Describe the orbital diagram of an atom with 16 electrons. Explain how this orbital diagram demonstrates Hund's rule.

The orbital notations shows the sequence of filling electrons into the orbitals of sublevels.

This filling is based on some certain principles.

For an atom with 16 electrons, the orbital diagram is shown below:

1s²2s²2p⁶3s²3p⁴

The maximum number of electrons in each sublevel of the orbitals are:

2 electrons for s-sublevel with one orbital

6 electrons for p-sublevel with three orbital

10 electrons for d-sublevel with five orbital

14 electrons for f-sublevel with seven orbital

According to the Aufbau's principle, sublevels with lower energy are filled before those with higher energy.

1s 2s 2p 3s 3p 4s 3d etc

Pauli's exclusion principle shows that no two electrons can have the same set of values for the four quantum numbers. Simply, no two electrons can spin in the same direction.

Hund's rule states that electrons go into degenerate orbitals of sub-levles (s, p, d and f) singly before pairing commences.

This rule shows that in each energy level, as the electron goes into the degenerate orbitals, they fill it one by one before they begin to pair up. As we know, each degenerate orbital can only accomodate 2 electrons.

From the orbital diagram 1s²2s²2p⁶3s²3p⁴, the 3p sublevel has 3 orbitals. In each of the orbitals, two electrons would occupy them to give a maximum capacity of 6. But the sublevel has just 4 electrons. Based on Hund's rule, an electron will go into each of the 3 orbitals first. The remaining electron will now pair with the first degenerate orbital. This makes a total of 4 electrons.