I am trying to rationalise a result that I have recently measured in the lab. I'm measuring the pKb of different compounds in anhydrous acetic acid, using the procedure outlined in this paper (https://pubs.acs.org/doi/abs/10.1021/ic50084a055). I'm able to re-produce their calibration plot for my pH probe set-up, so no issues there.
I notice that in their table however, that the pKb of diphenylamine is greater than that of triphenylamine (9.72 and 9.20, respectively). This is confusing to me, because I thought that when considering the pKb scale, a low pKb means that it is a stronger base. So the data recorded states triphenylamine is a stronger base than triphenylamine. But in triphenylamine, the nitrogen lone pair is less accessible for protonation, and therefore it should be a weaker base than diphenylamine (i.e: triphenylamine should have a higher pKb than diphenylamine).
I also see in anhydrous acetonitrile, you get the expected trend as you increase the number of phenyl rings attached to a nitrogen (shown below with pKa values). As more phenyl rings are introduced, the pKa decreases, which means that the pKb increases in that solvent. The rational in the literature agrees with what I describe above.
Why does glacial acetic acid act in the opposite way to other solvents when it comes to the pKb (and therefore the pKa)?
