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Zyklon-A
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Well I don't know, I've ordered tons of all sorts of chem related stuff online with no problems at all. Texas thinks it's its own country.
About boiling sulfuric acid, obviously wear long gloves to avoid splashing, goggles, a lab coat (if possible) and maybe more protection, Its always
worked for me , no accidents.
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Romain
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Nope. They don't have it in pool stores...
I never tried to buy sulfuric acid in a car shop, though I don't think they would even sell me a few ml. I'll try anyway in case I'm lucky.
HNO3 (though I don't need it) is my dream but I'll have to make it myself, assuming I can make/find conc. H2SO4.
Boiling sulfuric acid is not exactly something I want to try unless it's the only way to get a concentrated product.
I would probably set my hotplate to 300°C and backup a few meters away with the power cord handy in case it starts to boil.
I have a source of nitrates: 13% ammonium nitrate fertilizer. Pure NH4NO3 is 35% nitrogen content so mine should be ~37% (100/35*13).
I got my NaOH (1l of 30.5% solution) in France last year though I was affraid I would be controlled at the border, so I didn't take anything else with
me.
I never got searched when I crossed the border but I don't want to risk anything.
And for the gold bar, I just asked at my local bank for 1g and they ordered it for 54$ (49CHF). They said it would arrive Tuesday.
EDIT: typo
[Edited on 12-1-2014 by Romain]
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bismuthate
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Would a solution of sodium (or calcium) ethoxide react with acetlene to form sodium carbide?
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Nicodem
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See The literature searching guidelines (chapter 2.2) for a review of online pKa tables where you can easily find the answer to your question.
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thebean
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What glove brands can you confirm as having diethylhexyl phthalate as a plasticizer? (USA)
"You need a little bit of insanity to do great things."
-Henry Rollins
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papaya
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Hi comrades,
can anyone help me to find mutual solubilities (or is this the same as phase diagram?) in NaCL + NaOH + H2O system. In other words I want to know how
much NaCL will dissolve in some % NaOH aqueous solution at STP, or maybe how much NaOH in some % NaCl.
Thanks in advance.
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Brain&Force
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Товарищ papaya, it might help to know the Ksp's of each compound in question. The Ksp of NaCl is 37.7 and
the Ksp of NaOH is 7.7 (calculate by squaring the molar solubilities of each compound).
I have a similar question. If a mixture of an alkali metal carbonate and alkali metal hydroxide are reacted with an acid, will one react before the
other, or will they both react at the same time. If one reacts preferentially, why?
At the end of the day, simulating atoms doesn't beat working with the real things...
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Romain
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Hi,
I just tested gold electrodes for electrochemical sulfuric acid production from copper sulfate: It doesn't work!
I used a 1g gold bar (~8mm X ~16mm).
The electrode erodes very fast (-0.01g at 50ma/cm^2 for 30 minutes)...
I just posted my result in case anyone is interested.
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Zyklon-A
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Wow! That sucks, sorry for the bad advice, do you have a way to recover the gold?
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DraconicAcid
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Quote: Originally posted by Brain&Force  | | Товарищ papaya, it might help to know the Ksp's of each compound in question. The Ksp of NaCl is 37.7 and
the Ksp of NaOH is 7.7 (calculate by squaring the molar solubilities of each compound). |
Actually, I'm not sure that helps- Ksp values are only good for dilute aqueous solutions. In concentrated solutions, the ionic strength changes the
activity of the ions so that they are no longer equal to the concentrations, so the eq'n one normally uses for Ksp is no longer applicable.
ETA: I'm looking at a ternary phase diagram form sodium chloride and sodium nitrate here: http://www.phasediagram.dk/ternary/ternary4.htm
Sadly, it's been so long since I've looked at phase diagrams that I'm still not one hundred percent sure I understand what I'm reading.
| Quote: | | I have a similar question. If a mixture of an alkali metal carbonate and alkali metal hydroxide are reacted with an acid, will one react before the
other, or will they both react at the same time. If one reacts preferentially, why? |
The stronger base will react first. If the acid reacted with the carbonate to give bicarbonate ion, then the hydroxide that is still around would
react with the bicarbonate to give carbonate again.
