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I am doing the following question:

Calculate the pH of a $4.00\ \mathrm{mol\ L^{-1}}$ solution of citric acid. $\mathrm{pK_{a1}} = 3.09~~~~~ \mathrm{pK_{a2}} = 4.75~~~~~ \mathrm{pK_{a3}} = 6.40$

Usually for polyprotic acids, we can assume that the second, third and other dissociations of $\ce{H+}$ ions are negligible when compared to the first dissociation. However I believe that this is only valid when the pKa values are separated by around 4. So for citric acid, this would be an invalid assumption as the pKa values are pretty close to each other.

Now, I can find the pH of the solution by doing the long way by first finding the amount of $\ce{H+}$ ions formed by the first dissociation, then using that to find the amount of $\ce{H+}$ ions formed by the second dissociation and so on.

However the room that they have given for this question is pretty small and has only has enough room for a couple of lines of working out and I definitely wouldn't be able to fit all my working out in that box.

So am I missing some short cut method to finding the pH in cases like this, or is the above assumption that I have outline at the start still correct for citric acid?

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  • $\begingroup$ Looks mighty like they want you to assume the said approximation, despite its being not that great for this particular case. $\endgroup$ – Ivan Neretin Dec 7 '15 at 12:29
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The assumption that the second (and higher) deprotonation steps are negligible for the calculation of the pH of a polyprotic acid is valid for citric acid.

You can compare this by calculating the exact solution and then stepwise neglect deprotonation steps.

(For the sake of simplicity I write $c(\ce{H+})=x$ and $k_{a,n}=k_n$.)

Exact: $$x=\frac{k_W}{x}+c_0(\ce{H3A})\frac{x^2 k_1+2x k_1 k_2+3 k_1 k_2 k_3}{x^3+x^2 k_1+x k_1 k_2+k_1 k_2 k_3}$$ Neglecting $k_3$: $$x=\frac{k_W}{x}+c_0(\ce{H3A})\frac{x k_1+2 k_1 k_2}{x^2+x k_1+k_1 k_2}$$ Neglecting $k_2$ and $k_3$: $$x=\frac{k_W}{x}+c_0(\ce{H3A})\frac{k_1}{x+k_1}$$

Solving those equations give the following pH values (with super unrealistic high precision):

  • Exact: 1.2469291050971560
  • w/o $k_3$: 1.2469291065371300
  • w/o $k_2$ and $k_3$: 1.2470654318802320

As a pH can only be measured precisely within something like $\pm \left(10^{-2}\ldots10^{-3}\right)~\mathrm{pH}$, the calculated pH can be rounded to be 1.247 in every singly case. No matter which deprotonation step was neglected.

So for your case it is negligible. For more general cases I have to look it up and will add it to my answer within the next days.

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  • $\begingroup$ Thank you for the great answer! So is there any rule of thumb about knowing if you can assume the second dissociation is negligible? Before I thought the pKa values had to be different by 4 but that obviously doesn't work for this example. $\endgroup$ – Nanoputian Dec 7 '15 at 19:10
  • $\begingroup$ As always it depends and there is no clear border ... the pictures will follow. But another point ... is this just a "normal" homework excercise from school or is it more "real" from the lab? 4 M is a very high concentration and unless this is not just a "normal" homework excercise - which I assumed - there is no simple concentration=activity assumption. Afaik, this can actually only be made if the concentration is below ~0.001 M. $\endgroup$ – pH13 - Yet another Philipp Dec 8 '15 at 9:17
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Since the solution will be acid in water, the pH will be below all the pKas.

Given that the pH is below all the pKas, an upper limit is that the second pKa will contribute 10^(3.09-4.75) = 0.022 protons for each proton that the first pKa contributes.

The effect of these 2% more protons on pH would be about log (1.022) = 0.0009 pH units.

In real life, the problems calculating pH from the information given in the question, would be:

  1. 4M citric acid means 768 grams of citric acid in 1 liter of solution. Will that much really dissolve?

  2. If that much did dissolve, and concentration of H2O is down to about 22M instead of the usual 55M, how do you account for that in the equilibrium calculations?

  3. The 3 pKa values are calculated based upon extrapolating ionic strength to zero, what would these values be at 4 molar citric acid?

The pH of 50% citric acid is reported to be 0.6, and if you could make 4 M citric acid, it would be still lower pH. It doesn't matter if you use 1, 2, or 3 pKas, you wouldn't calculate the right value (to say 0.1 pH units) for a solution that concentrated.

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