How to separate NaHCO3 from NH4Cl in water and measure quantitiesbismuthate - 10-3-2014 at 03:59
You could add NaOH, evaporate it, then dissolve the NaCl in glycerin or propylene glycol. I'm not sure how you could get the NH4Cl except dissolving
the ammonia from step 1 in HCl.
EDIT changed messed up formula
[Edited on 10-3-2014 by bismuthate]shc25 - 10-3-2014 at 04:04
I need to separate and measure quantities of each
must be simple way to ppt out the sodi bic
[Edited on 10-3-2014 by shc25]TheChemiKid - 10-3-2014 at 07:18
You could add NaOH, evaporate it, then dissolve the NaCl in glycerin or propylene glycol. I'm not sure how you could get the NH3Cl except dissolving
the ammonia from step 1 in HCl.
Ammonium Chloride is NH4Cl, not NH3Clblogfast25 - 10-3-2014 at 09:48
How to separate NaHCO3 from NH4Cl in water and measure quantities
Separating them AND measuring in a non-destructive way will be near impossible.
Measuring should be possible with some crafty acid-base titrations, possibly two different ones.
A first, straightforward titration with HCl will give you the amount of NaHCO3.
The second: add a known excess of NaOH and back titration with HCl should give you the total of NH4Cl + NaHCO3. shc25 - 10-3-2014 at 12:05
how do you calculate the % composition of each separately
1. having added HCl
2. adding NaOHunionised - 10-3-2014 at 13:24
What problem are you actually trying to solve? What are you trying to achieve?shc25 - 10-3-2014 at 13:29
I'm attempting to separate NaHCO3 + NH4Cl + H2O and calculate the percentage composition of each component.
The problem with adding HCl is that it produces NaCl which is still souble in water - no?
Also, I was wondering if there is any process that would involve physical rather than chemical processes.hyfalcon - 10-3-2014 at 13:50
Look at your solubility tables then I suggest trying fractional crystallization.bismuthate - 10-3-2014 at 13:51
Well NH4Cl is slightly soluble in ethanol if your using small amounts.hyfalcon - 10-3-2014 at 14:11
He already states the two are in water. Ammonium Chloride is quit soluble in water where sodium bicarbonate isn't.blogfast25 - 10-3-2014 at 14:15
You may not be able to separate them 100% from each other. This is part of solvay process, and when saturated ammonium bicarbonate solution is
saturated with NaCl at 35-40C and cooled down to 5C, NH4Cl will salt out. Since solubility of NH4Cl is a lot higher than sodium carbonate(formed in
methathesis with NH4CO2 and NaCl), the precipitate could be sodium carbonate as well. By this way it is possible to obtain quite pure reagents anyway.shc25 - 11-3-2014 at 03:32
Is there a reason for heating the solution to 40C before cooling it to 5C in a water bath?
Will a precipitate form if you just cool the solution from room temp. to 5C?AJKOER - 11-3-2014 at 14:38
And upon adding ammonia to a cold concentrated carbonated brine solution, the NaHCO3 is observed to precipate out of the basic solution. This leaves
NH4Cl and some NaCl and aqueous ammonia.
Heat to remove water, ammonia and then further heat the dry salt to sublime out the NH4Cl.
[Edited on 11-3-2014 by AJKOER]packetforger - 11-3-2014 at 14:54
Quick "before I go to bed" thought:
Precipiate out the crystals and use sublimation to extract the NH4Cl from the sodium bicarb powder. It should be seperable that way at least, though
there might be some reaction or something that I am simply not thinking of at this hour...bismuthate - 11-3-2014 at 15:01
NH4Cl doesn't sublinate it just decomposes and NaHCO3 will decompose too so heating is out of the question.shc25 - 11-3-2014 at 15:02
I actually have NaHCO3 + NH4Cl in water (the 2 in solution) at room temp
Do you need to warm the solution before cooling to say 0-5 degrees C to precipitate out the NaHCO3
As a test just adding Baking Soda (presumable just NaHCO3) to water and cooling to 4 deg C did not ppt anything?
I need to calculate the percentage composition of each component
I need to only use a physical process not chemicalshc25 - 11-3-2014 at 15:09
I actually have NaHCO3 + NH4Cl in water (the 2 in solution) at room temp
Do you need to warm the solution before cooling to say 0-5 degrees C to precipitate out the NaHCO3
I understand heating above 50 deg decomposes it but much of the data on Solvay talks abou the hot solution after NaCl was added to make the NaHCO3 +
NH4Cl as if that was important before cooling to get the ppt
As a test just adding Baking Soda (presumable just NaHCO3) to water and cooling to 4 deg C did not ppt anything?
I need to calculate the percentage composition of each component
I need to only use a physical process not chemicalhyfalcon - 11-3-2014 at 19:26
Boil till you just start to see crystals floating on the top of your solution. This indicates a supersaturated solution. Cool to 0C and then filter.
Most if not all of the Ammonium Chloride should stay in solution. Most of the the sodium bicarbonate will fall out when cooled to 0C.Refinery - 12-3-2014 at 08:48
In the solvay the solution is heated to 35-40C because the decomp temp of ammonium bicarbonate is about 40C. Sodium chloride and amm. bicarb will
react to form ammonium chloride and sodium bicarbonate. When cooled down sodium carbonate will mostly precipitate out. The solution probably contains
some impurities of sodium chloride and sodium bicarbonate. The bicarb will decompose to sodium carbonate when heated(boiled). If it is evaporated to
dryness, ammonium chloride can be purified by sublimation to very high purity. It will actually decompose when heated, but HCl and NH3 will instantly
react to form ammonium chloride at the condensation surface.hyfalcon - 12-3-2014 at 11:04
You better check your decomposition temperature of sodium bicarbonate.shc25 - 12-3-2014 at 11:33
sodium bicarbonate decomposes at 50C.
I understand that in Solvay the 40C is needed for the initial production of NaHCO3 + NH4Cl but in this case I have them in solution at room temp.
Do I need a supersaturated solution before cooling (by heating) or can I just cool it to 0C from room temp and have the solid NaHCO3 to measure
Then I can I not just evaporate the water and measure the NH4Cl left hyfalcon - 12-3-2014 at 11:38
In solution, it actually begins to decompose at room temperature.
Pull a vacuum on it with gentle heating to drive the excess water from it and when I say gentle I mean gentle like in a beaker of hot water not so hot
that you couldn't stick your finger in it and hold it there.
[Edited on 12-3-2014 by hyfalcon]AJKOER - 13-3-2014 at 05:16
Here is an extract from my Wikipedia reference on the Solvay process:
"In industrial practice, the reaction is carried out by passing concentrated brine through two towers. In the first, ammonia bubbles up through the
brine (salt water) and is absorbed by it. In the second, carbon dioxide bubbles up through the ammoniated brine, and sodium bicarbonate (baking soda)
precipitates out of the solution. Note that, in a basic solution, NaHCO3 is less water-soluble than sodium chloride. The ammonia (NH3) buffers the
solution at a basic pH; without the ammonia, a hydrochloric acid byproduct would render the solution acidic, and arrest the precipitation.
The necessary ammonia "catalyst" for reaction (I) is reclaimed in a later step, and relatively little ammonia is consumed. "
which clearly stresses the role of the ammonia as a pH buffer enabling the precipitation of the NaHCO3.
