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ninhydric1
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[*] posted on 3-9-2017 at 17:45


Magnesium also works. I'm not sure if this works immediately, but HF reacts with glass. Maybe that could be used as a test?

[Edited on 9-4-2017 by ninhydric1]
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[*] posted on 3-9-2017 at 18:01


Calcium fluorite is amongst the most insoluble things there is. Other calcium halides are soluble and so a calcium salt makes for a sensible way of differentiating fluoride from other halides.

As for any test of this type, you presumably have a limited number of possibilities of what your compound contains and your test differentiates between those possibilities. So, context helps in answering this question. What other acid might it be if it is not hydrofluoric?

Testing a drop on a glass slide for 24 hours would seem to be a good idea. Even low concentrations will etch the glass in time. There really is not much else that will attack glass and so this would be pretty definitive.
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[*] posted on 4-9-2017 at 09:14


Here was my procedure. I acquired some 1,1 difluoroethane (DFE) as electronics duster. I did some research on hydrofluorocarbons and found that when these burn they often produce hydrogen fluoride and carbonyl fluoride. When carbonyl fluoride comes in contact with water it hydrolyses to HF and CO2.
My main dilemma was the carbonyl fluoride. I was pretty scared to mess around with carbonyl fluoride considering that it is toxic at 2 ppm. But I decided to do a small scale reaction anyway. I filled a test tube about halfway full of 1,1 DFE. I then put about 1ml of water in the bottom of the test tube. I bent a wooden splint so it would fit down in the test tube with my hand out of the way. I lit it on fire and stuck it down in the tube. The DFE took longer to burn than expected. About 5-7 seconds. Once it was done burning I quickly turned on the fume hood and stoppered it with a butyl rubber stopper. I vigorously shook the tube. Then I opened the tube and squirted a little bit on the walls of the test tube to get the excess HF of the walls. I was left with an acidic clear liquid.
The only things that I could imagine being in the liquid would be Hydrofluoric acid and a small amount of Carbonic acid. Obviously I would expect any carbonic acid to decompose shortly. Thus I would be left with reasonably pure Hydroflouric acid.
But I am not 100% sure if it is or not. That is why I asked. Like I said, I am working small scale. This means that I only have <1ml of solution at dilute concentrations so I don't have a lot to work with. I may try the magnesium method and as a secondary test the glass slide.
It would be much appreciated if I could get some input on the reaction going on here and tell me if I'm doing something wrong.
I cannot find any information online referring to the mechanism or the chemical equation to how this works so I will have to check out the library at my college to see if I can find any information.
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[*] posted on 4-9-2017 at 14:42


Given that, I would smear a tiny portion on a microscope slide and leave it for a number of hours. i would attempt to protect it from evaporation -- perhaps a cover slip and an evacuated zip-seal plastic bag. Then examine the surface of the slide under a microscope and see if any etching had taken place.
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[*] posted on 8-9-2017 at 11:34
Question: Solubility of multiple ions


I am working on making a safer version of a lead acid battery that uses sulfate salts instead of sulfuric acid for the electrolyte. The two salts I am considering are ammonium sulfate and magnesium sulfate, but neither are soluble enough on their own to put enough sulfate ions into solution for what I want to do. What limits how much of these salts can dissolve? Is it the magnesium/ammonium or the sulfate? What I am really asking is is the solubility of of ammonium sulfate and magnesium sulfate in the same solution the same as if they were in separate solutions?

I am away from my lab so I cannot test this experimentally.

[Edited on 8-9-2017 by Plunkett]
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[*] posted on 8-9-2017 at 13:27


Quote: Originally posted by Plunkett  
What I am really asking is is the solubility of of ammonium sulfate and magnesium sulfate in the same solution the same as if they were in separate solutions?

I am away from my lab so I cannot test this experimentally.


No. There's two factors at play here- first, an ionic compound will be more soluble in a solution with high ionic strength than in pure water. Secondly (and more importantly), two compounds with common ions will be less soluble in the same solution than in pure water, or a solution without common ions.

A saturated solution of MX (generic ionic compound) will have the equilibrium MX(s) <=> M+(aq) + X-(aq). If you already have some X- in the solution from the presence of another compound (say, NaX or KX), the equilibrium will shift to the left (less MX dissolves).




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[*] posted on 9-9-2017 at 21:42


Just out of curiosity, would copper(ii) form complexes with amphetamine or methamphetamine? It would be interesting if such complexes existed

Totally theoretical, I have no interest in drugs.




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[*] posted on 11-9-2017 at 03:05
Oxalate or Carbonate


I have a solution with in which iron is leached using oxalic acid. What will precipitate if I add NaHCO3 in the solution ?

it will be Sodium Oxalate or Iron Carbonate or both ?
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[*] posted on 11-9-2017 at 20:01


Depends on how saturated your solution is:
Sodium oxalate has a solubility of 3.7 g per 100 mL of water at 20 degrees Celsius, so it depends on how much iron oxalate is in solution. If you are able to keep all the sodium oxalate in solution, iron carbonate will be almost all of the precipitate that forms from the addition.
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[*] posted on 17-9-2017 at 16:05
Red P


Is it possible to get pure red phosphorus from red phosphorus fire retardant? Its added to some kind of resin. I wasnt sure if there would be a way to dissolve the resin leaving behind the red phosphorus?
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[*] posted on 17-9-2017 at 22:55


Quote: Originally posted by fabtasticwill  
Is it possible to get pure red phosphorus from red phosphorus fire retardant? Its added to some kind of resin. I wasnt sure if there would be a way to dissolve the resin leaving behind the red phosphorus?


