Biochem
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Potentiometric Titrations
I've done the experiment and now I'm on the calculations but i dunno if they are right...can someone please verify if I'm correct or in
error.
Experiment basically used unknown HCL-H3PO4 mixture titrated with known NaOH and took readings via pH meter with 2 trial runs. Plotted points graphed
pH vs vol, first deriv, and second deriv...used seocond deriv to find endpoints where it hit zero figured the second endpoint to determine H3PO4
concentration(1:1) and then used that and subtracted it from first endpoint where HCL and H3PO4 were still mixed so here are the calculations:
First Run
Volume at First Endpoint ~ 16.85 ml
Volume at Second Endpoint ~ 22.35 ml
Volume at 2nd Endpoint – Volume at 1st Endpoint= 5.5 ml, 5.5 ml /1000 = 0.0055 L
Concentration of NaOH = 0.1019 M
Moles of NaOH = Moles H3PO4 = 0.0055 L * 0.1019 M = 5.605e-4 M
16.85 ml /1000 = 0.1685 L
0.01685 L * 0.1019 M = Moles of NaOH reacted for both acids = 0.001717 M
Moles of HCl = Moles of NaOH reacted for both acids - Moles H3PO4
Moles of HCl = 0.001717 – 5.605e-4 = 1.167e-3 M
Second Run
Volume at First Endpoint ~ 16.80 ml
Volume at Second Endpoint ~ 22.24 ml
Volume at 2nd Endpoint – Volume at 1st Endpoint= 5.44 ml, 5.44 ml /1000 = 0.00544 L
Concentration of NaOH = 0.1019 M
Moles of NaOH = Moles H3PO4 = 0.00544 L * 0.1019 M = 5.543e-4 M
16.80 ml /1000 = 0.0168 L
0.0168 L * 0.1019 M = Moles of NaOH reacted for both acids = 0.001712 M
Moles of HCl = Moles of NaOH reacted for both acids - Moles H3PO4
Moles of HCl = 0.001712 – 5.543e-4 = 1.658e-3 M
Also we didn't calibrate the pH meters...but the instructor said we didn't have to any idea why? And it's strange how HCl and H3PO4
have 4 total protons but we only see two end points why not 4?
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unionised
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I'm not checking the arithmetic, but the idea looks right.
Imagine that you doubled the readings you got from your pH meter, then plotted the graph. It would be twice as high, but the sudden changes (the peaks
in the derivative) would still happen at the same volumes so the pH meter calibration doesn't matter.
You don't see all 4 protons because the first proton transfer from H3PO4 and the one from HCl happen at virtually the same pH. Both are strong
acids in water.
The third proton from H3PO4 is a very weak acid. You can take it off, but only at rather high pH. At that pH the solution picks up CO2 from the air
and makes the titration difficult also, the pH electrodes are less robust at high pH. Anyway, there's not much point- the two results you have
will give you the concentrations of HCl and H3PO4.
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sparkgap
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"...you don't see all 4 protons because the first proton transfer from H3PO4 and the one from HCl happen at virtually the same pH. Both are
strong acids in water..."
OK, I'm confused.  It has always been said that phosphoric acid is a weak
acid. What gives? 
Virtually the same... anyone know offhand the pKa's of both?
sparky (~.~)
"What's UTFSE? I keep hearing about it, but I can't be arsed to search for the answer..."
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Pyridinium
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| Quote: | Originally posted by Biochem
And it's strange how HCl and H3PO4 have 4 total protons but we only see two end points why not 4? |
I believe in mixed acid titrations, where 1 acid is strong and 1 is weaker and multiprotic, the titration does only the HCl and the first proton of
the weak multiprotic acid, provided the pKa's differ by at least 4 units.
I wish I could see your pH curve, but I'm guessing there was a small and not too well-defined pH break for HCl, followed by a better-defined
(steeper & taller) pH break for the first proton of H3PO4.
What was the value for your 1st derivative at the first equivalence point? If it was the equiv. pt. of HCl the slop probably didn't go vertical.
There's no sense in titrating all 3 protons of H3PO4, since you can get the number of moles just by titrating the first one.
I think HCl has a pKa around -2. First proton of H3PO4 is around 2 something, second is around 7 something.
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unionised
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OK, strictly, phosphoric isn't a strong acid but it's fairly strong.
It's pKa is about 2, so it would act as a buffer near pH2. On the other hand presumably, the solutions are fairly dilute. If they are about 0.1M
(ie comparable with the NaOH concn) then, even if they were fully dissociated, they would only have a pH of about 1, fairly near the pKa for H3PO4.
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