Here's a video discussion of types of bonding and polarity (10min): Crash Course Chemistry: Atomic Hook-ups, on YouTube
The facts described in the previous section, that almost all molecules have all their electrons paired, lead Lewis to the conclusion that electron pairs are of central importance in chemistry. He proposed that in atoms, and especially in molecules, electrons are usually paired. He proposed this theory a little after Bohr's theory of quantum mechanics, which proposed orbits for electrons, and that two electrons could fit in each orbit. Lewis imagined that when two H atoms form a molecule, the two electrons would share an orbit "between" the atoms. He didn't talk about what the orbits would really look like, but he did propose that the electrons could be imagined as sitting on the corners of tetrahedra. When atoms made bonds, they could share electrons on the points of the tetrahedra, as shown. Thus they could make single, double or triple bonds (such as in F2, O2 and N2).
In Lewis' theory, there is no fundamental difference between covalent and ionic bonding. (Actually, he invented the idea of covalent bonding, while Langmuir invented the word "covalent", but it had been known for a while that molecules made of only non-metals had different properties from those made of metals and non-metals together.) Electrons can be shared between atoms, and the sharing is more equal if the atoms are similar or the same, and more unequal if the atoms are different. Two shared electrons make one chemical bond.
| Molecule | bond strength (kJ/mol) | bond length (pm) |
|---|---|---|
| F2 | 155 | 141.8 |
| O2 | 494 | 120.7 |
| N2 | 942 | 109.8 |
Compare the Lewis dot structures for the three diatomic gases shown and the data, which suggests that the bonds get longer and weaker as we go from N to F. If we draw them with single, double and triple bonds, notice that each atom either has or shares eight electrons.
