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2.6 Resonance and Formal Charge

8 min read • december 27, 2022

Dalia Savy

Sometimes, there is not just one way to represent a molecule with a Lewis structure . If a structure can be drawn in multiple ways, the structure is known to have resonance . A good way of thinking of resonance is like mixing paint🎨.

When you draw out the two or three resonance structures of a molecule, those two/three structures make up the entire molecule. The actual structure is represented by an average of these 2-3 structures. This can lead to bond orders that are fractions, such as bond orders of 1/3 or 2/3.

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-Cxh232mXtCeI.png?alt=media&token=f95fd975-9476-4476-9643-c05c910cf81d

A common misconception is that structures that have resonance are truly able to exist with different bond connections. However, resonance is all about representing a molecule in varying ways just to make it clear that the molecule is actually an average of those representations in space. This will make more sense as we take a look at several examples.

How do you know when a structure has resonance?

Let's try drawing the lewis dot structure (LDS) of the polyatomic ion NO3-.

Recall some steps for this process that we discussed in the last study guide :

Count the number of valence electrons that the molecule has in total. Nitrogen has 5 and an oxygen atom has 6. Since there are three oxygen atoms, we must account for each: 5+6+6+6 = 23. But...NO3- is a polyatomic ion and there is a charge attached to the molecule as a whole. The -1 charge indicates that there is one more electron, so there must be a total of 24 valence electrons represented in the LDS.

Draw the central atom , which in this case, is nitrogen since there is only one nitrogen atom.

Draw the 3 oxygen atoms surrounding nitrogen and three single bonds connecting them to nitrogen. Draw the full octets :

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-2jzfNcIVILMW.JPG?alt=media&token=86dca1b2-df23-4a2f-8444-910e5cbf7f96

Count the current number of valence electrons . There are 26, so somehow we have to get rid of 2 of them🤔. Remember, when there are too many electrons, we can replace a single bond and a lone pair with a double bond. In this case, if we change them all to double bonds, we won't have enough electrons! This is where resonance comes in: you only change one to a double bond.

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-gvDuPWN5xvLe.JPG?alt=media&token=998172c8-cdf4-4401-9ce3-d42688469fb1

When drawing the lewis structure of polyatomic ions, make sure you include the brackets and charge.

Now we count: we have 24! This is one way to draw the lewis dot structure of NO3-. How do we know which bond to make a double bond?

Simple answer: any of them since each of the three possibilities simultaneously exist in reality. There are three ways to draw this structure on paper:

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-vrb4XWOXlsaN.JPG?alt=media&token=aae01727-c160-4ee7-819c-67c6e14a0443

Any of these would be acceptable. On the AP Exam itself, you would be required to draw all three side by side with these specific arrows between them ↔️.

In reality...

In reality, the structure doesn't have 2 single bonds and 1 double bond. It has a 4/3 bond order . There is just no way to represent 4/3s of a bond.

Bond Order?

How did I know it was a 4/3s bond? The easiest way to know the bond order is ( # of lines / # of spaces ). There are 4 lines or bonds and 3 spaces for bonds. A 4/3s bond is between a single bond and a double bond. Therefore, it is stronger than a single bond but weaker than a double bond. A 4/3 bond is what is represented by these three structures.

👉Now try drawing the LDS of nitrite (another polyatomic ion ). The answer is at the very end of the guide!

Formal Charge

Formal charge is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms. It reflects the electron count associated with the atom compared to the isolated neutral atom. It is used to predict the correct placement of electrons.

Here is an overall diagram, but let's get into the nitty-gritty.

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-zLraipHyVgMO.png?alt=media&token=a2486ed2-53fc-43a8-8fb8-ab48e49a3400

How do you calculate formal charges?

The easiest way to calculate the formal charge of an atom is to do (# of valence electrons - # of dots - # of dashes) . This might may not be the most technical but it's the easiest way to remember it!

Look at the image above at the SCN- ion to the left and the way they calculated formal charge :

For the sulfur atom, they took the number of valence electrons (6) and subtracted it by the total number of lone electrons (2) and single bonds with carbon (3 [but there really is a triple bond]). This tells you that sulfur has a formal charge of +1 when drawn in SCN- with that representation of electrons.

When do I check the formal charge of an atom?

It is good to always check your calculation of formal charge since it is easy to miscount something. If anything, just check it if an element past element 14 is involved (because that is when you can break the octet rule).

Example with Formal Charge

Let's draw the LDS of phosphate (PO4-3):

There should be 5+6+6+6+6+3 valence electrons , so 32 total.

Draw Phosphorus in the center and 4 oxygen atoms surrounding it with single bonds and full octets :

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-5WZNwWGezjLq.JPG?alt=media&token=d60e960b-1f01-4ed2-b001-aa22e833fb10

Looking at this lewis dot structure, we count 32 valence electrons and think it is perfect! However, phosphorus is element 15 so we should be looking at its formal charge to make sure the placement of electrons is correct.

Calculating the formal charge :

P: 5 valence electrons - 0 dots - 4 bonds = +1 formal charge

O: 6 valence electrons - 6 dots - 1 bond = -1 formal charge

When we are looking at the formal charges, we want the central atom to always have a formal charge of 0 or the most electronegative atom to have a negative formal charge . A formal charge of 0 means the electrons are localized, or not moving, and that representation of the molecule is the most stable.

One way we can lower the formal charge of Phosphorus to 0 is by adding a double bond:

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-LXlbevS60fbA.JPG?alt=media&token=607bde5c-3dd7-4b48-8a0c-5960bc66cceb

Now let's check formal charge once more:

P: 5 valence electrons - 5 bonds = 0 (Perfect🥳)

Top 3 Oxygen atoms: 6 valence electrons - 6 dots - 1 bond = -1

Botton oxygen atom: 6 valence electrons - 4 dots - 2 bonds = 0

Now, instead of having no atoms with a formal charge of 0, we have two atoms with a formal charge of 0. The electrons are placed correctly, resulting in a stable bond🎉.

The polyatomic ion also has resonance now, so you should draw all 4 ways of representing it and know the bonds are 5/4 order.

Since there were 3 oxygen atoms with a -1 charge, the sum is -3, which is equal to the -3 charge of the polyatomic ion .

AP FRQ Practice Questions

The following questions are pieces of free-response questions that were on previous AP Chemistry exams. These exams can be found on the College Board website .

AP Chemistry Exam 2016 - #2e

The HCO3− ion has three carbon-to-oxygen bonds. Two of the carbon-to-oxygen bonds have the same length and the third carbon-to-oxygen bond is longer than the other two. The hydrogen atom is bonded to one of the oxygen atoms. In the box below, draw a Lewis electron-dot diagram (or diagrams) for the HCO3 − ion that is (are) consistent with the given information.

They give you a large box for you to draw your answer in, but try this on your own!

AP Chemistry Exam 2017 - #1c

S2Cl2 is a product of the reaction. In the box below, complete the Lewis electron-dot diagram for the S2Cl2 molecule by drawing in all of the electron pairs.

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-gNvgW4wRtWW1.png?alt=media&token=5be96405-1955-4b5c-81d3-aaa528d51226

AP Chemistry Exam 2017 #2a

Two possible Lewis electron-dot diagrams for fulminic acid, HCNO, are shown below. Explain why the diagram on the left is the better representation for the bonding in fulminic acid. Justify your choice based on formal charges.

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-UA91wgDPzMcB.png?alt=media&token=b15e399f-e483-4a6d-8856-e252ea59027d

2016 #2e Scoring Guidelines

You are given the ion and its charge and asked to draw the Lewis dot structure. Taking a look at HCO3-, you should calculate that there are a total of 23 valence electrons present, plus the additional electron from the negative charge: 24.

When you draw the structure out, you should place carbon in the center as the central atom . From there, always attach the multiple oxygen atoms to the central atom . If there is a hydrogen atom, you can basically bond it to any of the oxygens.

Once you have the atoms drawn and the full octets filled in, you count 26 valence electrons . We have too many electrons, so we must replace one single bond and lone pair with a double bond. Since there are two single bonds that you could convert to a double bond, this structure has resonance !

Here is an acceptable answer:

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-KfGHUTCi6ysV.png?alt=media&token=09e94811-e86f-44c9-87bb-179139bb6621

This is what College Board had on their site for full credit, but their arrow is the same as the double-headed arrow we discussed. Just please make sure you do not draw an equilibrium arrow by accident.

If one were to get this question correct, two full points would be earned:

One point is given for drawing a correct Lewis structure with brackets and a charge.

One point is given for indicating that there is resonance by drawing both structures with an arrow in between.

2017 #1c Scoring Guidelines

The following answer is correct and would earn you one full point:

https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2F-8mtzSAUZhEgc.png?alt=media&token=19613f97-1ffe-486d-9c65-c603d51c75f5

2017 #2a Scoring Guidelines

The basis of this question is calculating formal charges, and they even give you that hint! The best representation of a structure is one where the most electronegative atom has a negative formal charge .

This part was worth two points, one for the correct calculation of formal charge and one for a correct explanation of which is the better representation.

The following is a sample response:

Point #1: In the diagram on the left, the C atom has a formal charge of zero and the O atom has a formal charge of -1. In the diagram on the right, the C atom has a formal charge of -1 and the O atom has a formal charge of zero.

Point #2: The diagram on the left is the better representation because it puts the negative formal charge on oxygen, which is more electronegative than carbon.

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PO43- Lewis structure

The information on this page is ✔ fact-checked.

PO43- Lewis Structure

PO 4 3- (phosphate) has one phosphorus atom and four oxygen atoms.

In the PO 4 3- Lewis structure, there is one double bond and three single bonds around the phosphorus atom, with four oxygen atoms attached to it. One oxygen atom with a double bond has two lone pairs, and three oxygen atoms with single bonds have three lone pairs.

