Sciencemadness Discussion Board - The Short Questions Thread (4)

The Short Questions Thread (4)

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alexleyenda - 8-1-2014 at 12:42

Then I guess it would be friction, cork is good at creating heat with friction so it makes sense, thought i'm not sure it can create enough heat to ignite H2. Anyways to me it's more plausible than static electricity as it takes quite a lot of static e to create a spark at my knowledge. Unless you rub your cork in your hair for luck before you experiment I doubt it explains the ignition :p

Pyro - 8-1-2014 at 15:12

is rectapur (chemically pure, http://www.surechem.co.uk/pdf/chemicalCompweb14509.pdf) anhydrous AlCl3 supposed to look like this? I can tell 3 different colors...

and is it really necessary to grind it to a powder? My lab is kind of damp and the fumes...

Storing magnesium.

Zyklon-A - 8-1-2014 at 15:56

I got some Mg ribbon a few months ago, and put it in a ziplock bag, and then forgot about it. When I took it out, it was coated with a white powder (MgO or Mg(OH)₂), anyway I put it in very dilute HCl for about 5 mins. It's mostly shiny now, but I don't want that to happen again, so I want to find a way to store it without it oxidizing anymore.
I put it in a glass jar and filled it with CO₂, I know Mg reacts with CO₂, but slower than air and only at high temps right?
Would this be affective for long time storage of magnesium?
(The jar of course was sealed.)

[Edited on 8-1-2014 by Zyklonb]

Pyro - 8-1-2014 at 16:02

Just protect it from damp air, in a jar or air tight bag. the oxidation is minimal, I have 2 rolls of Mg ribbon, one was lying around for a year, the other was in an airtight bag. there isn't much difference between the two.

Zyklon-A - 8-1-2014 at 16:45

Well I live in Texas, where it is very humid, and even though the bag was airtight ( I think), it still oxidized quite a bit.
But I guess the jar will work then?

thebean - 8-1-2014 at 17:49

I put my magnesium under mineral oil, it works fairly well.
My question would be about oxalic acid ester syntheses. Does the Fischer Esterification work? If I want to produce dimethyl oxalate or diethyl oxalate will a simple reflux of the oxalic acid the alcohol and a sulfuric acid catalyst work?

Crowfjord - 8-1-2014 at 18:35

Answer @thebean:

Yep, pretty much. There is a synthesis on orgsyn.org that does just that if I remember correctly. I don't have the link off hand, but a structure search for dimethyl oxalate brings it up.

thebean - 8-1-2014 at 19:20

Oh and one other question/conundrum.
I wanted to produce benzoic acid, so I took stoichiometric amounts of sodium benzoate and hydrochloric acid but I didn't see any sort of reaction. The odor is the same and the crystals look untouched. What happened, what can I do to fix it, and if it isn't fixable without a lot of work is the sodium benzoate salvageable?

Crowfjord - 8-1-2014 at 20:14

Benzoic acid is scarcely soluble in water, so the result isn't much of a surprise. The crystals are likely benzoic acid. I think you can extract with a non polar solvent, dry with desiccant of your choice and evaporate to recover.

DraconicAcid - 8-1-2014 at 21:07

You should have dissolved it in water, then added the acid to get a precipitate.

If what you have is benzoic acid (as it should be), dissolve it in hot water (not boiling, or you'll hate the smell), and let it cool slowly to crystallize. They should be white and feathery.

thebean - 9-1-2014 at 10:14

I added approximately 400mls of water to a beaker if the acid is included so I should think it would work. If I remove some precipitate and put it in hot water that should determine whether it is benzoic acid or not, yes?

DraconicAcid - 9-1-2014 at 10:35

400 mL of water should dissolve about 1.2 g of benzoic acid at RT, and 18 g at 90oC. How much did you add? Heat your solution, and if everything dissolves, cool it slowly.

thebean - 9-1-2014 at 13:35

I added .1 moles of sodium benzoate and .1 moles of HCl (actually a tiny bit more because excess HCl would be easier to deal with than benzoate).

