Net charge for phosphate and phosphite ion

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The discussion centers on the net charge of phosphate (PO4^3-) and phosphite (PO3^3-), both of which are -3. The oxidation state of oxygen is typically -2, while hydrogen carries a +1 charge. In phosphite, the presence of hydrogen alters its structure, leading to a net charge of -2 when calculated. The oxidation state of phosphorus in both ions is deduced to be +1, which accounts for the overall charge balance. The conversation highlights the unusual characteristics of phosphorus chemistry, including the stability of compounds like hypophosphite (H3PO2) and their behavior under heating. It also touches on the coordination preferences of phosphorus, noting that it often forms stable P-H or P-O bonds to achieve a coordination number of 4. The discussion emphasizes the complexity and unique aspects of phosphorus compounds in chemistry.
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Why is the net charge the same (-3) for phosphate (PO4) and (PO3)?
Why is the net charge the same (-3) for phosphate and phosphite ion?
PO4-3 and PO3 is also -3.
This is nor a homework question. I am 49 and love chemistry and helping my son with his studies.
Thank you in advance.
 
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Different oxidation states. By convention (and perhaps necessity), oxygen has a -2 oxidation state with notable exceptions being O2 (where it is 0) and H2O2 where it is -1.

Can you deduce the oxidation state of P in these two compounds and justify the identical net charge? Also remember that phosphite has an H on it! (+1 oxidation state typically)
 
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Surprising to find Phosphite ion does have a Hydrogen in it.
HPO2-2

Basic facts are, O has charge -2, and H has a charge of +1.
You were interested in what charge goes with the P.
The ion's charge is -2.

1(+1)+1(x)+2(-2)=-2

1+x-4=-2

x=-2-1+4

x=+1
 
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symbolipoint said:
Surprising to find Phosphite ion does have a Hydrogen in it.
HPO2-2
It is surprising and somewhat unusual but it is one of the relatively notable facts about chemistry of phosphorus. Likewise, phosphorus has hypophosphite H3PO2 which is importantly a monobasic acid.
But the underlying logic is that while the low coordination numbers of P (and S) prevail in lower halides, they are unstable for O and sometimes F compounds. P prefers to have either P-H bonds or else "saturate" P-O bonds by reaching the coordination number 4. Both phosphites and hypophosphites will dismute on heating - to PH3 and H3PO4.
This requires exchanging hydrogens between P atoms. At room temperature the compositions H3PO3 and H3PO2 are stable... but the molecules rearrange to add the P=O bond:
P(OH)3>O=P-H(OH)2
goes to practically completion and is not reversed by bases.
Realized a better comparison. You do not have "carbonous acid" C(OH)2 either. You do have acid H2CO2... but it is actually monobasic HCOOH, the formic acid.
 
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