The formal charge is a hypothetical charge calculated for atoms in a molecule with the assumption that all bonding electrons are equally shared by the two bonding atoms. Formal charges are only used to identify the most stable Lewis structure from various potential structures. For example, CO2 can in theory be represented by one of the following Lewis structures . . . .
C
O••••
C
O••••••
O
C••••
Although the 3rd structure can be eliminated because the central atom (O) is not the least electronegative atom (C), it is listed to show that it is also eliminated based on its formal charge.
Activity: click Show Formal Charges to display each atom's formal charge in the Lewis Structures for CO2. Note that the most stable arrangement of atoms and bonds is the structure whose atoms have a zero formal charge.
Calculating Formal Charge
Use the steps below to calculate the formal charge on each atom in a molecule or ion:
- Draw the Lewis structure to show how atoms are connected and how electrons are shared.
- Determine the number of valence electrons each atom contributed to the Lewis structure.
- H, Li, Na, K, Rb, Cs1 valence electron
- Be, Mg, Ca, Sr, Ba2 valence electrons
- B, Al3 valence electrons
- C, Si4 valence electrons
- N, P5 valence electrons
- O, S6 valence electrons
- F, Cl, Br, I7 valence electrons
- For each atom in the Lewis structure
- Count its number of lone (i.e. "non-bonding") electrons - a "lone pair" is 2 electrons, a radical is 1 electron. The "N" in the ••N•=O•••• radical has 3 lone electrons . . . . the "O" has 4 lone electrons.
- Count its number of bonding electrons - the number of electrons involved in bonds with other atoms . . . .
- a single bond is created from 2 bonding electrons
- a double bond is created from 4 bonding electrons
- a triple bond is created from 6 bonding electrons
The "N" in the ••N•=O•••• radical has 4 bonding electrons . . . . the "O" also has 4 bonding electrons.
- Calculate the formal charge on a specific atom using the formula below:
Atom's Formal Charge = Atom's Valence Electrons – Atom's Lone Electrons – 1
2 Atom's Bonding Electrons
Nitrogen's Formal Charge in ••N•=O•••• = 5 – 3 – ( 1
2 × 4 )
Nitrogen's Formal Charge in ••N•=O•••• = 0
Activity: determine the formal charge of the blue atom in structures below. Click the Show Answer text to check your answer.
| Molecule / Ion | Valence Electrons | Non-Bonding Electrons | Bonding Electrons | Formal Charge |
| ••O
C O•••••• |
6 | 2 | 6 |
Show Answer |
| O••••
C O•••• |
4 | 0 | 8 |
Show Answer |
| O••••
O C•••• |
6 | 0 | 8 |
Show Answer |
| ••C
N•• — |
4 | 2 | 6 |
Show Answer |
| ••C
N•• — |
5 | 2 | 6 |
Show Answer |
Observe that the sum of a structure's formal charges equals the charge on the structure.
The cyanide ion (:C≡N:–) has an overall charge of –1.
The C atom in :C≡N:– has a formal charge of -1.
The N atom in :C≡N:– has a formal charge of 0.
The sum of ‐1 and 0 equals ‐1 (the charge of the cyanide ion).
Using Formal Charges to Predict Molecular Structure
Three Lewis structures for Carbon dioxide were presented as the possible structure of CO2. The structure where all formal charges equal zero was deemed the most stable and the accepted Lewis structure of CO2.
C
O••••
Most Stable
C
O••••••
O
C••••
The following guidelines will help determine the most likely structure of a molecule when more than one Lewis structure can be written:
- A Lewis structure where all formal charges are zero is preferable to one in which some formal charges are not zero.
- If nonzero formal charges cannot be avoided, the Lewis structure with the smallest nonzero formal charges is preferable.
- A Lewis structure where adjacent formal charges are zero or of the opposite sign is more stable.
- If Steps 1-3 have not provided the most stable Lewis structure, structures where a negative formal charge is on the more electronegative atom is preferable.
Activity: complete the HW 7.4a: Formal Charges assignment. You will determine the formal charge of individual atoms in molecules and ions and then use the guidelines above to select the most stable Lewis structure from a list of possible structures.
Resonance
The Lewis structure of the nitrite ion, NO2‐, is correctly drawn with the least electronegative atom (Nitrogen) as the central atom. However, there are two ways to draw the NO2‐ ion and assigning formal charges does not reveal the most stable Lewis structure.
N••
O•••••• ]–
[
O••••••N••
O•••• ]–
Are both correct? Yes and No. Both are correct Lewis structures, but neither is the correct actual structure. The actual molecular structure is a . . . the average of the various resonance forms. of the two resonance forms.
A great analogy was given by George Wheland, a pioneer of resonance theory . . . .
A medieval traveler happened upon a rhinoceros and set about describing it in his journal. He chose two fictitious animals in his description because the rhinoceros had many attributes common to both. He wrote that the rhinoceros was a hybrid of a dragon and a unicorn.
Just as a rhinoceros does not alternate between a dragon and a unicorn, neither does a resonance hybrid switch between its resonance forms. A resonance form (or resonance structure) is as imaginary as is a dragon or a unicorn . . . . but, the resonance hybrid is as real as a rhinoceros. A resonance hybrid is not a mixture of its resonance forms, but rather a stable, real entity with attributes and properties of its resonance structures.
When a molecule or ion is represented by a single Lewis structure, we obtain a great "picture" of the actual molecule. When two or more Lewis structures can be drawn for a molecule or ion we can still get a great "picture" of the actual molecule, however, we must perform some mental gymnastics to merge the various resonance forms into a composite image that helps us envision the actual molecule.
Activity: complete the HW 7.4b: Drawing Resonance Structures assignment. Use the interactive Lewis Draw program to draw the resonance forms of selected molecules and ions.