Electron configuration refers to the arrangement
of the electrons in an atom into orbitals.
To determine the electron configuration of
an atom with multiple electrons, we follow
something called the Aufbau Principle.
Aufbau is a German word that means “to build
up,” and that’s what you do in this process
- you build up the electron configuration
from the bottom up in this diagram of orbitals.
Electrons seek the lowest energy state, starting
with the 1s orbital.
We use arrows to represent electrons, and
the direction they are pointing shows the
direction of their spin.
According to the Pauli exclusion principle,
each orbital holds a maximum of 2 electrons,
with opposite spin.
So there’s a maximum of 2 electrons in the
1s orbital.
If the atom has more than 2 electrons, they
start filling the subsequent, higher energy
orbitals.
Let’s see some examples.
Hydrogen has 1 electron, so it goes into the
1s orbital.
We write the electron configuration as 1s1.
The first number shows the energy level, and
the exponent on s shows the number of electrons
in the s orbital.
Helium has 2 electrons, so those 2 electrons
go into the 1s orbital, with opposite spins.
The electron configuration is written as 1s2.
Lithium has 3 electrons, so we know it will
start to fill a higher orbital, since the
1s orbital can only hold 2 electrons.
We write this as 1s22s1.
Again, remember that the large numbers represent
the energy levels of the orbitals, and the
superscripts represent the number of electrons
in that orbital.
Beryllium has 4 electrons.
We fill the diagram from the bottom up:
Next comes Boron, with 5 electrons.
We’re now starting to fill in the 2p subshell.
The p subshell has 3 orbitals (px, py, and
pz), which can each hold two electrons, for
a total of 6 electrons.
We put the first electron in px.
The electron configuration of Boron is written
as 1s22s22p1.
Carbon has 6 electrons.
Where does that second electron go in the
p orbitals?
According to Hund’s Rule, every orbital
in a subshell gets one electron before any
orbitals get two electrons; and each of the
single electrons have parallel spin.
So we put the next electron in 2py, with an
up arrow to show it has the same spin as the
electron in the 2px orbital.
After the second energy level, you really
have to pay attention to the energy level
diagram.
For instance, notice that the 4s orbital has
a lower energy than the 3d orbital, so you
should fill the 4s orbital first.
But these two orbitals are so close in energy,
that for a few elements, electrons go into
the 3d orbital first.
For example, in chromium, the electron configuration
ends in 3d54s1 instead of 3d44s2.
Similarly, copper’s electron configuration
ends with 3d104s1 rather than 3d94s1.
The rationale for this is that half-filled
or completely filled subshells are particularly
stable arrangements of electrons.
[We will discuss some tips and tricks to make
writing electron configurations go a little
faster, including referring to the periodic
table, in another video.
