Hund's Rules

In the Aufbau section, we learned that electrons will fill the lowest energy orbitals first, and then move up to higher energy orbitals only after the lower energy orbitals are full. If you think carefully, though, you'll realize that there's still a problem. Certainly, 1s orbitals should be filled before 2s orbitals, because the 1s orbitals have a lower value of n, and thus a lower energy. What about the three different 2p orbitals? In what order should they be filled? To answer this question, we need to turn to Hund's Rule.

Hund's Rule states that:

1. Every orbital in a sublevel is singly occupied before any orbital is doubly occupied.
2. All of the electrons in singly occupied orbitals have the same spin (to maximize total spin).

When assigning electrons in orbitals, each electron will first fill all the orbitals with similar energy (also referred to as degenerate) before pairing with another electron in a half-filled orbital. Atoms at ground states tend to have as many unpaired electrons as possible. When visualizing this processes, think about how electrons are exhibiting the same behavior as the same poles on a magnet would if they came into contact; as the negatively charged electrons fill orbitals they first try to get as far as possible from each other before having to pair up.

Example 1: Nitrogen Atoms

If we look at the correct electron configuration of the nitrogen (Z = 7) atom: 1s2 2s2 2p3

We can clearly see that p orbitals are half filled as there are three electrons and three p orbitals. This is because the three electrons in the 2p subshell will fill all the empty orbitals first before pairing with electrons in them.

Keep in mind that elemental nitrogen is found in nature typically as dinitrogen, $$N_2$$, which would require the filling in of molecular orbitals instead of atomic orbitals as demonstrated above.

Example 2: Oxygen Atoms

If we look at oxygen (Z = 8) atom, the element after nitrogen in the same period, its electron configuration is: 1s2 2s2 2p4

Oxygen has one more electron than nitrogen and as the orbitals are all half filled the electron must pair up.

Keep in mind that elemental oxygen is found in nature typically as dioxygen, $$O_2$$, which would require the filling in of molecular orbitals instead of atomic orbitals as demonstrated above.

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'.

Example: Carbon and Oxygen

Consider the electron configuration for carbon atoms: 1s22s22p2: The two 2s electrons will be in same orbitals, while the two 2p electrons will be in different orbital (and aligned the same direction) to obey Hund's Rule.

Consider the electron configuration of oxygen? Oxygen has 8 electrons. The electron configuration can be written as 1s22s22p4. 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.

23:16, 25 Aug 2014

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