Orbital diagrams are another way of visually showing electron configurations. They use boxes to represent orbitals, and arrows to represent electrons.

Here are the different subshells as orbital diagrams ↴

Note that these are empty orbitals, because there are no arrows. Remember that per energy level, there is 1 s-orbital, so there is 1 s-box. There are 3 p-orbitals in a p-subshell, so there are 3 p-boxes. The same reasoning explains why there are 5 d-boxes and 7 f-boxes.

Here are what the filled orbitals look like ↴

In each box, there is an upward arrow and downward arrow. Recall the mₛ quantum number, which represents the spin of the electron. Upward facing arrows have positive spin, and downward facing arrows have negative spin. In a filled orbital, there is one electron with positive spin and one with negative spin.

Here’s an example of orbital configuration for oxygen ↴

Pay attention to the incomplete 2p subshell. Below shows a common mistake ↴

When doing orbital configuration, fill out empty boxes within a subshell before double filling. This is illustrated by the diagram below, which shows the correct order of filling a d-subshell ↴

You will learn more about this rule in the “Principles” section.

Paramagnetic vs. Diamagnetic

Paramagnetic: An atom with unpaired electrons. Paramagnetic substances tend to act as magnets (attracted to magnetic fields).

Diamagnetic: An atom with no unpaired electrons. In other words, all electrons are paired.

Unpaired electrons are “alone” in their orbital diagram box, as shown below ↴

If an atom has any unpaired electrons, it is paramagnetic. The more unpaired electrons a substance has, the more paramagnetic it is.

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