Is Al2 paramagnetic according to the MO theory?

To determine if Al2 is paramagnetic according to the molecular orbital (MO) theory, we need to understand the electronic configuration and bonding in Al2. Here's how to approach this question:

1. Determine the electron configuration of each Al atom: Aluminum (Al) has atomic number 13, which means each Al atom has 13 electrons. The electron configuration of an isolated Al atom is 1s2 2s2 2p6 3s2 3p1.

2. Write the molecular orbital (MO) diagram: When two Al atoms come together to form Al2, their atomic orbitals interact to form molecular orbitals. To determine the electron configuration of Al2, we can fill up the molecular orbitals according to the Aufbau principle and Hund's rule.

a. Construct the MO diagram by combining the atomic orbitals. The two Al atoms each contribute three valence electrons for a total of six electrons.

b. Start by filling the molecular orbitals with the lowest energy, following the Aufbau principle. The molecular orbitals will be labeled as σ, σ*, π, π* and so on, based on their symmetries.

c. Fill the molecular orbitals with the electrons. Place the six electrons according to Hund's rule (one electron per orbital before pairing up) until all the electrons are distributed.

3. Determine if there are any unpaired electrons: If there are unpaired electrons in the molecular orbitals, the molecule is paramagnetic. If all electrons are paired, the molecule is diamagnetic.

So, based on the MO diagram and electron distribution, if Al2 has any unpaired electrons, it is paramagnetic; otherwise, it is diamagnetic. You can now analyze the filled molecular orbitals in your Al2 MO diagram to determine if there are any unpaired electrons.