Also, there is a negative (-1) charge on the three oxygen atoms with single bonds.

To properly draw the PO 4 3- Lewis structure, follow these steps:

#1 Draw a rough sketch of the structure #2 Next, indicate lone pairs on the atoms #3 Indicate formal charges on the atoms, if necessary #4 Minimize formal charges by converting lone pairs of the atoms #5 Repeat step 4 if necessary, until all charges are minimized

Let’s break down each step in more detail.

#1 Draw a rough sketch of the structure

  • First, determine the total number of valence electrons

how to calculate formal charge of po43

In the periodic table , phosphorus lies in group 15, and oxygen lies in group 16.

Hence, phosphorus has five valence electrons and oxygen has six valence electrons.

Since PO 4 3- has one phosphorus atom and four oxygen atoms, so…

Valence electrons of one phosphorus atom = 5 × 1 = 5 Valence electrons of four oxygen atoms = 6 × 4 = 24

Now the PO 4 3- has a negative (-3) charge, so we have to add three more electrons.

So the total valence electrons = 5 + 24 + 3 = 32

Learn how to find: Phosphorus valence electrons and Oxygen valence electrons

  • Second, find the total electron pairs

We have a total of 32 valence electrons. And when we divide this value by two, we get the value of total electron pairs.

Total electron pairs = total valence electrons ÷ 2

So the total electron pairs = 32 ÷ 2 = 16

  • Third, determine the central atom

We have to place the least electronegative atom at the center.

Since phosphorus is less electronegative than oxygen, assume that the central atom is phosphorus .

Therefore, place phosphorus in the center and oxygens on either side.

  • And finally, draw the rough sketch

PO43- Lewis Structure (Step 1)

#2 Next, indicate lone pairs on the atoms

Here, we have a total of 16 electron pairs. And four P — O bonds are already marked. So we have to only mark the remaining twelve electron pairs as lone pairs on the sketch.

Also remember that phosphorus is a period 3 element , so it can keep more than 8 electrons in its last shell. And oxygen is a period 2 element , so it can not keep more than 8 electrons in its last shell.

Always start to mark the lone pairs from outside atoms. Here, the outside atoms are oxygens.

So for each oxygen, there are three lone pairs, and for phosphorus, there is zero lone pair because all twelve electron pairs are over.

Mark the lone pairs on the sketch as follows:

PO43- Lewis Structure (Step 2)

#3 Indicate formal charges on the atoms, if necessary

Use the following formula to calculate the formal charges on atoms:

Formal charge = valence electrons – nonbonding electrons – ½ bonding electrons

For phosphorus atom, formal charge = 5 – 0 – ½ (8) = +1

For each oxygen atom, formal charge = 6 – 6 – ½ (2) = -1

Here, both phosphorus and oxygen atoms have charges, so mark them on the sketch as follows:

PO43- Lewis Structure (Step 3)

The above structure is not a stable Lewis structure because both phosphorus and oxygen atoms have charges. Therefore, reduce the charges (as below) by converting lone pairs to bonds.

#4 Minimize formal charges by converting lone pairs of the atoms

Convert a lone pair of the oxygen atom to make a new P — O bond with the phosphorus atom as follows:

PO43- Lewis Structure (Step 4)

In the above structure, you can see that the central atom (phosphorus) forms an octet. And the outside atoms (oxygens) also form an octet. Hence, the octet rule is satisfied.

Now there is still a negative (-1) charge on the three oxygen atoms.

This is okay, because the structure with a negative charge on the most electronegative atom is the best Lewis structure. And in this case, the most electronegative element is oxygen.

Also, the above structure is more stable than the previous structures. Therefore, this structure is the most stable Lewis structure of PO 4 3- .

And since the PO 4 3- has a negative (-3) charge, mention that charge on the Lewis structure by drawing brackets as follows:

PO43- Lewis Structure (Final)

Next: I 3 – Lewis structure

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External links

  • https://lambdageeks.com/po4-3-lewis-structure/
  • https://www.chemistryscl.com/general/lewis-structure-of-PO43/index.php
  • https://techiescientist.com/po43-lewis-structure/
  • https://www.thegeoexchange.org/chemistry/bonding/Lewis-Structures/PO4-3minus-lewis-structure.html
  • https://topblogtenz.com/po43-lewis-structure-molecular-geometry-hybridization-bond-angle/
  • https://geometryofmolecules.com/po34-lewis-structure/

how to calculate formal charge of po43

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  • 7.4 Formal Charges and Resonance
  • Introduction
  • 1.1 Chemistry in Context
  • 1.2 Phases and Classification of Matter
  • 1.3 Physical and Chemical Properties
  • 1.4 Measurements
  • 1.5 Measurement Uncertainty, Accuracy, and Precision
  • 1.6 Mathematical Treatment of Measurement Results
  • Key Equations
  • 2.1 Early Ideas in Atomic Theory
  • 2.2 Evolution of Atomic Theory
  • 2.3 Atomic Structure and Symbolism
  • 2.4 Chemical Formulas
  • 2.5 The Periodic Table
  • 2.6 Ionic and Molecular Compounds
  • 2.7 Chemical Nomenclature
  • 3.1 Formula Mass and the Mole Concept
  • 3.2 Determining Empirical and Molecular Formulas
  • 3.3 Molarity
  • 3.4 Other Units for Solution Concentrations
  • 4.1 Writing and Balancing Chemical Equations
  • 4.2 Classifying Chemical Reactions
  • 4.3 Reaction Stoichiometry
  • 4.4 Reaction Yields
  • 4.5 Quantitative Chemical Analysis
  • 5.1 Energy Basics
  • 5.2 Calorimetry
  • 5.3 Enthalpy
  • 6.1 Electromagnetic Energy
  • 6.2 The Bohr Model
  • 6.3 Development of Quantum Theory
  • 6.4 Electronic Structure of Atoms (Electron Configurations)
  • 6.5 Periodic Variations in Element Properties
  • 7.1 Ionic Bonding
  • 7.2 Covalent Bonding
  • 7.3 Lewis Symbols and Structures
  • 7.5 Strengths of Ionic and Covalent Bonds
  • 7.6 Molecular Structure and Polarity
  • 8.1 Valence Bond Theory
  • 8.2 Hybrid Atomic Orbitals
  • 8.3 Multiple Bonds
  • 8.4 Molecular Orbital Theory
  • 9.1 Gas Pressure
  • 9.2 Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law
  • 9.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions
  • 9.4 Effusion and Diffusion of Gases
  • 9.5 The Kinetic-Molecular Theory
  • 9.6 Non-Ideal Gas Behavior
  • 10.1 Intermolecular Forces
  • 10.2 Properties of Liquids
  • 10.3 Phase Transitions
  • 10.4 Phase Diagrams
  • 10.5 The Solid State of Matter
  • 10.6 Lattice Structures in Crystalline Solids
  • 11.1 The Dissolution Process
  • 11.2 Electrolytes
  • 11.3 Solubility
  • 11.4 Colligative Properties
  • 11.5 Colloids
  • 12.1 Chemical Reaction Rates
  • 12.2 Factors Affecting Reaction Rates
  • 12.3 Rate Laws
  • 12.4 Integrated Rate Laws
  • 12.5 Collision Theory
  • 12.6 Reaction Mechanisms
  • 12.7 Catalysis
  • 13.1 Chemical Equilibria
  • 13.2 Equilibrium Constants
  • 13.3 Shifting Equilibria: Le Châtelier’s Principle
  • 13.4 Equilibrium Calculations
  • 14.1 Brønsted-Lowry Acids and Bases
  • 14.2 pH and pOH
  • 14.3 Relative Strengths of Acids and Bases
  • 14.4 Hydrolysis of Salts
  • 14.5 Polyprotic Acids
  • 14.6 Buffers
  • 14.7 Acid-Base Titrations
  • 15.1 Precipitation and Dissolution
  • 15.2 Lewis Acids and Bases
  • 15.3 Coupled Equilibria
  • 16.1 Spontaneity
  • 16.2 Entropy
  • 16.3 The Second and Third Laws of Thermodynamics
  • 16.4 Free Energy
  • 17.1 Review of Redox Chemistry
  • 17.2 Galvanic Cells
  • 17.3 Electrode and Cell Potentials
  • 17.4 Potential, Free Energy, and Equilibrium
  • 17.5 Batteries and Fuel Cells
  • 17.6 Corrosion
  • 17.7 Electrolysis
  • 18.1 Periodicity
  • 18.2 Occurrence and Preparation of the Representative Metals
  • 18.3 Structure and General Properties of the Metalloids
  • 18.4 Structure and General Properties of the Nonmetals
  • 18.5 Occurrence, Preparation, and Compounds of Hydrogen
  • 18.6 Occurrence, Preparation, and Properties of Carbonates
  • 18.7 Occurrence, Preparation, and Properties of Nitrogen
  • 18.8 Occurrence, Preparation, and Properties of Phosphorus
  • 18.9 Occurrence, Preparation, and Compounds of Oxygen
  • 18.10 Occurrence, Preparation, and Properties of Sulfur
  • 18.11 Occurrence, Preparation, and Properties of Halogens
  • 18.12 Occurrence, Preparation, and Properties of the Noble Gases
  • 19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds
  • 19.2 Coordination Chemistry of Transition Metals
  • 19.3 Spectroscopic and Magnetic Properties of Coordination Compounds
  • 20.1 Hydrocarbons
  • 20.2 Alcohols and Ethers
  • 20.3 Aldehydes, Ketones, Carboxylic Acids, and Esters
  • 20.4 Amines and Amides
  • 21.1 Nuclear Structure and Stability
  • 21.2 Nuclear Equations
  • 21.3 Radioactive Decay
  • 21.4 Transmutation and Nuclear Energy
  • 21.5 Uses of Radioisotopes
  • 21.6 Biological Effects of Radiation
  • A | The Periodic Table
  • B | Essential Mathematics
  • C | Units and Conversion Factors
  • D | Fundamental Physical Constants
  • E | Water Properties
  • F | Composition of Commercial Acids and Bases
  • G | Standard Thermodynamic Properties for Selected Substances
  • H | Ionization Constants of Weak Acids
  • I | Ionization Constants of Weak Bases
  • J | Solubility Products
  • K | Formation Constants for Complex Ions
  • L | Standard Electrode (Half-Cell) Potentials
  • M | Half-Lives for Several Radioactive Isotopes

Learning Objectives

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

  • Compute formal charges for atoms in any Lewis structure
  • Use formal charges to identify the most reasonable Lewis structure for a given molecule
  • Explain the concept of resonance and draw Lewis structures representing resonance forms for a given molecule

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.

