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I have learnt that for a weak acid — strong base titration, $\mathrm{p}K_\mathrm{a} = \mathrm{pH}$ at the half equivalence point.

However, the same conclusion is not drawn when discussing strong acid — strong base titrations. Why does the above only hold for weak acid — strong base and not strong acid — strong base?

Similarly, why is $\mathrm{pOH} = \mathrm{p}K_\mathrm{b}$ at half-equivalence point only true for strong acid weak base and not for strong acid strong base?

andselisk
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vgupt
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    What do you know about buffers? – Mithoron Dec 19 '20 at 20:40
  • I know that a buffer solution is one which resists changes to pH when small amounts of acid/alkali are added and I know the two ways in which a buffer solution can be formed. I also know that a buffer is most effective when the concentration of the ions is very high so that additional ions added (H+ or OH-) won't have a large relative impact...is that useful? – vgupt Dec 19 '20 at 20:47
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    How about pKa = pH vs buffers then and why strong acid won't work? – Mithoron Dec 19 '20 at 21:23
  • https://chemistry.stackexchange.com/questions/133853/why-does-the-ph-before-the-equivalence-point-of-a-titration-depend-on-the-initia – Mithoron Dec 19 '20 at 21:47

1 Answers1

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The reason is that strong acids have $\mathrm{p}K_\mathrm{a}$ values that are poorly known. These $\mathrm{p}K_\mathrm{a}$s cannot be determined directly, because the exact concentration of the ions is difficult to know with precision: the electrodes do not react with the concentration of $\ce{H^+}$ or $\ce{H3O^+}$ ions. They react with the activity of $\ce{H+}$ or $\ce{H3O+}$, which may be quite different from the concentration in concentrated solutions.

So, the $\mathrm{p}K_\mathrm{a}$ values of strong acids can only be determined indirectly, for example by extrapolation of values obtained in organic solvents, and mixtures of organic solvents plus water. This extrapolation is never precise enough.

andselisk
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Maurice
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  • Thanks Maruice - I still don't get. Surely the Ka of strong acids is almost entirely to the right (complete dissociation) - why don't we know it? Also, is the topic of buffers important at all, as mentioned by Mithoron above in the comments? Thank you! – vgupt Dec 19 '20 at 20:55
  • @ vgupt. I don't know what Mithoron wanted to say with his short question ! – Maurice Dec 19 '20 at 21:00
  • So, can the pKa value of a strong acid not be determined from it titration curve? – vgupt Dec 19 '20 at 21:02
  • are you there? Would appreciate any help, if possible! – vgupt Dec 19 '20 at 21:11
  • No. The pH cannot be determined in a concentrated solution. Or, if you prefer : the measured pH values has no relation with the concentration of the ion $\ce{H^+}$. – Maurice Dec 19 '20 at 21:13
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    Just to explain my purpose. If you measure the pH of a 0.01 M solution of HCl, you obtain 2. If now you dissolve enough NaCl to saturate this solution, you measure pH 1.1, as if the solution were more concentrated. This shows that the pH is the log of the concentration calculated NOT by unit of volume, but by unit of volume of "free water". Free water is the water that is not adsorbed around the dissolved ions. In a concentrated solution, there is practical no free water any more. The concentration calculated by volume of free water seems to increase. That is the same for pure strong acids. – Maurice Dec 19 '20 at 21:26