In this section, we will cover two important aspects before starting an in-depth investigation of covalent bonds. The first aspect is a property of electrons, known as electron spin. The second topic is electron configuration, which is basically the art of designating each electron its "address."

Think about a spinning top. It can only spin in two different directions, namely clockwise and counterclockwise. Likewise, in any orbital, the two electrons (assuming both electrons exist) must have opposite spins. One will spin in a "counterclockwise" direction while the other spins in a "clockwise" direction.
This is known as the Pauli exclusion principle, named after Wolfgang Pauli. Look at the first figure, which is the Hydrogen 1s orbital. Because it has only one electron, the electron's spin is designated by an up arrow. The second figure is helium, which has two electrons, and has two arrows in opposite directions. This is in compliance with the Pauli exclusion principle.

Now that we understand electron spin, lets look at the quantum numbers, which will lead to a better understanding of electron configuration.

Quantum numbers are simply a set of numbers used to represent the location of any one electron. For example, your address, which is a set of numbers and letters, represents only your house out of all the millions of houses in the world. There are four different quantum numbers: the principal, azimuthal, magnetic, and the spin quantum numbers. Look at the following chart.

 Quantum Number Symbol of quantum number Description of quantum number principalazimuthalmagneticspin nlms energy level of electronsublevel (s, p, d, f)specific orbitalspin direction

The arrangement of electrons in an atom is called the electron configuration of the atom. In a special notation, all the quantum numbers of a single electron can be shown. Lets look at two different examples, helium and oxygen. In helium, which has only two electrons, the 1s orbital is filled up. Thus, we write 1s². The 1 designates the principal quantum number, which is the first energy level of the electron. The s is the only sublevel possible for the first energy level. The 2 is the number of electrons. This number can be no more than 2 and no less than 0, since only two electrons can be in the same orbital. Now lets look at oxygen. It has eight electrons, since its atomic number is 8. We can use the first two electrons for the 1s orbital, giving them opposite spins. Now we have six electrons left. The 2s orbital is the next lowest in energy after the 1s and we can put two electrons there, also with opposite spins. We have four electrons left, which we can put in the 2px, 2py, and 2pz orbitals. There are some possible choices: we could fill up one 2p orbital at a time, or we could fill up separate orbitals, one at a time, as illustrated below. What should we do? Before we continue, we need to shed some light on Hund's rule of maximum multiplicity. It states that the lowest energy configuration possible (remember, atoms love lowest energy states) is that in which as many electrons as possible reside in different orbitals with all their spins in the same direction.
What does that mean? It means we should put one electron in the 2px orbital, another in the 2py orbital, and one more in the 2px orbital, all with the same spins. But wait, we still have one electron left! The final electron will go back to the 2px orbital, but with a spin opposite to the first one. Note - You cannot put any electron in a higher energy orbital, like a 3s orbital, until the lower energy orbitals have been filled up. The electron configuration, therefore, for oxygen is 1s² 2s² 2p4.

We have come a long way through the quantum numbers. On the next page is a list of the elements from boron to sodium. It is merely a reinforcement to what you learned now. Take some time to review and understand how and why each atom has the given electron configuration. Then, when you are ready, the next page will introduce covalent bonds. Press 'Next Section' to continue.