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UnintentionalChaos
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Quote: Originally posted by Panache  | | Quote: Originally posted by UnintentionalChaos | I'd suspect that the per-citric acid may form, but decomposes, ejecting O2 and CO2. I would suspect that the citric acid is completely consumed in the
process with things like the very oxidation-prone acetone dicarboxylic acid being generated as intermediates.
Keep in mind that the HCl will generate HClO and Cl2 with the H2O2. Perhaps you should test it with dilute H2SO4 instead.
[Edited on 5-13-09 by UnintentionalChaos] |
Apologies i should have mentioned the evolved gas was colourless and without odour. |
I had meant more along the lines that hypochlorite or chlorine would be a powerful intermediate oxidizer that might have screwed with your results.
Perhaps straight citric acid and H2O2 acidified with H3PO4 or H2SO4 would behave differently.
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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Lambda-Eyde
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First off, let me put something straight. I have absolutely no experience in practical "large-scale" chemistry, such as distillations and such.
So tonight I attempted my first distillation. I decided on distilling 50 ml of 30% technical grade hydrochloric acid.
As I suspected, HCl is a real bitch to distill.  But let's not blame the acid
alone.
Knowing that aqueous HCl forms an azeotrope at 20,2% HCl I confered with ScienceGeek for some advice on the distilliation. He adviced me to start the
distillation slowly, so that as little HCl gas as possible escaped from the solution, increasing the yield.
A happy first year student then set out on this adventure, plugged in the heating mantle, assembling the apparatus, greasing the joints, rigging
cooling and covering the still head in aluminium foil.
It took me two hours to reach the boiling point of 30% HCl, 90*C. Don't laugh!  I see that I was waaaay too careful in the beginning, and I could have sped it up many times.
Yes, I know a distillation isn't done in 15 minutes, but 2 hours to reach the boiling point is just ridiculous.
During the distillation I observed lots of white smoke at the vacuum port which obviously was HCl gas escaping from solution, so the concentration in
the distillation flask must have changed.
I have approx. 5 ml of nice and clear HCl solution in the receiver flask of unknown concentration, but about 20 minutes ago the distillation came to a
halt. The temperature is now 98*C and I suspect that the concentration in the distillation flask has reached something around 20% (b.p. 108*C).
I don't dare to take this distillation any further as my thermometer only goes to 110*C. 
But at least I got a few milliliters of nice reagent grade HCl!
So, do any kind souls feel like sharing any advice on such procedures? What did I do wrong, what could I do to speed it up without getting a horrible
yield and so forth. Any advice is appreciated. 
(Yes, I'm gonna buy another thermometer that goes beyond 110*C!)
[Edited on 14-5-2009 by Lambda-Eyde]
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Jor
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You should just heat the flask with a hotplate/mantle/burner.
When you heat concentrated hydrochloric acid, pure HCl will evolve until the concentration of the solution has reached about 20% (azeotrope). In the
same matter, when you heat dilute 10% acid, you will drive off water until the concentration will be about 20%. Then the temperature of the
thermometer will rise, and the azeotrope comes over.
I see no benefit in doing the distillation slowly. Just lead the evolved HCl in water for extra hydrochloric acid. You do this by attaching a tube at
the vaccuum port and leading this into water.
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woelen
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Jor, what you are saying is dangerous. You should not simply attach the vacuum port to a tube and put this in water. As you already stated, with 30%
HCl, you first will get fairly pure HCl gas when the solution is heated. When the pure gas is hitting the water, it will dissolve and then at once,
the water flashes back into the apparatus. This is going fast, amazingly fast and the cold water will suck back into the distillation setup in a
fraction of a second and the apparatus will crack 
I once was surprised by suckback of water in an HCl-filled system and this is going fast, soooo fast! I was really shocked to see this happening.
What you can do is take an inverted funnel, which just is a few mm below the water surface, with the water in a beaker, which is just somewhat wider
than the funnel. When the HCl is absorbed, then the water is sucked into the funnel, but this causes the water surface to lower and then the funnel
looses contact with the water surface and no suckback occurs into the distillation setup.
Another option is to dilute the acid somewhat before distilling (take 2 parts of acid and 1 part of water) and then distill this acid. The thermometer
must go well beyond 120 C if you want to distill the water/HCl azeotrope.
