Formal charge practice

Get some practice with these formal charge practice problems! Click here to download the PDF version .

Assign a formal charge to all atoms and determine the overall charge of the molecule. All lone pairs and hydrogens attached to carbon are shown. (Click on the picture to zoom in!)

Part 1 formal charge practice problem. Hand-drawn molecules with lone pairs shown but formal charge missing.

Click here to reveal the answers

Draw in any lone pairs and any hydrogens attached to carbon. If the formal charge for an atom is not indicated, it is assumed to be zero. (Click on the picture to zoom in!)

Part 2 formal charge practice problem. Hand-drawn molecules drawn correctly but lone pairs not shown.

If you're seeing this message, it means we're having trouble loading external resources on our website.

If you're behind a web filter, please make sure that the domains *.kastatic.org and *.kasandbox.org are unblocked.

To log in and use all the features of Khan Academy, please enable JavaScript in your browser.

AP®︎/College Chemistry

Course: ap®︎/college chemistry   >   unit 2, formal charge.

  • Worked example: Using formal charges to evaluate nonequivalent resonance structures
  • Resonance and formal charge

Want to join the conversation?

  • Upvote Button navigates to signup page
  • Downvote Button navigates to signup page
  • Flag Button navigates to signup page

Video transcript

  • 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 Molecular and Ionic 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 Salt Solutions
  • 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 Multiple Equilibria
  • 16.1 Spontaneity
  • 16.2 Entropy
  • 16.3 The Second and Third Laws of Thermodynamics
  • 16.4 Free Energy
  • 17.1 Balancing Oxidation-Reduction Reactions
  • 17.2 Galvanic Cells
  • 17.3 Standard Reduction Potentials
  • 17.4 The Nernst Equation
  • 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

  • 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: CNS – , NCS – , or CSN – . The formal charges present in each of these molecular structures can help us pick the most likely arrangement of atoms. Possible Lewis structures and the formal charges for each of the three possible structures for the thiocyanate ion are shown here:

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

You may have noticed that the nitrite anion in Example 7.8 can have two possible structures with the atoms in the same positions. The electrons involved in the N–O double bond, however, are in different positions:

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. Thus, the electronic structure of the NO 2 − NO 2 − ion is shown as:

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

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

As an Amazon Associate we earn from qualifying purchases.

This book may not be used in the training of large language models or otherwise be ingested into large language models or generative AI offerings without OpenStax's permission.

Want to cite, share, or modify this book? This book uses the Creative Commons Attribution License and you must attribute OpenStax.

Access for free at https://openstax.org/books/chemistry/pages/1-introduction
  • Authors: Paul Flowers, William R. Robinson, PhD, Richard Langley, Klaus Theopold
  • Publisher/website: OpenStax
  • Book title: Chemistry
  • Publication date: Mar 11, 2015
  • Location: Houston, Texas
  • Book URL: https://openstax.org/books/chemistry/pages/1-introduction
  • Section URL: https://openstax.org/books/chemistry/pages/7-4-formal-charges-and-resonance

© Feb 15, 2022 OpenStax. Textbook content produced by OpenStax is licensed under a Creative Commons Attribution License . The OpenStax name, OpenStax logo, OpenStax book covers, OpenStax CNX name, and OpenStax CNX logo are not subject to the Creative Commons license and may not be reproduced without the prior and express written consent of Rice University.

FORMAL CHARGES QUIZ Assign the correct formal charge to the specified atom in the molecules below

3. What is the formal charge on the oxygen

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.

  • Pingback: Como posso calcular a cobrança formal? – CorujaSabia

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

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

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

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

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

Leave a Reply

Your email address will not be published. Required fields are marked *

Save my name, email, and website in this browser for the next time I comment.

Notify me via e-mail if anyone answers my comment.

This site uses Akismet to reduce spam. Learn how your comment data is processed .

Module 7: Chemical Bonding and Molecular Geometry

Formal charges and resonance, learning outcomes.

  • 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:

[latex]\text{formal charge}=\text{# valence shell electrons (free atom)}-\text{# lone pair electrons}-\dfrac{1}{2}\text{# bonding electrons}[/latex]

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 1: Calculating Formal Charge from Lewis Structures

Assign formal charges to each atom in the interhalogen ion [latex]{\text{ICl}}_{4}^{-}.[/latex]

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

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

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

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

Example 2: Calculating Formal Charge from Lewis Structures

Assign formal charges to each atom in the interhalogen molecule BrCl 3 .

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

  • Assign the lone pairs to their atom. Now each Cl atom has seven electrons and the Br atom has seven electrons.
  • Subtract this number from the number of valence electrons for the neutral atom. This gives the formal charge:Br: 7 – 7 = 0Cl: 7 – 7 = 0All 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.

Determine the formal charge for each atom in NCl 3 .

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

Using Formal Charge to Predict Molecular Structure

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

  • 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:

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

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

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

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

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

Example 3: Using Formal Charge to Determine Molecular Structure

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

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

Determining formal charge yields the following:

Two Lewis structures are shown with the word “or” in between them. The left structure depicts a nitrogen atom with two lone pairs of electrons double bonded to a nitrogen atom that is double bonded to an oxygen atom with two lone pairs of electrons. The numbers negative one, positive one, and zero are written above this structure. The right structure shows a nitrogen atom with two lone pairs of electrons double bonded to an oxygen atom that is double bonded to a nitrogen atom with two lone pairs of electrons. The numbers negative one, positive two, and negative one are written above this structure.

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

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

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

Which is the most likely molecular structure for the nitrite [latex]\left({\text{NO}}_{2}^{-}\right)[/latex] ion?

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

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 [latex]{\text{NO}}_{2}^{-}[/latex] have the same strength and length, and are identical in all other properties.

It is not possible to write a single Lewis structure for [latex]{\text{NO}}_{2}^{-}[/latex] 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 [latex]{\text{NO}}_{2}^{-}[/latex] is the average of a double bond and a single bond. We call the individual Lewis structures resonance forms . The actual electronic structure of the molecule (the average of the resonance forms) is called a resonance hybrid of the individual resonance forms. A double-headed arrow between Lewis structures indicates that they are resonance forms. Thus, the electronic structure of the [latex]{\text{NO}}_{2}^{-}[/latex] ion is shown as:

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

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).

