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Zyklon-A
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You could use a syringe or something like that. Or you could tintrate it against a know amount of barium nitrate solution.
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Fulmen
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What sort of accuracy can one expect from a disposable plastic syringe? I tested the first batch this way with a 1ml syringe, it gave me a density of
1.8. But once you reach 80-90% you really need a resolution of 0,01g/ml to determine concentration, and I just don't believe this is achievable with
make-shift equipment.
Right now I'm pondering if it's possible to use the weird freezing points to both concentrate and determine concentration. A combination of boiling
and fractionate freezing could possibly give a useful result without having to breathe in too much noxious gas.
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blogfast25
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Quote: Originally posted by Fulmen  | What sort of accuracy can one expect from a disposable plastic syringe? I tested the first batch this way with a 1ml syringe, it gave me a density of
1.8. But once you reach 80-90% you really need a resolution of 0,01g/ml to determine concentration, and I just don't believe this is achievable with
make-shift equipment.
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A glass measuring cylinder of 10 ml, used relative to deionised water, should give fairly good accuracy. You can't beat a decent pycnometer with mg
scales of course, for 4 or more significant digit densities.
And there's always good old acidometry, easier and faster than you might think if you're well organised.
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Fulmen
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The acid is undergoing some freezing experiments now, so it'll have to wait until morning. I'm just so used to quality precision equipment I have a
real time trusting these improvised methods. A titration seems like the best option, I'll have to dig around to find a suitable indicator.
After a bit of reheating the acid now freezes partially (perhaps 3/4) at -20°C. This should indicate appr 90%, decanting the liquid should give me
92% according to this: http://www.generalchemical.com/assets/pdf/Sulfur_Trioxide_Wa...
This seems to be the limit with the lab hotplate, I'll try again with another hotplate and a 500ml flat-bottom florence flask. Hopefully the long stem
will act like a crude column and add a few % of separation without too much loss. Would be nice to reach 05%, but I'm not complaining abut the results
so far.
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Fulmen
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This is weird. After freezing at -20°C I got a waxy solid precipitation at the bottom that would not melt even at room temperature. So I did a couple
of titrations using 1 & 10% NaOH (w/w) and litmus indicator, the NaOH was technical grade (drain opener, not sure what purity one can expect).
According to this the liquid phase was 95% while the solid was 87%. I'll have to recheck my results, but it could indicate that fractional freezing
could work. Only problem is the fact that sulfuric acid should be liquid at room temperature, regardless of concentration. Does anybody have a better
theory?
I also found some K2CO3 puriss (min 98%) and haematoxylin indicator (pH 5-7). Would titrating the acid with the carbonate work reliable? I can't seem
to find any examples of this.
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bbartlog
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Sulfuric acid has a monohydrate (corresponding to 84.5% concentration) which apparently has a fairly high melting point - above freezing, anyway.
Phase diagram: http://image.bayimg.com/aakdgaabi.jpg
I'm guessing that's what you have, though 'room temperature' still seems like it should be high enough to melt this.
The less you bet, the more you lose when you win.
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Fulmen
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Actually this is something else completely. Heating to 80C partially dissolved the solid, but it solidified again upon cooling. I have no idea what
this is, perhaps a salt of some sorts? It's soluble in water but not cold concentrated acid.
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Metacelsus
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Sounds like a bisulfate of some sort.
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Fulmen
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Maybe. Not sure where the cation would come from, but it's pretty obvious it's some unwanted contaminate. Don't think I'll be wasting more time on it,
it's just a matter of freezing it out and decanting.
I'd like to find a better method for determining strength, I don't trust the purity of the sodium hydroxide. I'll have to try titrating K2CO3 with
diluted acid to see if that gives a different result.
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blogfast25
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| Quote: Originally posted by Fulmen |
I also found some K2CO3 puriss (min 98%) and haematoxylin indicator (pH 5-7). Would titrating the acid with the carbonate work reliable? I can't seem
to find any examples of this. |
For titrating the salt of strong base and a weak acid (carbonic acid) with a strong acid you need an indicator like Methyl Orange because at end point
the pH will be slightly acidic, due to carbonic acid.
Recrystallise the K2CO3 at least once, then dry it at 200 C or so for 2 h. Prepare a solution 0.1 N (i.e. 0.05 M) K2CO3 using accurate scales and a
calibrated volumetric flask. Note the titer of the solution, calculated from the precise weight of K2CO3 used. This is your Standard Solution.
Prepare a solution of the sulphuric acid with an accurately weighed sample of the acid. Using your expected value of H2SO4 concentration in your acid,
aim for a solution strength of 0.1 N (0.05 M). Use a closed weighing boat (H2SO4 is hygroscopic), accurate scales (1 mg, 0.1 mg is better of course)
and a calibrated volumetric flask.
Titrate, repeatedly, 20.0 ml (pipette!) of Standard Solution with the H2SO4 solution and Methyl Orange, from yellow to orange (pH = about 3.8).
A relative measuring error of 1 - 2 % is within reach if executed properly.
You can use low grade NaOH but NOT without standardising the titrant solution, for instance against pure potassium hydrogenphthalate
(KHP).
[Edited on 15-11-2014 by blogfast25]
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