Previous Next

Formal Charge

It is possible to get an estimate of where charges in a molecule are likely to lie by inspecting its Lewis structure and assigning formal charges to specific atoms. To obtain the formal charge of an atom:

  1. Take the atom's group number.
  2. Subtract the number of lone pairs on the atom.
  3. Subtract half the number of bonding electrons.
For example, let's examine the molecule methane, seen in . Carbon is group IV; the carbon atom in methane has zero lone pairs and eight electrons involved in bonds around it. Following the steps listed above: 4 - 0 - 1/2(8) = 0. In this Lewis structure Carbon has no Formal charge.

Remember that formal charges are just a bookkeeping tool and do not necessarily represent actual charges. However, the sum of formal charges on a molecule must equal its net charge.

How to Write Lewis Structures

As you have seen throughout this section, the simplest way to represent and describe molecules is to use a Lewis structure. The Lewis structure model generally follows the octet rule and provides a framework to understand covalent bonding. Lewis structures represent valence electrons as dots and bonding electrons as lines. Lewis structures do not represent inner electrons; only valence electrons are shown.

Here we give a step-by-step procedure for writing valid Lewis structures for any given molecular formula:

  1. Count the total number of valence electrons by summing the group numbers of all the atoms. If there is a net positive charge, subtract that number from the total electron count. If there is a net negative charge, add that number to the total electron count.
  2. Draw single bonds to form the desired connectivity.
  3. Add lone pairs and multiple bonds, keeping the octet rule in mind.
  4. Add formal charges as needed.
An example:
Figure %: Drawing a Lewis Structure for [C2H3O]-
Notice that it is possible to write two separate valid Lewis structures for the given connectivity. This phenomenon is discussed in the next section, Resonance.

Some Common Bonding Motifs in Organic Molecules

You have seen that carbon tends to form four bonds, nitrogen three, oxygen two, and hydrogen/halogens one (remember also: as the number of bonds of an atom decreases, the number of its lone pairs increases). The number of bonds that a neutral atom forms is called its valence. Hence carbon is tetravalent, nitrogen is trivalent, oxygen is divalent, and so on. However, a carbon atom, for example, can be tetravalent in a number of different ways. The following chart shows a number of common bonding motifs for carbon, nitrogen, oxygen, and hydrogen.

Figure %: Bonding patterns
The majority of motifs in the table above obey the octet rule, with one exception. Carbon can form a trivalent species with a positive charge. This phenomenon is known as a carbocation. Such a carbon atom is extremely unstable due to its lack of an octet, but its reactivity makes it a source of a great deal of fascinating chemistry that we will later discuss.

While atoms can occasionally be short of a full octet, elements in the first two rows of the periodic table can never exceed the octet. Students often make the dreaded mistake of drawing pentavalent carbons. Never do this! (However, if you do, rest assured that all organic students make this mistake at sometime or other.) Elements in row 3 and above can exceed the octet by using d-orbitals.

Previous page Covalent Bonds and Lewis Structures page 1 Next section

Take a Study Break