What is the difference between ion charge and formal charge

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 — multiple different bonds 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:. To see how these guidelines apply, consider some possible structures for carbon dioxide, CO 2. It is known that the less electronegative atom typically occupies the central position, but formal charges help understand why this occurs.

Three possibilities for the structure can be drawn: carbon in the center with two double bonds, the carbon in the center with a single and triple bond, and oxygen in the center with double bonds.

On comparing the three formal charges, the structure on the left can be identified as preferable because it has only formal charges of zero. 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: NCS—, CNS—, or CSN—. The formal charges present in each of these molecular structures can help 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 - carbon in the center with double bonds, nitrogen in the center with double bonds, and sulfur in the center with double bonds.

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 with carbon in the center is preferred because it has the lowest number of atoms with non-zero formal charges. Also, it places the least electronegative atom in the center and the negative charge on the more electronegative element. This text is adapted from Openstax, Chemistry 2e, Chapter 7. To learn more about our GDPR policies click here.

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Chapter 4: Chemical Quantities and Aqueous Reactions. Chapter 5: Gases. The formal charges for the two Lewis electron structures of CO 2 are as follows:.

Thus the symmetrical Lewis structure on the left is predicted to be more stable, and it is, in fact, the structure observed experimentally. Remember, though, that formal charges do not represent the actual charges on atoms in a molecule or ion. They are used simply as a bookkeeping method for predicting the most stable Lewis structure for a compound. The Lewis structure with the set of formal charges closest to zero is usually the most stable. Draw two possible structures, assign formal charges on all atoms in both, and decide which is the preferred arrangement of electrons.

Asked for: Lewis electron structures, formal charges, and preferred arrangement. B Calculate the formal charge on each atom using Equation 2. C Predict which structure is preferred based on the formal charge on each atom and its electronegativity relative to the other atoms present.

B We must calculate the formal charges on each atom to identify the more stable structure. If we begin with carbon, we notice that the carbon atom in each of these structures shares four bonding pairs, the number of bonds typical for carbon, so it has a formal charge of zero.

Continuing with sulfur, we observe that in a the sulfur atom shares one bonding pair and has three lone pairs and has a total of six valence electrons.

In b , the sulfur atom has a formal charge of 0. Continuing with the nitrogen, we observe that in a the nitrogen atom shares three bonding pairs and has one lone pair and has a total of 5 valence electrons.

C Which structure is preferred? Note: N is the central atom. In each case, use the method of calculating formal charge described to satisfy yourself that the structures you have drawn do in fact carry the charges shown.

Give the formal charges for all non-hydrogen atoms in the following moelcules:. Steven Farmer Sonoma State University. Objectives After completing this section, you should be able to calculate the formal charge of an atom in an organic molecule or ion.

Key Terms Make certain that you can define, and use in context, the key term below. Study Notes It is more important that students learn to easily identify atoms that have formal charges of zero, than it is to actually calculate the formal charge of every atom in an organic compound. Determining the Formal Charge on an Atom A formal charge compares the number of electrons around a "neutral atom" an atom not in a molecule versus the number of electrons around an atom in a molecule.

To calculate formal charges, we assign electrons in the molecule to individual atoms according to these rules: Non-bonding electrons are assigned to the atom on which they are located. Bonding electrons are divided equally between the two bonded atoms, so one electron from each bond goes to each atom. Example 2. Determining the Charge of Atoms in Organic Structures The calculation method reviewed above for determining formal charges on atoms is an essential starting point for a novice organic chemist, and works well when dealing with small structures.

Carbon Carbon, the most important element for organic chemists. Carbon usually makes four bonds Carbon is tetravalent in most organic molecules, but there are exceptions. Hydrogen The common bonding pattern for hydrogen is easy: hydrogen atoms in organic molecules typically have only one bond, no unpaired electrons and a formal charge of zero. Neutral Hydrogen: one bond, no lone pair. This type of reaction can be recognized because it involves a change in oxidation number of at least one element.

More information on these reactions is found in the section on redox reactions. Oxidation numbers are also used in the names of compounds. The internationally recommended rules of nomenclature involve roman numerals which represent oxidation numbers. Oxidation numbers can sometimes also be useful in writing Lewis structures, particularly for oxyanions. Since sulfur has six valence electrons, we conclude that two electrons are not involved in the bonding, i.

With this clue, a plausible Lewis structure is much easier to draw:. Formal Charge On the page discussing the covalent bond , it is shown that the density of electrons in a covalent bond is shared between both atoms. Formal Charge "Rules" Here are some rules for determining the Formal Charge on each atom in a molecule or polyatomic ion: Electrons within a Lone Pair on an atom are assigned exclusively to that atom. Half of the electrons in each bond around an atom are assigned to that atom.

The Formal Charges on all atoms in a molecule must sum to zero; for a polyatomic ion the Formal Charges must sum to the charge on the ion which may be positive or negative. Oxidation Numbers We have also discussed electronegativity, which gives rise to polarity in bonds and molecules. 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. 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. 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. Draw all possible resonance structures for each of the compounds below. Determine the formal charge on each atom in each of the resonance structures:. The structure with formal charges of 0 is the most stable and would therefore be the correct arrangement of atoms.

The structure that gives zero formal charges is consistent with the actual structure:. The empirical formula is NF 3 and its molar mass is The Lewis structure is.

Skip to main content. Chemical Bonding and Molecular Geometry. Search for:. Formal Charges and Resonance Learning Objectives By the end of this section, you will be able to: 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.

Each Cl atom now has seven electrons assigned to it, and the I atom has eight. Show Answer Assign one of the electrons in each Br—Cl bond to the Br atom and one to the Cl atom in that bond: Assign the lone pairs to their atom.