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Author: Subject: Temperature for sodium acetate and bisulfate reaction?
ZeulMos
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[*] posted on 4-8-2016 at 21:51
Temperature for sodium acetate and bisulfate reaction?


Hi Sciencemadness, this is my first post here, so please forgive me if I do something wrong.
I wanted to know the required temperature for the reaction between sodium acetate and sodium bisulfate in order to yield acetic acid?
I have anhydrous sodium bisulfate and sodium acetate trihydrate.
Peace!
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ahill
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[*] posted on 4-8-2016 at 22:54


GDay ZeulMos,

The reaction gets going at the melting point of sodium hydrogen sulphate .. 58.5C (according to wikipedia) - but assuming you are distilling the acetic acid off - you'd need to go to 119C to boil that.

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woelen
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[*] posted on 4-8-2016 at 23:48


Sodium hydrogen sulfate has a much higher melting point than 58.5 C. This temperature of 58.5 C is the temperature of decomposition of the mono-hydrate of sodium hydrogen sulfate: NaHSO4.H2O

On heating, NaHSO4.H2O liquefies. It is no real melting, it is breaking down of the hydrate and you get free water, in which the NaHSO4 dissolves.
Sodium acetate tri-hydrate has a similar behavior. It also breaks apart and forms a solution of NaCH3COO in water.

These two solutions react with each other and indeed you can drive off acetic acid, together with water. This process, however, will not be really smooth. You get a cake of solid Na2SO4, mixed with unreacted NaHSO4 and NaCH3COO. It will be hard to nicely and evenly heat this solid cake. You will have a lot of sputtering and cracking and if you are not really careful with heating you will crack the flask in which this solid cake is formed.

So, theoretically, the reaction is simple. A practical execution of the reaction and the isolation of the acid may be quite cumbersome.




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Tdep
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[*] posted on 5-8-2016 at 06:01


I've tried this and it works alright. Whenever you use bisulfate in the place of sulfuric acid, you get lower yields and cleanup of the reaction flask is more annoying. But sodium bisulfate here is massively cheaper than sulfuric acid, to the point where it's almost cheaper to use bisulfate and break a few flasks and replace them, rather than use sulfuric acid. That's probably not the case everywhere, but don't hate on bisulfate ok?

I use it to make anhydrous nitric and acetic acids, and dry HCl gas. It might not be the smoothest, but it's great. I wouldn't say cumbersome.
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ZeulMos
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[*] posted on 5-8-2016 at 11:41


So if I use this method will I get at least 70% acetic acid?
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woelen
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[*] posted on 7-8-2016 at 04:58


Yes, I think so, taking Tdep's experience into account. But be prepared to have a hard solid cake of Na2SO4 in your glassware and the issue of non-uniform heating, leading to the risk of cracks. But apparently this method of getting acids is doable and it is good to know this, especially if obtaining H2SO4 becomes more and more difficult, due to all kinds of regulations.



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AJKOER
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[*] posted on 7-8-2016 at 08:57


Mix in water solid Sodium hydrogen sulfate and Sodium acetate. If this reaction is endothermic, apply heat as needed.

Then, freeze and separate out the acetic acid and water from the Na2SO4.10H2O precipitate.

If this is a homework problem to calculate the amount of heat, please ask.
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ahill
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[*] posted on 7-8-2016 at 15:29


..in retrospect, I dont know that my sodium hydrogen sulphate is anhydrous - it is swimming pool conditioner from the hardware. I calculated expected yields based on the quantity of sodium acetate I used, and used stacks of excess sodium hydrogen sulphate. Yeilds were very reasonable, and I dont recall cleanup being an issue. While I did have some doubts as to _how_ dry the resulting acetic acid actually was - it displayed the passivating effect some concentrated acids show. (like it got more reactive as I diluted it). I've still got maybe 25ml left lying around - if someone has a good suggestion for how I could measure the concentration, I'll give it a go. (my scales can only do .1g)

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[Edited on 8-8-2016 by ahill]
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zwt
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[*] posted on 7-8-2016 at 15:48


Quote: Originally posted by ahill  
how do you measure the diff between 95 and 99% acetic acid ?

If you have a digital scale and some volumetric glassware (and you should), density is easy to determine and a reliable way of finding concentration within a percent or so. Here is a link to the specific gravity of acetic acid solutions at various concentrations at 15*C.

Edit:
Thanks for editing your post after I replied. That's not disruptive at all. Good thing I quoted the original. You should really consider getting a centigram or milligram balance and some good volumetric glassware.

[Edited on 8-8-2016 by zwt]
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clearly_not_atara
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[*] posted on 8-8-2016 at 11:55


Can you use e.g. anhydrous ethanol as a cosolvent to make life easier?

IIRC sodium acetate is soluble in dry ethanol. Dropping liquid NaHSO4*H2O into this solution at 60 C might be a much easier version of this rxn. Na2SO4 is highly insoluble in ethanol after all. Maybe methanol would be better. Since Na2SO4 is an excellent drying agent this should work even with sodium acetate trihydrate.

I'm not sure how easy it is to distill acetic acid from ethanol but I don't think it should be that hard since their bps are quite far apart (118 C vs 78 C), as long as there isn't any extra acid present (adding a little NaHCO3 should not generate any sulfate precipitate)

[Edited on 8-8-2016 by clearly_not_atara]
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ahill
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[*] posted on 1-10-2016 at 17:50


..I finally got around to weighing my acid - (thanks for the table zwt) - I only had 13ml left - and it weighed 13.8g (remember .1g resolution on the scales) - indicating a surprising and disappointing 50%.

Anyone recommend a cheap high res scale on ebay ?
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