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Author: Subject: Making Sodium Bisulfate
Runningbear
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[*] posted on 13-10-2009 at 03:44
Making Sodium Bisulfate


Hi guys,

A quick question. If I have 33% sulfuric acid and 100% sodium hydroxide, can I produce sodium bisulfate without having to concentrate the sulfuric acid? And if so, what quantities of each are required?

thanks in advance,

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[*] posted on 13-10-2009 at 04:38


Yes you can produce sodium bisulfate without concentrating the acid. Just slowly add sodium hydroxide to the sufuric acid until a 1:1 molar ratio has been acheived (this requires fairly basic calculations, you should be able to figure it out).



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[*] posted on 13-10-2009 at 06:43


It requires basic calculations AND a table of densities of sulphuric acid of different concentrations. You need to know the density of 33% sulphuric acid if you use volume measurements for the acid.

If you measure weight for the acid and no volume, then no such table is needed, but usually such liquid chemicals like 33% H2SO4 are not weighed, but volume is measured (e.g. 100 ml of liquid is taken, using a good volumetric flask).




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entropy51
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[*] posted on 13-10-2009 at 07:22


Runningbear, making bisulfate is a nice experiment, but you can buy it rather cheaply at a place that stocks pool supplies. It is often sold as "pH Down". You probably knew this, but I didn't want you to "waste" H2SO4 and NaOH unless you just especially wanted to roll your own.
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[*] posted on 13-10-2009 at 08:35


I've made it several times (also K bisulphate) from 50 % H2SO4 ('Draino') and technical caustic soda (NaOH) but I never seem to obtain a dry product. Rather a tough, moist, crystalline mass that fumes like mad when I try and get it to anhydrous. Any tips?


I'm guessing 'pH down' may also be available from 'Skunk 'r us' sites, tt's where I get my nitric acid from.
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[*] posted on 13-10-2009 at 09:28


Quote: Originally posted by blogfast25  
Rather a tough, moist, crystalline mass that fumes like mad when I try and get it to anhydrous. Any tips?
Could it be an excess of H2SO4? I've recovered it from nitric distillations, recrys from H2O and dry in a vacuum dessicator, which works OK. Probably not as pure as from NaOH, though.
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[*] posted on 13-10-2009 at 14:36


Thanks for the feedback guys.

Entrophy, I know you can buy it quite cheaply, but I thought I might try and make it just for the hell of it. A little bit more knowledge can't hurt.

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[*] posted on 13-10-2009 at 14:51


I thought that might be your plan, and it is a nice experiment. I guess you know too that if your H2SO4 and NaOH are not very pure, you may not hit it just right. NaOH is prone to absorb CO2 and contain Na2CO3, which has a different molecular weight than NaOH (obviously). According to the Merck Index, the pH of a 0.1 M solution of NaHSO4 is 1.4, so if you have a way to determine pH this might be a better way to exactly get to NaHSO4 than by a straight calculation. But you may be willing to accept a slight error, I probably would. Let us know how it turns out.
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[*] posted on 14-10-2009 at 06:11


Quote: Originally posted by entropy51  
Quote: Originally posted by blogfast25  
Rather a tough, moist, crystalline mass that fumes like mad when I try and get it to anhydrous. Any tips?
Could it be an excess of H2SO4? I've recovered it from nitric distillations, recrys from H2O and dry in a vacuum dessicator, which works OK. Probably not as pure as from NaOH, though.


I always thought that that was the cause but having calculated the quantities precisely, I was at the same time sceptical of that hypothesis. I guess next time I'll aim for a small excess of NaOH (perhaps 1 % ?) That way the product will contain a small amount of Na2SO4, which can't hurt...
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[*] posted on 14-10-2009 at 15:16


Quote: Originally posted by entropy51  
I thought that might be your plan, and it is a nice experiment. I guess you know too that if your H2SO4 and NaOH are not very pure, you may not hit it just right. NaOH is prone to absorb CO2 and contain Na2CO3, which has a different molecular weight than NaOH (obviously). According to the Merck Index, the pH of a 0.1 M solution of NaHSO4 is 1.4, so if you have a way to determine pH this might be a better way to exactly get to NaHSO4 than by a straight calculation. But you may be willing to accept a slight error, I probably would. Let us know how it turns out.


Entrophy, thanks for the advice on PH. That's a handy tip.
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[*] posted on 17-10-2009 at 12:31


Hmmm... careful with that pH plan. Remember the basics of acid-base titrations? The S-shape of the pH curve? Near the point of equivalency the pH of the solution changes very sharply: difficult to stop titrating at precisely pH = 1.4 (that value is concentration dependent, BTW).
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[*] posted on 17-10-2009 at 14:10


Quote: Originally posted by blogfast25  
Near the point of equivalency the pH of the solution changes very sharply: difficult to stop titrating at precisely pH = 1.4 (that value is concentration dependent, BTW).


That's the reason you use calculated quantities but add dropwise when adding the last portion.

That's the reason I said the pH is 1.4 at 0.1 M. You can remove 0.1 mL aliquots, dilute to 0.1 M, stick a pH electrode in it.

Even better is to titrate a test run on a small scale to determine the quantities that bring you to the equivalence point, then weigh out those quantitites for the large scale prep.

[Edited on 17-10-2009 by entropy51]
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