Sciencemadness Discussion Board

sodium bisulfate monohydrate synthesis

cyberzed - 12-2-2007 at 10:12

Hi i was wondering if anyone could get me on the right track with a sodium bisulfate synthesis.

According to Wikipedia there are two syntheses which are actually very easy:

* NaOH + H2SO4 → NaHSO4 + H2O

and

* NaCl + H2SO4 → NaHSO4 + HCl

Both are very exothermic reactions, best to try on a small scale under controlled conditions !!
I understand that those reactions would create NaHSO4, but suppose that u want to create the monohydrate version what should you modify to this synthesis?

[Edited on 12-2-07 by cyberzed]

12AX7 - 12-2-2007 at 10:16

Or Na2SO4 + H2SO4 = 2NaHSO4, or...

The first is a big fat waste of NaOH, and the second requires removal of HCl (efforvescent in conc. H2SO4).

To make whatever hydrate, recrystallize from H2O.....

I noticed KHSO4 readily crystallizes in irregular octahedrons, which are apparently rhombic somethings. Anhydrous, I think. I don't know how sodium responds, but it's probably more soluble.

Tim

cyberzed - 12-2-2007 at 10:22

Why is the first reaction a big fat waste of NaOH, it is a chemical which is very otc, easy to come by and cheap, so i wouldn't mind a little loss on NaOH .....

I don't like the reaction with NaCl and H2SO4 because of the production of HCL in the reaction, for this you better have a decent venting system.

If i understand correctly if i would create NaHSO4 by the reaction of NaOH with H2SO4 which leaves a H2O molecule wouldnt the end product be monohydrate already?

Since those reactions are very exothermic, shouldnt it be possible to work with diluted solutions of either component and distill off the water afterwards?



[Edited on 12-2-07 by cyberzed]

woelen - 12-2-2007 at 10:35

Don't do the reaction with concentrated materials. It will lead to excessive production of heat, and non-homegeneous mixes, because all of it solidifies.

Dissolve H2SO4 in water, make an appr. 20% solution (careful: add acid to water, slowly, while stirring).
Dissolve NaOH in water, also make an appr. 20% solution, but concentrations are not critical at all.

Take as precisely as you can 10.00 ml of the H2SO4 and add an indicator (e.g. phenolphtaleine).
Now titrate with the solution of NaOH, until the indicator just changes color (e.g. for phenolphtaleine the pink color just persists on stirring).
Now you know how much solution of NaOH is needed for full neutralization of the H2SO4.

Now take a precisely measured amount of solution of H2SO4 and compute how much solution of NaOH is needed to neutralize that. Take half that amount and carefully and slowly add to the H2SO4 (this produces a lot of heat). If your volumetric measurements are somewhat uncertain, then it is best to take a slight excess of NaOH. Too little results in excess H2SO4 making the stuff very hygroscopic, it does not dry well. A little too much of NaOH results in Na2SO4, but that has no adverse effects on drying.

Now let the half-neutralized liquid evaporate on a dry and dust-free place at room temperature in an evaporation dish. Crystals of NaHSO4.H2O will separate from the liquid.

cyberzed - 12-2-2007 at 11:07

Woelen, thank you for this info.
This might seem like a dumb question, but in your post you mention that only half of the NaOH solution required to neutralize the acid will be used for neutralization?
Why are we only taking half the solution?
I was thinking that according to the equation 1 mole of NaOH and 1 mole of H2SO4, this diluted in water should react together to produce NaHSO4 + some more water.


Also another question:
Suppose that you have anhydrous sodium bisulfate, and you want to convert this to the monohydrate, would it be sufficient to dissolve this in water and the evaporate over a water bath, and afterwards dry this in the dessicator?




[Edited on 12-2-07 by cyberzed]

[Edited on 12-2-07 by cyberzed]

Rosco Bodine - 12-2-2007 at 14:26

Sodium Bisulfate Monohydrate is commonly available
cheap as a swimming pool chemical called " pH Minus " .

Aurum - 13-2-2007 at 05:54

Quote:
Originally posted by Rosco Bodine
Sodium Bisulfate Monohydrate is commonly available
cheap as a swimming pool chemical called " pH Minus " .


Never use it though as it severely reduces your total alkalinity by precipitating calcium.

woelen - 13-2-2007 at 10:26

Quote:
Originally posted by cyberzed
Woelen, thank you for this info.
This might seem like a dumb question, but in your post you mention that only half of the NaOH solution required to neutralize the acid will be used for neutralization?
Why are we only taking half the solution?
I was thinking that according to the equation 1 mole of NaOH and 1 mole of H2SO4, this diluted in water should react together to produce NaHSO4 + some more water.

The titration gives the result for full neutralization to Na2SO4, but you want it neutralized only halfway, so you need half the amount of NaOH.
If you were titrating 1M H2SO4 with 1 M NaOH, then you would find the indicator to switch color when 2 volumes of NaOH are added to 1 volume of NaOH.
Quote:
Also another question:
Suppose that you have anhydrous sodium bisulfate, and you want to convert this to the monohydrate, would it be sufficient to dissolve this in water and the evaporate over a water bath, and afterwards dry this in the dessicator?

That would be sufficient. Just let it evaporate at room temperature.

cyberzed - 15-2-2007 at 12:59

Ok, thanks for the info, but evaporation is such a slow process, isn't there the possibility to separate this by distillation, since water boils at 100° and Na2SO4 decomposes somewhere above 300° while never boiling?

cyberzed - 24-2-2007 at 15:12

is there a conclusive test to see if sodium bisulfate is the monohydrate form?

12AX7 - 24-2-2007 at 18:12

Weigh it, dehydrate (without overheating, producing sulfuric fumes) and weigh again.

Tim

unionised - 25-2-2007 at 05:39

Or titrate it but since it crystalises from water as the monohydrate, that's what you will get.