Drawing the Lewis Structure for PO43-

Viewing Notes:

  • PO43- has a charge of -3 (that makes it a negative ion or anion). That means that it has an three extra electrons that needs to be taken into account.
  • Since Phosphorous is in Period Three on the periodic table it can have more than eight valence electrons.
  • The Lewis structure for PO43- has 32 valence electrons available to work with.
  • You should check the formal charges to make sure that the total charge of the atom is -3.
  • Since it's an ion you need to put brackets around the Lewis structure and then put a -3 outside.

Transcript: This is Dr. B. Let's do the Lewis structure for PO4 3-. Phosphorus has 5 valence electrons. Oxygen has 6, we've got 4 Oxygens. This negative 3 up here means we have three additional electrons. Five plus 24 plus 3 gives you 32. So those are our valence electrons. Put Phosphorus at the center and the Oxygens around it, all 4 of them. We'll draw bonds between the Oxygen and the Phosphorus. We're forming chemical bonds right here, two electrons each. So we've used 8 and then let's fill the octets for the Oxygen. So we've used 2, 4, 6, 8, 10, 12, and 32.

So we've used all 32 of the valence electrons. Each of the atoms has an octet and it really feels like we're done with this one. But Phosphorus is in period 3. It can hold more than 8 valence electrons. So let's check our formal charges to see if this is the best structure. For the Phosphorus, on the periodic table, it has 5 valence electrons. All of these right here are involved in bonds, so we have zero nonbonding; but we have 2, 4, 6, 8 bonding. We'll divide that by 2. That gives us a +1 formal charge for Phosphorus. For Oxygen, on the periodic table, 6 valence electrons. Each Oxygen has 6 of these nonbonding valence electrons and then 2 bonding, which we divide by 2. Six minus 6 minus 1 gives us -1 as the formal charge on the Oxygen. And they're all the same.

So we know we have to end up with a 3- as the total formal charge. And that's not working right now. And I can see that, on the Phosphorus, there's this +1 right here. What I can do, if I see a +1 like that, if I move two electrons from one of these atoms and form a double bond, that should resolve the +1. So let's see if that works. So what I've done is, I've taken these two electrons that were out here and I've formed a double bond. Now, when I check my formal charges, the Phosphorus ends up being zero, which is much better. And this Oxygen right here, 6 minus the 4 nonbonding, and then the 4 bonding divided by 2, that's zero. And the other Oxygens remain at a -1, so we end up with a total charge of minus one, minus two, minus 3, which makes a whole bunch of sense.

So this is the best structure for the phosphate ion—Lewis structure for PO4-. We do need to put brackets around it and a 3- out here so that everyone knows that it is, indeed, the phosphate ion. That's the Lewis structure for PO4 3-. A bit tougher: watch your formal charges, you'll get the correct structure. This is Dr. B., and thanks for watching.