Example 7.6

Calculating formal charge from lewis structures.

  • Step 2. 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.
  • Step 3. Subtract this number from the number of valence electrons for the neutral atom: I: 7 – 8 = –1 Cl: 7 – 7 = 0 The sum of the formal charges of all the atoms equals –1, which is identical to the charge of the ion (–1).

Check Your Learning

Example 7.7.

  • Step 2. Assign the lone pairs to their atom. Now each Cl atom has seven electrons and the Br atom has seven electrons.
  • Step 3. Subtract this number from the number of valence electrons for the neutral atom. This gives the formal charge: Br: 7 – 7 = 0 Cl: 7 – 7 = 0 All atoms in BrCl 3 have a formal charge of zero, and the sum of the formal charges totals zero, as it must in a neutral molecule.

N: 0; all three Cl atoms: 0

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:

  • A molecular structure in which all formal charges are zero is preferable to one in which some formal charges are not zero.
  • If the Lewis structure must have nonzero formal charges, the arrangement with the smallest nonzero formal charges is preferable.
  • Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign.
  • When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable.

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:

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: NCS – , CNS – , 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:

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).

Example 7.8

Using formal charge to determine molecular structure.

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

The number of atoms with formal charges are minimized (Guideline 2), there is no formal charge with a magnitude greater than one (Guideline 2), the negative formal charge is on the more electronegative element (Guideline 4), and the less electronegative atom is in the center position.

Notice that the more likely structure for the nitrite anion in Example 7.8 may actually be drawn in two different ways, distinguished by the locations of the N-O and N=O bonds:

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 NO 2 − NO 2 − have the same strength and length, and are identical in all other properties.

It is not possible to write a single Lewis structure for NO 2 − NO 2 − 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 NO 2 − NO 2 − 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.

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, CO 3 2− , CO 3 2− , provides a second example of resonance:

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.

Link to Learning

Use this online quiz to practice your skills in drawing resonance structures and estimating formal charges.

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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: February 21st, 2024 |

How To Calculate Formal Charge

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

  • evaluating the number of valence electrons ( VE ) the neutral atom has (e.g. 3 for boron, 4 for carbon, 5 for nitrogen, and so on).  (note: this is also equivalent to the effective nuclear charge Z eff , the number of protons that an electron in the valence orbital “sees” due to screening by inner-shell electrons.)
  • counting the number of  non-bonded valence electrons ( NBE ) on the atom. Each lone pair counts as  2 , and each unpaired electron counts as 1.
  • counting the number of  bonds ( B ) to the atom, or alternatively, counting the number of bonding electrons and dividing this by 2 .

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

  • Formal Charge
  • Simple Examples For First-Row Elements
  • Formal Charge Calculations When You Aren’t Given All The Details
  • Some Classic Formal Charge Problems
  • Formal Charges and Curved Arrows

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:

FC = VE – NBE – B

  • VE corresponds to the number of electrons around the neutral atom (3 for boron, 4 for carbon, 5 for nitrogen, 6 for oxygen, 7 for fluorine)
  • NBE corresponds to the number of non-bonded electrons around the atom (2 for a lone pair, 1 for a singly-occupied orbital, 0 for an empty orbital)
  • B is the number of  bonds around the atom (equivalent to half the number of bonding electrons)

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.

  • In the hydronium ion (H 3 O) the central atom is oxygen , which has 6 valence electrons in the neutral atom
  • The central atom has 2 unpaired electrons and 3 bonds
  • The formal charge on oxygen is [6 – 2 – 3 = +1 ] giving us H 3 O +

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.

Become a member to see the clickable quiz with answers on the back.

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.

  • With carbon, we often omit drawing hydrogens . You’re still supposed to know that they are there, and add as many hydrogens as necessary to give a full octet (or sextet, if it’s a carbocation). 
  • If there is a lone pair or unpaired electron on a carbon, it’s always drawn in .

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.

  • Bonds to hydrogen are always drawn in.
  • The lone pairs that are often omitted.
  • Nitrogen and oxygen will always have full octets. Always. [ Note 2 – OK, two exceptions ]

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.

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:

  • Positive charges on oxygen and nitrogen do not represent an empty orbital. Assume that oxygen and nitrogen have full octets! [ Note 2 ]
  • In contrast, positive charges on carbon do represent empty orbitals.

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.

  • Formal charge can be calculated using the formula FC = VE – NBE – B
  • Line drawings often omit lone pairs and C-H bonds. Be alert for these situations when calculating formal charges.
  • Positively charged carbon has an empty orbital, but assume that positively charged nitrogen and oxygen have full octets.
  • The example of the hydronium ion H 3 O+ shows the perils of relying on formal charge to understand reactivity. Pay close attention to the differences in electronegativity between atoms and draw out the dipoles to get a true sense of their reactivity.

Related Articles

  • Partial Charges Give Clues About Electron Flow
  • How To Use Electronegativity To Determine Electron Density (and why NOT to trust formal charge)
  • How to apply electronegativity and resonance to understand reactivity
  • Maybe they should call them, “Formal Wins” ?
  • Common Mistakes: Formal Charges Can Mislead

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

  • Lewis Structures
  • Ionic and Covalent Bonding
  • Chemical Kinetics
  • Chemical Equilibria
  • Valence Electrons of the First Row Elements
  • How Concepts Build Up In Org 1 ("The Pyramid")

01 Bonding, Structure, and Resonance

  • How Do We Know Methane (CH4) Is Tetrahedral?
  • Hybrid Orbitals and Hybridization
  • How To Determine Hybridization: A Shortcut
  • Orbital Hybridization And Bond Strengths
  • Sigma bonds come in six varieties: Pi bonds come in one
  • The Four Intermolecular Forces and How They Affect Boiling Points
  • 3 Trends That Affect Boiling Points
  • Introduction to Resonance
  • How To Use Curved Arrows To Interchange Resonance Forms
  • Evaluating Resonance Forms (1) - The Rule of Least Charges
  • How To Find The Best Resonance Structure By Applying Electronegativity
  • Evaluating Resonance Structures With Negative Charges
  • Evaluating Resonance Structures With Positive Charge
  • Exploring Resonance: Pi-Donation
  • Exploring Resonance: Pi-acceptors
  • In Summary: Evaluating Resonance Structures
  • Drawing Resonance Structures: 3 Common Mistakes To Avoid
  • Bond Hybridization Practice
  • Structure and Bonding Practice Quizzes
  • Resonance Structures Practice

02 Acid Base Reactions

  • Introduction to Acid-Base Reactions
  • Acid Base Reactions In Organic Chemistry
  • The Stronger The Acid, The Weaker The Conjugate Base
  • Walkthrough of Acid-Base Reactions (3) - Acidity Trends
  • Five Key Factors That Influence Acidity
  • Acid-Base Reactions: Introducing Ka and pKa
  • How to Use a pKa Table
  • The pKa Table Is Your Friend
  • A Handy Rule of Thumb for Acid-Base Reactions
  • Acid Base Reactions Are Fast
  • pKa Values Span 60 Orders Of Magnitude
  • How Protonation and Deprotonation Affect Reactivity
  • Acid Base Practice Problems

03 Alkanes and Nomenclature

  • Meet the (Most Important) Functional Groups
  • Condensed Formulas: Deciphering What the Brackets Mean
  • Hidden Hydrogens, Hidden Lone Pairs, Hidden Counterions
  • Don't Be Futyl, Learn The Butyls
  • Primary, Secondary, Tertiary, Quaternary In Organic Chemistry
  • Branching, and Its Affect On Melting and Boiling Points
  • The Many, Many Ways of Drawing Butane
  • Wedge And Dash Convention For Tetrahedral Carbon
  • Common Mistakes in Organic Chemistry: Pentavalent Carbon
  • Table of Functional Group Priorities for Nomenclature
  • Summary Sheet - Alkane Nomenclature
  • Organic Chemistry IUPAC Nomenclature Demystified With A Simple Puzzle Piece Approach
  • Boiling Point Quizzes
  • Organic Chemistry Nomenclature Quizzes

04 Conformations and Cycloalkanes

  • Staggered vs Eclipsed Conformations of Ethane
  • Conformational Isomers of Propane
  • Newman Projection of Butane (and Gauche Conformation)
  • Introduction to Cycloalkanes (1)
  • Geometric Isomers In Small Rings: Cis And Trans Cycloalkanes
  • Calculation of Ring Strain In Cycloalkanes
  • Cycloalkanes - Ring Strain In Cyclopropane And Cyclobutane
  • Cyclohexane Conformations
  • Cyclohexane Chair Conformation: An Aerial Tour
  • How To Draw The Cyclohexane Chair Conformation
  • The Cyclohexane Chair Flip
  • The Cyclohexane Chair Flip - Energy Diagram
  • Substituted Cyclohexanes - Axial vs Equatorial
  • Ranking The Bulkiness Of Substituents On Cyclohexanes: "A-Values"
  • Cyclohexane Chair Conformation Stability: Which One Is Lower Energy?
  • Fused Rings - Cis-Decalin and Trans-Decalin
  • Naming Bicyclic Compounds - Fused, Bridged, and Spiro
  • Bredt's Rule (And Summary of Cycloalkanes)
  • Newman Projection Practice
  • Cycloalkanes Practice Problems

