A Quick Review of Chemical Bonds
A Crystal Clear Chemistry Tutorial
- Molecular Orbitals
- Electronegativity
Chemical bonds are what hold atoms in molecules together, and they are formed by the interactions of the electrons in each atom. In almost all circumstances, only the valence electrons (that is, the electrons in the uppermost energy level of an atom) interact with each other to form the bonds.
Pi-Bonds and Sigma-Bonds
We can think of a bond as the overlap of atomic orbitals after they've been hybridized. For instance, in water (H2O), oxygen has four sp3-hybridized orbitals, and the bond between oxygen and hydrogen can be viewed as the overlap between an sp3-hybridized orbital and an s orbital from hydrogen. You can see the overlap in the figure below, which is shaded grey.
There are essentially two basic ways that orbitals can overlap with each other to form a bond. Consider two atoms with one p-orbital each that will overlap to form a bond. The orbitals can either bond head-on, so that they touch only once, or they can bond side-by-side, touching in two places. The first type, the head-on bond, is called a sigma-bond (usually denoted with the lowercase Greek letter sigma, σ), and the second type, the side-by-side bond, is called a pi-bond (denoted with π). In the figures below, an example of the sigma bond is on the left and the pi-bond is on the right.
A sigma bond is almost always stronger than a pi bond, because the overlap of the orbitals is bigger with sigma bonds. Thus, pi bonds tend to be more reactive (they're easier to break). Double bonds are composed of one sigma and one pi bond, while triple bonds are composed of one sigma and two pi bonds. Thus, double and triple bonds are stronger overall than signle bonds (because they have additional pi bonds involved) but they also tend to be much more reactive than single bonds (the pi bonds often break to leave just the sigma bond there).
Molecular Orbitals
When two orbitals overlap, they actually combine and form two new orbitals, called "molecular orbitals" (or MOs): one is called the "bonding" orbital, and the other is called the "antibonding" orbital (the symbol for the antibonding orbital usually has a star next to it). The full explanation as to why this occurs needs some heavy duty quantum physics, but all you need to know is that 2 new molecular orbitals form for any two overlapping orbitals, and that bonding orbitals tighten the bond between atoms while the antibonding orbitals pull the atoms apart. You can see a diagram of the two orbitals formed from two p orbitals in a pi-bond. The lower energy bonding orbital is called π (pi) and the higher energy anti-bonding orbital is called π* (pi star). For a more in-depth explanation, read the box below, though it is by no means necessary for the understanding of the rest of the article.
The overlap of the atomic orbitals to form molecular orbitals is essentially the interference of waves to form new waves. In quantum mechanics, orbitals have something called phase, which can be positive or negative. Don't confuse this with charge; the positive and negative phases have nothing to do with the positive or negative charges. Instead, they're just ways to label the different phases, and in fact, most organic chemists just denote the phases as "dark" and "light" in pictures by coloring in orbitals. The p orbital, as you can see, has one half dark and one half light, denoting that there is a change of phase between the different lobes of the orbital.
When two orbitals overlap and "interfere" with each other, if the phases match, then there is "constructive" interference, and the overlapping regions add to each other. When two orbitals are different in phase, then there is "destructive" interference and the overlapping regions substract from each other. You can get a visual sense of this interference concept from the picture of the two waves below adding and substracting from each other depending on the phase. This is exactly analogous to the formation of bonding and antibonding molecular orbitals from the overlap of two atomic orbitals.
Because the phases of orbitals can rapidly shift to just about anything, both bonding and antibonding orbitals form. If there are two electrons in the bond, then these electrons both occupy the lower energy orbital (the bonding orbital) and contribute to the stability of the chemical bond. When two more electrons from somewhere else fill the antibonding orbital, then the bond will break.