Tsjerk
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percentage of deprotonated HCl in concentrated solution
Dear all,
Some years ago, during a first year course of my study biology, I was asked to calculate the pH of 37% HCl. Obviously the calculation turned into a
negative pH. The answer was said to be wrong, as it would not be possible to calculate the pH of 37% as not all HCl de-protonates, and cannot not be
measured either, because there is no pH sensor capable of measuring a pH that low.
My question to you, does anyone know to what extend HCl in a 37% solution is de-protonated? If that number is known, a real pH could be calculated.
The reason for this question is the fact I'm going to assist in the said course, and don't want to be wrong on the matter 
Thanks in advance.
[Edited on 8-6-2014 by Tsjerk]
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Nicodem
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Quote: Originally posted by Tsjerk  | | My question to you, does anyone know to what extend HCl in a 37% solution is de-protonated? If that number is known, a real pH could be calculated.
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The pKa = -7 of HCl was measured in water. This value is useless for non-water such as 37% HCl(aq) (at such deviation from the model, it cannot even
be approximated any more). I have never seen the pKa of HCl in 37% HCl, though I would not be surprised it has been published.
Also, as a rough orientation, keep in mind that water is a base with the pKa(H2O) -1.7 while the pKa(H2O) of HCl is -7 (5
magnitudes apart from water). So, even at high HCl concentration and the deviations from these pKa values due to the change of solvent, the amount of
non-dissociated HCl is still likely to be low.
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blogfast25
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According to this source (which unfortunately provides no references and isn’t very authorative):
http://www.newton.dep.anl.gov/askasci/chem99/chem99230.htm
… the pH of 37 w% HCl is -1.1. That would mean the deprotonation is near 100 %, even at that high concentration.
That the pH must almost certainly be negative is clear: pH = 0 is only [H3O+] = 1 M. With a molar concentration of HCl 37 w% of about 12 M, less than
10 % would be deprotonated if the pH was not negative. That just doesn't make any sense.
[Edited on 8-6-2014 by blogfast25]
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aga
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For the course work, you have three choices if faced with this question from a student :-
1. say you do not know.
2. with Great Authority say the first number that pops into your head, and stick to it.
3. press the fire alarm button
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Tsjerk
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Haha Aga,
Thanks for the suggestions, but I prefer just to be right
I remember being quite upset when they told me my answer was wrong, I don't want to tell a student they are wrong if they are not.
So I will just assume that more than 10% is de-protonated, meaning the pH should be somewhere between 0 and -1,2-ish.
They also claimed a strong acids not to be strong, and sulfuric acid to be a mono-hydrogen acid in the same syllabus... not to worse for a biologist,
but with a bit of a chemistry background I can't afford to overlook these things, if it only were for my own conscious.
The acid was phosphoric acid, I think the miss assumption came from the fact we only use phosphate in biology as a buffer, so as the mono-hydrogen or
the di-hydrogen ion, in those cases it is weakly acidic or basic. The same for hydrogensulfate.
[Edited on 8-6-2014 by Tsjerk]
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kmno4
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| Quote: Originally posted by Tsjerk |
So I will just assume that more than 10% is de-protonated, meaning the pH should be somewhere between 0 and -1,2-ish.
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It was pointed out by Nicodem, that such a large concentration is far beyond theory of solute-solvent interaction for diluted
solutions.
Diluted means less than 0,1 M but even at this retively low conc., deviations from theory are easily observed. In another words: molal fraction of
solute should be low (<1/100). This is region where (almost) all components (ionic or not) of solution are fully coordinated, structure of solvent
is not significantly changed, ion association is small.
Let us take ~40% HCl. It is ~12 M solution, it corrensponds to molal fraction ~1/3: one molecule of HCl per 3 molecules of water.
No conditions for diluted solutios are fulfiled.
Even familiar K=[OH][H3O]/[H2O] loses its meaning, because of above mentioned reasons.
That is why you cannot use "pH" concept for such solutions, because it "fits to nothing".
One can calculate value = -log12, it gives ~1,1 and say "I calculated pH of 40% HCl", what is the problem ?????"
"Ph" concept was created for purely practical reasons, and in theory it is defined by unmeaurable activity of H+.
That is why, practical standard of pH is defined by electrochemical methods.
"Ph" is also a measure of acidity - the lower pH, the greater acidity. But - as concentration of HCl is rising (to saturation), solvation of protons
becomes more and more limited.
It causes very important thing: acidity of such solution rises, because less solvated proton is more active("acidic") !!!!!
So acidity - ability to protonate bases - rises faster than concentration, it is valid for other concentrated strong asids too.
For weak electrolytes, above certain concentration, acidity remains (almost) constant.
That is why various "acidity funtions" were introduced. For decreasing concentrations, these functions pass to "pH" function and give the same
(almost) values.
| Quote: | | The answer was said to be wrong, as it would not be possible to calculate the pH of 37% as not all HCl de-protonates, and cannot not be measured
either, because there is no pH sensor capable of measuring a pH that low. |
Of cours it is not true, there are electrochemical methods allowing measurements in such solutions.
In reality, however, no-one measures pH but only some voltage 
Theory simply allows calculation of "pH" from voltage.
On the picture - values of Ho acidity function for some acids in water. See that ~8 M sol. of HCl is about 100-fold more acidic than ~2 M solution.
See also that below 0,1 M concentrations , Ho passes to pH. (table taken from Chem. Rev., 1957, 57 (1), pp 1–45)
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