05 A Primer On Organic Reactions

  • The Most Important Question To Ask When Learning a New Reaction
  • Learning New Reactions: How Do The Electrons Move?
  • The Third Most Important Question to Ask When Learning A New Reaction
  • 7 Factors that stabilize negative charge in organic chemistry
  • 7 Factors That Stabilize Positive Charge in Organic Chemistry
  • Nucleophiles and Electrophiles
  • Curved Arrows (for reactions)
  • Curved Arrows (2): Initial Tails and Final Heads
  • Nucleophilicity vs. Basicity
  • The Three Classes of Nucleophiles
  • What Makes A Good Nucleophile?
  • What makes a good leaving group?
  • 3 Factors That Stabilize Carbocations
  • Equilibrium and Energy Relationships
  • What's a Transition State?
  • Hammond's Postulate
  • Learning Organic Chemistry Reactions: A Checklist (PDF)
  • Introduction to Free Radical Substitution Reactions
  • Introduction to Oxidative Cleavage Reactions

06 Free Radical Reactions

  • Bond Dissociation Energies = Homolytic Cleavage
  • Free Radical Reactions
  • 3 Factors That Stabilize Free Radicals
  • What Factors Destabilize Free Radicals?
  • Bond Strengths And Radical Stability
  • Free Radical Initiation: Why Is "Light" Or "Heat" Required?
  • Initiation, Propagation, Termination
  • Monochlorination Products Of Propane, Pentane, And Other Alkanes
  • Selectivity In Free Radical Reactions
  • Selectivity in Free Radical Reactions: Bromination vs. Chlorination
  • Halogenation At Tiffany's
  • Allylic Bromination
  • Bonus Topic: Allylic Rearrangements
  • In Summary: Free Radicals
  • Synthesis (2) - Reactions of Alkanes
  • Free Radicals Practice Quizzes

07 Stereochemistry and Chirality

  • Types of Isomers: Constitutional Isomers, Stereoisomers, Enantiomers, and Diastereomers
  • How To Draw The Enantiomer Of A Chiral Molecule
  • How To Draw A Bond Rotation
  • Introduction to Assigning (R) and (S): The Cahn-Ingold-Prelog Rules
  • Assigning Cahn-Ingold-Prelog (CIP) Priorities (2) - The Method of Dots
  • Enantiomers vs Diastereomers vs The Same? Two Methods For Solving Problems
  • Assigning R/S To Newman Projections (And Converting Newman To Line Diagrams)
  • How To Determine R and S Configurations On A Fischer Projection
  • The Meso Trap
  • Optical Rotation, Optical Activity, and Specific Rotation
  • Optical Purity and Enantiomeric Excess
  • What's a Racemic Mixture?
  • Chiral Allenes And Chiral Axes
  • Stereochemistry Practice Problems and Quizzes

08 Substitution Reactions

  • Introduction to Nucleophilic Substitution Reactions
  • Walkthrough of Substitution Reactions (1) - Introduction
  • Two Types of Nucleophilic Substitution Reactions
  • The SN2 Mechanism
  • Why the SN2 Reaction Is Powerful
  • The SN1 Mechanism
  • The Conjugate Acid Is A Better Leaving Group
  • Comparing the SN1 and SN2 Reactions
  • Polar Protic? Polar Aprotic? Nonpolar? All About Solvents
  • Steric Hindrance is Like a Fat Goalie
  • Common Blind Spot: Intramolecular Reactions
  • The Conjugate Base is Always a Stronger Nucleophile
  • Substitution Practice - SN1
  • Substitution Practice - SN2

09 Elimination Reactions

  • Elimination Reactions (1): Introduction And The Key Pattern
  • Elimination Reactions (2): The Zaitsev Rule
  • Elimination Reactions Are Favored By Heat
  • Two Elimination Reaction Patterns
  • The E1 Reaction
  • The E2 Mechanism
  • E1 vs E2: Comparing the E1 and E2 Reactions
  • Antiperiplanar Relationships: The E2 Reaction and Cyclohexane Rings
  • Bulky Bases in Elimination Reactions
  • Comparing the E1 vs SN1 Reactions
  • Elimination (E1) Reactions With Rearrangements
  • E1cB - Elimination (Unimolecular) Conjugate Base
  • Elimination (E1) Practice Problems And Solutions
  • Elimination (E2) Practice Problems and Solutions

10 Rearrangements

  • Introduction to Rearrangement Reactions
  • Rearrangement Reactions (1) - Hydride Shifts
  • Carbocation Rearrangement Reactions (2) - Alkyl Shifts
  • Pinacol Rearrangement
  • The SN1, E1, and Alkene Addition Reactions All Pass Through A Carbocation Intermediate

11 SN1/SN2/E1/E2 Decision

  • Identifying Where Substitution and Elimination Reactions Happen
  • Deciding SN1/SN2/E1/E2 (1) - The Substrate
  • Deciding SN1/SN2/E1/E2 (2) - The Nucleophile/Base
  • SN1 vs E1 and SN2 vs E2 : The Temperature
  • Deciding SN1/SN2/E1/E2 - The Solvent
  • Wrapup: The Quick N' Dirty Guide To SN1/SN2/E1/E2
  • Alkyl Halide Reaction Map And Summary
  • SN1 SN2 E1 E2 Practice Problems

12 Alkene Reactions

  • E and Z Notation For Alkenes (+ Cis/Trans)
  • Alkene Stability
  • Addition Reactions: Elimination's Opposite
  • Stereoselective and Stereospecific Reactions
  • Regioselectivity In Alkene Addition Reactions
  • Stereoselectivity In Alkene Addition Reactions: Syn vs Anti Addition
  • Hydrohalogenation of Alkenes and Markovnikov's Rule
  • Hydration of Alkenes With Aqueous Acid
  • Rearrangements in Alkene Addition Reactions
  • Halogenation of Alkenes and Halohydrin Formation
  • Oxymercuration Demercuration of Alkenes
  • Hydroboration Oxidation of Alkenes
  • m-CPBA (meta-chloroperoxybenzoic acid)
  • OsO4 (Osmium Tetroxide) for Dihydroxylation of Alkenes
  • Palladium on Carbon (Pd/C) for Catalytic Hydrogenation of Alkenes
  • Cyclopropanation of Alkenes
  • A Fourth Alkene Addition Pattern - Free Radical Addition
  • Alkene Reactions: Ozonolysis
  • Summary: Three Key Families Of Alkene Reaction Mechanisms
  • Synthesis (4) - Alkene Reaction Map, Including Alkyl Halide Reactions
  • Alkene Reactions Practice Problems

13 Alkyne Reactions

  • Acetylides from Alkynes, And Substitution Reactions of Acetylides
  • Partial Reduction of Alkynes With Lindlar's Catalyst
  • Partial Reduction of Alkynes With Na/NH3 To Obtain Trans Alkenes
  • Alkyne Hydroboration With "R2BH"
  • Hydration and Oxymercuration of Alkynes
  • Hydrohalogenation of Alkynes
  • Alkyne Halogenation: Bromination, Chlorination, and Iodination of Alkynes
  • Alkyne Reactions - The "Concerted" Pathway
  • Alkenes To Alkynes Via Halogenation And Elimination Reactions
  • Alkynes Are A Blank Canvas
  • Synthesis (5) - Reactions of Alkynes
  • Alkyne Reactions Practice Problems With Answers

14 Alcohols, Epoxides and Ethers

  • Alcohols - Nomenclature and Properties
  • Alcohols Can Act As Acids Or Bases (And Why It Matters)
  • Alcohols - Acidity and Basicity
  • The Williamson Ether Synthesis
  • Ethers From Alkenes, Tertiary Alkyl Halides and Alkoxymercuration
  • Alcohols To Ethers via Acid Catalysis
  • Cleavage Of Ethers With Acid
  • Epoxides - The Outlier Of The Ether Family
  • Opening of Epoxides With Acid
  • Epoxide Ring Opening With Base
  • Making Alkyl Halides From Alcohols
  • Tosylates And Mesylates
  • PBr3 and SOCl2
  • Elimination Reactions of Alcohols
  • Elimination of Alcohols To Alkenes With POCl3
  • Alcohol Oxidation: "Strong" and "Weak" Oxidants
  • Demystifying The Mechanisms of Alcohol Oxidations
  • Protecting Groups For Alcohols
  • Thiols And Thioethers
  • Calculating the oxidation state of a carbon
  • Oxidation and Reduction in Organic Chemistry
  • Oxidation Ladders
  • SOCl2 Mechanism For Alcohols To Alkyl Halides: SN2 versus SNi
  • Alcohol Reactions Roadmap (PDF)
  • Alcohol Reaction Practice Problems
  • Epoxide Reaction Quizzes
  • Oxidation and Reduction Practice Quizzes

15 Organometallics

  • What's An Organometallic?
  • Formation of Grignard and Organolithium Reagents
  • Organometallics Are Strong Bases
  • Reactions of Grignard Reagents
  • Protecting Groups In Grignard Reactions
  • Synthesis Problems Involving Grignard Reagents
  • Grignard Reactions And Synthesis (2)
  • Organocuprates (Gilman Reagents): How They're Made
  • Gilman Reagents (Organocuprates): What They're Used For
  • The Heck, Suzuki, and Olefin Metathesis Reactions (And Why They Don't Belong In Most Introductory Organic Chemistry Courses)
  • Reaction Map: Reactions of Organometallics
  • Grignard Practice Problems

