Chemical Forums
Chemistry Forums for Students => High School Chemistry Forum => Topic started by: ydude03 on April 04, 2008, 09:27:10 AM
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Hi all,
I've spent about an hour trying to solve this question I have as homework...
So here is the question:
When placing pyridinium chloride (C5H5NH+Cl-) in water, the pyridinium ion's dissociation degree is equal to 0.0237.
A) What is the pH of the solution?
B) Find the concentration of the pyridinium ion.
(I have no other info, exept that I already know that Kb for C5H5N is 1.5E-9 - im not sure if it's useful, I just took it from a table of Kb's)
I have an ideao, though I'm not sure if that will work:
maybe I could use
pH = Ka + log([products]/[reactants])
which is the same to say : pH = Ka + log(0.0237), findind Ka from the internet (as it is not given...)
Am I right or ...
If anyone could help me find this out, it would be really appreciated,
Thanks all
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You're close with your equation but that isn't the correct henderson-hasselbalch equation. Look it up and use that. You're on a great track. Keep going with it. a question:
how do you convert Kb to Ka?
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Hi
We know that:
Kb*Ka=Kwater=E-14
so we can easily find Ka using Kb and Kwater.
I had this idea, but wasn't sure if it would work due to the fact that the Kb is C5H5N's, not the one for C5H5NCl or so...
Does this cause a problem? or is it really 'convert' Kb to Ka and use:
pH = Ka + Log[diccociation degree]
Thanks again dude.
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Ah, just realized what's going on here. forgot to follow rule 1 of solving chemistry problems:
write out a balanced equation for the reaction
Do that and then write out the equilibrium equation. Post what you have and if you've realized what I have...you'll know if you have.
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No concentration, no answer IMHO.
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When you set up the equilibrium equation I bet you get to cancel out the base/conjugate acid making the need for concentrations moot. Haven't tried it but I have a feeling...