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Lambda-Eyde
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Thanks for your answers!
| Quote: Originally posted by woelen | When the pure gas is hitting the water, it will dissolve and then at once, the water flashes back into the apparatus. This is going fast, amazingly
fast and the cold water will suck back into the distillation setup in a fraction of a second and the apparatus will crack
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I would of course have used a trap in the line. The water would anyways only be sucked into the cold end of the distillation setup, wouldn't it?
Yes, I thought about diluting the HCl to 20,2%, but in the end I decided against it for some reason. Azeotropes are cumbersome!
Wikipedia says (yes, I know. I'm saving up for a CRC) the boiling point for a 20% solution is 108*C, but at 10% it drops to 103*C. No data is provided
of the boiling points between those two, but you say that it will go beyond 120*C ? Sigh... I hate limitations.
[Edited on 15-5-2009 by Lambda-Eyde]
[Edited on 16-5-2009 by Lambda-Eyde]
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querjek
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Can anybody give me a quick rundown as to what a spectrophotometer does? I bought one (cheaply) today, thinking, "oh, yeah, this'll be really useful",
but now can't remember whether it measures concentration or (roughly) identifies samples.
it's all about chemistry.
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Paddywhacker
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| Quote: Originally posted by querjek | | Can anybody give me a quick rundown as to what a spectrophotometer does? I bought one (cheaply) today, thinking, "oh, yeah, this'll be really useful",
but now can't remember whether it measures concentration or (roughly) identifies samples. |
See http://en.wikipedia.org/wiki/Spectrophotometer
In order to make any use of it you will need the ability to weigh accurately, and to measure volumes of liquid accurately.
You get it at a good price? Going to resell? What make and model?
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Panache
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Does anyone have a neat trick for reducing iodate back to iodine. I know of the common methods (KI, H2S, Hydrazine, SO2, ascorbic) but these all fail
the 'i don't have any at hand nor am inclined to make some' test. The iodate is a potassium/sodium with about 20% citric acid and a few of salts of
PO4 mixed in for good measure.
Many many references have said it undergoes oxygen loss reverting to KI with heat but i have not found this happens to any significant extent. However
i only raised it to around 600C for some minutes. I can heat it heavily up to 1200, this would also remove the citrate. Is this foolish?
Oh it also now has some chromium in it as it turned green when i was evaporating off the water (stainless dish).
Oh also does anyone know of illegal/diversionary uses for benzotriazole, a sales rep asked me because he had a dodgy looking customer asking for it
and when he asked the customer how much he was after the customer said 'how much have you got! Ha!'
[Edited on 16-5-2009 by Panache]
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Lambda-Eyde
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querjek: http://www.chemguide.co.uk/analysismenu.html
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UnintentionalChaos
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Panache- if you have some sulfite/bisulfite/metabisulfite, you have a source of SO2 sitting around, waiting for a little acid. That'll get you back to
iodide, but getting it out may be a bit tricky.
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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S.C. Wack
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A way to avoid an iodine clock reaction is hot aq. oxalic acid. It works.
[Edited on 16-5-2009 by S.C. Wack]
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woelen
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| Quote: Originally posted by Lambda-Eyde | Thanks for your answers!
I would of course have used a trap in the line. The water would anyways only be sucked into the cold end of the distillation setup, wouldn't it?
Yes, I thought about diluting the HCl to 20,2%, but in the end I decided against it for some reason. Azeotropes are cumbersome!
Wikipedia says (yes, I know. I'm saving up for a CRC) the boiling point for a 20% solution is 108*C, but at 10% it drops to 103*C. No data is provided
of the boiling points between those two, but you say that it will go beyond 120*C ? Sigh... I hate limitations. |
Sorry for my mistake, I accidently exchanged the BP of nitric acid azeotrope with that of HCl azeotrope.
But even then, if your thermometer is going up to 110 C then you indeed should not attempt this distillation, you will be on the edge with that and
just a little overheating will cause destruction of your thermometer. What can be done though is just putting in a stopper in the hole, where you
normally put the thermometer. Best of course is a glass stopper, ground joint, but if you don't have that, then you also can use a rubber stopper,
which is covered in a few layers of white teflon tape (from hardware stores, used for fixing water taps and the like), before it is pushed in the
hole. The teflon tape will give sufficient protection, certainly for the duration of the distillation. With HCl you can do that, it is not that
corrosive. Even a cork stopper can be used, but again, this must be well covered with teflon tape, no cork may be exposed directly to the acid vapors.
