# As atomic orbitals form they require how many electrons to fill the orbital?

##### 1 Answer
Jul 18, 2016

I think you're confusing atomic orbitals with molecular orbitals.

When molecular orbitals form, two valence electrons are required to be situated in between two atoms for a molecular orbital, forming a chemical bond.

So for this hypothetical ${\text{C}}_{2}$ molecule, the molecular orbital that forms between them has two valence electrons.

Here's the thing - atomic orbitals don't actually require any electrons in them.

Case in point: $\text{Be}$ doesn't have any electrons in its $2 p$ orbitals, and that's totally normal. A simple way of talking about this is that there can be $0$, $1$, or $2$ electrons in a single atomic orbital, but there is no requirement for a specific number.

When atomic orbitals overlap, they will form molecular orbitals (bonding and antibonding, such that the number of molecular orbitals is equal to the number of atomic orbitals). That is...

The electron population within molecular orbitals constitute the chemical bonds made between atoms!

• Any sigma ($\sigma$) bond requires the head-on overlap of atomic orbitals.
• Any pi ($\pi$) bond requires the sidelong overlap of atomic orbitals.

(The dots in the diagram represent atoms, so those are not electrons.)

Actually, each $\sigma$ and $\pi$ overlap (and $\delta$, which is more complicated) requires two valence electrons to be in between the two atoms when the chemical bond forms.

• For $\sigma$ bonding molecular orbitals, the electrons are right smack dab in between the atoms.
• For $\pi$ bonding molecular orbitals, the electrons are in between the atoms, but in the lobe above the atoms.

(How the electrons are shared or transferred is another story, but what matters is that there are two.)