GreenD
National Hazard
Posts: 623
Registered: 30-3-2011
Member Is Offline
Mood: Not really high anymore
|
|
pH in IPA
Hello. I've got a problem that I'm not sure how to look at.
I need to create a solution, 20 mL of a certain acid (.2% by weight) in IPA (and H2O if necessary).
The solution must be basic via ammonium hydroxide to a high pH to fully deprotonate all of the acid moieties. Does simply calculating the
concentration of NH4OH in IPA give you the pH (i.e. OH- concentration).
How would I go about finding the "pH"?
If you're ambitious;
20mL of a .2% solution of an acid fully deprotonated with (preferably) 25% ammonium hydroxide.
I was fairly poor at calculating pH until one day I had a eureka I understand it, then it disappeared.
ʃ Ψ*Ψ
Keepin' it real.
Check out my new collaborated site: MNMLimpact.com
|
|
|
unionised
International Hazard
Posts: 5126
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
Measuring pH isn't easy in non-aqueous solutions.
You would probably do better to calculate the number of moles of acid and add a bit more than that number of moles of ammonia.
|
|
|
GreenD
National Hazard
Posts: 623
Registered: 30-3-2011
Member Is Offline
Mood: Not really high anymore
|
|
Quote: Originally posted by unionised  | Measuring pH isn't easy in non-aqueous solutions.
You would probably do better to calculate the number of moles of acid and add a bit more than that number of moles of ammonia. |
Ok, yeah because it isn't actually pH anymore, right?
But still, if you had an excess of base, you would have a fully deprotonated acid in IPA, right? There aren't some weird occurences where the
equivalents of base needs to be really high or something?
Since carboxylic acids can be easily deprotonated by ammonium hydroxide a ~1.2:1 base:acid should fully deprotonate, right?
ʃ Ψ*Ψ
Keepin' it real.
Check out my new collaborated site: MNMLimpact.com
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