16 Spectroscopy

  • Degrees of Unsaturation (or IHD, Index of Hydrogen Deficiency)
  • Conjugation And Color (+ How Bleach Works)
  • Introduction To UV-Vis Spectroscopy
  • UV-Vis Spectroscopy: Absorbance of Carbonyls
  • UV-Vis Spectroscopy: Practice Questions
  • Bond Vibrations, Infrared Spectroscopy, and the "Ball and Spring" Model
  • Infrared Spectroscopy: A Quick Primer On Interpreting Spectra
  • IR Spectroscopy: 4 Practice Problems
  • 1H NMR: How Many Signals?
  • Homotopic, Enantiotopic, Diastereotopic
  • Diastereotopic Protons in 1H NMR Spectroscopy: Examples
  • C13 NMR - How Many Signals
  • Liquid Gold: Pheromones In Doe Urine
  • Natural Product Isolation (1) - Extraction
  • Natural Product Isolation (2) - Purification Techniques, An Overview
  • Structure Determination Case Study: Deer Tarsal Gland Pheromone

17 Dienes and MO Theory

  • What To Expect In Organic Chemistry 2
  • Are these molecules conjugated?
  • Conjugation And Resonance In Organic Chemistry
  • Bonding And Antibonding Pi Orbitals
  • Molecular Orbitals of The Allyl Cation, Allyl Radical, and Allyl Anion
  • Pi Molecular Orbitals of Butadiene
  • Reactions of Dienes: 1,2 and 1,4 Addition
  • Thermodynamic and Kinetic Products
  • More On 1,2 and 1,4 Additions To Dienes
  • s-cis and s-trans
  • The Diels-Alder Reaction
  • Cyclic Dienes and Dienophiles in the Diels-Alder Reaction
  • Stereochemistry of the Diels-Alder Reaction
  • Exo vs Endo Products In The Diels Alder: How To Tell Them Apart
  • HOMO and LUMO In the Diels Alder Reaction
  • Why Are Endo vs Exo Products Favored in the Diels-Alder Reaction?
  • Diels-Alder Reaction: Kinetic and Thermodynamic Control
  • The Retro Diels-Alder Reaction
  • The Intramolecular Diels Alder Reaction
  • Regiochemistry In The Diels-Alder Reaction
  • The Cope and Claisen Rearrangements
  • Electrocyclic Reactions
  • Electrocyclic Ring Opening And Closure (2) - Six (or Eight) Pi Electrons
  • Diels Alder Practice Problems
  • Molecular Orbital Theory Practice

18 Aromaticity

  • Introduction To Aromaticity
  • Rules For Aromaticity
  • Huckel's Rule: What Does 4n+2 Mean?
  • Aromatic, Non-Aromatic, or Antiaromatic? Some Practice Problems
  • Antiaromatic Compounds and Antiaromaticity
  • The Pi Molecular Orbitals of Benzene
  • The Pi Molecular Orbitals of Cyclobutadiene
  • Frost Circles
  • Aromaticity Practice Quizzes

19 Reactions of Aromatic Molecules

  • Electrophilic Aromatic Substitution: Introduction
  • Activating and Deactivating Groups In Electrophilic Aromatic Substitution
  • Electrophilic Aromatic Substitution - The Mechanism
  • Ortho-, Para- and Meta- Directors in Electrophilic Aromatic Substitution
  • Understanding Ortho, Para, and Meta Directors
  • Why are halogens ortho- para- directors?
  • Disubstituted Benzenes: The Strongest Electron-Donor "Wins"
  • Electrophilic Aromatic Substitutions (1) - Halogenation of Benzene
  • Electrophilic Aromatic Substitutions (2) - Nitration and Sulfonation
  • EAS Reactions (3) - Friedel-Crafts Acylation and Friedel-Crafts Alkylation
  • Intramolecular Friedel-Crafts Reactions
  • Nucleophilic Aromatic Substitution (NAS)
  • Nucleophilic Aromatic Substitution (2) - The Benzyne Mechanism
  • Reactions on the "Benzylic" Carbon: Bromination And Oxidation
  • The Wolff-Kishner, Clemmensen, And Other Carbonyl Reductions
  • More Reactions on the Aromatic Sidechain: Reduction of Nitro Groups and the Baeyer Villiger
  • Aromatic Synthesis (1) - "Order Of Operations"
  • Synthesis of Benzene Derivatives (2) - Polarity Reversal
  • Aromatic Synthesis (3) - Sulfonyl Blocking Groups
  • Birch Reduction
  • Synthesis (7): Reaction Map of Benzene and Related Aromatic Compounds
  • Aromatic Reactions and Synthesis Practice
  • Electrophilic Aromatic Substitution Practice Problems

20 Aldehydes and Ketones

  • What's The Alpha Carbon In Carbonyl Compounds?
  • Nucleophilic Addition To Carbonyls
  • Aldehydes and Ketones: 14 Reactions With The Same Mechanism
  • Sodium Borohydride (NaBH4) Reduction of Aldehydes and Ketones
  • Grignard Reagents For Addition To Aldehydes and Ketones
  • Wittig Reaction
  • Hydrates, Hemiacetals, and Acetals
  • Imines - Properties, Formation, Reactions, and Mechanisms
  • All About Enamines
  • Breaking Down Carbonyl Reaction Mechanisms: Reactions of Anionic Nucleophiles (Part 2)
  • Aldehydes Ketones Reaction Practice

21 Carboxylic Acid Derivatives

  • Nucleophilic Acyl Substitution (With Negatively Charged Nucleophiles)
  • Addition-Elimination Mechanisms With Neutral Nucleophiles (Including Acid Catalysis)
  • Basic Hydrolysis of Esters - Saponification
  • Transesterification
  • Proton Transfer
  • Fischer Esterification - Carboxylic Acid to Ester Under Acidic Conditions
  • Lithium Aluminum Hydride (LiAlH4) For Reduction of Carboxylic Acid Derivatives
  • LiAlH[Ot-Bu]3 For The Reduction of Acid Halides To Aldehydes
  • Di-isobutyl Aluminum Hydride (DIBAL) For The Partial Reduction of Esters and Nitriles
  • Amide Hydrolysis
  • Thionyl Chloride (SOCl2)
  • Diazomethane (CH2N2)
  • Carbonyl Chemistry: Learn Six Mechanisms For the Price Of One
  • Making Music With Mechanisms (PADPED)
  • Carboxylic Acid Derivatives Practice Questions

22 Enols and Enolates

  • Keto-Enol Tautomerism
  • Enolates - Formation, Stability, and Simple Reactions
  • Kinetic Versus Thermodynamic Enolates
  • Aldol Addition and Condensation Reactions
  • Reactions of Enols - Acid-Catalyzed Aldol, Halogenation, and Mannich Reactions
  • Claisen Condensation and Dieckmann Condensation
  • Decarboxylation
  • The Malonic Ester and Acetoacetic Ester Synthesis
  • The Michael Addition Reaction and Conjugate Addition
  • The Robinson Annulation
  • Haloform Reaction
  • The Hell–Volhard–Zelinsky Reaction
  • Enols and Enolates Practice Quizzes
  • The Amide Functional Group: Properties, Synthesis, and Nomenclature
  • Basicity of Amines And pKaH
  • 5 Key Basicity Trends of Amines
  • The Mesomeric Effect And Aromatic Amines
  • Nucleophilicity of Amines
  • Alkylation of Amines (Sucks!)
  • Reductive Amination
  • The Gabriel Synthesis
  • Some Reactions of Azides
  • The Hofmann Elimination
  • The Hofmann and Curtius Rearrangements
  • The Cope Elimination
  • Protecting Groups for Amines - Carbamates
  • The Strecker Synthesis of Amino Acids
  • Introduction to Peptide Synthesis
  • Reactions of Diazonium Salts: Sandmeyer and Related Reactions
  • Amine Practice Questions

24 Carbohydrates

  • D and L Notation For Sugars
  • Pyranoses and Furanoses: Ring-Chain Tautomerism In Sugars
  • What is Mutarotation?
  • Reducing Sugars
  • The Big Damn Post Of Carbohydrate-Related Chemistry Definitions
  • The Haworth Projection
  • Converting a Fischer Projection To A Haworth (And Vice Versa)
  • Reactions of Sugars: Glycosylation and Protection
  • The Ruff Degradation and Kiliani-Fischer Synthesis
  • Isoelectric Points of Amino Acids (and How To Calculate Them)
  • Carbohydrates Practice
  • Amino Acid Quizzes

25 Fun and Miscellaneous

  • A Gallery of Some Interesting Molecules From Nature
  • Screw Organic Chemistry, I'm Just Going To Write About Cats
  • On Cats, Part 1: Conformations and Configurations
  • On Cats, Part 2: Cat Line Diagrams
  • On Cats, Part 4: Enantiocats
  • On Cats, Part 6: Stereocenters
  • Organic Chemistry Is Shit
  • The Organic Chemistry Behind "The Pill"
  • Maybe they should call them, "Formal Wins" ?
  • Why Do Organic Chemists Use Kilocalories?
  • The Principle of Least Effort
  • Organic Chemistry GIFS - Resonance Forms
  • Reproducibility In Organic Chemistry
  • What Holds The Nucleus Together?
  • How Reactions Are Like Music
  • Organic Chemistry and the New MCAT

