Hund's Rule Explained
According to the first rule, electrons will always occupy an empty orbital before they pair up. This should make sense given what you know about electrons. Electrons are negatively charged and, as a result, they repel each other. Electrons tend to minimize repulsion by occupying their own orbital, rather than sharing an orbital with another electron. Further, quantum-mechanical calculations have shown that the electrons in singly occupied orbitals are less effectively screened or shielded from the nucleus. You will learn more about electron shielding in the next section.
For the second rule, unpaired electrons in singly occupied orbitals have the same spins. Technically speaking, the first electron in a sublevel could be either "spin-up" or "spin-down". Once you've chosen the spin of the first electron in a sublevel, though, the spins of all of the other electrons in that sublevel depend on the spin you chose for the first. To avoid confusion, scientists always draw the first electron, and any other unpaired electron, in an orbital as 'spin-up'.
Applying Hund's Rule
Take a look at the electron configuration for carbon: (Figure 0). Notice how the two 2p electrons in the orbital diagram on the left are in separate orbitals, while the two 2p electrons in the orbital diagram on the right are sharing a single orbital. The orbital diagram on the left is the correct orbital diagram, because it obeys Hund's Rule.
What about oxygen? Oxygen has 8 electrons. The electron configuration can be written as 1s22s22p3. To draw the orbital diagram we diagram the following: the first two electrons will pair up in the 1s orbital; the next two electrons will pair up in the 2s orbital. That leaves 4 electrons, which must be placed in the 2p orbitals. According to Hund’s Rule, all orbitals will be singly occupied before any is doubly occupied. Therefore, we know that two p orbital get one electron and one will get two electrons. Hund's Rule also tells us that all of the unpaired electrons must have the same spin. Keeping with convention, we draw all of the unpaired electrons as "spin-up", which gives (Figure 1)
Purpose of Electron Configurations
What do electron configurations tell us? Soon you will learn that when atoms come into contact with one another, it is the outermost electrons of these atoms, or valence shell, that will interact first. An atom is least stable (and therefore most reactive) when its valence shell is not full. The valence electrons are largely responsible for an element's chemical behavior. Elements that have the same number of valence electrons often have similar chemical properties.
Electron configurations can also predict stability. An atom is at its most stable (and therefore unreactive) when all its orbitals are full. The most stable configurations are the ones that have full energy levels. These configurations occur in the noble gases. The noble gases are very stable elements that do not react easily with any other elements.
Electron configurations can help you to make predictions about the ways in which certain elements will react, and the chemical "compounds" or "molecules" that different elements will form. The principles that you're learning now will help you to understand the behavior of all chemicals, from the most basic elements like hydrogen and helium, to the most complex proteins (huge biological chemicals made of thousands of different atoms bound together) found in your body.