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

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

You can view the transcript for “Resonance” here (opens in new window) .

Key Concepts and Summary

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

Key Equations

  • [latex]\text{formal charge}=\text{# valence shell electrons (free atom)}-\text{# one pair electrons}-\dfrac{1}{2}\text{# bonding electrons}[/latex]
  • selenium dioxide, OSeO
  • nitrate ion, [latex]{\text{NO}}_{3}^{-}[/latex]
  • nitric acid, HNO 3 (N is bonded to an OH group and two O atoms)

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

  • sulfur dioxide, SO 2 
  • carbonate ion, [latex]{\text{CO}}_{3}^{2-}[/latex]
  • hydrogen carbonate ion, [latex]{\text{HCO}}_{3}^{-}[/latex] (C is bonded to an OH group and two O atoms)

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

  • Write the resonance forms of ozone, O 3 , the component of the upper atmosphere that protects the Earth from ultraviolet radiation.
  • Sodium nitrite, which has been used to preserve bacon and other meats, is an ionic compound. Write the resonance forms of the nitrite ion, [latex]{\text{NO}}_{\text{2}}^{-}.[/latex]

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

  • Toothpastes containing sodium hydrogen carbonate (sodium bicarbonate) and hydrogen peroxide are widely used. Write Lewis structures for the hydrogen carbonate ion and hydrogen peroxide molecule, with resonance forms where appropriate.
  • [latex]{\text{SO}}_{4}^{2-}[/latex]
  • [latex]{\text{O}}_{2}^{2-}[/latex] (e) H 2 O 2
  • Calculate the formal charge of chlorine in the molecules Cl 2 , BeCl 2 , and ClF 5 .
  • [latex]{\text{BF}}_{4}^{-}[/latex]
  • [latex]{\text{SnCl}}_{3}^{-}[/latex]
  • [latex]{\text{PO}}_{4}^{\text{3-}}[/latex]
  • [latex]{\text{NO}}_{2}^{-}[/latex]
  • [latex]{\text{NO}}_{3}^{-}[/latex]
  • Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in nitrosyl chloride: ClNO or ClON?
  • Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in hypochlorous acid: HOCl or OClH?
  • Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in sulfur dioxide: OSO or SOO?
  • Draw the structure of hydroxylamine, H 3 NO, and assign formal charges; look up the structure. Is the actual structure consistent with the formal charges?
  • Write the Lewis structure and chemical formula of the compound with a molar mass of about 70 g/mol that contains 19.7% nitrogen and 80.3% fluorine by mass, and determine the formal charge of the atoms in this compound.

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

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

2. The resonance forms are as follows:

Two Lewis structures are shown with a double-headed arrow in between. The left structure shows a sulfur atom with a lone pair of electrons single bonded to the left to an oxygen atom with three lone pairs of electrons. The sulfur atom is also double bonded on the right to an oxygen atom with two lone pairs of electrons. The right structure depicts the same atoms, but this time the double bond is between the left oxygen and the sulfur atom. The lone pairs of electrons have also shifted to account for the change of bond types. The sulfur atom in the right structures, also has a third electron dot below it.

6. The Lewis structures are as follows:

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

CO has the strongest carbon-oxygen bond, because there are is a triple bond joining C and O. CO 2 has double bonds, and carbonate has 1.3 bonds.

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

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

The structure with formal charges of 0 is the most stable and would therefore be the correct arrangement of atoms.

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

18. There are 19.7 g N and 80.3 g F in a 100.0-g sample:

[latex]\begin{array}{l}\dfrac{19.7\text{g}}{14.0067\text{ g}{\text{ mol}}^{-1}}=1.406\text{ mol}\\ \dfrac{1.406\text{ mol}}{1.406\text{ mol}}=1\text{ N}\\ \dfrac{80.3\text{ g}}{18.9984\text{ g}{\text{ mol}}^{-1}}=4.2267\text{ mol}\\ \dfrac{4.2267\text{ mol}}{1.406\text{ mol}}=3\text{ F}\end{array}[/latex]

The empirical formula is NF 3 and its molar mass is 71.00 g/mol, which is consistent with the stated molar mass.

  • Oxidation states: N = +3, F = –1.

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

formal charge:  charge that would result on an atom by taking the number of valence electrons on the neutral atom and subtracting the nonbonding electrons and the number of bonds (one-half of the bonding electrons)

molecular structure:  arrangement of atoms in a molecule or ion

resonance:  situation in which one Lewis structure is insufficient to describe the bonding in a molecule and the average of multiple structures is observed

resonance forms:  two or more Lewis structures that have the same arrangement of atoms but different arrangements of electrons

resonance hybrid:  average of the resonance forms shown by the individual Lewis structures

  • Chemistry 2e. Provided by : OpenStax. Located at : https://openstax.org/ . License : CC BY: Attribution . License Terms : Access for free at https://openstax.org/books/chemistry-2e/pages/1-introduction
  • Resonance. Authored by : Khan Academy. Located at : https://youtu.be/6XOm3Km7r30 . License : Other . License Terms : Standard YouTube License

MCAT and Organic Chemistry Study Guides, Videos, Cheat Sheets, tutoring and more

Formal Charge Formula and Shortcut for Organic Chemistry

April 6, 2015 By Leah4sci 53 Comments

formal charge equation

Why do I call it pesky? Because the equation in your textbook is long, confusing, and needlessly annoying.

How can you afford that kind of time when working through a multi-step synthesis?

You Can’t! You shouldn’t have to.

Yet understanding the nature of Formal Charge is a critical component when it comes to mastering organic chemistry reactions and mechanisms. Formal charge helps you understand reaction patterns by showing you why a specific atom attacks, and why its ‘victim’ accepts the attack.

So what exactly IS formal charge?

Formal charge is the actual charge on an individual atom within a larger molecule or polyatomic ion.

The sum of formal charges on any molecule or ion results in the net overall charge.

This concept is simple enough for small ions. Chloride obviously has a negative charge. Even the negative charge on the hydroxide oxygen is simple to understand.

NO3- lewis structure

But what if you have a much larger group of bound atoms with an overall net charge? For example, the negative nitrate or triple negative phosphate.

Where does the charge come from?

Which specific atom is responsible for the overall charge?

Can it be more than one atom?

