Electron configuration describes the positions (orbital shapes and energy levels) of electrons in a certain atom or ion.
Notation
Written in the form of nlᵐ, where:
– n is the principle energy level
– l is the orbital letter (s, p, d, or f)
– m is the number of electrons in that orbital
Ex. 2s² represents an s subshell in the second principal energy level that contains two electrons.
Ex. 3d⁵ represents a d subshell in the third principal energy level that contains five electrons.
Ex. 3d⁶ represents a d subshell in the third principal energy level that contains six electrons.
When writing electron configuration, we use this notation to describe all electrons in an atom. For example, a neutral sodium atom has an electron configuration of ↴
1s² 2s² 2p⁶ 3s¹
This means that sodium has ↴
– 2 electrons in an s-subshell in the first principal energy level
– 2 electrons in an s-subshell in the second principal energy level
– 6 electrons in a p-subshell in the second principal energy level
– 1 electron in an s-subshell in the third principal energy level
Notice how when you add the superscripts, they add up to 11, which is how many electrons sodium has.
The superscripts in electron configuration must add up to the total number of electrons in the atom because the superscripts represent the number of electrons.
Keep reading to understand how to actually write these electron configurations.
Orbitals
Each subshell can only hold a certain amount of electrons. The shape of orbitals in that subshell (s, p, d, or f) determine how many electrons can be held at max.
s-subshells can hold 2 electrons at max. When you see an s in electron configuration, the superscript cannot be greater than 2.
p-subshells can hold 6 electrons at max. When you see a p in electron configuration, the superscript cannot be greater than 6.
d-subshells can hold 10 electrons at max. When you see a d in electron configuration, the superscript cannot be greater than 10.
f-subshells can hold 14 electrons at max. When you see an f in electron configuration, the superscript cannot be greater than 14.
This is extremely important to remember, as you will see below.
Pattern
Electrons fill out the lowest energy positions first.
This makes life easy, because all elements follow the same order of electron configuration. That sounds confusing, but let’s illustrate through an example.
Sodium (11 electrons), has an electron configuration of
1s² 2s² 2p⁶ 3s¹
Magnesium (12 electrons), has an electron configuration of
1s² 2s² 2p⁶ 3s²
Notice how they are the exact same, except magnesium has one more electron the 3s subshell. This is because magnesium has one more electron than sodium.
Aluminum (13 electrons), has an electron configuration of
1s² 2s² 2p⁶ 3s² 3p¹
Again, the same as magnesium, with the addition of one electron in the 3p subshell.
Essentially, you start with the lowest energy subshell (1s), and “fill out” each subshell until you have the correct amount of electrons.
When filling out subshells, remember that s-subshells fill up after 2 electrons, p-subshells fill up after 6 electrons, d-subshells fill up after 10 electrons, and f-subshells fill up after 14 electrons.
You also need to know the order of the subshells. Remember, electrons fill out the lowest energy subshells first. Recall that s<p<d<f in terms of energy, and as you increase the principal energy level (n), you increase the energy.
For example 1s < 2s < 3s and 3s < 3p < 3d.
Also note ↴
-There is no 1p subshell
-The first d-subshell is 3d (so no 1d or 2d).
-You don’t see an f-subshell until 4f (so no 1f, 2f, or 3f).
Sometimes principal energy level (n) and orbital shape conflict. For example, 3d is higher energy than 4s. Even though 4s has a greater principal energy level, d > s, so 3d is higher energy. Therefore, 4s is filled before 3d because it is lower energy.
It is impossible for you to mentally determine when principal energy level beats orbital shape, so you will have to somewhat memorize the order. Below is a diagram showing the order that you should “fill” out subshells.
Essentially, just follow the arrows from the top arrow to the bottom arrow. “Fill out” the subshells in that order until you have the desired number of electrons.
Check out the example at the bottom if you are at all confused.
Noble Gas Configuration
We said before that electron configurations represent all electrons in an atom. Especially for larger atoms, it can be tedious to write the entire electron configuration. Noble gas configuration is a shortcut method.
Teachers may ask you to write the “complete electron configuration” of an atom, in which case you cannot use noble gas configuration.
First, find the nearest noble gas that is before your desired element in the periodic table. Then, remove the first x electrons from your element’s electron configuration (where x is the number of electrons that the noble gas has). Replace the omitted electrons with the noble gas in square brackets.
Ex. Calcium has a complete electron configuration of: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². The nearest noble gas before Ca is Argon. Argon has 18 electrons, so we remove the first 18 electrons from calcium’s configuration. 1s² 2s² 2p⁶ 3s² 3p⁶ are the first 18, so we remove and replace them with [Ar]. This gives us the noble gas configuration [Ar] 4s². See how much simpler that looks.
Notice how the removed electrons are the electron configuration of Argon. Therefore, adding the “[Ar]” represents those electrons.
Chemistutor 