lone pairs of electrons single bonded to a nitrogen atom with one lone pair of electrons that is double bonded to an oxygen atom with two lone pairs of electrons. Brackets surround this structure, and there is a superscripted negative sign.

We should remember that a molecule described as a resonance hybrid never possesses an electronic structure described by either resonance form. It does not fluctuate between resonance forms; rather, the actual electronic structure is always the average of that shown by all resonance forms. George Wheland, one of the pioneers of resonance theory, used a historical analogy to describe the relationship between resonance forms and resonance hybrids. A medieval traveler, having never before seen a rhinoceros, described it as a hybrid of a dragon and a unicorn because it had many properties in common with both. Just as a rhinoceros is neither a dragon sometimes nor a unicorn at other times, a resonance hybrid is neither of its resonance forms at any given time. Like a rhinoceros, it is a real entity that experimental evidence has shown to exist. It has some characteristics in common with its resonance forms, but the resonance forms themselves are convenient, imaginary images (like the unicorn and the dragon).

The carbonate anion, ${\text{CO}}_{3}{}^{\text{2−}},$ provides a second example of resonance:

Three Lewis structures are shown with double headed arrows in between. Each structure is surrounded by brackets, and each has a superscripted two negative sign. The left structure depicts a carbon atom bonded to three oxygen atoms. It is single bonded to two of these oxygen atoms, each of which has three lone pairs of electrons, and double bonded to the third, which has two lone pairs of electrons. The double bond is located between the lower left oxygen atom and the carbon atom. The central and right structures are the same as the first, but the position of the double bonded oxygen has moved to the lower right oxygen in the central structure and to the top oxygen in the right structure.

One oxygen atom must have a double bond to carbon to complete the octet on the central atom. All oxygen atoms, however, are equivalent, and the double bond could form from any one of the three atoms. This gives rise to three resonance forms of the carbonate ion. Because we can write three identical resonance structures, we know that the actual arrangement of electrons in the carbonate ion is the average of the three structures. Again, experiments show that all three C–O bonds are exactly the same.

The online Lewis Structure Make includes many examples to practice drawing resonance structures.

Key Concepts and Summary

In a Lewis structure, formal charges can be assigned to each atom by treating each bond as if one-half of the electrons are assigned to each atom. These hypothetical formal charges are a guide to determining the most appropriate Lewis structure. A structure in which the formal charges are as close to zero as possible is preferred. Resonance occurs in cases where two or more Lewis structures with identical arrangements of atoms but different distributions of electrons can be written. The actual distribution of electrons (the resonance hybrid) is an average of the distribution indicated by the individual Lewis structures (the resonance forms).

Key Equations

$\text{formal charge}\phantom{\rule{0.2em}{0ex}}=\text{# valence shell electrons (free atom)}\phantom{\rule{0.1em}{0ex}}-\phantom{\rule{0.2em}{0ex}}\text{# one pair electrons}\phantom{\rule{0.2em}{0ex}}-\phantom{\rule{0.2em}{0ex}}\frac{1}{2}\phantom{\rule{0.3em}{0ex}}\text{# bonding electrons}$

Chemistry End of Chapter Exercises

Write resonance forms that describe the distribution of electrons in each of these molecules or ions.

(a) selenium dioxide, OSeO

(b) nitrate ion, ${\text{NO}}_{3}{}^{\text{−}}$

(d) benzene, C 6 H 6 :

A Lewis structure shows a hexagonal ring composed of six carbon atoms. They form single bonds to each another and single bonds to one hydrogen atom each.

(e) the formate ion:

A Lewis structure shows a carbon atom single bonded to two oxygen atoms and a hydrogen atom. The structure is surrounded by brackets and there is a superscripted negative sign.

(a) sulfur dioxide, SO 2

(b) carbonate ion, ${\text{CO}}_{3}{}^{\text{2−}}$

(d) pyridine:

A Lewis structure depicts a hexagonal ring composed of five carbon atoms and one nitrogen atom. Each carbon atom is single bonded to a hydrogen atom.

(e) the allyl ion:

A Lewis structure shows a carbon atom single bonded to two hydrogen atoms and a second carbon atom. The second carbon atom is single bonded to a hydrogen atom and a third carbon atom. The third carbon atom is single bonded to two hydrogen atoms. The whole structure is surrounded by brackets, and there is a superscripted negative sign.

Write the resonance forms of ozone, O 3 , the component of the upper atmosphere that protects the Earth from ultraviolet radiation.