If you have the 10% mix, then the boiling point may be 103 C, but don't expect to get a 10% mix boiling off. At 10% concentration, you get almost pure
water, and I think that it certainly will contain less than 1% of HCl.
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Lambda-Eyde
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Thanks again for a helpful answer, woelen.
My distillation apparatus is all 24/40 ground glass, however, my thermometer adapter has a rubber fitting. I'm going to purchase a teflon adapter from
UGT along with a vigreux column very soon (and of course a new thermometer).
I didn't get 103 C on the thermometer when I distilled, that number was also taken from the wikipedia article. I'm curious as to what the bp's between
10% (103 C) and 20% (108 C) are, more specifically if they exceed 110 C. But like you say, slight overheating could reduce my thermometer to pieces.
By the way, the graduation of the thermometer ends at 110 C, but the inner tube doesn't come to an end until about 3 centimeters above the end of the
graduation. I understand that I shouldn't go above 110 C just because of that, but it doesn't mean that I shouldn't go to 108 C?
Regarding the concentration of the HCl, I don't know what its concentration is right now, either in the still pot nor the receiver. And I have no way
of determining it either, as I lack a burette and a suitable titrant (not to mention an indicator...).
The starting concentration was 30% HCl, and I'm pretty sure it has reached the azeotropic concentration. At least it has according to my logics.
And a slight digression (yes, even more questions!): I have a one liter solvent can containing 70-100% dichloromethane, 5-10% formic acid, and 1-5%
"anionic surfactants". Obviously, I want the CH2Cl2 for lab purposes, and I'm thinking of distilling it. The boiling point of CH2Cl2 is 40 C, and for
HCOOH it is 101 C. I suppose I won't have to do a fractional distillation, a simple distillation will suffice, right?
And those surfactants, how will I get rid of those? I don't know anything about their chemical composition or physical properties.
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UnintentionalChaos
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Lambda-Eyde- If you add a base, you will make formate salts and not need to worry about the volatility of formic acid. The anionic surfactacants
probably have a relatively high molecular weight and will be virtually nonvolatile under the temperatures needed to distill DCM. You will just need to
let the DCM sit over a dessicant after distillation since neutralizing the formic acid will generate a small amount of water.
[Edited on 5-17-09 by UnintentionalChaos]
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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Lambda-Eyde
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And I suppose neither NaOH or CaCl2 will react with the DCM?
Thanks!
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stateofhack
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Hello,
Quick thing that i am not sure about, i need to buffer 1.5 mls of 70 % ethylamine sol with Glacial acetic acid, how much would it take?
edit: found it.
[Edited on 18-5-2009 by stateofhack]
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manimal
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Some solubility values for substances list two different values: one for 'anhydrous' substance, and one for a 'hydrated' substance.
My question is this: How do you have an aqueous solution of an anhydrous compound? For example, wikipedia lists the solubility for calcium sulfate
(anhydrous) as .0021g/100ml and calcium sulfate (dihydrate) as .24g/100ml.
[Edited on 19-5-2009 by manimal]
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Maja
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Does sodium sulfite(Na2SO3) forms bisulfite adduct with aldehyde and ketone groups ?
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woelen
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| Quote: Originally posted by manimal | Some solubility values for substances list two different values: one for 'anhydrous' substance, and one for a 'hydrated' substance.
My question is this: How do you have an aqueous solution of an anhydrous compound? For example, wikipedia lists the solubility for calcium sulfate
(anhydrous) as .0021g/100ml and calcium sulfate (dihydrate) as .24g/100ml.
[Edited on 19-5-2009 by manimal] |
There indeed can be quite a strong difference between solubilities of anhydrous compounds and hydrated compounds. I myself made anhydrous NiSO4 and
that hardly dissolves in water, while NiSO4.6H2O dissolves very well. Another example is CrCl3 vs. CrCl3.6H2O.
This has to do with other structure of the compounds. Many of these anhydrous compounds are not really salts, but form covalent (sometimes polymeric)
units. Some of these anhydrous salts very slowly are hydrated and if that is the case, then indeed you can distinguish two different solubility
figures for the two. E.g. nickel sulfate anhydrous only dissolves with hundreds of mg in 100 ml of water, while the hydrated salt dissolves with tens
of grams in 100 ml of water. If you wait for several days, then the anhydrous nickel sulfate also dissolves to a much larger extent and this is
because it slowly is hydrated and then the solubility of the hydrated salt comes into play again.
This distinguishment only is important for anhydrous salts which slowly are hydrated in water. When the anhydrous salt immediately is hydrated (such
as CuCl2 or CuSO4) then of course one cannot speak of different solubilities, then only the hydrated compound is present in aqueous solution. Calcium
sulfate also slowly hydrates. Think of gypsum sculptures, which are mixed with water and which take a day or so to harden.
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smuv
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| Quote: | | Does sodium sulfite(Na2SO3) forms bisulfite adduct with aldehyde and ketone groups ? |
No. The reaction would be reversible, forming 1eq hydroxide for every eq of adduct. Just buy sodium bisulfite....or if you are really desperate try
halfway neutralizing a solution of it.
"Titanium tetrachloride…You sly temptress." --Walter Bishop
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Formula409
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At school we were tasked to analyse an unknown white powder. Nitrate ions were confirmed with the "brown ring" test, and an emerald green colour was
observed with a flame test. My first thought was that it was Zinc Nitrate, so I added a solution of sodium carbonate which garnered a white
precipitate of Zinc Carbonate (so I thought). The teacher, however, suggested that I may be wrong and that I should look at other metal ions which
give a green colour in the flame. I tested for barium with the addition of a sulfate which yielded nothing - I don't know what else to look for. This
is only high school so it would not be anything TOO exotic.
Formula409.
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DJF90
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According to this site:
http://chemistry.about.com/library/weekly/aa110401a.htm
Emerald green is copper (not a halide). Which makes perfect sense if you tested positive for nitrate. However one thing doesnt add up. You said it was
a white powder... copper nitrate would be blue/green. Unless of course its copper (I) nitrate... that would be colourless right? But how stable would
that be in aqueous solution?
Either that or its Thallium nitrate
EDIT: It doesnt look as if copper (I) nitrate exists... well at least not as a compound that is usable in the lab. But surely your teachers wouldnt
give you thallium salts to vaporise in the flame... I would expect anhydrous Copper (II) nitrate to be "white", but surely upon adding water the Cu
(II) ion would hydrate to the blue aqua ion [Cu(OH2)6]2+ ?
[Edited on 24-5-2009 by DJF90]
[Edited on 24-5-2009 by DJF90]
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Formula409
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| Quote: Originally posted by DJF90 | According to this site:
http://chemistry.about.com/library/weekly/aa110401a.htm
Emerald green is copper (not a halide). Which makes perfect sense if you tested positive for nitrate. However one thing doesnt add up. You said it was
a white powder... copper nitrate would be blue/green. Unless of course its copper (I) nitrate... that would be colourless right? But how stable would
that be in aqueous solution?
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Exactly what I said too! Plus the addition of sodium carbonate would yield a coloured precipitate!
Formula409.
[Edited on 24-5-2009 by Formula409]
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UnintentionalChaos
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| Quote: Originally posted by DJF90 | | I would expect anhydrous Copper (II) nitrate to be "white", but surely upon adding water the Cu (II) ion would hydrate to the blue aqua ion
[Cu(OH2)6]2+ ? |
Anhydrous copper nitrate is also blue, so it definetly isn't that, plus a blue aqueous solution is pretty noticable.
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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Sedit
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Two quick questions one more of an urgent nature..
First one. I have alot of good cloths of mine that when we put in our new well they forced us to go 125 ft down when our last one was only 50. Well
the water sucks back and is loaded with iron. Almost all of my good cloths are ruined with iron hardwater stainds and a dingy orange yellow color. I
have tried everything from bleach to hard water rust remover to peroxide soaks and nothing touches it. I am not considering a dilute muratic acid soak
but thats a last resort.
Any suggestions as to what I could use to turn my whites white again?
Another question thats bugs me a bit.
Earlyer I made pasta. After it was done I started eating it with garlic bread and ate a decent amount as did my children until someone else came in
and tasted the meat and said it was bad... Damned if it wasn't to. I couldn't tell the difference over the taste of garlic and my grandfather ate two
plates... It was cooked very well but I was woundering if I should worry for myself/kids and more important my grandfather.
Knowledge is useless to useless people...
"I see a lot of patterns in our behavior as a nation that parallel a lot of other historical processes. The fall of Rome, the fall of Germany — the
fall of the ruling country, the people who think they can do whatever they want without anybody else's consent. I've seen this story
before."~Maynard James Keenan
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