Q. Orbital diagram for Mg.

An orbital diagram for magnesium (Mg) shows how its electrons fill orbitals according to the aufbau principle, the Pauli exclusion principle, and Hund’s rule.

First, identify the atomic number of magnesium.

\( \text{Mg} \) has atomic number \(12\), so a neutral Mg atom has \(12\) electrons.

Write the electron configuration by filling subshells in order of increasing energy.

The electrons fill in this order (lowest energy first): \(1s\), \(2s\), \(2p\), \(3s\).

Count and place electrons into each subshell.

• \(1s\): holds up to 2 electrons, so Mg gets \(2\).
• \(2s\): holds up to 2 electrons, so Mg gets \(2\).
• \(2p\): holds up to 6 electrons, so Mg gets \(6\).
• \(3s\): holds up to 2 electrons, so Mg gets the remaining \(2\).

This gives the electron configuration:

\[
\text{Mg: } 1s^2\, 2s^2\, 2p^6\, 3s^2
\]

Now construct the orbital diagram by drawing each subshell as boxes and placing electrons as arrows (one arrow in each box can represent one electron; direction is used to show opposite spins).

Orbital diagram for Mg (showing individual orbitals):

\[
\begin{array}{c}
1s: \quad \uparrow\downarrow \\
\\
2s: \quad \uparrow\downarrow \\
\\
2p: \quad \uparrow\downarrow \quad \uparrow\downarrow \quad \uparrow\downarrow \\
\\
3s: \quad \uparrow\downarrow
\end{array}
\]

Alternatively, written in a single compact orbital-box style (one set of three boxes for \(2p\)):

\[
\begin{array}{c}
1s:\ [\uparrow\downarrow] \\
2s:\ [\uparrow\downarrow] \\
2p:\ [\uparrow\downarrow]\ [\uparrow\downarrow]\ [\uparrow\downarrow] \\
3s:\ [\uparrow\downarrow]
\end{array}
\]

Key check: Count the electrons to ensure it matches \(12\).

\(1s^2\) gives \(2\), \(2s^2\) gives \(2\), \(2p^6\) gives \(6\), and \(3s^2\) gives \(2\).
Total \(2+2+6+2 = 12\).