Sodium nitrite, which has been used to preserve bacon and other meats, is an ionic compound. Write the resonance forms of the nitrite ion, ${\text{NO}}_{\text{2}}{}^{\text{–}}\text{.}$

Two pairs of Lewis structures are shown with a double-headed arrow in between each pair. The left structure of the first pair shows a nitrogen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. It is also double bonded to an oxygen with two lone pairs of electrons. The right image of this pair depicts the mirror image of the left. Both images are surrounded by brackets and a superscripted negative sign. They are labeled, “For N O subscript two superscript negative sign.” The left structure of the second pair shows an oxygen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. It is also double bonded to an oxygen atom with two lone pairs of electrons. The right structure appears as a mirror image of the left. These structures are labeled, “For O subscript three.”

In terms of the bonds present, explain why acetic acid, CH 3 CO 2 H, contains two distinct types of carbon-oxygen bonds, whereas the acetate ion, formed by loss of a hydrogen ion from acetic acid, only contains one type of carbon-oxygen bond. The skeleton structures of these species are shown:

Two Lewis structures are shown with a double headed arrow in between. The left structure shows a carbon atom single bonded to three hydrogen atoms and a second carbon atom. The second carbon is single bonded to two oxygen atoms. One of the oxygen atoms is single bonded to a hydrogen atom. The right structure, surrounded by brackets and with a superscripted negative sign, depicts a carbon atom single bonded to three hydrogen atoms and a second carbon atom. The second carbon atom is single bonded to two oxygen atoms.

Write the Lewis structures for the following, and include resonance structures where appropriate. Indicate which has the strongest carbon-oxygen bond.

This structure shows a carbon atom double bonded to two oxygen atoms, each of which has two lone pairs of electrons.

Toothpastes containing sodium hydrogen carbonate (sodium bicarbonate) and hydrogen peroxide are widely used. Write Lewis structures for the hydrogen carbonate ion and hydrogen peroxide molecule, with resonance forms where appropriate.

Determine the formal charge of each element in the following:

(a) H: 0, Cl: 0; (b) C: 0, F: 0; (c) P: 0, Cl 0; (d) P: 0, F: 0

(a) H 3 O +

(b) ${\text{SO}}_{4}{}^{\text{2−}}$

(d) ${\text{O}}_{2}{}^{\text{2−}}$

(e) H 2 O 2

Calculate the formal charge of chlorine in the molecules Cl 2 , BeCl 2 , and ClF 5 .

Cl in Cl 2 : 0; Cl in BeCl 2 : 0; Cl in ClF 5 : 0

Calculate the formal charge of each element in the following compounds and ions:

(d) ${\text{SnCl}}_{3}{}^{\text{−}}$

(e) H 2 CCH 2

(h) ${\text{PO}}_{4}{}^{\text{3−}}$

Draw all possible resonance structures for each of these compounds. Determine the formal charge on each atom in each of the resonance structures:

(d) ${\text{NO}}_{3}{}^{\text{−}}$

Two Lewis structures are shown with a double-headed arrow in between. The left structure shows an oxygen atom with one lone pair of electrons single bonded to an oxygen atom with three lone pairs of electrons. It is also double bonded to an oxygen atom with two lone pairs of electrons. The symbols and numbers below this structure read, “( 0 ), ( positive 1 ), ( negative 1 ).” The phrase, “Formal charge,” and a right-facing arrow lie to the left of this structure. The right structure appears as a mirror image of the left and the symbols and numbers below this structure read, “( negative 1 ), ( positive 1 ), ( 0 ).”

Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in nitrosyl chloride: ClNO or ClON?

Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in hypochlorous acid: HOCl or OClH?

Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in sulfur dioxide: OSO or SOO?

Draw the structure of hydroxylamine, H 3 NO, and assign formal charges; look up the structure. Is the actual structure consistent with the formal charges?

The structure that gives zero formal charges is consistent with the actual structure:

A Lewis structure shows a nitrogen atom with one lone pair of electrons single bonded to two hydrogen atoms and an oxygen atom which has two lone pairs of electrons. The oxygen atom is single bonded to a hydrogen atom.

Iodine forms a series of fluorides (listed here). Write Lewis structures for each of the four compounds and determine the formal charge of the iodine atom in each molecule:

Write the Lewis structure and chemical formula of the compound with a molar mass of about 70 g/mol that contains 19.7% nitrogen and 80.3% fluorine by mass, and determine the formal charge of the atoms in this compound.

A Lewis structure shows a nitrogen atom with one lone pair of electrons single bonded to three fluorine atoms, each with three lone pairs of electrons.

Which of the following structures would we expect for nitrous acid? Determine the formal charges:

Two Lewis structures are shown, with the word “or” in between. The left structure shows a nitrogen atom single bonded to an oxygen atom with three lone pairs of electrons. It is also single bonded to a hydrogen atom and double bonded to an oxygen atom with two lone pairs of electrons. The right structure shows a hydrogen atom single bonded to an oxygen atom with two lone pairs of electrons. The oxygen atom is single bonded to a nitrogen atom which is double bonded to an oxygen atom with two lone pairs of electrons.