And, what if you’re studying an uncharged molecule – Is it still possible for individual atoms to carry charge despite the net neutrality?

The answers are yes, yes, yes, and more yes.

Think back to general chemistry when you studied ion formation. An ion is simply an atom or molecule that gained or lost electrons to get a net charge.

If the atom has just one more negative electron, than protons, it will have a net negative charge.

If the atom loses negative electrons and therefore has as few as one extra POSITIVE protons in its nucleus, it will carry a net positive charge.

Now that you know WHAT we’re studying, let’s see HOW! – That’s where I can help you.

How to Calculate Formal Charge

When I first studied formal charge I was lost. The formula in my textbook was long, tedious and brutal. A long, annoying trip to my professor’s office hours enabled me to figure it out, but it was still overwhelmingly TEDIOUS!

Let’s start with the proper formula and then learn a shortcut.

Formal Charge Equation

Formal Charge Formula

Formal Charge = [# valence electrons on neutral atom] – [(# lone electron pairs) + (½ # bonding electrons)]

  • Valence electrons = corresponds to the group number of the periodic table (for representative elements).
  • Lone Pairs = lone electrons sitting on the atom. Each electron counts as one and so a pair counts as two.
  • ½ # bonding electrons = Since a bond is formed by sharing 2 electrons between 2 atoms, every atom in the bond can only take credit for one of the 2 electrons. A double bond containing four electrons allows each atom to claim 2. A triple bond of 6 electrons allows each atom to claim three.

Are you confused?

Do you dread running through this every time you finish a reaction?

I’m with you.

Yes the equation works, but it’s far too tedious and annoying! There are just too many steps and calculations.

Here another option, MY version – one that is easier, faster, and comes out with the same result!

Formal Charge Calculation Shortcut

Here’s my version of this equation.

Formal Charge = Should – Has

Formal Charge Formula and Shortcut

This shortcut is guaranteed to save precious seconds on your exam IF AND ONLY IF you understand how to apply it.

But when you understand it you’ll be able to solve formal charge in your head, in under 8 seconds per atom.

Let’s make sure you understand this shortcut

Formal Charge Electron Demonstration

Should = the number of valence electrons that a neutral atom SHOULD have.

Has = the number of electrons an atom HAS directly attached, touching the atom in question.

Lone pairs represent 2 electrons sitting on the atom so that Has = 2

Each bond only counts for a single electron since the second electron in the bond is touching the other atom.

Want to see this shortcut brought to life? See my formal charge video below

When you first study formal charge it helps to draw out the Lewis Structure for every molecule in question. As you get more comfortable with this topic you’ll be able to pick out bonds on skeletal bond-line drawings.

For now, let’s stick with the basics (or better – Let’s keep it simple), and look at a handful of examples utilizing the checklist shared in the Lewis Structures video below. This is a must-learn checklist – so if you didn’t watch yet, review this video:

periodic table to memorize for organic chemistry

  • Group numbers as written on top
  • Electronegativity increases up and to the right
  • Size increases going down and to the left

Let’s start with an easier example. We’ve already discussed hydroxide so let’s take a look at the OH- Lewis Structure.

Following the checklist in the Lewis Structures video we have oxygen bound to hydrogen, with 3 lone pairs around oxygen.

hydroxide OH- lewis structure

Let’s apply the shortcut:

Hydroxide OH- formal charge and lewis structure

Neutral oxygen should have 6 valence electrons

Oxygen has 7 attached electrons as shown in red

Should – Has = 6 – 7 = -1 Hydroxide has a negative formal charge on the oxygen atom.

What about Nitrate?

Let’s analyze the NO3- Lewis Structure and Formal Charge

Following the checklist we draw our atoms, bonds, and electrons.

NO3 Lewis Structure

Now for formal charge

NO3- Formal charge and lewis structure

Neutral nitrogen should have 5 valence electrons but our drawing only shows 4 attached.

Next we’ll look at the double bound oxygen. Should – has = 6 – 6 = 0

The double bound oxygen is happy, stable, and has a net neutral charge.

Finally we have 2 single bound oxygen atoms with 3 lone pairs each. Since they are identical we have to calculate just one to get the answer for both. Should – Has = 6 – 7 = -1 Oxygen should have 6 valence electrons, each of the remaining oxygen atoms have 7 attached electrons for a net negative charge.

Let’s try one more simple yet tricky example. Carbon monoxide.

Simple because it only has 2 atoms.

Tricky because carbon monoxide is neutral, but don’t let this fool you. Many a hemoglobin has been fooled by CO making it such a deadly poison.

There are 2 types of molecules with a net zero charge

  • molecules with no formal charge
  • molecules with formal charge that cancel out for a net zero formal charge.

CO Lewis Structure and Formal Charge

CO lewis structure

Now let’s tackle the individual atoms.

Carbon should have 4, has 5 attached, formal charge = -1 Oxygen should have 6, has 5 attached, formal charge = +1 +1 and -1 cancel for a net formal charge of zero

CO formal charge and lewis structure

What do you think? Do you still find formal charge as tedious and daunting as you did when you first learned it? Hopefully the shortcut ‘should – has’ given you the confidence to simply stare at a molecule and come up with instant formal charges.

Or, are you needing more practice with Lewis Structures, Shape, Hybridization and more? Watch below!

I'd love to read your feedback in the comments below!

Ready to test your resonance skills try my free resonance practice quiz ..

' srcset=

March 26, 2019 at 8:19 pm

Thank you soooo much…

' srcset=

August 24, 2018 at 8:38 pm

Its so helpful!!! Thank you

' srcset=

August 12, 2018 at 8:43 am

Thank you Leah, you’ve been so helpful!!!

' srcset=

August 11, 2018 at 1:10 pm

Literally awsum trick for solving formal charge in fraction of seconds…thank you

' srcset=

August 9, 2018 at 2:18 am

Amazing. Thank you very much for your great help. Nobody ever taught me that. It is an excellent time saver. Please keep on posting new shortcuts for different topics if you have.

' srcset=

July 25, 2018 at 9:57 am

Namaste Mam,

Thanks for your help, it’s easy to understand now.

' srcset=

July 22, 2018 at 11:43 pm

i got it just now by watching it several times thanks

' srcset=

July 19, 2018 at 6:02 pm

Why don’t they teach this at school if it’s such a good equation? Why are u teaching things that rnt in the curriculum

' srcset=

April 21, 2018 at 11:34 am

How do we treat a coordinate covalent bond?

