The electron configuration is the distribution of the electrons in an atom. It is how the
electrons are arranged that helps us understand chemical bonding and chemical reactions.
So, if you know how electrons are arranged, especially the valence
electrons, you will better understand chemical formulas, equations and reactions. But, first a few rules need
to be stated concerning the placement of electrons into energy levels.
Electron Arrangement Rules
Now that the quantum numbers have been introduced, the possible electron arrangements for the elements in an
atom can be described.
The number of electrons equals the number of protons in a stable atom.
The number of electrons each energy level can hold is 2n2.
The number of sub shells in an energy level is equal to n.
The s sub shell has only one possible position, the p sub shell
has three, the d sub shell has five and the f sub shell has seven.
Each possible position is an orbital.
Only two electrons can occupy each orbital.
Hund's Rule: When a p, d, or f sublevel is being filled, one electron will occupy
each orbital before pairing.
The maximum number of electrons is two in any s sub shell, six in any p sub shell, ten in any d subshell,
and fourteen in any f sub shell.
Pauli Exclusion Principle: No two electrons in an atom have the same four quantum
Aufbau Principle: An electron occupies the lowest energy level available, filling in
orbitals of higher energy levels until all electrons are distributed.
Using the above rules, you can easily diagram the distribution of the electrons in an atom. Just determine how
electrons the atom has
energy levels to be used
sub energy levels to be used
orbitals there will be in each sub energy level
Let's look at carbon as an example.
According to the first rule, carbon has six electrons and according to the second rule would use two energy
levels. Using rule three, you can determine that the first energy level will have only one sub energy level and
only one orbital (s). The second energy level will have two sub energy levels and four orbitals ( one s and three
p). Now, using Hund's rule and Aufbau's principle distribute carbon's six electrons:
two electrons go into the first energy level's only orbital (s orbital)
then two go into the second energy level's first orbital (s orbital)
then one in the second energy level's first p orbital
finally one in the second energy level's second p orbital
carbon = 1s2 2s2 2p1 2p
Electron and Orbital Notations
Using our knowledge of quantum numbers and the distribution rules, let's see how electron configurations and
orbital notations of the elements are represented.
If you’re thinking this is too easy to be true, you’re right. There are a few complications as the atoms get
larger. As the energy levels get farther from the nucleus, the distance between the energy levels decreases.
As a matter of fact, it is believed that the energy levels actually overlap. Therefore, some energy levels start
filling orbitals before the previous energy level is finished filling its sublevels.
The first time this is encountered is with potassium, in which the 4s starts to fill
before the 3d.
There's More ...
The second complication has to do with a variation of Hund’s Rule that takes into account the minimizing of the
It states, the most stable arrangement of electrons is the arrangement with the maximum number of unpaired
electrons. So, when the transition metals’ orbitals are filling with electrons, at
d4 and d9, an electron from the
s jumps up into the d5 or
Writing out electron configurations and orbital notations can become awkward as the atoms increase in the number
of electrons. So, scientists have agreed on a type of shorthand to help make writing electron configurations and
orbital notations less cumbersome.
The shorthand involves using the abbreviation of the last noble gas (placed in brackets) to indicate that all
the orbitals to that point are full. Then the notation is continued as usual.