26 Organic Chemistry Tips and Tricks

  • Draw The Ugly Version First
  • Organic Chemistry Study Tips: Learn the Trends
  • The 8 Types of Arrows In Organic Chemistry, Explained
  • Top 10 Skills To Master Before An Organic Chemistry 2 Final
  • Common Mistakes with Carbonyls: Carboxylic Acids... Are Acids!
  • Planning Organic Synthesis With "Reaction Maps"
  • Alkene Addition Pattern #1: The "Carbocation Pathway"
  • Alkene Addition Pattern #2: The "Three-Membered Ring" Pathway
  • Alkene Addition Pattern #3: The "Concerted" Pathway
  • Number Your Carbons!
  • The 4 Major Classes of Reactions in Org 1
  • How (and why) electrons flow
  • Grossman's Rule
  • Three Exam Tips
  • A 3-Step Method For Thinking Through Synthesis Problems
  • Putting It Together
  • Putting Diels-Alder Products in Perspective
  • The Ups and Downs of Cyclohexanes
  • The Most Annoying Exceptions in Org 1 (Part 1)
  • The Most Annoying Exceptions in Org 1 (Part 2)
  • The Marriage May Be Bad, But the Divorce Still Costs Money
  • 9 Nomenclature Conventions To Know
  • Nucleophile attacks Electrophile

27 Case Studies of Successful O-Chem Students

  • Success Stories: How Corina Got The The "Hard" Professor - And Got An A+ Anyway
  • How Helena Aced Organic Chemistry
  • From a "Drop" To B+ in Org 2 – How A Hard Working Student Turned It Around
  • How Serge Aced Organic Chemistry
  • Success Stories: How Zach Aced Organic Chemistry 1
  • Success Stories: How Kari Went From C– to B+
  • How Esther Bounced Back From a "C" To Get A's In Organic Chemistry 1 And 2
  • How Tyrell Got The Highest Grade In Her Organic Chemistry Course
  • This Is Why Students Use Flashcards
  • Success Stories: How Stu Aced Organic Chemistry
  • How John Pulled Up His Organic Chemistry Exam Grades
  • Success Stories: How Nathan Aced Organic Chemistry (Without It Taking Over His Life)
  • How Chris Aced Org 1 and Org 2
  • Interview: How Jay Got an A+ In Organic Chemistry
  • How to Do Well in Organic Chemistry: One Student's Advice
  • "America's Top TA" Shares His Secrets For Teaching O-Chem
  • "Organic Chemistry Is Like..." - A Few Metaphors
  • How To Do Well In Organic Chemistry: Advice From A Tutor
  • Guest post: "I went from being afraid of tests to actually looking forward to them".

Comment section

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

Hello, thanks for your wonderful posts on organic chemistry. It reallys helps me to recap org chem and I really like how you explain all these topics with a bit of humor.

That said, I think in this posts may be some typos: I think there are two typos in the solution of the last quiz of chapter 3 (ID 2310): (a) In the third task [C3H7N] of the quiz, there is just one electron on the negative charged carbon. Shouldn’t there be two electrons? (b) And in the fourth task [O-CH2] the sign of the formal charge of the carbon atom should be +1 (in the calculation). (c) Note 4 says “(…) of various compounds of first-row elements.” Aren’t the shown elements in the picture from the second row of the periodic table?

Thank you very much!

  • Pingback: A Key Skill: How to Calculate Formal Charge | Straight A Mindset

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

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!

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

Nice simple explanation

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

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.

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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.

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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.

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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.

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.

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.

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 was so helpful n the best explanation about the topic…

Thanks for the easy approach.

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

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!

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

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 .

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

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

It’s neutral!

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.

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!

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

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!

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?

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?

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.

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

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.

thank you for collaboration of formal charge

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 excellent explanation

Glad you found it useful Peter!

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

Glad you found it helpful.

nice, concise explanation

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

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How to calculate formal charge

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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:

  • Determining the appropriate number of valence electrons for an atom – This can be accomplished by inspecting the periodic table. The group number indicates the appropriate number of valence electrons for each atom
  • Determining whether the atom exhibits the appropriate number of electrons – In the Lewis structure, determine whether some of the atoms show an unexpected number of electrons

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

formal-charge

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.

formal-charge-2

CH 4 , methane

CH4-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

CH3-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

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 .

If you have any questions or would like to share your reviews on the How to calculate formal charge , then comment down below. I would love to hear what you have to think.

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Lewis structure of PO43-

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Lewis structure of PO43-

The Lewis structure of PO 4 3- contains one double bond and three single bonds, with phosphorus in the center, and four oxygens on either side. The left oxygen atom has two lone pairs. The right oxygen atom, top oxygen atom, and bottom oxygen atom have three lone pairs, and the phosphorus atom does not have any lone pair.

Plus, there is a negative (-1) charge on the right oxygen atom, top oxygen atom, and bottom oxygen atom.

By using the following steps, you can easily draw the Lewis structure of PO 4 3- :

#1 Draw skeleton #2 Show chemical bond #3 Mark lone pairs #4 Calculate formal charge and check stability (if octet is already completed on central atom) #5 Convert lone pair and calculate formal charge again (if formal charges are not closer to zero)

Let’s one by one discuss each step in detail.

#1 Draw skeleton

In this step, first calculate the total number of valence electrons. And then, decide the central atom.

  • Let’s calculate the total number of valence electrons

We know that… phosphorus is a group 15 element and oxygen is a group 16 element. Hence, phosphorus has five valence electrons and oxygen has six valence electrons.

Now PO 4 3- has one phosphorus atom and four oxygen atoms.

So the total number of valence electrons = valence electrons of phosphorus atom + (valence electrons of oxygen atom × 4)

And PO 4 3- has a negative (-3) charge, so we have to add three more electrons.

Therefore, the total number of valence electrons = 5 + 24 + 3 = 32

  • Now decide the central atom

The atom with the least electronegative value is placed at the center. By looking at the periodic table, we get the electronegativity values for phosphorus and oxygen as follows:

Electronegativity value of phosphorus = 2.19 Electronegativity value of oxygen = 3.44

Obviously, phosphorus is less electronegative than oxygen. Hence, assume that phosphorus is the central atom .

So now, put phosphorus in the center and oxygens on either side. And draw the rough skeleton structure for the Lewis structure of PO 4 3- something like this:

how to calculate formal charge of po43

Also read: How to draw Lewis structure of HCl (4 steps)

#2 Show chemical bond

Place two electrons between the atoms to show a chemical bond. Since phosphorus is surrounded by four oxygens, use eight electrons to show four chemical bonds as follows:

how to calculate formal charge of po43

#3 Mark lone pairs

As calculated earlier, we have a total of 32 valence electrons. And in the above structure, we have already used eight valence electrons. Hence, twenty-four valence electrons are remaining.

Two valence electrons represent one lone pair. So twenty-four valence electrons = twelve lone pairs .

Note that phosphorus is period 3 element, so it can keep more than 8 electrons in its last shell. And oxygen is a period 2 element, so it can not keep more than 8 electrons in its last shell.

Also, make sure that you start marking these lone pairs on outside atoms first. And then, on the central atom.

The outside atoms are oxygens, so each oxygen will get three lone pairs. And the central atom (phosphorus) will not get any lone pair, because all twelve lone pairs are used.

So the Lewis structure of PO 4 3- looks something like this:

how to calculate formal charge of po43

In the above structure, you can see that the octet is completed on the central atom (phosphorus), and also on the outside atoms. Therefore, the octet rule is satisfied.

Now calculate the formal charge and check the stability of the above structure.

Also read: How to draw Lewis structure of ClO 2 – (5 steps)

#4 Calculate formal charge and check stability

The following formula is used to calculate the formal charges on atoms:

Formal charge = valence electrons – nonbonding electrons – ½ bonding electrons

Collect the data from the above structure and then, write it down below as follows:

  • For phosphorus atom

Valence electrons = 5 Nonbonding electrons = 0 Bonding electrons = 8

Formal charge = 5 – 0 – ½ (8) = +1

  • For each oxygen atom

Valence electrons = 6 Nonbonding electrons = 6 Bonding electrons = 2

Formal charge = 6 – 6 – ½ (2) = -1

Mention the formal charges of atoms on the structure. So the Lewis structure of PO 4 3- looks something like this:

how to calculate formal charge of po43

In the above structure, you can see that the formal charges of atoms are not closer to zero. Therefore, convert lone pair and calculate formal charge again.

Also read: How to draw Lewis structure of H 2 O 2 (4 steps)

#5 Convert lone pair and calculate formal charge again

As mentioned earlier, phosphorus is a period 3 element, so it can keep more than 8 electrons in its last shell.

So convert one lone pair from one oxygen atom to make a new bond with the phosphorus atom. And then, the Lewis structure of PO 4 3- looks something like this:

how to calculate formal charge of po43

Now one last thing we need to do is, calculate the formal charge again and check the stability of the above structure.

Valence electrons = 5 Nonbonding electrons = 0 Bonding electrons = 10

Formal charge = 5 – 0 – ½ (10) = 0

  • For left oxygen atom

Valence electrons = 6 Nonbonding electrons = 4 Bonding electrons = 4

Formal charge = 6 – 4 – ½ (4) = 0

  • For right oxygen , top oxygen , and bottom oxygen atom

how to calculate formal charge of po43

In the above structure, you can see that the formal charges of atoms are closer to zero. Therefore, this is the most stable Lewis structure of PO 4 3- .

And each horizontal line drawn in the above structure represents a pair of bonding valence electrons.

