When summed the overall charge is zero, which is consistent with the overall charge on the NH 3 molecule. Typically, the structure with the most charges on the atoms closest to zero is the more stable Lewis structure. In cases where there are positive or negative formal charges on various atoms, stable structures generally have negative formal charges on the more electronegative atoms and positive formal charges on the less electronegative atoms.
The next example further demonstrates how to calculate formal charges. The nitrogen atom shares four bonding pairs of electrons, and a neutral nitrogen atom has five valence electrons. Each hydrogen atom in has one bonding pair.
The formal charge on each hydrogen atom is therefore. Adding together the formal charges on the atoms should give us the total charge on the molecule or ion. If an atom in a molecule or ion has the number of bonds that is typical for that atom e.
As an example of how formal charges can be used to determine the most stable Lewis structure for a substance, we can compare two possible structures for CO 2. Both structures conform to the rules for Lewis electron structures. C is less electronegative than O, so it is the central atom. C has 4 valence electrons and each O has 6 valence electrons, for a total of 16 valence electrons. Dividing the remaining electrons between the O atoms gives three lone pairs on each atom:. This structure has an octet of electrons around each O atom but only 4 electrons around the C atom.
No electrons are left for the central atom. To give the carbon atom an octet of electrons, we can convert two of the lone pairs on the oxygen atoms to bonding electron pairs. There are, however, two ways to do this. We can either take one electron pair from each oxygen to form a symmetrical structure or take both electron pairs from a single oxygen atom to give an asymmetrical structure:. Both Lewis electron structures give all three atoms an octet.
How do we decide between these two possibilities? 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 4. 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.
Again, move a lone pair from oxygen to make a double bond with the carbons. This gives both atoms a share in a complete octet. The double bond could be between the carbon and any one of the equivalent oxygen atoms.
There is more electron density between the nuclei. However, it has been shown that all three carbon-oxygen bonds are identical. This can be explained if we use all three configurations to explain the structure. When two or more Lewis Structures can be used to describe the same molecule, each is called a resonance structure.
Each has an identical atomic structure , only the position of the electrons has been changed. A double arrow is used to represent resonance structures. Another way of representing these structures is:.
The atoms can be arranged 2 different ways:. This leaves 6 more electrons to complete the octets around C and O. Again, we have to move lone pair electrons to create double bonds. Which of these two structures represents formaldehyde? To determine this, we need to look at the concept of formal charge.
This is a comparison of the number of valence electrons on an isolated atom versus the number of electrons assigned to an atom in a given Lewis Structure.
The number of electrons in an isolated atom is simply the Group number of the element. The number of electrons assigned by a Lewis Structure is the number of lone pair electrons unshared plus one-half of the shared, bonding electrons. In the figure below, the electrons "belonging" to carbon are shown in red, and the electrons "belonging" to oxygen are shown in blue.
For the structure on the left , the formal charges will be:. For the structure on the right , the formal charges will be:. For a neutral molecule, a Lewis Structure with no formal charges is preferred. Small formal charges are preferable to large formal charges. Like charges on adjacent atoms are undesirable. A more negative formal charge should reside on a more electronegative element. The structure on the left puts a positive formal charge on oxygen, an electronegative element.
The structure on the right has no formal charges and is the best structure for formaldehyde. The Lewis dot structures of the individual, non-metal atoms give a good indication of the bonding possibilities for the atoms.
This same method can be used to calculate the number of electrons that are not participating in bonding. The number of non-bonding electrons is equal to the the number of electrons in a full valence shell minus the number electrons which are participating in bonding which is 2 x the typical number of bonds.
The number of lone pairs is the number of non-bonding electrons divided by two. For example, hydrogen typically has 0 non-bonding electrons. The full valence shell for hydrogen is 2 and the number of electrons in bonds is also 2.
The difference is zero. In 5-coordinated molecules containing lone pairs, these non-bonding orbitals which are closer to the central atom and thus more likely to be repelled by other orbitals will preferentially reside in the equatorial plane. Substituting nonbonding pairs for bonded atoms reduces the triangular bipyramid coordination to even simpler molecular shapes. Boundless vets and curates high-quality, openly licensed content from around the Internet.
This particular resource used the following sources:. Skip to main content. Advanced Concepts of Chemical Bonding. Search for:. Lone Electron Pairs.
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