Covalent Bonds: The Covalent Bond

So far we have only dealt with very simple, uncharged molecules. For more complex molecules and molecular ions, it becomes important to keep an accurate count of the number of electrons in the molecule. For example, let us make a Lewis structure for NO₂^-. We have five electrons from N, twelve from the oxygen (six from each O), and one extra electron because the molecule has a negative charge. Therefore, NO₂^- has a total of eighteen electrons and we should draw the following Lewis structure:

If we had tried to draw the above structure without taking the charge of the ion into account, we could not have produced a full octet around at least one atom. If the ion had been positively charged, as in NO₂⁺, we would count the electrons as follows: five from N, twelve from O, and minus one due to the charge. The total number of electrons is sixteen for NO₂⁺, and the molecule will have a Lewis structure different from that of NO₂^- because it has a different number of electrons.

To improve your skills in writing Lewis structures, you should draw as many molecules as possible until you feel confident in your ability to draw Lewis structures.

Formal Charge

When trying to draw the Lewis structures of charged molecules like NO₂^-, we encounter the problem of trying to tell where the negative charge is located. Is it on nitrogen or on one of the oxygens? To combat these troubles, chemists have devised the notion of formal charge. Using the Lewis structure and the rules for assigning formal charges, we can assign a formal charge to each atom in a Lewis structure to determine where the charges are located.

Using NO₂^- as an example, let's discuss how to determine the formal charges on atoms in molecules. First, we must draw the correct Lewis structure. Then, we break all bonds around each atom giving half the electrons in the bond to each bonded atom. All lone pairs remain on the atom to which they belong in the molecule. This process serves to count the number of electrons each atom has in the molecule and is shown in the figure below.

Figure %: How to determine the electron count for formal charges

Once we have counted the number of electrons assigned to each atom, we compare the number to the number of valence electrons in the free atom. For example, oxygen has six electrons in the free atom, and it has six electrons in the right-hand oxygen in the . Therefore, the right-hand oxygen has no formal charge because it has the same number of electrons in the NO₂^- molecule as it does as an atom. The left-hand, singly bonded oxygen has seven electrons-- one more electron than has the free atom. Therefore, this oxygen has a -1 formal charge because it has one more electron in the molecule than oxygen has as a free atom. The nitrogen has five electrons around it and five valence electrons in the free atom, so the N has no formal charge. In general, formal charge equals the difference between the number of valence electrons of the atom and the number of electrons around the atom in a molecule as assigned by the rules for drawing Lewis structures.