[Edited on 14-1-2014 by DraconicAcid]
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Zyklon-A
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I started to decompose Ba(NO3), it worked well at first, but then some of the mixture melted, it is not BaO for sure as the melting point is way to
high.
As for nitrate, wiki just says it decomposes.
I stopped it about half way done and let it cool, when I heated it up again, some water condensed on the test tube.
How do I know when it's done?
Edit: Damnit, my tube cracked as it cooled.
Transferred to a vile:
[Edited on 14-1-2014 by Zyklonb]
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Romain
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No, I know of no way to recover the gold but I didn't loose much just 0.01g so it's ok...
I thought it would be inert enough to make some H2SO4 though. Sad.
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Zyklon-A
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The gold is a powder on the bottom right? Not in solution. If so, I think powdered gold would be awesome for an element collection!
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elementcollector1
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What was the purity? If it was anything less than 24k, the base metals in the structure will dissolve while the gold collects as a sludge at the
anode.
Powdered gold would admittedly be cool, but regular gold was hard enough (I had to pan for all of mine, and even then it's barely anything!)...
Although I suppose you could use one of those gold flake containers in tourist and rock shops.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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alexleyenda
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I made KCLO3 boiling down bleach and adding the solution to a saturated solution of KCl. Basically at this point it was a solution of water, KCl,
KCLO3, NaCLO3 and NaCl. After my first filtration, I decided to put it in the freezer just to see how it behaves and if I could get some more chlorate
out of it. 20h laters, I took it out and only 2/3 of the solution was frozen. Why? I know that salts makes the water freeze at lower temperatures but
I can't explain why a part of it was frozen and the rest was liquid.
[Edited on 14-1-2014 by alexleyenda]
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elementcollector1
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I believe this is 'freeze distillation' in action: More concentrated solutions of potassium chlorate freeze at lower temperatures. The ice you see is
relatively pure ice, and the solution should now be quite concentrated in salts, more so than earlier. Have you checked for any more precipitate?
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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alexleyenda
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hmm there were particles in suspension when I poured the first part of it after letting it melt 2-3 minutes. I'm not sure they were there at first,
it was hard to see. I guess I'll search on "freeze distillation", thanks. If anyone knows exactly what happenned let me know !
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DraconicAcid
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What's happened is this- you have a concentrated solution of salts, which will have a much lower melting point than pure water. When you cool it
down, some of the water will precipitate as ice (freezing), and the solution which is left as liquid will become more concentrated, thus having an
even lower melting point. If you take that liquid solution and put it back in the freezer at the same temperature, it will not freeze. If you cool
it down even further, then more water will crystallize out.
If you look at the phase diagram for NaCl-water (there will be similar diagrams for water-KCl, water-NaClO3, etc), if you have a solution
that is 10% salt, it will remain liquid until you cool it to about -8oC. Cooling it further will result in the formation of ice- you will then have
two phases (one is pure water (ice), and the other is brine, the composition of which is given by that black curve; at -10oC, it will be about 12%
salt, and at -20oC, it will be about 20% salt, which means about half of your water will have precipitated as ice). If you cool it down to -21 oC or
-22 oC, then the whole thing will solidify as a mixture of sodium chloride dihydrate and ice.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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elementcollector1
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I have a 2-hole stopper (size 0) that is current occupied by 2 ~3/16" diameter test tubes. The stopper is now too large to fit the test tube intended
for it. What do I do? I don't have any larger test tubes, and the reaction is only meant for a test tube - not a beaker or flask.
