Formal charge is a bookkeeping tool that helps chemists track how electrons are assigned in a Lewis structure. It does not always equal the real electrical charge on an atom, but it helps compare possible structures for the same molecule or ion. Learning formal charge is important because many molecules can be drawn in more than one way, and not all drawings are equally reasonable.
A good Lewis structure usually keeps formal charges small and places negative charge on more electronegative atoms.
To calculate formal charge, compare an atom's valence electrons to the electrons it owns in the Lewis structure. Lone pair electrons count fully for the atom, while bonding electrons are split equally between bonded atoms. Chemists use these values to choose the best structure, check resonance forms, and predict reactive sites.
Formal charge is especially useful for ions, expanded octets, resonance structures, and molecules with multiple possible bonding patterns.
Understanding Chemistry: Formal Charge
Start formal charge work only after the basic Lewis structure is complete. First count every valence electron available, including electrons added for a negative ion or removed for a positive ion. Choose a reasonable central atom, connect the atoms, and fill outer octets before changing single bonds into multiple bonds.
This order matters. A student can get a correct formal charge for a drawing that used the wrong total number of electrons. The final drawing must conserve electrons, give the correct overall ionic charge, and follow normal octet patterns where possible.
Formal charge becomes especially useful when several bond arrangements seem possible. Consider the nitrate ion. One drawing can show nitrogen with one double bond to oxygen and single bonds to two other oxygens.
Moving the double bond creates two more valid drawings. These are resonance forms, not molecules switching back and forth between separate shapes. The real ion has electron density spread across all three oxygen atoms.
Experiments show that its nitrogen to oxygen bonds have the same average length. Formal charges help identify these related drawings and explain why no single one shows the whole electron arrangement.
The usual rules need careful use because some elements behave differently. Second period atoms such as carbon, nitrogen, oxygen, and fluorine generally cannot exceed an octet. Third period elements such as phosphorus and sulfur can appear with more than eight electrons in some Lewis structures.
Boron often has fewer than eight electrons in compounds such as boron trifluoride. A structure with a small formal charge is not automatically best if it forces carbon or oxygen beyond an octet. Some stable species must contain charge separation.
In ozone, for example, one oxygen carries a positive formal charge while another carries a negative formal charge in each resonance form. The charges are unavoidable within the Lewis model.
Formal charge gives useful clues about chemical behavior. An atom shown with negative formal charge often has electron density available for bonding to an electron poor atom. An atom shown with positive formal charge can be a likely target for electron rich particles.
This helps students later with acids, bases, reaction mechanisms, and organic functional groups. Still, formal charge is not a map of exact physical charge. Bond polarity, molecular shape, and resonance can spread charge across several atoms.
When checking a structure, track the total electron count, the octet rule, the overall charge, and every atom with a nonzero formal charge. Then compare alternative structures rather than trusting the first drawing that seems to fit.
Key Facts
- Formal charge = valence electrons - nonbonding electrons - 1/2 bonding electrons.
- For a single bond, each bonded atom is assigned 1 electron from the bond.
- For a double bond, each bonded atom is assigned 2 electrons from the bond.
- The sum of all formal charges must equal the overall charge of the molecule or ion.
- The best Lewis structure usually has the smallest possible formal charge magnitudes.
- When formal charges cannot be avoided, negative formal charge is usually best placed on the more electronegative atom.
Vocabulary
- Formal charge
- The charge assigned to an atom in a Lewis structure by assuming bonding electrons are shared equally.
- Valence electrons
- Electrons in the outer shell of an atom that are available for bonding and lone pairs.
- Lone pair
- A pair of valence electrons localized on one atom and not shared in a bond.
- Lewis structure
- A diagram that shows atoms, bonds, lone pairs, and sometimes formal charges in a molecule or ion.
- Resonance structure
- One of two or more valid Lewis structures that differ only in electron placement, not atom positions.
Common Mistakes to Avoid
- Counting all bonding electrons for one atom is wrong because formal charge splits each bond equally between the bonded atoms.
- Forgetting lone pairs is wrong because nonbonding electrons are fully assigned to the atom when calculating formal charge.
- Choosing a structure only because every atom has an octet can be wrong because formal charges and electronegativity also affect which structure is best.
- Ignoring the total charge is wrong because the sum of formal charges must match the molecule or ion charge exactly.
Practice Questions
- 1 Calculate the formal charge on nitrogen in NH4+ if nitrogen has four single bonds and no lone pairs.
- 2 In CO2 drawn as O=C=O, calculate the formal charge on each oxygen and on carbon.
- 3 Two Lewis structures for NCO- place the negative formal charge either on nitrogen or oxygen. Which placement is usually more favorable, and why?