Summary and Terms

Summary

It's impossible to understate the importance of understanding atomic bonding, so much that Richard Feynman, a giant in physics, thought it was the most important piece of information. It is certainly true that without a clear understanding of the way atoms interact, we wouldn't be able to work the scientific wonders we all take for granted every day. Atomic Bonding is an incredibly complicated phenomenon with nuance far beyond what we can cover in this class, but there are several simple models we used in this chapter to navigate the phenomenon in useful ways.

A simple way of thinking about atomic bonds is the electrostatic interpretation which you might have seen in prior courses - it's an approachable and useful model for defining bonds. Atoms consist of positively charged nuclei and negatively charged electrons which do not fly around like little projectiles but constitute a probabilistic cloud that surrounds the nucleus. As you saw in the simulations, if we simply introduce the concept that like charges repel one another and opposite charges attract one another, we see a rudimentary form of bond emerge!

To add further complexity to our simple model, we considered the concept of electronegativity, a measure of how strongly an atom attracts electrons toward itself. The force that the atom's outermost electrons feel is known as the effective nuclear charge, and can be roughly approximated by subtracting the number of inner shell electrons from the atom's nuclear charge. In considering effective nuclear charge and the distance an electron is away from the nucleus, we can back out important information like periodic trends in atomic size.

All atomic bonds are the result of increased electron density between two nuclei, but with electronegativity under our belt, we defined three general types of atomic bonds: Covalent, Ionic and Metallic.

Covalent bonds tend to occur between atoms of similar, but overall high, electronegativities, causing increased electron density between the nuclei which is shared roughly equally. Ionic bonds form between nuclei with very different electronegativities, so the electron density is pulled more strongly to the more electronegative element. Finally, metallic bonds occur between atoms of similar but low electronegativities which, like in the covalent case, results in roughly equal sharing of electron density. However, metallic bond electrons are much more delocalized than in the covalent case due to the lower electronegativities of the involved atoms.