DraconicAcid - 9-1-2014 at 13:41

That should give you about 12 g of benzoic acid, so getting it hot should dissolve it all nicely. Once it's dissolved, let it cool slowly to room temp, then chill it in an ice bath for twenty minutes or so, then filter your acid out.

thebean - 9-1-2014 at 15:03

Yeah my theoretical yield was just barely over 12 grams. I did the hot water test and it is definitely benzoic acid.

plante1999 - 9-1-2014 at 17:29

Anyone got a good chemical engineering book? I already have perry.

I mean, a book that would be about design of plants. For example pipes and pumps and all the other chemical plant goodies.

Thanks

Brain&Force - 9-1-2014 at 17:47

Also, how can you tell if an oxide is nonstochiometric? I know MnO₂ is and Tb₄O₇ may be, and they're both brownish-black. I've also seen the color of Ho₂O₃ explained as being caused by stoichometric defects. Can such defects create colors that subtle?

DraconicAcid - 9-1-2014 at 17:52

Quote: Originally posted by Brain&Force

Also, how can you tell if an oxide is nonstochiometric? I know MnO₂ is and Tb₄O₇ may be, and they're both brownish-black. I've also seen the color of Ho₂O₃ explained as being caused by stoichometric defects. Can such defects create colors that subtle?

Yes, they can. One class of defect is called an f-centre; the f stands for "farbe" or "colour", because it can turn an otherwise colourless compound blue.

How to tell if it's nonstoichiometric? No idea....

Romain - 11-1-2014 at 04:32

Hi,
I'm wondering if you can use gold electrodes to electrochemically make sulfuric acid form copper sulfate.
I don't have access (at all) to H2SO4, HCl, HNO3,... where I live so that would be very useful if I could make my own sulfuric acid.
I'd like to etch circuit boards with it (+ H2O2) so if it works I could restore my etchant for almost nothing.

I asked for platinum at my local bank but they only have gold... though I ordered a 1g gold bar anyway because I need some for my element collection.

Zyklon-A - 11-1-2014 at 10:06

I think so. HCl is sold at Home Depot FYI.

Romain - 11-1-2014 at 11:13

Ok thanks!

Also I live in Switzerland and very few chemicals are available here.

HCl, H2SO4, NaOH, acetone, 35% H2O2,... are readily available in France (and I live only 60km away from the border) though I need to bring them back to Switzerland and I don't think it's legal...

elementcollector1 - 11-1-2014 at 11:15

You can order 1g gold bars? How much do they cost?

alexleyenda - 11-1-2014 at 11:20

He lives in switzerland, I don't know if they have any home depot. Anyways, usually you can find HCl in pool stores to lower the PH, you can also find it in hardware stores near the paint thinners. Conc. Sulfuric acid is sold in hardware stores as drain cleaner and 35% sulfuric acid is sold in car pieces shop as battery acid, you have to ask for it. You can concentrate it yourself ( very dangerous however).

HNO3 can be more tricky. At least here in Canada it's impossible to find , you have to make it yourself with nitrate salt like KNO3 (also banned but obtainable from farms) and conc. H2SO4.

edit after I saw your post: Even where I told you they don't have it? That would be real chimicalphobia. And I know what you live, I often think about going in the states to buy chemicals and bring them back but if they catch me at the border I'll be in deep trouble :p

[Edited on 11-1-2014 by alexleyenda]

Zyklon-A - 11-1-2014 at 16:02

Not that dangerous to concentrate H2SO4, as long as you do it outside and don't breath in the fumes, it takes a lot of energy though, boil it >300C.
I never realized how lucky I am, in Texas I can get almost any chemical I want, (mostly online).
Does the border patrol search you when you cross?

[Edited on 12-1-2014 by Zyklonb]

alexleyenda - 11-1-2014 at 16:29

Really? I thought it heard you could not even have chemistry glassare in Texas. Anyways the danger with H2SO4 is that it boils at 320+ °C and that it tends to splash as it bubbles, even with boiling chips. If 98% H2SO4 splashes on you, it's very bad. If 320°C 98% H2SO4 splashes on you, you're in deep _____.

