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.