' srcset=

March 23, 2018 at 5:52 pm

What if some charge is present on compound I mean lyk if compound is HClO4–?? What about this negative charge?

' srcset=

March 14, 2018 at 12:25 pm

Very helpful ..thanx a lot

' srcset=

February 3, 2018 at 11:18 am

Wow… I never thought that I could easily calculate the formal charge without even remembering the formula.Superb trick !!!

' srcset=

January 7, 2018 at 6:47 am

I love you so so much.. But I, have to take entry test for the medical field.. We are given only 42seconds to solve the question and to fill up the circle.. Moreover, calculator is also not allowed in the test centre. I need short cut methods to calculate no of moles, no of atoms, no of ions and charges on them, volume of gas from the given mass or mass from given volume or given no of particles .also formulae to calculate tackle with mass to mass, mass to volume, volume to mass,no of particles to volume…. Similarly, molarity, molality, mole fraction, percentage w/w,w/v,v/w,v/v, and out of them find out molarity or molality etc…

' srcset=

November 23, 2017 at 2:30 am

Thank you so much for this excellent trick! You made it so easy!!

' srcset=

October 11, 2017 at 3:08 am

Thank uu it is very easy …..

' srcset=

September 18, 2017 at 11:28 am

Hi leah thank u so much for the tricks! 🙂 was wondering how then, would we know when to use dative bonding?

' srcset=

June 30, 2017 at 10:48 am

i want to know formal charge for O3

' srcset=

February 21, 2017 at 10:54 am

Thank you for this kind of help.. Could you discuss question I want to send on …?swap method

' srcset=

February 15, 2017 at 3:35 am

Please write about the Lewis structure of NO2—

' srcset=

January 27, 2017 at 11:32 am

thanks leah for your wonderful technique

' srcset=

January 17, 2017 at 7:41 am

When I first come to your site and read the tutorial,it was large and I was feeling boring but after reading it with concentration, I got the trick and It got worked.Really Very Helpful Trick.Kindly update some more tricks. Regards, Sagar Gupta

' srcset=

December 15, 2016 at 9:21 pm

Thank you so much! This literally saved my life for the MCAT Gen Chem and Organic Chem sections! I struggled so much previously using the old equation and spent way too much time trying to go through the equation. This was so helpful!

' srcset=

December 13, 2016 at 12:06 am

Leah kindly do share more shortcuts.

December 13, 2016 at 12:03 am

Leah you helped me a lot after a long struggle in organic chemistry. You’re superb. Thanks a lot

' srcset=

December 8, 2016 at 1:10 pm

Formal charge concept was bouncing over my head and now it’s over.Thank you for this easy trick.

' srcset=

December 29, 2016 at 4:17 pm

You’re welcome!

' srcset=

August 8, 2017 at 1:25 pm

leah madam namaste. i am so glad that u are highly knowledgable and i found no teacher as u in india where i study in a coaching institute to clear entrance exam for under graduation in mbbs. thank u leah maa’m. one day my success is sure in which u have a bit worthy role.

August 8, 2017 at 1:30 pm

namaste leah madam. i have found no teacher as u in india where i am studying in a coaching institute to clear an entrance exam for undergraduation in mbbs. i hope that with ur simple n easy tricks i will be strong n bold enogh to solve the questions easily and get a good rank. thank u madam, u are so knowledgable . god will be always with u . keep on teaching like this .

' srcset=

July 22, 2019 at 12:11 pm

Formal charge of carbon in benzene

' srcset=

December 6, 2016 at 9:36 am

Leah, can u pls find formal chrg of P in H3PO4, cz by ur formula the answer is zero by in my book it’s given +1.. Thnx

' srcset=

November 2, 2016 at 6:55 pm

This really would have saved me some time on my last quiz…. I had no idea about this. Very cool.

' srcset=

July 28, 2016 at 10:49 am

Wow Leah, I didn’t know about this quick, short and time saving method. Thank you for such good work

August 10, 2016 at 9:27 am

Glad you found this helpful!

' srcset=

May 12, 2016 at 2:34 am

This should – has formula is great. But i’m having some problems with NO3-. In NO3-, one oxygen atom is bonded with nitrogen atom – one sigma and one pie bond. But there is a coordinate bond between N and the other O atom. so the formal charge on O should be= 6-(6+2). It should be -2. So how come it is -1? Can you clear my confusion pls.

July 3, 2016 at 11:05 am

You should only have 2 numbers in your equation NOT 3. Oxygen would be 6 – 7 for -1 or 6-6 = 0 for the other

' srcset=

April 1, 2016 at 3:52 pm

You have done a fantastic job! Thanks a lot for all these, You are amazing…

April 4, 2016 at 5:54 pm

Thanks Shubham

' srcset=

December 1, 2015 at 8:19 am

Formal charge was frustrating me for a long time,but after this “should-has” technique I feel like a burden is off my head. Thank you very much for your fantastic trick.

December 6, 2015 at 5:22 pm

You’re very welcome. I felt the same way but now it’s second nature to me. Glad you feel the same way

' srcset=

November 5, 2015 at 1:34 pm

The reason it’s normally not taught this way is that this just leads to a different confusion. The students response is “but I thought we need octets. Why does carbon only have 5 electrons in CO?”

November 10, 2015 at 9:32 pm

I don’t understand your confusion. Carbon has 8 electrons in CO 5 attached and 3 more in the bonds for a total of 8

' srcset=

August 30, 2015 at 2:59 pm

so one atom can be negative and the other positive, such as in (CH3)3NO. what would the overall charge be on that compound?

August 30, 2015 at 10:21 pm

To find the overall or net charge, simply add up the individual formal charges.

' srcset=

July 26, 2015 at 9:52 am

This is very helpful! Thanks!

August 1, 2015 at 11:59 am

You’re welcome

' srcset=

April 6, 2015 at 10:19 pm

Thanks now am going well with this series of tutorial but I need more different questions to make me practise on lewis structures and formal charge as well.

April 8, 2015 at 9:37 am

I agree Ally. I’m planning to add a quiz to this post down the line. For now refer to your textbook and look for as many problems as you can find to drive this topic home.