Now PO 4 3- is an ion having a negative (-3) charge, so draw brackets around the above Lewis structure and mention that charge on the top right corner. And then, the Lewis structure of PO 4 3- looks something like this:

how to calculate formal charge of po43

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  • Lewis structure of HCl
  • Lewis structure of ClO 2 –
  • Lewis structure of H 2 O 2
  • Lewis structure of ClO 3 –
  • Lewis structure of CH 2 Cl 2

External links

  • PO43- Lewis Structure (Phosphate ion) – Chemistry School
  • PO43- Lewis Structure: Drawings, Hybridization, Shape, Charges, Pair and Detailed Facts – Lambda Geeks
  • Drawing the Lewis Structure for PO43 – The University of Maryland
  • PO43- lewis structure, molecular geometry, hybridization, and bond angle – Topblogtenz
  • PO43- Lewis Structure, Hybridization, Polarity, and Molecular Geometry – Techiescientist
  • How many lone pairs of electrons are represented in the Lewis structure of a phosphate ion (po43-)? – Quora
  • PO43- Lewis Structure in 5 Steps (With Images) – Pediabay
  • Chemical Bonding: PO43- Lewis Structure – The Geoexchange
  • For PO43 -, phosphate ion, draw the Lewis structure (by counting valence electrons of each atom) – Studocu
  • Which of the following phosphate, PO43- Lewis structures is the best, most valid resonance structure? – Pearson
  • What is the formal charge on phosphorus in a Lewis structure for the phosphate ion (PO43-) that satisfies the octet rule? – Homework.Study.com
  • PO34- Lewis Structure, Properties and more – Geometry of Molecules
  • Draw a Lewis structure for PO43- in which the central P atom obeys the octet rule – Chegg
  • Lewis structure of phosphate ion – AceOrganicChem
  • How many lone pairs of electrons are represented in the Lewis structure of a phosphate ion (PO43-)? – Brainly

how to calculate formal charge of po43

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In PO 4 3 - , the formal charge on each oxygen atom and the P - O bond order respectively are

- 0 . 75 , 0 . 6

- 0 . 75 , 1 . 0

- 0 . 75 , 1 . 25

- 3 , 1 . 25

The correct option is C - 0 . 75 , 1 . 25 The explanation for the correct answer: Option (C) - 0 . 75 , 1 . 25 Step 1 Bond order of PO 4 3 - Bond order = Number of bonds Number of resonating structures The resonating structures of phosphate ions are, The total number of bonds is 5 The total number of resonating structures is 4 Hence, the Bond order is 5 4 = 1 . 25 Step 2 Formula charge of PO 4 3 - Formula Charge = Valence electrons − Number of non bonding electrons − No . of bonding electrons 2 For oxygen atom that forms a double bond with Phosphorus atom, The valence electrons are 6 The number of the non-bonding electrons is 4 The number of bonding electrons is 4 Formula charge = 6 - 4 + 4 2 = 6 - 4 + 2 = 0 For oxygen atom that forms a single bond with Phosphorus atom, The valence electrons are 6 The number of the non-bonding electrons is 6 The number of bonding electrons is 2 Formula charge = 6 - 6 + 2 2 = 6 - 6 + 1 = - 1 In the resonance structure, a total of - 3 charge is distributed over four oxygen atoms. Thus, the formal charge of each oxygen atom is - 3 4 = - 0 . 75 Therefore, in PO 4 3 - , the formal charge on each oxygen atom is - 0 . 75 and the P - O bond order is 1 . 25 . Hence, option (C) is the correct answer.

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PO 4 3- (Phosphate ion) Resonance Structures

Resonance structures of PO 4 3- ion can be drawn by using lewis structure of phosphate ion. four stable resonance structures can be drawn for PO 4 3- ion. These resonance structures are used to build resonance hybrid.

Lewis structure of phosphate ion (PO 4 3- )

lewis structure of PO43-

Lewis structure of PO 4 3- ion is important because it is required to draw resonance structures of phosphate ion.

Resonance structures of PO 4 3- ion

Let's draw four stable four resonance structures for the phosphate anion (NO 3 - ).

resonance structures of phosphate ion

Lone pairs, charges and bonds of PO 4 3- ion

When we draw resonance structures, we convert lone pairs to bonds and bonds to lone pairs when it is possible. When we do this, you should be careful to protect stability of ion, octal rule.

In lewis structure of PO 4 3- ion, there are three lone pairs (in the last shell) in three oxygen atoms and that oxygen atoms. Also, those oxygen atoms have a -1 charge on each atoms.

There is another oxygen atom. That oxygen atom is connected to the phosphorous atom by a double bond has two lone pairs in its last shell. Also, there is no charge in that oxygen atom.

On phosphorous atom, there are no lone pair or a charge.

Steps to draw resonance structures for PO 4 3-

You can convert a lone pair of one oxygen atom which already has three lone pairs to make a bond with phosphorous atom. With that, total electrons around phosphorous atom is going to be twelve . It is not a problem because phosphorous can keep more than eight electrons in its valence shell (phosphorous has 3d orbitals which help to keep more than eight electrons).

steps of PO43- resonance structure

With that electron transferring, number of bonds around the phosphorous atom become 6. Phosphorous atom gains -1 charge. Therefore, this is not a best resonance structure because oxygen should gain the negative charges (electronegativity of oxygen is higher than phosphorous).

To obtain a stable resonance structure, convert a bond of former double bond as in the figure.

how to draw stable resonance structures of phosphate ion

Now, phosphorous do not have a charge and three oxygen atom has charges. As like this, we can convert lone pairs to bonds and vice versa to draw four resonance structures of phosphate ion.

Four resonance structure of PO 4 3-

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What is the Charge on PO4 (Phosphate ion)? And Why?

Charge on PO4 (Phosphate ion)

The Charge of PO4 (Phosphate ion) is 3- .

But the question is how can you find the charge on PO4 ( phosphate ion )?

Well there are 2 methods by which you can find the charge of PO4.

Lets dive right into these methods one by one.

If you are a visual learner like me, then here is a short one minute video for you.

Method 1: By looking at what it is bonded to

The charge of PO4 (Phosphate ion) can be found out by looking at what it is bonded to.

So let’s take some examples of compounds that contain PO 4 ; like H 3 PO 4 , K 3 PO 4 , etc.

Example 1: H 3 PO 4 In H 3 PO 4 , the PO 4 is bonded to Hydrogen (H). You know that the ionic charge of H is 1+. So you can easily say that the charge of PO 4 should be 3-, then only it will get canceled out. Hence the charge of PO 4 in H 3 PO 4 is 3-.

Example 2: K 3 PO 4 In K 3 PO 4 , the PO 4 is bonded to Potassium (K). And again, you know that the ionic charge of K is 1+. So here also you can easily say that the charge of PO 4 should be 3-, then only it will get canceled out. Hence the charge of PO 4 in K 3 PO 4 is 3-.

As seen from the above examples, The charge of PO4 is 3- .

In this way, you can easily find the charge of PO4 by looking at what it is bonded to.

Method 2: By calculating the formal charge using lewis structure

In order to calculate the formal charge on PO4 (Phosphate ion), you should know the Lewis dot structure of PO4 (Phosphate ion) .

Here is the lewis structure of PO4.

how to calculate formal charge of po43

Now using the above lewis structure of PO4, you have to find the formal charge on each atom that is present in the PO4 molecule.

For calculating the formal charge, you need to remember this formula;

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

You can see the bonding and nonbonding electrons of PO4 from the image given below.

how to calculate formal charge of po43

So now let’s calculate the formal charge on each individual atom present in PO4.

Formal charge on Phosphorus atom: Valence electrons = 5 (as it is in group 15 on periodic table) [1] Nonbonding electrons = 0 Bonding electrons = 10

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

Formal charge on Phosphorus = Valence electrons – Nonbonding electrons – (Bonding electrons)/2 = 5 – 0 – (10/2) = 0

how to calculate formal charge of po43

So the formal charge on phosphorus atom is 0.

Formal charge on double bonded Oxygen: Valence electron = 6 (as it is in group 16 on periodic table) [2] Nonbonding electrons = 4 Bonding electrons = 4

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

how to calculate formal charge of po43

So the formal charge on double bonded oxygen atom is 0.

Formal charge on single bonded Oxygen : Valence electron = 6 (as it is in group 16 on periodic table) Nonbonding electrons = 6 Bonding electrons = 2

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

how to calculate formal charge of po43

So the formal charge on single bonded oxygen atom is 1-.

Now let’s put all these charges on the lewis dot structure of PO4.

how to calculate formal charge of po43

So there is overall 3- charge left on the entire molecule.

This indicates that the PO4 (Phosphate ion) has 3- charge .

I hope you have understood the above calculations of PO4 (Phosphate ion). But for your tests, you don’t need to remember the entire calculations. You should just try to remember that PO4 has 3- charge.

Check out some other related topics for your practice.

Related topics: Charge of Silver (Ag) Charge of Zinc (Zn) Charge of Aluminum (Al) Charge of Sodium (Na) Charge of Chlorine (Cl)  

how to calculate formal charge of po43

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4.3: Formal Charge and Oxidation State (Problems)

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PROBLEM \(\PageIndex{1}\)

Determine the formal charge and oxidation state of each element in the following:

a. HCl b. CF 4 c. PCl 3

FC: H: 0, Cl: 0

OX: H: +1, Cl: -1

FC: C: 0, F: 0

OX: C: +4, F: -1

FC: P: 0, Cl: 0

OX: P: +3, Cl: -1

PROBLEM \(\PageIndex{2}\)

a. H 3 O + b. \(\ce{SO4^2-}\) c. NH 3 d. \(\ce{O2^2-}\) e. H 2 O 2

FC: H: 0, O: +1

OX: H: +1, O: -2

FC: S: +2, O: -1

OX: S: +6, O: -2

FC: N: 0, H: 0

OX: N: -3, H: +1

FC: O: 0, H: 0

OX: O: -1, H: +1

PROBLEM \(\PageIndex{3}\)

Calculate the formal charge and oxidation state of chlorine in the molecules Cl 2 and CCl 4 .