Zyklon-A - 11-1-2014 at 20:00

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. :cool:

Romain - 12-1-2014 at 03:23

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]

bismuthate - 13-1-2014 at 04:45

Would a solution of sodium (or calcium) ethoxide react with acetlene to form sodium carbide?

Nicodem - 13-1-2014 at 09:18

Quote: Originally posted by bismuthate

Would a solution of sodium (or calcium) ethoxide react with acetlene to form sodium carbide?

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.

thebean - 13-1-2014 at 09:51

What glove brands can you confirm as having diethylhexyl phthalate as a plasticizer? (USA)

papaya - 13-1-2014 at 10:00

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.

Brain&Force - 13-1-2014 at 18:24

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?

Romain - 14-1-2014 at 11:10

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.

Zyklon-A - 14-1-2014 at 11:15

Wow! That sucks, sorry for the bad advice, do you have a way to recover the gold?

DraconicAcid - 14-1-2014 at 11:23

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]

Zyklon-A - 14-1-2014 at 12:40

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]

Romain - 14-1-2014 at 12:48

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.

Zyklon-A - 14-1-2014 at 13:24

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!

elementcollector1 - 14-1-2014 at 13:56

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.

alexleyenda - 14-1-2014 at 14:34

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]

elementcollector1 - 14-1-2014 at 14:39

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?

alexleyenda - 14-1-2014 at 15:08

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 !

DraconicAcid - 14-1-2014 at 17:00

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.

elementcollector1 - 14-1-2014 at 17:05

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.

Metacelsus - 14-1-2014 at 18:09

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.

elementcollector1 - 14-1-2014 at 18:34

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.

Zyklon-A - 15-1-2014 at 07:56

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.

alexleyenda - 15-1-2014 at 13:35

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.

[Edited on 15-1-2014 by alexleyenda]

alexleyenda - 15-1-2014 at 13:42

Quote: Originally posted by Cheddite Cheese

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]

alexleyenda - 16-1-2014 at 10:49

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.

elementcollector1 - 16-1-2014 at 13:18

1) You can't concentrate H2O2 like that, it just decomposes.
2) You don't pour a cold liquid into a hot beaker. Period. Even regular ones can't take it.

Zyklon-A - 16-1-2014 at 13:43

How can I test for benzene? I have some zippo lighter fluid, that I want to use as a nonpolar solvent, for a substance that needs to be uncontaminated to the point of edibility. Obviously, I don't want to get cancer. The bottle says nothing about benzene, but I want to be sure. All my other nonpolar solvents won't work in this situation... (I have very few nonpolar solvents)

Sorry if the answer is really obvious, I know very little about organic chemistry.

[Edited on 16-1-2014 by Zyklonb]

Pyro - 16-1-2014 at 13:50

how old is it? what does it smell like?
I wouldn't use lighter fluid for anything that needs to be edible, benzene or not.

about the H2O2
but in his case it shouldn't have broken because:
1)RT is 20*C and BP of water (what H2O2 becomes is you boil it) is 100*C, 80*C difference in temp. shouldn't crack a beaker

2) the temp after addition will be 92,727*C, that means a drop of 7,273*C, this shouldn't break any glass.

[Edited on 16-1-2014 by Pyro]

Zyklon-A - 16-1-2014 at 13:55

It's almost brand new, I will allow it to evaporate completely before use. It smells fine, I guess.
Edit: It's good quality, premium Zippo brand, it says very clean burning, and from what I understand benzene produces some smoke when burned,

[Edited on 16-1-2014 by Zyklonb]

Mailinmypocket - 16-1-2014 at 13:58

I'm pretty sure lighter fluid is naphtha, and it should be fairly pure. Don't quote me on it, perhaps check an msds? As for a test for benzene, it would likely require somewhat advanced equipement since it likely would be present (if at all) in small concentrations. What are you trying to extract? Perhaps we could suggest something besides lighter fluid...

alexleyenda - 16-1-2014 at 13:59

Quote: Originally posted by elementcollector1

1) You can't concentrate H2O2 like that, it just decomposes.
2) You don't pour a cold liquid into a hot beaker. Period. Even regular ones can't take it.