I'm not sure about how to actually isolate the phosphorus, you'll need to check what the other components are to get an idea of how to separate them, try checking an msds?

But I think it's really interesting that they ate using phosphorus to put out fires considering it's history... it seems the mechanism is the creation of phosphoric acid which is able to subdue plastic fires.




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[*] posted on 18-9-2017 at 12:31


It says that there is a layer of polymer protective on the surface, though others seem to be in some sort of resin. Would it be simple to dissolve the polymer without damaging the red phosphorus?
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[*] posted on 18-9-2017 at 15:15



Fire retardant doesn't attempt to put out wildfires or even necessarily halt flames in their advance. Consisting primarily of ammonium phosphate — fertilizer, basically — fire retardant is formulated to slow down the combustion of trees, brush and grass. The idea is to give firefighters time to mount a ground attack. Excerp wikipedia
There is no red phosphorous it is a dye
Aluminum and ammonium phosphate in a exothermic reaction
Will produce phosphorous but it will be white phosphorous.
Just use fertilizer it is a more pure source




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[*] posted on 19-9-2017 at 12:10


My goal is actually to get white phosphorus. So if I burn it in an inert atmosphere so the phosphorus doesnt burn it will produce white phosphorus?
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[*] posted on 21-9-2017 at 00:27


Quote: Originally posted by gluon47  
Just out of curiosity, would copper(ii) form complexes with amphetamine or methamphetamine? It would be interesting if such complexes existed

Totally theoretical, I have no interest in drugs.


Should do, at a guess, I imagine however that the freebase coordinated to the copper would still be susceptible to reaction with atmospheric CO2 as all such amines are, I would doubt you could isolate the complexes as solids though but really I actually wouldn't really know some stranger things form stable complexes, I think silver forms a volatile complex with pyridine if I recall.

edit - the qualifier 'I have no interest in drugs'...bit tedious, bit irrelevant, no one cares anyway and if they do, they are unlikely to believe your caveat anyway.

[Edited on 21-9-2017 by Panache]




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[*] posted on 21-9-2017 at 06:31


What does fluorine smell like?



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[*] posted on 21-9-2017 at 07:29


Probably similar to chlorine as in it attacks the nose.



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[*] posted on 1-10-2017 at 05:26


Would bubbling chlorine through a solution of KNO3 yield NO3 gas which would react with water to make HNO3?



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[*] posted on 1-10-2017 at 05:34


Quote: Originally posted by xfusion44  
Would bubbling chlorine through a solution of KNO3 yield NO3 gas which would react with water to make HNO3?


No.




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[*] posted on 1-10-2017 at 05:36


Quote: Originally posted by xfusion44  
Would bubbling chlorine through a solution of KNO3 yield NO3 gas which would react with water to make HNO3?


No.

First NO3 is a radical and extremely unstable. But i guess you mean N2O5, which is also unstable but much more stable than NO3. It will react with water however so it will not be formed.

You would get a solution of K+, H3O+, NO3-, Cl- and ClO-.
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[*] posted on 1-10-2017 at 06:06


Oh well, just looking for a cheap and easy way to make HNO3

Thank you both for your answers.

I once tried NurdRage's method of making it and it works but the acid obtained this way is of rather low concentration and a lot of NO2 is wasted.




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[*] posted on 1-10-2017 at 06:19


Quote:
Oh well, just looking for a cheap and easy way to make HNO3

Don't give up on chlorine, just yet ─ the dry gas will displace NO3 from AgNO3 as the anhydride, N2O5?




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[*] posted on 6-10-2017 at 19:44


Which sulfamate salts are insoluble in water?
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[*] posted on 14-10-2017 at 19:11
Stopcock size


Are the sizes of glass stockcocks standardized? I need one that has a 2mm bore, ~30mm long, ~13mm widest, ~10mm narrowest. This eBay listing has dimensions that are close, but not exactly the same (.4mm off). Is this just error of measurement or the wrong size?

Some pictures of what needs the stopcock:

IMG_0070.JPG - 1.2MB IMG_0071.JPG - 1.2MB

As always, thanks for any help :)




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[*] posted on 17-10-2017 at 09:03


Does PTFE adsorb much water?

I want to render an auger-type powder dispensing funnel rigorously dry, but the PTFE auger has a phenolic looking handle that isn't obviously removable so I would hesitate to heat it at all.

Is the PTFE already dry if it looks dry, or is there some procedure for drying it that doesn't involve the ovens or torches used to dry glass?

This is a pretty expensive piece of glassware for me, and I don't want to ruin it by doing something stupid, but I would really like to be sure of its dryness.
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