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How to convert P2O5 concentration to H3PO4 concentration?
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Marry Harry
How to convert P2O5 concentration to H3PO4 concentration?
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.
Second part: I will let you solve the rest of the problem: Start from here
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.
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.
Second part: I will let you solve the rest of the problem: Start from here
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.
Thanks for the answer I believe I solved it based on first poster response as follows (in 1 kg). 700g P2O5 x 1 mol/ 142 g = 4.9 mol P2O5. 300 g H2O is 16.7 mol H2O. 4.9 mol P2O5 + 17.78 mol H2O --> 9.8 mol H3PO4 Mass of 9.8 mol H3PO4 * 98 g mol = 960.4 g H3PO4More
Seems right, when dealing with percentage concentrations, it is better to use 100g as the unit. So using your numbers, it is 96% w/w phosphoric acid.More
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$\ce{P2O5}$ 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).
$\ce{P2O5}$ 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).
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.
Second part: I will let you solve the rest of the problem: Start from here
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.
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.
Second part: I will let you solve the rest of the problem: Start from here
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.
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$\ce{P2O5}$ 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).
$\ce{P2O5}$ 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).
More
VOTE