FC: Cl in Cl 2 : 0; Cl in CCl 4 : 0

OX: Cl in Cl 2 : 0; Cl in CCl 4 : -1

PROBLEM \(\PageIndex{4}\)

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

a. F 2 CO b. NO – c. \(\ce{BF4-}\) d. \(\ce{SnCl3-}\) e. H 2 CCH 2 f. \(\ce{PO4^3-}\)

FC: C: 0, F :0, O: 0

OX: C: +4, F: -1, O: -2

FC: N: -1, O: 0

OX: N: +1, O: -2

FC: B: -1, F: 0

OX: B: +3, F: -1

FC: Sn: -1; Cl: 0

OX: Sn: +2, Cl: -1

FC: C: 0, H: 0

OX: C: -2, H: +1

FC: P: +1, O: -1

OX: P: +5, O: -2

PROBLEM \(\PageIndex{5}\)

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

PROBLEM \(\PageIndex{6}\)

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

PROBLEM \(\PageIndex{7}\)

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

PROBLEM \(\PageIndex{8}\)

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:

PROBLEM \(\PageIndex{9}\)

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

The first structure is the best structure. the formal charges are closest to 0 (and also the second structure does not give a complete octet on N)

Contributors

Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors.  Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/[email protected] ).

  • Adelaide Clark, Oregon Institute of Technology

Think one of the answers above is wrong? Let us know here .

IMAGES

  1. PO43- Formal charge, How to calculate it with images?

    how to calculate formal charge of po43

  2. PO43- Formal charge, How to calculate it with images?

    how to calculate formal charge of po43

  3. Calculating PO43- Formal Charges: Formal Charges for the Phosphate Ion

    how to calculate formal charge of po43

  4. PO43- Formal charge, How to calculate it with images?

    how to calculate formal charge of po43

  5. How to Calculate the Formal Charges for PO4 3- (Phosphate ion)

    how to calculate formal charge of po43

  6. PO43- Formal charge, How to calculate it with images?

    how to calculate formal charge of po43

VIDEO

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  5. Formal charge of carbonate ion #shorts #short

  6. Formal charge. calculate formal charge

COMMENTS

  1. How to Calculate the Formal Charges for PO4 3- (Phosphate ion)

    Subscribed 252 38K views 5 years ago In order to calculate the formal charges for PO4 3- we'll use the equation: Formal charge = [# of valence electrons] - [nonbonding val electrons] -...

  2. Calculating PO43- Formal Charges: Formal Charges for the ...

    In order to calculate the formal charges for PO43-- we'll use the equationFormal charge = [# of valence electrons] - [nonbonding val electrons] - [bonding ...

  3. Formal Charge

    The formal charge of any atom in a molecule can be calculated by the following equation: FC = V − N − B 2 (1) (1) F C = V − N − B 2. where V is the number of valence electrons of the neutral atom in isolation (in its ground state); N is the number of non-bonding valence electrons on this atom in the molecule; and B is the total number ...

  4. 2.3: Formal Charges

    Adding together the formal charges on the atoms should give us the total charge on the molecule or ion. In this case, the sum of the formal charges is 0 + 1 + 0 + 0 + 0 = 1+, which is the same as the total charge of the ammonium polyatomic ion. Exercise 2.3.1 2.3. 1. Write the formal charges on all atoms in BH−4 BH 4 −.

  5. PO43- Lewis Structure (Phosphate ion)

    Total valance electrons pairs = σ bonds + π bonds + lone pairs at valence shells Total electron pairs are determined by dividing the number total valence electrons by two. For, PO 43- ion, Total pairs of electrons are 16. Center atom of PO 43- ion To be the center atom, ability of having greater valance is important.

  6. AP Chem Unit 2.6 Resonance & Formal Charge

    The easiest way to calculate the formal charge of an atom is to do (# of valence electrons - # of dots - # of dashes). This might may not be the most technical but it's the easiest way to remember it! ... Example with Formal Charge. Let's draw the LDS of phosphate (PO4-3): There should be 5+6+6+6+6+3 valence electrons, so 32 total.

  7. PO43- Lewis structure

    Steps To properly draw the PO 43- Lewis structure, follow these steps: #1 Draw a rough sketch of the structure #2 Next, indicate lone pairs on the atoms #3 Indicate formal charges on the atoms, if necessary #4 Minimize formal charges by converting lone pairs of the atoms #5 Repeat step 4 if necessary, until all charges are minimized

  8. 6.5: Formal Charges and Resonance

    Each Cl atom now has seven electrons assigned to it, and the I atom has eight. Subtract this number from the number of valence electrons for the neutral atom: I: 7 - 8 = -1. Cl: 7 - 7 = 0. The sum of the formal charges of all the atoms equals -1, which is identical to the charge of the ion (-1). Exercise 6.5.1 6.5. 1.

  9. 7.4 Formal Charges and Resonance

    Step 1. We divide the bonding electron pairs equally for all I-Cl bonds: Step 2. 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. Step 3. Subtract this number from the number of valence electrons for the neutral atom: I: 7 - 8 = -1.

  10. How To Calculate Formal Charge

    How To Calculate Formal Charge. To calculate the formal charge of an atom, we start by:. evaluating the number of valence electrons (VE) the neutral atom has (e.g. 3 for boron, 4 for carbon, 5 for nitrogen, and so on). (note: this is also equivalent to the effective nuclear charge Z eff, the number of protons that an electron in the valence orbital "sees" due to screening by inner-shell ...

  11. How to calculate formal charge

    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.

  12. 7.54

    Calculate the formal charge of each element in the following compounds and ions:(a) F2CO(b) NO -(c) BF4 -(d) SnCl3 -(e) H2CCH2(f) ClF3(g) SeF6(h) PO4 3−OpenS...

  13. PO43- Formal charge, How to calculate it with images?

    Bonding electrons around Phosphorus = 3 single bonds + 1 double bond = 3 (2) + 1 (4) = 10 electrons Non-bonding electrons on Phosphorus = no lone pair = 0 electrons Formal charge on the Phosphorus atom = 5 - 0 - 10/2 = 5 - 0 - 5 = 5 - 5 = 0 ∴ The formal charge on the Phosphorus (P) atom in [PO4]3- is 0. For single-bonded oxygen atoms

  14. Lewis structure of PO43- Root Memory

    Steps. By using the following steps, you can easily draw the Lewis structure of PO 4 3-: #1 Draw skeleton #2 Show chemical bond #3 Mark lone pairs #4 Calculate formal charge and check stability (if octet is already completed on central atom) #5 Convert lone pair and calculate formal charge again (if formal charges are not closer to zero). Let's one by one discuss each step in detail.

  15. In PO43 , the formal charge on each oxygen atom and the P O ...

    Byju's Answer Standard XII Chemistry Lewis Kossel Structure In PO43-, the... Question In PO 4 3 -, the formal charge on each oxygen atom and the P - O bond order respectively are A - 0. 75, 0. 6 B - 0. 75, 1. 0 C - 0. 75, 1. 25 D - 3, 1. 25 Solution The correct option is C - 0. 75, 1. 25 The explanation for the correct answer:

  16. Lewis Structure of PO4 3- (With 5 Simple Steps to Draw!)

    5 Steps to Draw the Lewis Structure of PO43- Step #1: Calculate the total number of valence electrons Here, the given ion is PO4 3-. In order to draw the lewis structure of PO4 3- ion, first of all you have to find the total number of valence electrons present in the PO4 3- ion.

  17. 2.3: Formal Charges

    Chemistry 350: Organic Chemistry I 2: Polar Covalent Bonds; Acids and Bases 2.3: Formal Charges

  18. Phosphate Ion (PO₄³⁻)

    Oxidation-Reduction. Phosphate is a very weak oxidizing agent. Since the phosphorus is in its highest oxidation state in phosphate ion, this ion cannot act as a reducing agent. This page titled Phosphate Ion (PO₄³⁻) is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by James P. Birk. Halide Ions (Cl⁻, Br ...

  19. PO43- (Phosphate ion) Resonance Structures

    There is another oxygen atom. That oxygen atom is connected to the phosphorous atom by a double bond has two lone pairs in its last shell. Also, there is no charge in that oxygen atom. On phosphorous atom, there are no lone pair or a charge. Steps to draw resonance structures for PO 43-

  20. What is the Charge on PO4 (Phosphate ion)? And Why?

    As seen from the above examples, The charge of PO4 is 3-. In this way, you can easily find the charge of PO4 by looking at what it is bonded to. Method 2: By calculating the formal charge using lewis structure In order to calculate the formal charge on PO4 (Phosphate ion), you should know the Lewis dot structure of PO4 (Phosphate ion).

  21. Formal Charges in Lewis Structures

    To find formal charges in a Lewis structure, for each atom, you should count how many electrons it "owns". Count all of its lone pair electrons, and half of its bonding electrons. The difference between the atom's number of valence electrons and the number it owns is the formal charge. For example, in NH 3, N has 1 lone pair (2 electrons) and 3 ...

  22. Formal charge on oxygen (video)

    Note I'm using a different method to calculate formal charge from Jay, I feel this one shows you where the numbers come from better. Remember each bond is 2 electrons, and each lone pair is 2 electrons. Formal charge = # of valence electrons - # of lone pair electrons - # of bonding electrons/2 2 bonds and 2 lone pairs = 6 - 4 - 4/2 = 0 formal ...

  23. 4.3: Formal Charge and Oxidation State (Problems)

    The first structure is the best structure. the formal charges are closest to 0 (and also the second structure does not give a complete octet on N) Contributors Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors.