Formal charge is a bookkeeping value used to show how the electrons in a Lewis structure are assigned to each atom. To calculate it, compare an atom’s valence electrons with the nonbonding electrons shown around it and the share of bonding electrons assigned to it.
The result helps you check a Lewis structure, compare possible structures, and identify where charge appears in a molecule or ion. You will need a Lewis structure with lone pairs and bonds drawn clearly, plus the periodic table’s group information for each element.
What You’ll Need
- A Lewis structure showing every bond and lone pair. Include brackets and the overall charge when drawing a polyatomic ion.
- A periodic table, so you can find the number of valence electrons for each atom. For common main-group elements, this usually follows the group number.
- A way to keep track of three quantities for the atom you are checking: its valence electrons, its nonbonding electrons, and its bonding electrons.
- Paper or a calculator if you want to record the arithmetic. The calculation is simple, but writing each term helps prevent counting errors.
Step-by-Step
- Choose one atom. Calculate formal charge separately for each atom in the structure. Do not assign the molecule’s entire charge to the central atom by default. A polyatomic ion’s charge is distributed among its atoms according to the particular Lewis structure.
- Find the atom’s valence-electron count. Use the neutral atom’s number of valence electrons, not its atomic number and not a number changed to match the molecule’s charge. For typical main-group atoms, carbon has four valence electrons, nitrogen has five, oxygen has six, and a halogen such as fluorine has seven. Hydrogen has one. These are useful examples, but check the element’s group when in doubt.
- Count nonbonding electrons on that atom. These are the electrons in lone pairs drawn on the atom. Count individual dots, not pairs: one lone pair contains two nonbonding electrons. If the atom has two lone pairs, it has four nonbonding electrons. Bonding electrons are not included in this count.
- Count the atom’s bonds, then determine its share of the bonding electrons. A single bond contains two electrons, a double bond contains four, and a triple bond contains six. For formal charge, assign half of each bond’s electrons to each bonded atom. Equivalently, count one electron for each bond line connected to the atom: a single bond contributes one, a double bond contributes two, and a triple bond contributes three.
- Substitute into the formal-charge formula. Use formal charge = valence electrons − nonbonding electrons − half of bonding electrons. In the bond-line shortcut, this is formal charge = valence electrons − nonbonding electrons − number of bond lines counted by bond order. Bond order means one for a single bond, two for a double bond, and three for a triple bond.
- Work through a neutral oxygen example. In the usual Lewis structure of water, oxygen has six valence electrons, four nonbonding electrons in two lone pairs, and two single bonds. Its formal charge is 6 − 4 − 2 = 0. Each hydrogen has one valence electron, no lone-pair electrons, and one single bond, so each hydrogen’s formal charge is 1 − 0 − 1 = 0.
- Check a charged atom in a familiar ion. In a Lewis structure for hydroxide, oxygen has six valence electrons, six nonbonding electrons in three lone pairs, and one single bond. Its formal charge is 6 − 6 − 1 = −1. Hydrogen’s formal charge is 1 − 0 − 1 = 0. The charges add to the ion’s overall charge, −1.
- Calculate every atom and add the results. The sum of all formal charges in a Lewis structure must equal the structure’s overall charge. A neutral molecule must have a total of zero; a charged ion must have a total matching its stated charge. This is a key check, but it does not by itself prove that the structure is the best Lewis structure.
- Compare plausible structures when needed. If more than one Lewis structure is possible, calculate the formal charges in each. In general, a useful structure has small formal charges, places negative formal charge on a more electronegative atom when reasonable, and satisfies applicable valence rules. Treat these as comparison guidelines, not a substitute for the full Lewis-structure rules.
For example, a nitrogen atom with one lone pair and three single bonds has five valence electrons, two nonbonding electrons, and three electrons assigned from bonds. Its formal charge is 5 − 2 − 3 = 0. If the same nitrogen is shown with four single bonds and no lone pair, its formal charge is 5 − 0 − 4 = +1. The atom’s formal charge changes because the Lewis structure assigns its electrons differently.
Common Mistakes
- Counting a lone pair as one electron. A lone pair contains two electrons. Count the dots, not the pair symbols.
- Subtracting all bonding electrons. Formal charge assigns half of the bonding electrons to the atom. Subtracting the full number makes the result too negative.
- Forgetting that double and triple bonds count more than once. A double bond gives an atom a share of two electrons; a triple bond gives it a share of three. Counting a double bond as just one bond gives the wrong result.
- Using the ion’s charge as the atom’s valence-electron count. Start with the neutral element’s valence electrons. The structure’s bonds and lone pairs determine the atom’s formal charge.
- Assuming a formal charge is the same as an oxidation state. They are different bookkeeping methods. Formal charge divides each bonding pair equally between the bonded atoms; oxidation-state assignments treat bonding electrons differently based on relative electronegativity.
- Thinking formal charge is a measured partial charge. It is an integer assigned by a Lewis-structure convention. It is not, by itself, a direct measurement of the actual distribution of electron density.
- Checking only the central atom. A correct total charge requires accounting for every atom. Terminal atoms often carry formal charges that are easy to overlook.
- Accepting a structure because its charges add up. The total-charge check is necessary, but a structure can pass it and still break an expected valence rule or represent a less favorable arrangement.
Tips
- Write the three inputs above the formula before calculating: valence electrons, nonbonding electrons, and bond-order count. This makes it easy to see which term caused an unexpected result.
- Use the shortcut only after you understand what it represents. “One per bond line” means one per bond order, not one per connection regardless of whether the bond is single, double, or triple.
- Mark lone pairs clearly on the Lewis structure before counting. If a pair is omitted, the formal charge can change even when the atoms and bonds appear unchanged.
- For ions, put the whole structure in brackets and label the overall charge. Then compare that charge with the sum of the atom-by-atom formal charges.
- Keep the distinction between formal charge and actual electron density in mind when interpreting a structure. Formal charges are most useful for consistent Lewis-structure bookkeeping and comparison.
- When comparing resonance structures, calculate each one separately. The positions of atoms stay the same while electron placement changes, so formal-charge locations may differ between contributing structures.
Troubleshooting
The atom has a formal charge that seems unexpected. Recount its dots and bond orders. A common source of error is missing a lone pair or treating a double bond as one electron assigned to the atom. Then verify that you used the neutral element’s valence-electron count.
The charges do not add up to the ion’s stated charge. Recalculate each atom rather than changing one result to force the total. Check the overall electron count used to draw the Lewis structure, whether the brackets show the correct ion charge, and whether every bond and lone pair has been included. A mismatch usually signals an incomplete or incorrectly drawn structure.
An atom appears to have too many bonds. Check whether the element can be represented with that bonding pattern under the rules used in your course, and confirm that the Lewis structure has the correct total number of electrons. Do not try to repair the result by altering formal charges alone. Some elements have bonding patterns that need additional context, so follow the conventions taught for the molecule and element in question.
Several structures seem possible. First ensure that each candidate uses the correct total number of valence electrons and has the right overall charge. Then compare formal-charge magnitudes and locations alongside octet or other applicable valence requirements. If a class or problem specifies a preferred convention, use that convention consistently.
The answer is a fraction or does not match the expected integer. In an ordinary Lewis-structure calculation, formal charge is an integer for each atom. Check that you divided bonding electrons equally, and that your bond-order count matches the number of shared pairs. A fractional value may mean you have mixed formal charge with another way of describing charge distribution.
Frequently Asked Questions
What is the formula for formal charge?
Formal charge equals valence electrons minus nonbonding electrons minus half the bonding electrons. In a Lewis structure, the last term can be counted as one per single bond, two per double bond, and three per triple bond.
Can formal charge be positive, negative, or zero?
Yes. A positive value means the atom is assigned fewer electrons than the neutral atom has as valence electrons; a negative value means it is assigned more. Zero means the assigned count matches its neutral valence-electron count.
Does a neutral molecule have to give every atom a formal charge of zero?
No. A neutral molecule’s formal charges must sum to zero, but individual atoms can have nonzero formal charges that balance one another.
Why do I count half of a bond’s electrons?
Formal charge uses an equal-sharing convention for bonding electrons. Each atom in a bond is assigned one electron from a single shared pair, even if the atoms differ in electronegativity.
Is formal charge the same as the charge written on an ion?
No. The ion’s overall charge belongs to the whole species, while formal charge is assigned atom by atom in a particular Lewis structure. The individual formal charges should add up to the overall charge.
What should I do if two Lewis structures have different formal charges?
Check that both structures use the correct total electrons and satisfy relevant valence rules. Then compare their formal charges; structures with smaller charges and sensible charge placement are generally preferred, though other chemical considerations can matter.
In short, count an atom’s valence electrons, subtract its lone-pair electrons and its share of bonding electrons, and repeat for every atom. Finally, confirm that the formal charges add to the overall charge and that the Lewis structure itself is reasonable.