' srcset=

August 4, 2016 at 12:21 am

How would you work that out for N NH2 ? Suppose I take 5 electrons as ‘should’ and 1 electron from each Hydrogen which makes ‘Has’ 7. So, the formal charge on Nitrogen becomes -2 but it should be -1. Please tell me what I am missing?

September 2, 2016 at 10:42 am

I’m not sure I understand your question. Is your molecule N2H2? Or the N in NH2, and if so, first give everything their complete octet. Then determine that nitrogen should have five, hydrogen should have one, that’s twelve total. Count how many you have around each. Twelve will include bonds. But whatever is touching directly is considered in your formal charge calculation.

' srcset=

April 6, 2015 at 3:21 pm

I’ve never seen a trivalent oxygen …how do u explain this?

April 8, 2015 at 9:35 am

Sheila: Oxygen with 3 bonds is very common in both organic and inorganic molecules. Think about ozone and hydronium, and so many more

' srcset=

July 5, 2016 at 10:15 am

Please find me the Bond order of ozone???| Thanks!

' srcset=

March 3, 2018 at 3:59 am

Wow! That’s a cool method!! Its so easy to find formal charge now…Thank you so much….does that trick work for all molecules without a single exeption? If so please do let me know.

Leave a Reply Cancel reply

Your email address will not be published. Required fields are marked *

Save my name, email, and website in this browser for the next time I comment.

Organic Chemistry Reference Material and Cheat Sheets

calculate formal charge practice

Alkene Reactions Overview Cheat Sheet – Organic Chemistry

The true key to successful mastery of alkene reactions lies in practice practice practice. However, … [Read More...]

Click for additional cheat sheets

MCAT Tutorials

mcat math without a calculator 1 play

Introduction To MCAT Math Without A Calculator

While the pre-2015 MCAT only tests you on science and verbal, you are still required to perform … [Read More...]

Click for additional MCAT tutorials

Organic Chemistry Tutorial Videos

KET Keto enol tautomerization reaction and mechanism leah4sci

Keto Enol Tautomerization Reaction and Mechanism

Keto Enol Tautomerization or KET, is an organic chemistry reaction in which ketone and enol … [Read More...]

Click for additional orgo tutorial videos

In order to be most effective for you, try to answer these questions before you look at the answers!

Talk to our experts

1800-120-456-456

  • Formal Charge

ffImage

What is Formal Charge?

A formal charge (F.C. or q) is the charge assigned to an atom in a molecule in the covalent view of bonding, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity.

The formal charge is the difference between an atom's number of valence electrons in its neutral free state and the number allocated to that atom in a Lewis structure.

When choosing the optimum Lewis structure (or predominant resonance structure) for a molecule, it is important to keep the formal charge on each of the atoms as low as feasible.

The following equation can be used to compute the formal charge of an atom in a molecule:

F = V - L - \[\frac{B}{2}\]

F = Formal Charge

V = Valence Electron of the neutral atom in isolation

L = Number of non-bonding valence electrons on this atom in the molecule

B = Total number of electrons shared in bonds with other atoms in the molecule

Formula, Calculation, Importance, and Example

The formula for computing a formal charge is:

(Number of valency electrons in neutral atom)-(electrons in lone pairs + 1/2 the number of bonding electrons)

The number of bonding electrons divided by two equals the number of bonds that surround the atom, hence this expression can be reduced to:

Formal Charge = (number of valence electrons in neutral atom)-(non-bonded electrons + number of bonds)

Take the compound BH 4 or tetrahydrdoborate.

Boron (B) possesses three valence electrons, zero non-bonded electrons, and four bonds around it.

This changes the formula to 3-(0+4), yielding a result of -1.

Let us now examine the hydrogen atoms in BH4. One valence electron, zero non-bonded electrons, and one bond make up hydrogen.

In BH4, the formal charge of hydrogen is 1-(0+1), resulting in a formal charge of 0.

Calculate the formal charge on the following:

O atoms of O3

Cl atom in HClO4- ion

S atom in HSO4- ion

Ans: We are showing how to find a formal charge of the species mentioned.

Formal charge on O1: 6 – 6/2 – 2 = +1

Formal charge on O2: 6 – 4/2 – 4 = 0

Formal charge on O3: 6 – 2/2 – 6 = -1

Formal charge on Cl atom of HClO4 ion: 7 – 8/2 – 0 = 3

Formal charge on S atom of HSO4- ion: 6 – 8/2 – 0 = 2

Significance

Molecular Structure

An atom in a molecule should have a formal charge of zero to have the lowest energy and hence the most stable state. If there are numerous alternatives for a molecule's structure, this gives us a hint: the one with the least/lowest formal charges is the ideal structure.

While formal charge can indicate a molecule's preferred structure, the problem becomes more complicated when numerous equally preferred structures exist. This condition could point to resonance structures, especially if the structures have the same atom arrangement but different types of arrangements of bonds.

The formal charge of a molecule can indicate how it will behave during a process. A negative formal charge indicates that an atom is more likely to be the source of electrons in a reaction (a nucleophile). If it has a positive one, on the other hand, it is more likely to take electrons (an electrophile), and that atom is more likely to be the reaction's site.

It's also worth noting that an atom's formal charge differs from its actual charge. Formal charge ignores electronegativity and assumes that electrons in a bond are uniformly distributed.

It's only a courtesy that's utilized to make molecular structures and reaction mechanisms more understandable. The actual charge, on the other hand, is based on the electronegativities of the atoms and the polarity of the bonds and looks at the actual electron density.

Importance Of Formal Charge  

Now that we know what is the formal charge and we are familiar with the process for calculating a formal charge, we will learn about its importance. 

The formal charge is a theoretical concept, useful when studying the molecule minutely. It does not indicate any real charge separation in the molecule. This concept and the knowledge of ‘what is formal charge' is vital.

The formal charge is crucial in deciding the lowest energy configuration among several possible Lewis structures for the given molecule. Therefore, calculating formal charges becomes essential.

Knowing the lowest energy structure is critical in pointing out the primary product of a reaction. This knowledge is also useful in describing several phenomena.

The structure of least energy is usually the one with minimal formal charge and most distributed real charge. 

Besides knowing what is a formal charge, we now also know its significance. 

