Electron configurations article

The cells in our bodies are masters of quantum physics---they’ve figured out the complicated dance of atoms and electrons, and they use this knowledge to build an endlessly complex series of signalling pathways and genetic circuits. We, however, are not nearly as good at quantum mechanics, and so we must adopt careful notation if we want to keep track of all the complicated situations that the electrons in our bodies find themselves in.

Electron configurations are a simple way of writing down the locations of all of the electrons in an atom. As we know, the positively-charged protons in the nucleus of an atom tend to attract negatively-charged electrons. While these electrons all stick within the atom because of their attraction to the protons, they also mutually repel each other, causing them to spread out around the nucleus in regular patterns. This results in beautiful geometric structures called orbitals that represent the distinct regions around the nucleus that each electron traces out. The reason that electrons tend to stay in their separate orbitals rather than piling on top of one another is the Pauli Exclusion Principle, a theorem from quantum mechanics that dictates that no two electrons can ever be in the same place. The Pauli Exclusion Principle arises from more than just the electrostatic repulsion of negative electrons: it comes from fundamental physical principles that constrain all subatomic particles.

The orbitals represent identifiable “addresses” for each electron around an atom. Think of the electrons as being tenants in one of several blocks of studio apartments located near a nice park. The electrons all want to live near the park (nucleus), but they can’t all live in the same place. Instead, some electrons get to live in the apartments closest to the nucleus, but as the number of electrons that want to live near a given nucleus increases, the further out some of them need to move since apartments closer to the nucleus fill up. This describes a trend observed in the periodic table: elements with small atomic number (and thus fewer electrons) tend to have most of their electrons living in orbitals near the nucleus. As we move further down the periodic table, orbitals and energy levels further out from the nucleus begin to fill up with electrons.

In order to track down where a given electron lives in an atom, you need to know not only how far from the nucleus it is found (which determines its energy level, since electrons further out from the nucleus tend to have higher energy) but also the type of orbital that it can be found in. Think of this as knowing not only which apartment building (energy level) the electron lives in, but also its apartment number. If we drew a map of the apartment buildings where the electrons live, as well as the types of apartments available in each one, it would look something like this:

Notice how each block (energy level) includes one new type of apartment (orbital) with a larger number of residents, in addition to a copy of all of the smaller apartment types.

Because chemists are really interested in keeping track of where all the electrons in a given atom live, they write down a series of symbols called an electron configuration that keeps track of all of this information for a given atom. For example, the electron configuration of oxygen looks like:

The tiny superscripts say how many electrons live in each orbital, the letters represent the orbitals that are available, and the big numbers say which energy level the orbitals are found in. Remember that the total number of electrons just equals the total number of protons, and so the superscripts add up to 8, the atomic number of oxygen. Each energy level (apartment building) has up to four different orbitals (apartments) available, but the number of electrons that can live in each orbital is different depending on the type:

It turns out that there are even more complicated rules for determining which electrons end up in which orbitals as you move further down the periodic table, since subtle effects due to quantum mechanics can encourage electrons to choose apartments further out from the nucleus even when there are openings available in closer apartments. But two general trends remain the same: electrons want to live close to the nucleus, and electrons want to fill their apartments to complete occupancy.

The easiest way to create electron configurations is using an electron configuration table, which is a way of writing down the various orbitals available to electrons. This table is easy to remember, and it makes it possible to generate the electron configuration table for any given element. It looks something like this

The way to use this is to first draw the table, which should be pretty easy to remember because the row numbers correspond to energy levels and the columns correspond to orbital types. The reason that the first two rows have fewer columns involves crowding effects due to the proximity of the electrons to the nucleus. Once you’ve remembered and drawn the table, just follow the arrow (starting from the top) until the subscripts add up to the total number of electrons in your atom. So if you wanted to write down the electron configuration for beryllium (4 electrons), you’d start at the top and pass through 1s, and then you’d loop around until you reached 2s. The sum of the subscripts of the bubbles you’ve passed through is now 4, and so you’re done. Your path spells out the correct configuration for beryllium:

One cool detail that you might notice is that the electron configurations of successive elements (ordered by their periodic number) contain each other. For example, the electron configurations of the first four elements, hydrogen, helium, lithium, and beryllium, look like

and so on for the remaining elements. When chemists write down the really long electron configuration of an atom with a large atomic number (and thus lots of electrons), they’ll sometimes abbreviate the first couple of terms corresponding to the heaviest noble gas with an atomic number smaller than the one they’re studying. For example, sodium looks like

which can then be abbreviated as

[Ne] $3s^1$3, s, start superscript, 1, end superscript

Because Neon has the electron configuration $1s^22s^22p^6$1, s, squared, 2, s, squared, 2, p, start superscript, 6, end superscript. This trick just allows chemists to avoid writing down too many terms, since often only the outer couple of energy levels in an element are actually involved in chemical reactions.

The electron configuration of carbon makes it really good at forming a wide variety of molecules necessary to sustain life. The configuration looks something like

Recall that there are usually six electrons living in an orbital of the type “p,” and so there is a lot of space available (4 openings) for electrons in carbon atoms to share their orbitals with electrons in other atoms. This makes carbon extremely versatile, since it can form stable bonds with a wide variety of atoms including hydrogen, oxygen, and other essential atoms for various metabolic processes. Thus carbon occurs in a wide variety of biochemical compounds, including energy sources, structural building blocks, and essential digestive enzymes. This versatility means that a living organism can “recycle” carbon atoms by easily converting them between various compounds and purposes.