Organic Chemistry Tutorials

What is Formal Charge

structures of O3 and NO3, with the formal charges labeled

Core Concepts

Formal charge is an essential, basic concept to master in order to better understand molecular structures and reactions. In this tutorial, you will learn what is formal charge, how to calculate it, and its significance in practice.

Topics Covered in Other Articles

Vocabulary

  • Ion: an atom or molecule with a net ionic charge, due to the presence or lack of electrons.
  • Molecule: a group of atoms bonded together
  • Resonance: a term used to describe the hybrid bonding in a molecule with multiple possible structures.

What is Formal Charge?

Formal Charge is a charge assigned to an atom under the assumption that all electrons in bonds are shared equally. This is a hypothetical measure, not a real representation of the actual charge on an atom, which looks at the ways electrons are actually shared between atoms in a bond. But more on that later!

How to Calculate:

Formal Charge (FC) = (# of valence electrons) – (½)(number of bonded electrons) – (number of unbonded electrons)

Examples:

calculate formal charge from the lewis dot structure of NH3

NH3: what is the formal charge on the nitrogen?

Nitrogen has 5 valence electrons, 6 bonded electrons (as there are 3 single bonds, each containing 2 electrons), and 2 unbonded electrons in this configuration. Thus:

FC = (5) – (½)(6) – (2)
FC = 0

calculate formal charge from the lewis dot structure of CH3O

CH3O: what is the formal charge on the carbon?

Carbon has 4 valence electrons, 8 bonded electrons (two single bonds and one double bond), and no unbonded electrons. Thus:

FC = (4) – (½)(8) – 0
FC = 0

Note: though the formal charge in these two examples has been zero, that will not always be the case. We will explore some examples of nonzero charge below.

Significance

1. Molecular Structure

Ideally, an atom in a molecule wants to have a formal charge of zero: this is the lowest energy, and thus the most stable state for it to be in. This clues us into the structure of a molecule if there are multiple options: the one with the least/lowest formal charges is the preferred structure. There are even specific guidelines to help you figure this out:

  1. The preferred molecular structure is one where all formal charges are zero, as opposed to one where some formal charges are not zero.
  2. If there is no possible structure where all formal charges are zero, then the preferred structure is one with the least number of nonzero charges.
  3. Adjacent atoms in a molecule should have opposite signs if charges are present.
  4. If there are multiple structures that satisfy requirements 1-3, then the structure with negative formal charges on the more electronegative atoms is preferred.

Example: shown below are three possible structures for N2O. Let’s figure out which structure is correct.

different possible structures of N2O
  1. The top structure:
    First, we calculate the formal charge of the nitrogen on the left. Nitrogen has 5 valence electrons, this atom has 6 bonded electrons (a triple bond), and 2 unbonded electrons, thus the formal charge is (5) – (½)(6) – (2) = 0.
    Next, we calculate the formal charge of the nitrogen in the middle. This one has 8 bonded electrons and no unbonded, thus the formal charge is (5) – (½)(8) – (0) = +1.
    Finally, we calculate the formal charge of the oxygen. Oxygen has 6 valence electrons, and this atom has 2 bonded electrons and 6 unbonded, thus the formal charge is (6) – (½)(2) – (6) = -1.

  2. The middle structure:
    Similarly, we calculate the formal charge of the nitrogen on the left: (5) – (½)(4) – (4) = -1.
    Next, the formal charge of the nitrogen in the middle: (5) – (½)(8) – (0) = +1.
    Finally, the formal charge of the oxygen: (6) – (½)(4) – (4) = 0.

  3. The bottom structure:
    Again, first we calculate the formal charge of the nitrogen on the left: (5) – (½)(2) – (6) = -2.
    Next, the formal charge of the nitrogen in the middle: (5) – (½)(8) – (0) = +1.
    Finally, the formal charge of the oxygen: (6) – (½)(6) – (2) = +1.

Given these calculated formal charges, let’s consult the guidelines discussed above. First, are there any structures possible where all the formal charges are zero? There are not, so we move on to rule #2. This eliminates the bottom structure, as it has a greater number of nonzero charges than the top two (it also has greater charges, as it contains a -2 charge, whereas the other two only contain +/-1). Both the top and middle structures have adjacent atoms with opposite charges, so both satisfy rule #3. This leaves rule #4, meaning the preferred structure is the one with the negative charge on the more electronegative atom. Oxygen is more electronegative than nitrogen, meaning the preferred structure is the one with a negative charge on the oxygen—the top structure!

It is also worth noting that the sum of all the formal charges of the atoms in a molecule must equal the overall charge on the molecule/ion. That is, they should sum up to zero if its an neutral molecule, and should sum up to the ion’s charge if it is not.

Example: the ion BH4 has an overall charge of -1. This means that the formal charges of all the individual atoms in it should add up to -1. Let’s see if this is true.

BH4 - ion

Boron has three valence electrons, eight bonded electrons, and zero unbonded electrons. This makes its formal charge: (3) – (½)(8) – (0) = -1.

The four hydrogens in this molecule are all identical, thus we can calculate all of their formal charges at once. Hydrogen has one valence electron, two bonded electrons, and zero unbonded electrons. This makes its formal charge: (1) – (½)(2) – (0) = 0.

As we can see, the formal charges add up to 0 + 0 + 0 + 0 + (-1) = -1. This sum does equal the overall charge on the ion, which is -1.

2. Resonance

While formal charge can indicate the preferred structure of a molecule, as discussed above, the situation gets a bit more complicated when there are multiple equally preferred structures. This situation may indicate resonance structures, particularly when the structures have the same arrangement of atoms, but different types of or arrangement of bonds.

Example: the diagram below shows three possible structures for the ion CO32-. We can see that the arrangement of atoms is the same in all three structures (with the carbon in the center, connected to the three oxygens), but the placement of the double bond differs in each of the three.

resonance structures of CO3 2-

In each of them, the formal charge on the center carbon is 0, the double bonded oxygen is 0, and the two single bonded oxygens are each -1. See if you can calculate these yourself correctly! Note that as discussed above, 0 + 0 + (-1) + (-1) adds up to -2, which is the overall charge on the ion.

Since these bonds are the same in all three structures, their placement in the molecule is just different, the formal charges and distributions of it in each structure are the exact same, meaning that each of them is equally likely to occur. This means that all three are correct structures, and in reality, the molecule forms a hybrid of all three structures.

Read more about resonance structures here, and see more examples here!

3. Reactivity

Finally, the formal charge can give an indication as to how a molecule will behave during a reaction. If an atom has a negative formal charge, it is more likely to be the source of electrons in a reaction (a nucleophile). Conversely, if it has a positive one, then it is more likely to be accept electrons (an electrophile), and that atom specifically is most likely to be the site of the reaction.

Formal Charge vs. Actual Charge

It is also important to not that formal charge is different from the actual charge of an atom. Formal charge does not take electronegativity into mind: it assumes that electrons in a bond are shared equally. It’s merely a formality, used to help make sense of molecular structures and reaction mechanisms. Actual charge, on the other hand, looks at the actual electron density, based on the atoms’ electronegativities and polarity of the bonds. To read more on these topics, check out these tutorials on: ion-dipole forces, periodic trends, and polarity!

Formal Charge – Further Reading

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