A Typical PH Problem
Calculate the pH and percentage protonation of a .20 M aqueous solution of pyridine, C5H5N. The Kb for C5H5N is 1.8 x 10−9.
First, write the proton transfer equilibrium:
The equilibrium table, with all concentrations in moles per liter, is
| C5H5N | C5H6N+ | OH- | |
|---|---|---|---|
| initial normality | .20 | 0 | 0 |
| change in normality | -x | +x | +x |
| equilibrium normality | .20 -x | x | x |
| Substitute the equilibrium molarities into the basicity constant | |
| We can assume that x is so small that it will be meaningless by the time we use significant figures. | |
| Solve for x. | |
| Check the assumption that x << .20 | ; so the approximation is valid |
| Find pOH from pOH = -log with =x | |
| From pH = pKw - pOH, | |
| From the equation for percentage protonated with = x and initial = .20, |
This means .0095% of the pyridine is in the protonated form of C5H5NH+.
Read more about this topic: Weak Base
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