Structural Formula of Polyatomic ions
 

Hello, I'm very new to chemistry so bear with me if I seem to be too much of a newbie..hehe.

I seem to be having trouble understanding the structural formulas of polyatomic ions. Can somebody please explain to me how to draw the structural formula of a phosphate ion, PO4^3-. <-----(hope that makes sense; please correct me if it doesn't). I have attempted this and my solution is the phosphorus atom bonded to four oxygen atoms each with a single covalent bond. The answer I have retrieved off of various online databases however, suggests the phosphorus atom bonded to four oxygen atoms, three with a single covalent bond, the fourth O atom with a double covalent bond. Which one is right? Please help!! I get the impression that my error has something to do with the fact that one or more of the oxygen atoms are bonded as a coordinate covalent bond in which case I still have no clue of what's going on.

Thanks in advance,
Kevin

Its been awhile since I did this, so please forgive me of any inconsistances. A very nice way to grasp covalant bonding is to make a lewis dot diagrams of the molecule ( :shock: !). Heh, enjoy:

1- You have 4 oxygen atoms, each with a valence of 6eˉ so total 4x6=24eˉ
2- You have one phosphorus with 5eˉ
3- Since it is an ion, with charge of 3eˉ you have grand total of 24+5+3=32eˉ

4- Now write the letter P on a peice of paper and an O on top, right, bottom, and left sides of it. Start adding dots first (rememb, total of 32) to where the oxygens bond with phosphorus; between their symbols. Next start dotting the phosphorus around their outsides (2 each side) one by one until you have used up 32 dots.

You should find that all 32 dots are used up and make all the atoms' valences 'full', and between each oxygen and the phosphorus is two dots denoting a single bond. Now you can convert to a structural formula consisting of single bonds.

Repeat the above with something like CO (carbon monoxide), you'll find C is 4 dots short from a full valence, therefore requires a triple bond.

hmmm....
 

That's the same solution I got. So then I was right?
The other solution I found on the internet can be seen at the following link:
http://pubchem.ncbi.nlm.nih.gov/summary ... i?cid=1061

Is the solution posted in the link correct?

Yeah... after seeing another source that seems to be what the scientific community agrees on. I too think it has something to do with cordinate covalent bonds. Check: http://www.chem.purdue.edu/gchelp/gloss/ccbond.html

Try this, draw 3 O's with 7 dots and one O with 6 dots. The phosphorus in the middle has 5 dots you can draw to make it a total of 32. For the three O's with 7 dots, phosphorus only lines up 1 dot for each. 2 dots are left which it shares with the O with 6 dots. I guess that counts as a double bond.

The problem is, marking an oxygen atom with 7 valence electrons is not allowed. The maximum number of valences an oxygen atom can have is 6. The bonding capacity is 2.

The standard method of figuring out structures of simple molecules that is taught to beginning chemistry students is not sufficient for this ion. Phosphorus has access to d orbital hybridization, which is one way to explain why it can form more than four bonds and break the octet rule.

Are you asking because you need help on homework/tests, or out of simple curiousity?

I may have solved it.
 

d orbital hybridization, what is this?

In answer to your query, yes this is for homework. The textbook states nothing about d orbital hybridization. I did finally work out a solution that correlates with the solutions given on the internet. I'll describe the method:

1. Draw 4 o's and one p in the centre.
2. Draw a pair of electrons in between all four Os and the P.
3. Draw the remaining electrons around each O leaving 6 electrons around each Oxygen atom, and a bond between each oxygen-phosphorus pair.
This is where I got confused:
4. Since eight electrons were placed around the P, and P only brought 5 electrons to begin with...the P has a charge of 3-...since each O now has 7 electrons associated with its charge and it originally brought 6, the charge on each O is 1-.
8. To simplify this, we need to carry one pair of electrons from one of the O's and place it with the already existent bond. This results in a double bond.
9. Now when we figure out the charges....we get....5 electrons on the P that are associated with the charge (i.e. not including the other valence in the pair)...since P originally brought 5, this atom is neutral. each single-bonded O is still 1- since we haven't changed anything with them...the double bonded O now has 6 electrons associated with its charge (i.e. not including the other valence in the pair) which is exactly what it had originally....thus it too is neutral.
10. So we have one phosphorus atom in the centre, surrounded by three oxygen anions with a charge of 1- single bonded, and 1 oxygen atom double bonded with no charge.

Is this correct? It may also be incorrect due to typos, grammatical errors etc.

What is d orbital hybridization? Sounds interesting.

btw, I have no idea why skipped four steps.
:lol:

Ok...I looked up d orbital hybridization on the net...it appears to be way out of my league for the moment. Why would my gr.11 textbook have this problem if it's beyond the scope of the text itself?

Isn't that what i just said? (heh). Don't worry, you have it right. As far as I know, d-orbitals and the like are only covered from grade 12.