No reference, but you may calculate it by yourself (although you will be probably forced to use some software, like my BATE, as these calculations can be challenging). The whole idea of pKa determination by titration is based on the
Henderson-Hasselbalch equation and fact that pH of weak acid titrated 50% equals its pH. But this approach is based on simplyfying assumption that concentrations of HA (acid) and A
- (conjuagted base from neutralized acid) are defined by the neutralization stoichiometry. In the case of stronger acids HA will tend to dissociate "on its own", in the case of weaker acids A
- will hydrolize "on its own" and the difference between pH at 50% and pKa will be larger. Perhaps 3-11 is slightly conservative, perhaps you can extend it to 2.5-11.5 - it all depends on the error you are ready to accept.
| pKa | pH at 50% |
| 0.00 | 1.50 |
| 1.00 | 1.67 |
| 2.00 | 2.18 |
| 3.00 | 3.02 |
| 4.00 | 4.00 |
| 5.00 | 5.00 |
| 6.00 | 6.00 |
| 7.00 | 7.00 |
| 8.00 | 8.00 |
| 9.00 | 9.00 |
| 10.00 | 10.00 |
| 11.00 | 10.97 |
| 12.00 | 11.82 |
| 13.00 | 12.32 |
| 14.00 | 12.48 |
You may try to use some correction tables - like pH of 2.18 at 50% means pKa of 2.00 - but still, the stronger the acid the larger the error, as the dependence becomes more and more flat (ie differences in acid strength give smaller differences in pH).
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Topic: Chemical Identification of an unknown alkali (Read 109009 times)