Fun Facts On Formal Charge 

In organic chemistry, convention governs that formal charge is essential for depicting a complete and correct Lewis-Kekulé structure. However, the same does not apply to inorganic chemistry.  

The structure variation of a molecule having the least amount of charge is the most superior. 

The differences between formal charge and oxidation state led to the now widely followed and much more accurate valence bond theory of Slater and the molecular orbital theory of Mulliken.

arrow-right

Advertisement

The Civil Fraud Ruling on Donald Trump, Annotated

By Kate Christobek

Former President Donald J. Trump was penalized $355 million , plus millions more in interest, and banned for three years from serving in any top roles at a New York company, including his own, in a ruling on Friday by Justice Arthur F. Engoron. The decision comes after the state's attorney general, Letitia James, sued Mr. Trump, members of his family and his company in 2022.

The ruling expands on Justice Engoron’s decision last fall , which found that Mr. Trump’s financial statements were filled with fraudulent claims. Mr. Trump will appeal the financial penalty and is likely to appeal other restrictions; he has already appealed last fall’s ruling.

The New York Times annotated the document.

Download the original PDF .

Page 1 of undefined PDF document.

New York Times Analysis

This ruling by Justice Arthur F. Engoron is a result of a 2022 lawsuit filed by New York’s attorney general, Letitia James , against Donald J. Trump and the Trump Organization; his adult sons, Donald Trump Jr. and Eric Trump; the company’s former chief financial officer Allen Weisselberg and former controller Jeffrey McConney; and several of their related entities. Mr. Trump’s daughter, Ivanka Trump, was also initially a defendant until an appeals court dismissed the case against her.

Page 2 of undefined PDF document.

The law under which Ms. James sued, known by its shorthand 63(12), requires the plaintiff to show a defendant’s conduct was deceptive . If that standard is met, a judge can impose severe punishment, including forfeiting the money obtained through fraud. Ms. James has also used this law against the oil company ExxonMobil, the tobacco brand Juul and the pharma executive Martin Shkreli.

Page 4 of undefined PDF document.

Justice Engoron is now providing a background of this case. This ruling comes after a three-year investigation by the attorney general’s office and the conclusion of a trial that ended last month. But this likely won’t be Mr. Trump’s last word on the matter — he will appeal the financial penalty and is likely to appeal other restrictions, as he has already appealed other rulings.

In late 2022, Justice Engoron assigned a former federal judge, Barbara Jones, to serve as a monitor at the Trump Organization and tasked her with keeping an eye on the company and its lending relationships. Last month, she issued a report citing inconsistencies in its financial reporting, which “may reflect a lack of adequate internal controls.”

Page 5 of undefined PDF document.

Here, Justice Engoron is laying out the laws he considered in his ruling beyond 63(12). The attorney general’s lawsuit included allegations of violations of falsifying business records, issuing false financial statements, insurance fraud and related conspiracy offenses.

Justice Engoron is explaining the decision, issued a week before the trial, in which he found that Mr. Trump’s financial statements were filled with fraud , fundamentally shaping the rest of the trial.

Page 6 of undefined PDF document.

For over 50 pages, Justice Engoron describes his conclusions about the testimony of all of the witnesses who spoke during the trial.

Page 8 of undefined PDF document.

Justice Engoron discusses Mr. McConney’s important role in preparing Mr. Trump’s financial statements. The judge points out that Mr. McConney prepared all the valuations on the statements in consultation with Mr. Weisselberg.

Page 24 of undefined PDF document.

In his discussion of Mr. Weisselberg, Justice Engoron calls his testimony in the trial “intentionally evasive.” Justice Engoron then brings up Mr. Weisselberg’s separation agreement from the Trump Organization, which prohibited him from voluntarily cooperating with any entities “adverse” to the organization. Justice Engoron says that this renders Mr. Weisselberg’s testimony highly unreliable.

Page 27 of undefined PDF document.

When Donald Trump Jr. testified in court, he disavowed responsibility for his father’s financial statements despite serving as a trustee of the Donald J. Trump Revocable Trust while his father was president. But Justice Engoron specifically cites here that Donald Trump Jr. certified that he was responsible for the financial statements, and testified that he intended for the banks to rely on them and that the statements were “materially accurate.”

Page 30 of undefined PDF document.

During his testimony, Eric Trump, the Trump Organization’s de facto chief executive, initially denied knowing about his father’s financial statements until this case. As Justice Engoron points out here, Eric Trump eventually conceded to knowing about them as early as 2013. As a result, Justice Engoron calls Eric Trump’s credibility “severely damaged.”

Page 33 of undefined PDF document.

Justice Engoron points to Mr. Trump’s testimony when he took the witness stand in November when Mr. Trump acknowledged that he helped put together his annual financial statements. Mr. Trump said he would see them and occasionally have suggestions.

Page 35 of undefined PDF document.

After four pages of describing Mr. Trump’s testimony, Justice Engoron says Mr. Trump rarely responded to the questions asked and frequently interjected long, irrelevant speeches, which all “severely compromised his credibility.”

Page 38 of undefined PDF document.

For several pages, Justice Engoron provides background on specific assets that Mr. Trump included in his annual financial statements.

Page 61 of undefined PDF document.

The judge is clarifying that Ms. James had to prove her claims by a “preponderance of the evidence,” meaning she had to demonstrate it was more likely than not that Mr. Trump and the co-defendants should be held liable. This is a lower standard than that of a criminal trial, which requires that evidence be proven “beyond a reasonable doubt.”

Page 76 of undefined PDF document.

During the trial, Mr. Trump and his legal team tried to shift the blame for any inaccuracies in his financial statements onto his outside accountants. But Justice Engoron criticizes that argument here.

Page 77 of undefined PDF document.

During the monthslong trial, Mr. Trump, his legal team and several witnesses stressed that real estate appraisals are an art, not a science. But here it’s clear Justice Engoron, while agreeing with that sentiment, also believes it’s deceptive when different appraisals rely on different assumptions.

Page 78 of undefined PDF document.

Justice Engoron is now going through the defendants one by one and articulating the evidence that shows each of their “intent to defraud,” which is required by the statute against falsifying business records. Notably, his first paragraph describing the former president’s intent provides examples including Mr. Trump’s awareness that his triplex apartment was not 30,000 square feet and his valuation of Mar-a-Lago as a single-family residence even though it was deeded as a social club.

