To draw the Lewis structure for the phosphate ion (PO3³⁻), we start by determining the total number of valence electrons. Phosphorus (P) has 5 valence electrons, and each oxygen (O) has 6 valence electrons. Since there are three oxygen atoms, we have:
5 (from P) + 3 x 6 (from O) + 3 (for the overall charge) = 5 + 18 + 3 = 26 valence electrons.
In the Lewis structure, the phosphorus atom is the central atom surrounded by three oxygen atoms. We can form double bonds with one of the oxygen atoms to satisfy the octet rule for the oxygen. The structure will look like this:
O−
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O = P - O−
Now let’s count the bonds and non-bonding pairs around the central phosphorus atom:
- There are three bonds (two single bonds with two oxygen atoms, and one double bond with one oxygen).
- There are no lone pairs on the phosphorus atom.
Therefore, around phosphorus, there are:
- Bonds: 3
- Non-bonding pairs: 0
The shape of the PO3³⁻ molecule is trigonal pyramidal due to the arrangement of the three oxygen atoms around the phosphorus atom. The lone pairs present on the oxygen atoms also influence the overall geometry but do not affect the shape around the phosphorus itself.