Lewis Octet Theory
Compounds consisting of nontransition elements usually involve sharing of electrons between atoms so that each atom has a completed outer shell. In the case of hydrogen the outer shell is the 1s orbital which is complete with two electrons. For larger atoms it means that the s and p orbitals are filled with eight electrons, two in the s orbital, six in the p orbital (s2p6). These are the electronic structures of the noble gases. Thus the atoms, in bonding together to form a molecule, are approaching the stability of the noble (or inert) gases. This applies to covalent bonding where the electrons are shared (such as in H2) as opposed to highly ionic bonding, such as in NaCl, where the sodium ion has donated its outer electron the the chloride ion.
Below are examples of atoms with the s and p outer electrons indicated with dots.
Two chlorine atoms can combine to form Cl2 and share two electrons as shown below.
Once you get familiar with writing structures with all the electrons shown you can substitute a line for a shared electron pair and use the dots to represent unshared pairs. In cases where it is unimportant to indicate unshared pairs you can simply indicate shared pairs with lines. Below are three ways to write NH3.
In some compounds two electron pairs are shared in "double" bonds. CO2 is an example.
The double bond is usually shown with two lines.
Formal charges are assigned to an atom by assuming that the bonding electrons are equally divided between the two bonded atoms. This can be accomplished quickly by considering the "bond order". This involves remembering the bond order for a formal charge of zero for an atom. This is 1 for the halides, 2 for the oxygen group, 3 for the nitrogen group, etc. If the bond order is 1 greater than this the formal charge is +1. If the bond order is 2 greater than that for formal charge of zero, the formal charge is +2, etc. If the bond order is 1 less than for zero, the formal charge is -1, and so on. All three of the following bond arrangements for nitrogen result in a formal charge of zero as all have three bonds to nitrogen.
The hydronium ion H3O+ is shown with the formal charge of +1 on oxygen because it has three bonds rather than the two for a zero formal charge.
The neutral molecule below has four bonds to nitrogen and boron. The characteristic bond order of nitrogen is three and that of boron is five. Thus the nitrogen has a formal charge of +1 and boron a charge of -1.
Another way to determine the formal charge is to assign one electron form each bond to the atom, count the total electrons including unshared pairs, and compare this with the number of electrons in the neutral atom. The ammonium ion has four single bonds to the nitrogen. There are no unshared pairs. One electron from each single bond gives four electrons. The nitrogen atom has five electrons. Thus the nitrogen in the ammonium ion is one negative charge short of the neutral atom. The formal charge is +1.
In some cases there are more than one way to write a structure for a molecule. Ozone can be drawn in two ways. Obviously they are both equivalent, but they show two different types of bonds. It is known that both bonds are the same length.
These are known as "resonance" structures. A more accurate way to draw the structure is with partial "hybrid" bonds and fractional formal charges.
Non-equivalent resonance structures can be drawn for some compounds. An example is the cyanate ion.
The nitrate ion NO3- has three resonance structures.
The magic number eight does not apply to some structures containing larger atoms such as sulfur and phosphorus. The oxyacids are examples. Sulfuric acid is best drawn with twelve rather than eight shared electrons.
If the two hydrogen ions are removed from the sulfuric acid structure we have the sulfate anion. The sulfate ion has six possible arrangements of the two single and two double bonds, that is, six resonance structures.