Page 79 of undefined PDF document.

Among the defendants, Justice Engoron finds only Allen Weisselberg and Jeffrey McConney liable for insurance fraud. Here, he doesn’t provide an explanation for why the other defendants, including Mr. Trump and his adult sons, were not found liable, and he says that both Mr. Weisselberg and Mr. McConney made false representations to insurance companies about Mr. Trump’s financial statements.

While Mr. Trump and his adult sons were not found liable for insurance fraud, here Justice Engoron finds them liable for conspiracy to commit insurance fraud, explaining that they all “aided and abetted” the conspiracy to commit insurance fraud by falsifying business records.

Page 82 of undefined PDF document.

Justice Engoron here adopts the approximations of Michiel McCarty, the attorney general’s expert witness. Justice Engoron says Mr. McCarty testified “reliably and convincingly,” and finds that the defendants’ fraud saved them over $168 million in interest.

Page 83 of undefined PDF document.

In finding that the defendants were able to purchase the Old Post Office in Washington, D.C., through their use of the fraudulent financial statements, Justice Engoron rules that the defendants’ proceeds from the sale of the post office in 2022 should be considered “ill-gotten gains.” He penalizes Donald Trump and his companies over $126 million, and Donald Trump Jr. and Eric Trump $4 million each, for this one property.

Page 84 of undefined PDF document.

Justice Engoron blasts the defendants for failing to admit that they were wrong in their valuations — adding that “their complete lack of contrition and remorse borders on pathological.” He says that this inability to admit error makes him believe they will continue their fraudulent activities unless “judicially restrained.”

Page 88 of undefined PDF document.

The judge cites other examples of Mr. Trump’s “ongoing propensity to engage in fraud,” bringing up lawsuits against Trump University and the Donald J. Trump Foundation. He also notably raises two criminal cases brought by the Manhattan district attorney’s office: one against Mr. Weisselberg, who pleaded guilty to tax fraud and falsifying business records , and another against the Trump Organization, which was convicted of 17 criminal counts including tax fraud .

Justice Engoron states that Judge Barbara Jones, who has been serving as an independent monitor at the Trump Organization since 2022, will continue in that role for at least three years. He clarifies that going forward, her role will be enhanced and she will review Trump Organization financial disclosures before they are submitted to any third party, to ensure that there are no material misstatements.

Page 89 of undefined PDF document.

In addition to extending the monitor’s tenure and strengthening her powers, Justice Engoron also took the unusual step of ordering that an independent compliance director be installed inside The Trump organization, and that they report directly to the monitor.

— William K. Rashbaum

In his pre-trial order, Justice Engoron ordered the cancellation of some of Mr. Trump’s business licenses . But here, he pulls back on that decision and instead says that any “restructuring and potential dissolution” would be up to Ms. Jones, the independent monitor.

Page 90 of undefined PDF document.

Justice Engoron lays out his bans against the defendants, ruling that Mr. Trump, Mr. Weisselberg and Mr. McConney cannot serve as officers or directors of any corporation or legal entity in New York for the next three years, and bans his sons Donald Trump Jr. and Eric Trump for two years from the same. He also prohibits Mr. Trump from applying for any loans from any New York banks for the next three years. The ruling goes further in the cases of Mr. Weisselberg and Mr. McConney, permanently barring them from serving in the financial control function of any New York business.

Page 91 of undefined PDF document.

Justice Engoron also ordered that Mr. Trump and his sons pay the interest, pushing the penalty to $450 million, according to Ms. James.

Page 92 of undefined PDF document.

An earlier version of this article misstated how long the adult sons of former President Donald J. Trump — Donald Trump Jr. and Eric Trump — were barred by Justice Arthur F. Engoron from serving as officers or directors of any corporation or legal entity in New York. It was two years, not three. The article also misstated the number of pages in which Justice Engoron describes his conclusions about the testimony of all of the non-defendant witnesses. It was under 50 pages, not over 50 pages. The article also misstated the number of pages in the section in which Justice Engoron provides background on specific assets that Mr. Trump included in his annual financial statements. It was several pages, not more than a dozen pages.

  • Share full article

Library homepage

  • school Campus Bookshelves
  • menu_book Bookshelves
  • perm_media Learning Objects
  • login Login
  • how_to_reg Request Instructor Account
  • hub Instructor Commons
  • Download Page (PDF)
  • Download Full Book (PDF)
  • Periodic Table
  • Physics Constants
  • Scientific Calculator
  • Reference & Cite
  • Tools expand_more
  • Readability

selected template will load here

This action is not available.

Chemistry LibreTexts

Formal Charges

  • Last updated
  • Save as PDF
  • Page ID 5129
  • Not all atoms within a neutral molecule have to be neutral.
  • Atoms within a molecule with formal charges are often the site of chemical reactions.
  • You should be able to determine the formal charge for any atom in any molecule.

Knowing the formal charges on specific atoms in a molecule is an important step in keeping tract of the electrons and determine the chemical reactivity of the molecule. Formal charges can be calculated mathematically, but they can also be determined by intuition. The instinctive method is faster but requires more skill and knowledge of common structures.

Mathematical Method

Formal charge equation is based on the comparing the number of electrons in the individual atom with that in the structure. For each atom, we then compute a formal charge:

\[ \text{formal charge}= \underbrace{valence\; e^-}_{free\; atom } - \underbrace{ \left ( nonbonding\; e^{-}+\dfrac{bonding\;e^{-}}{2} \right ) }_{atom\; in\; Lewis\; structure } \label{8.5.1} \]

where the group number is equal to the number of electrons of the neutral atom.

Intuitive Method

The instinctive method relies on experience and the understanding of common, known neutral structures. To do this you need to know the common bonding patterns for C, N, O, and the halides:

  • C: 4 bonds;
  • N: 3 bonds, 1 lone pair;
  • O: 2 bonds, 2 lone pairs;
  • F: 1 bond, 3 lone pairs (or any other halide).