Elements Collected:52/87
Latest Acquired: Cl
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Metacelsus
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Related question:
One of my stoppers has fallen into a big (2L) Erlenmeyer flask (I accidentally pushed it through the neck), and now I can't get it out. I can't use
any pressure differential schemes because the stopper has a hole in it. Does anyone know of a good way to get the stopper out? Preferably, it would be
non-destructive, but I am willing to sacrifice the stopper to save the flask.
@ec1: is it that the tubing stretches the stopper so that it doesn't fit any more? You could put in the stopper first, and then put in the tubing, if
this is the case.
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elementcollector1
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| Quote: Originally posted by Cheddite Cheese | Related question:
One of my stoppers has fallen into a big (2L) Erlenmeyer flask (I accidentally pushed it through the neck), and now I can't get it out. I can't use
any pressure differential schemes because the stopper has a hole in it. Does anyone know of a good way to get the stopper out? Preferably, it would be
non-destructive, but I am willing to sacrifice the stopper to save the flask.
@ec1: is it that the tubing stretches the stopper so that it doesn't fit any more? You could put in the stopper first, and then put in the tubing, if
this is the case. |
Do you have one of those things where you press or slide a button and three prongs extend outward? If not, something along the lines of a miniature
crowbar might work. Screwdriver, maybe?
As an update, I managed to get the test tubes and the stopper in - however, the arrangement is quite haphazard: One of the tubes broke the outside of
the stopper, and is now touching the test tube wall. But hey, it works. Hopefully. Probably going to tape (if not glue) that just to make sure.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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Zyklon-A
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| Quote: Originally posted by elementcollector1 | | I believe this is 'freeze distillation' in action: More concentrated solutions of potassium chlorate freeze at lower temperatures. The ice you see is
relatively pure ice, and the solution should now be quite concentrated in salts, more so than earlier. Have you checked for any more precipitate?
|
Next time boil the bleach a lot more, I made the mistake of not boiling it enough, and there was to much water so the chlorate would not precipitate.
You shouldn't have to cool it down past 1C. If nothing precipitates at 1C, boil it down some more and try again.
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alexleyenda
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Thank you very much Draconic, it really helped, I never heard of that before! I guess I had frozen water with KClO3 cristals trapped in it that
precipitated before it froze and the remaining brine was mainly NaCl and KCl in soluition.
I already had 4 grams out of it before I did that Zyk,
| Quote: Originally posted by alexleyenda | | After my first filtration, I decided to put it in the freezer just to see how it behaves and if I could get some more chlorate out of it.
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[Edited on 15-1-2014 by alexleyenda]
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alexleyenda
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| Quote: Originally posted by Cheddite Cheese | Related question:
One of my stoppers has fallen into a big (2L) Erlenmeyer flask (I accidentally pushed it through the neck), and now I can't get it out. I can't use
any pressure differential schemes because the stopper has a hole in it. Does anyone know of a good way to get the stopper out? Preferably, it would be
non-destructive, but I am willing to sacrifice the stopper to save the flask.
@ec1: is it that the tubing stretches the stopper so that it doesn't fit any more? You could put in the stopper first, and then put in the tubing, if
this is the case. |
I got a quite random Idea, but it could work. Find a screw with a large head, put it in the hole of the stopper and pull it out of the flask with
pliers attached to the rod of the screw.
[Edited on 16-1-2014 by alexleyenda]
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alexleyenda
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I was boiling H2O2 in a beaker to concentrate it and when it was almost done I added some more 3% H2O2 ( I only added around 40 mL, room temp, poured
in the beaker directly from the bottle) in the 2000 mL beaker containing around 400 mL boiling 30% H2O2 and the bottom of the beaker cracked because
of the thermal shock. I was wondering if that was due to the fact that I used a cheap chinese beaker (borosillicate and not bubbled but still...) or
if this thermal shock was too big and even a pyrex beaker would have cracked.
In other words, do I need to buy better beakers, or to pour slower with like a separatory funnel.
Sorry for 3 posts in a row but this one is like 24h later, if I just edit it won't show up as new.
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