A 50.0 mL sample of 0.10 M pyridine, C5H5N, is titrated with 0.2 M HBr. Calculate the pH to one decimal place when the following volumes of titrant have been added.
0.00 mL
a) 0ml
b)17ml
c)25ml
d)40ml
This is a titration problem. How much do you know how to do? I'll help you but I need t know where to start?
To calculate the pH at different volumes of titrant added, we need to determine the reaction that occurs and then use the stoichiometry to find the concentration of the resulting species. Here's how you can approach each case:
a) 0 mL of titrant added:
At this point, no HBr has been added, so the reaction has not begun. The solution consists only of the pyridine and its conjugate base, the pyridinium ion (C5H5NH+). Pyridine acts as a weak base, so we can use the Henderson-Hasselbalch equation to calculate its pH:
pH = pKa + log ([A-]/[HA]), where pKa is the acid dissociation constant of pyridine.
The pKa of pyridine is usually given as 5.2.
b) 17 mL of titrant added:
At this point, some HBr has been added, and a reaction occurs between pyridine and HBr. The balanced chemical equation for this reaction is:
C5H5N + HBr -> C5H5NH+ + Br-
Since HBr is a strong acid, it completely ionizes in water. So, we can assume that the concentration of Br- is equal to the volume of HBr added in moles.
c) 25 mL of titrant added:
Similar to case b, a reaction has occurred between pyridine and HBr. Use the balanced chemical equation to determine the change in concentration of pyridine, pyridinium ion, and Br-.
d) 40 mL of titrant added:
At this point, a significant amount of HBr has been added. We again use the balanced chemical equation to determine the change in concentrations.
Once you have determined the concentrations of the relevant species in each case, you can calculate the pH using the appropriate equations for weak acids and bases.