Once mastered, this is much easier and quicker. The diagram below shows the common bonding situations for C, N, O, and F. The central column is the neutral atom, in the left most column the atom has gained an electron by having a bonding pair converted to a lone pair of electrons. The right most column has a formal charge of +1 due to either: a) using a lone pair of electrons as a bonding pair (N and O), b) losing a bonding pair of electrons (C), or c) losing a lone pair of electrons (F).

common formal charges

IMAGES

  1. Formal Charge Formula

    calculate formal charge practice

  2. Formal Charge

    calculate formal charge practice

  3. Formal Charge Practice Worksheets

    calculate formal charge practice

  4. Formal Charge Practice Worksheet

    calculate formal charge practice

  5. How To Calculate The Formal Charge of an Atom

    calculate formal charge practice

  6. Formal Charges and How To Calculate Them

    calculate formal charge practice

VIDEO

  1. ||How to calculate formal charge|| #trick #chemistry #mhtcet #mhtcet2022 #exam

  2. Assigning formal charge on intermediate and ions bsc 1st year |organic chemistry

  3. Formal Charge H2SO4 ।। How to calculate formal Charge H2SO4 ।। How to calculate formal charge

  4. How to calculate Formal charge |class-XI

  5. Formal Charge Why do we study Formal Charge Class 11 Chemical bonding CONFIRMED EXAM QUESTION

  6. Numerical 1

COMMENTS

  1. Formal charge practice problems with answers (PDF)

    Part 1 Assign a formal charge to all atoms and determine the overall charge of the molecule. All lone pairs and hydrogens attached to carbon are shown. (Click on the picture to zoom in!) Click here to reveal the answers Part 2 Draw in any lone pairs and any hydrogens attached to carbon.

  2. Calculating Formal Charge Practice

    Calculating Formal Charge AP Chemistry Skills Practice 1. What is the formal charge of boron in B C l 3 ? 2. What is the formal charge of silicon in S i O 2 ? 3. What is the formal...

  3. Resonance and formal charge (practice)

    Choose 1 answer: All of the bonds in CO A 3 A 2 − are identical in length and strength. A All of the bonds in CO A 3 A 2 − are identical in length and strength. The bonds in CCl A 4 are more polar than the bonds in CH A 4 . B The bonds in CCl A 4 are more polar than the bonds in CH A 4 . Both of the bonds in BeH A 2 have a bond order of 1 . C

  4. 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.

  5. Formal charge (video)

    We can calculate an atom's formal charge using the equation FC = VE - [LPE - ½ (BE)], where VE = the number of valence electrons on the free atom, LPE = the number of lone pair electrons on the atom in the molecule, and BE = the number of bonding (shared) electrons around the atom in the molecule. Created by Sal Khan. Questions Tips & Thanks

  6. 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 ...

  7. 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 −.

  8. 7.4 Formal Charges and Resonance

    Solution 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 Cl: 7 - 7 = 0

  9. Formal Charge Practice Problems with Explanations

    A video of formal charge practice problems (from easy to difficult) with clear, concise answers and explanations.Calculating the formal charges for a molecul...

  10. How To Calculate The Formal Charge of an Atom

    This chemistry video tutorial provides a basic introduction into how to calculate the formal charge of an atom or element in a lewis structure. This video i...

  11. Formal Charges Quiz

    1. What is the formal charge on the oxygen (note: double bonds count as four shared electrons) +1 -1 zero 2. What is the formal charge on the carbon atom +1 -1 zero 3. What is the formal charge on the oxygen +1 -1 zero

  12. Formal Charges Video Tutorial & Practice

    Suzuki Reaction 25m. Sonogashira Coupling Reaction 17m. Fukuyama Coupling Reaction 15m. Kumada Coupling Reaction 13m. Negishi Coupling Reaction 16m. Buchwald-Hartwig Amination Reaction 19m. Eglinton Reaction 17m. Learn Formal Charges with free step-by-step video explanations and practice problems by experienced tutors.

  13. Formal Charges: Calculating Formal Charge

    ...more Join this channel and unlock members-only perks A step-by-step description on how to calculate formal charges. Formal charges are important because they allow us to predict which...

  14. How To Calculate Formal Charge

    To obtain the formal charge of an atom, we start by counting the number of , and then subtract from it the number of electrons that it " i.e. electrons in lone pairs, or singly-occupied orbitalshalf the number of bonding electrons, which is equivalent to the number of bonds The simplest way to write the formula for formal charge (

  15. 4.3: Formal Charge and Oxidation State

    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 4.3.1 4.3. 1. Calculate the formal charge for each atom in the carbon monoxide molecule:

  16. Formal Charges and Resonance

    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.

  17. Formal Charge Formula and Shortcut for Organic Chemistry

    Formal Charge = [# valence electrons on neutral atom] - [ (# lone electron pairs) + (½ # bonding electrons)] Valence electrons = corresponds to the group number of the periodic table (for representative elements). Lone Pairs = lone electrons sitting on the atom. Each electron counts as one and so a pair counts as two.

  18. Ch 1 : Formal charge questions

    Ch 1 : Formal charge questions. Formal Charge Questions. In order to be most effective for you, try to answer these questions before you look at the answers! You might need a periodic table to help you here. 1. For each of the structures shown below, identify the formal charge of any atoms that are not neutral.

  19. 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.

  20. Lewis Structures and Formal Charges Practice Problems

    Practice drawing these lewis structures and don't worry we will go over all the answers step by step. This video will explain how to find the formal charges ...

  21. Quiz & Worksheet

    This worksheet and quiz will let you practice the following skills: Reading comprehension - ensure that you draw the most important information from the related how to calculate formal charge ...

  22. Formal Charge

    In BH4, the formal charge of hydrogen is 1-(0+1), resulting in a formal charge of 0. Example 2: Calculate the formal charge on the following: O atoms of O3. Cl atom in HClO4- ion. S atom in HSO4- ion. Ans: We are showing how to find a formal charge of the species mentioned. Formal charge on O1: 6 - 6/2 - 2 = +1. Formal charge on O2: 6 - 4 ...

  23. The Civil Fraud Ruling on Donald Trump, Annotated

    Former President Donald J. Trump was penalized $355 million plus interest and banned for three years from serving in any top roles at a New York company, including his own, in a ruling on Friday ...

  24. Formal Charges

    Formal charge equation is based on the comparing the number of electrons in the individual atom with that in the structure. For each atom, we then compute a formal charge: formal charge = valence e− free atom −(nonbonding e− + bonding e− 2) atom in Lewis structure (1) (1) formal charge = v a l e n c e e − ⏟ f r e e a t o m − ( n o ...