Why does the MO theory cause lone pairs to bond?

[amk0713]

Hi there. I'm new to the forums, a freshman in college, and mainly asking this for the sake of understanding.

Why does the MO theory cause lone pairs to bond? For example, in the Lewis sturcture of O₂ there are 2 pairs of unbonded lone pair electrons which makes sense since both oxygen atoms completed their octet. In addition, the 2p_y and 2p_z orbitals are used for overlapping in the valence bond theory, staying consistent with Lewis. But then in the MO theory, the (sigma)2s, (sigma)2s*, and (pi)2p_x are used to bond the electrons that were lone pairs. Why is this?

I really want to understand this, but I'm having a lot of trouble putting together how the valence bond theory and MO theory fit together.

Thank you for any input.

 

[Navin A S]

The MO theory allows for more accurate predictions regarding the structure of a molecule than the Lewis structure or valence bond theories. In the case of O2, the two lone pairs that are present in the Lewis structure remain, but the MO theory suggests that these lone pairs can actually interact strongly with each other and form a bond. This is because when the two oxygen atoms are close enough together, the (sigma)2s, (sigma)2s*, and (pi)2px orbitals can overlap, creating a strong bonding interaction between the two atoms. This type of overlap is not possible in the valence bond theory, which is why the MO theory is often used to better explain the structure of molecules.

 

[quicknote]

The MO (molecular orbital) theory is a more advanced theory that explains the bonding in molecules by taking into account the wave nature of electrons. In this theory, the electrons are not confined to specific atomic orbitals, but rather they are described by molecular orbitals that extend over the entire molecule. These molecular orbitals are formed by the combination of atomic orbitals from the participating atoms.

In the case of O2, the two oxygen atoms each have two unpaired electrons in their 2p orbitals. In the valence bond theory, these electrons are paired to form a sigma bond and a pi bond between the two oxygen atoms. However, in the MO theory, these electrons are described by molecular orbitals that extend over both oxygen atoms, resulting in a stronger bond between the two atoms.

The MO theory also takes into account the presence of lone pair electrons. These electrons, which do not participate in bonding, occupy molecular orbitals that are closer to the nucleus of the atom. This makes them more stable and less likely to participate in bonding. However, in some cases, these lone pair electrons can overlap with the molecular orbitals involved in bonding, resulting in a stronger bond.

In the case of O2, the lone pair electrons in the 2s orbital can overlap with the antibonding molecular orbital (sigma 2s*) and contribute to the bonding between the two oxygen atoms. This helps explain why O2 is a stable molecule even though it has two lone pairs of electrons.

In summary, the MO theory provides a more comprehensive explanation of bonding in molecules by considering the wave nature of electrons and their interactions with each other. This theory helps us understand why lone pairs can sometimes participate in bonding, and how they contribute to the overall stability of a molecule.