[rayznack]

Hello anyone reading.

Today at work, I measured the pH of phosphoric acid at 0.84.

The acid was 85% phosphoric acid and 15% water.

Just now, I did a quick and dirty calculation of what pH I should have measured.

Using a Ka value of 7.25x10^(-3) and molarity of 14.8 - based on the links below - I calculated a pH of 0.48.

I didn't take into account the pH of the water or the dissociation of the secondary and tertiary protons (or adjust for temperature - the temperature was around 72 F, however).

https://www.erowid.org/archive/rhodium/chemistry/equipment/molarity.html

https://en.wikipedia.org/wiki/Phosphoric_acid#Reactions

I was wondering if anyone agreed with my results or perhaps knows where I can find the pH of 85% phosphoric acid.

Thanks.

[Arkcon]

The measurement of pH really only works for dilute solutions, so measurement of mostly phosphoric acid in a bit of water is going to be off.

That's really the least of the problems with your particular application.  For example, how did you standardize your pH meter?  How did you flank your expected pH?  What slope did the pH meter report?

[kriggy]

Also, how did you calculate the pH?

[rayznack]

The measurement of pH really only works for dilute solutions, so measurement of mostly phosphoric acid in a bit of water is going to be off.

That's really the least of the problems with your particular application.  For example, how did you standardize your pH meter?  How did you flank your expected pH?  What slope did the pH meter report?

This was a basic pH meter.  I have no way to answer your questions.

[rayznack]

Also, how did you calculate the pH?

Basic calculation

x = (14.8*(7.25x10^-3))^0.5

x = 0.33 = [H30+]

-log(0.33) = 0.48

[Borek]

In such a concentrated solutions you are way outside the applicability of [itex][H^+] = \sqrt{C_A K_a}[/itex].