# How to convert P2O5 concentration to H3PO4 concentration?

Why are some elements such as phosphorus reported as $$\ce{P2O5}$$ concentrations when the materials have no $$\ce{P2O5}$$ in them. I would like an explanation of how the conversion is done between real life and this idealized state.

For example. the Wikipedia page on phosphoric acid states that a $$\ce{P2O5}$$ concentration of 70% is equivalent to nearly 100% $$\ce{H3PO4}$$. Can someone explain this conversion?

• What's to explain? P2O5 reacts with water. Do you know what the product is? Can you balance the reaction? – Ivan Neretin Sep 7 at 16:25
• okay I didnt realize the concentration was in reference to water because its often discussed as purity etc. makes sense thanks. – dlight Sep 7 at 16:39
• @IvanNeretin Im reading a paper which states energy used to process phosphoric acid is "6500 kwh to make a ton of phosphoric acid (P2O5)". I want to express this in kWh/ton H3PO4. So I get 4710 kWh/ton H3PO4. Agreed? The paper is: A Technical Review of the Improved Hard Process – dlight Sep 7 at 16:54
• It is a common practice to express content of the component A in the equivalent content of the component B. There is an older unit of the water "hardness" dGH ( degree of General/German Hardness ) where 1 dGH means equivalent of 10 mg CaO / 1 L. But there is no CaO in water. Similarly, in fertilizers, there is AFAIK historical custom to use content of oxides, even if there are none. – Poutnik Sep 7 at 17:32
• P2O5 + 3H2O --> 2H3PO4. Using molecular weights P2O5/2H3PO4 = 142/198 = 0.71 = 71% – user55119 Sep 7 at 22:33

Your question is a valid analytical chemistry question which has deep roots in history. No, the reason is not that $$\ce{P2O5}$$ reacts with water to from phosphoric acid and that is why $$\ce{P2O5}$$ values are quoted as suggested in the comments below the question. This is a very old historical tradition. In older analytical chemistry, chemists had only two main tools to analyse something namely, titration and gravimetry. Oxygen was also used to determine atomic weights, and people were interested in the ratios of element:oxygen ratio. One person even got a Nobel Prize as well. Generally, when the precipitates were heated, what was left was oxides. In really old books (at least 100 year old) you will notice that many analyses were quoted as oxides. For example, this is true in fertilizer industry to quote potassium as $$\ce{K2O}$$ and so on. Now if you strongly heat a precipitate phosphate, one will not get a $$\ce{P2O5}$$ but the convention to report elements as oxides continues.
1 mol $$\ce{P2O5}$$ = 2 mol $$\ce{H3PO4}$$
and do wt percentage conversions to moles. A 70% $$\ce{P2O5}$$ solution in water means 70 g $$\ce{P2O5}$$ is present in 100 grams of solution. Do mol conversions. You will also need the density of the solution.
$$\ce{P2O5}$$ content is useful as it is the phosphorous content that a scientist often cares about when using phosphoric acids. This might seem silly given that as you pointed out anything from $$0-70\% \ \ce{P2O5}$$ could just be expressed as a percentage of $$\ce{H3PO4}$$ but there are times when a higher $$\ce{P2O5}$$ concentration is desired. For example polyphosphoric acids are useful reagents that have more phosphorous content that phosphoric acid. It would be ambiguous to state the concentration as $$115\%\ \ce{H3PO4}$$ (what does 115% mean) so instead procedures are reported in terms of $$\ce{P2O5}$$ concentration with the balance assumed to be water (for those in the know).