Sulfuric acid is the industrial chemical produced in greatest quantity worldwide. About 90 billion pounds are produced each year in the United States alone. Write the Lewis structure for sulfuric acid, H 2 SO 4 , which has two oxygen atoms and two OH groups bonded to the sulfur.

A Lewis structure shows a hydrogen atom single bonded to an oxygen atom with two lone pairs of electrons. The oxygen atom is single bonded to a sulfur atom. The sulfur atom is double bonded to two oxygen atoms, each of which have three lone pairs of electrons, and single bonded to an oxygen atom with two lone pairs of electrons. This oxygen atom is single bonded to a hydrogen atom.

Formal Charges and Resonance by OpenStaxCollege is licensed under a Creative Commons Attribution 4.0 International License , except where otherwise noted.

Organic Chemistry

Formal Charges

Formal charges in organic chemistry is, perhaps, one of the most fundamental bookkeeping devices which is often misunderstood or neglected by students.

Why Formal Charges are Important in Organic Chemistry?

Knowing formal charges can help us understand the reactivity patterns in reactions, find reactive centers, and make sense out of electron flow in the mechanisms. For instance, negatively charged species tend to be the sources of electron density in reactions, while the positively charged species—accept those electrons.

Here is a couple of examples of molecules with formal charges:

The top species has a negative charge. We call such species anions. Since it has a negative charge, it means that it has an excess of electron density. Thus, it is likely to be the source of those electrons in a reaction. Sources of electron density in an organic reaction are acting as nucleophiles or bases.

The second species bears a positive charge, thus it is a cationic species. In a reaction, a cationic species will be an acceptor of electrons acting as an acid (either Brønsted or Lewis) or as an electrophile.

It’s important to keep in mind that a formal charge is not the same thing as an actual charge ! I’ll talk about it a little later in this post.

How to Calculate a Formal Charge

The “official” way is to subtract 1/2 bonding electrons and nonbonding electrons from possible valence electrons that an atom can have. In other words:

calculating formal charge by the definition

What are all those terms?

Bonding electrons are those that make up bonds. Each bond contains 2 electrons. So, 1/2 of bonding electrons equals to the number of bonds and atom has.

Nonbonding electrons are those that are not participating in any bonding. In other words, nonbonding electrons are the spare electrons (usually electron pairs) on an atom. Using the examples from above we have:

The definitions and the “official” method looks a little ugly. As a professional chemist I can talk all day about the “official rules” and “proper names” and bore you to death. Instead, I’d rather you use a simple “trick” that always works and, essentially, is the same thing. The trick is:

Formal charge = Valence Electrons – Sticks – Dots

The number of valence electrons equals to the element’s group (column) in the periodic table. This way, carbon has 4, oxygen has 6, and hydrogen has 1 valence electrons. The bonds and the spare electrons will be indicated (or can be easily found from) the molecule’s Lewis structure. So, for as long as you have a complete Lewis structure and periodic table handy, you can quickly find the formal charge of any atom in a molecule.

The Difference Between the Formal and Actual Charge

Now, I’ve mentioned earlier that there’s a difference between the formal and the actual charge. Formal charge is a bookkeeping tool that is important to help us keep track of the electron flow in the reaction.

The actual charge , however, is the actual electron density that is present on the atom. For instance, let’s take a look at borohydride anion:

The electronegativity of boron is 2.0 while electronegativity of hydrogen is 2.2. So, the hydrogen is more electronegative (not by much but still) and will polarize the bond. This means that hydrogen actually “pulls” the electron density towards itself. Thus:

Formal charge in borohydride ion vs the actual charge distribution in the anion

While formal charges are merely a “formality,” they are very important for the reactions mechanisms understanding. Thus you need to make sure you master the skill of quickly finding the formal charge.

You also notice that I’ve indicated my real electron densities with the delta-minus (𝛿-) symbol. That denotes that I only have a partial negative charge on each of the hydrogens. How much of that partial charge we have on them? Well, we could calculate it using fancy quantum chemical calculations, but that’s utterly unnecessary for the purposes of a typical organic chemistry course. What’s more important, is to realize that the boron is not actually negatively charged in this molecule. So, when we write a reaction with a borohydride anion, we won’t be showing electrons coming from boron!

Notice how in the example above my arrow starts at the H-B bond and not at the boron atom! That’s because it’s not the boron that is a source of electron density here! You remember from earlier in this post, it’s the hydrogens are the atoms with the 𝛿-. So now it makes sense that the arrow doesn’t start at boron.

how to calculate the formal charge of o3 electron

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