1) false, I already made it many times and made density test for concentrations and it almost did not decompose at all from 3% to 30% according to its density. I also tested it's reaction and it was muuuch more strong, showing again that it works quite well.

2) Alright, then I guess I have to do it slowly with a separatory funnel right?

Zyrk : Then I guess I'll pay the extra cost for pyrex for my experiments that deals with heat.

[Edited on 16-1-2014 by alexleyenda]

Mailinmypocket - 16-1-2014 at 14:13

1- let it cool between boilings, top it off, repeat.
2- use two beakers, and swap them, when one is boiled down let it cool and bottle it, add your next beaker of fresh peroxide, boil, swap... Etc etc

Just don't pour cold peroxide into a hot beaker. Simple.

Zyklon-A - 16-1-2014 at 14:20

@mailinmypocket, I'm helping my brother extract LSA (not LSD,) from from Hawaiian Baby Woodrose Seeds, in case you're wondering, it's completely legal.
Also it's for my brother.
@alexleyenda, maybe you can concentrate H2O2 by boiling a little, but you'll lose a lot, what you should do, is put it in the oven at ~80 to 90C, the decomposition is MUCH slower and yeilds will be MUCH better. It will just take a little longer.

[Edited on 16-1-2014 by Zyklonb]

alexleyenda - 16-1-2014 at 14:30

Quote: Originally posted by Zyklonb

@mailinmypocket, I'm helping my brother extract LSA (not LSD,) from from Hawaiian Baby Woodrose Seeds, in case you're wondering, it's completely legal.
Also it's for my brother.
@alexleyenda, maybe you can concentrate H2O2 by boiling a little, but you'll lose a lot, what you should do, is put it in the oven at ~80 to 90C, the decomposition is MUCH slower and yeilds will be MUCH better.

I'll try, I already tried keeping it around 80°C and bublling air through it, it was really slower and I did not see much difference in the final product. I should do a density comparison, but by boiling 90% of a 3% H2O2 solution I got 30% H2O2 according to my density test with 4 numbers. Either you overestimate the decomposition, either I made a mistake, either the 3% peroxide is in reallity a bit more concentrated to compensate for the decomposition.

[Edited on 16-1-2014 by alexleyenda]

Mailinmypocket - 16-1-2014 at 14:32

Hmm I don't really care if it's legal or who it's for, was just a question. Assuming you are going to dry what you extract, given how volatile benzene is you could simply let it evaporate off along with the naphtha IF there even was benzene present. Otherwise buy something from a supplier like elemental scientific such as hexane.

[Edited on 16-1-2014 by Mailinmypocket]

Zyklon-A - 16-1-2014 at 14:36

Ok, thanks.

thebean - 16-1-2014 at 16:52

Am I insane for considering cleaning my matchbox red phosphorus with aqua regia? It sounds a little nuts to me but the thought occurred when I remembered spilling a small drop of aqua regia on paper and watching it be destroyed. I haven't tried it but I'm worried about losses of RP or potentially dangerous redox reactions occurring.

Zyklon-A - 16-1-2014 at 16:56

What are you hoping the aqua regia will do?

alexleyenda - 16-1-2014 at 17:10

Quote: Originally posted by Zyklonb

What are you hoping the aqua regia will do?

I guess he wants it to dissolve the paper and leave the phosphorus behind.

Zyklon-A - 16-1-2014 at 17:18

Ok, that makes sense, I guess, I don't know if it will work though.

A quick google search: 'does aqua regia react with phosphorus' gave no answers.

[Edited on 17-1-2014 by Zyklonb]

thebean - 16-1-2014 at 17:22

I was hoping that the aqua regia would be more efficient than acetone washes, H2SO4 washes, HCl washes and NaOH washes like I normally do.

Zyklon-A - 16-1-2014 at 17:26

Do you know if aqua regia reacts with phosphorus?

alexleyenda - 16-1-2014 at 17:52

Quote: Originally posted by thebean

I haven't tried it but I'm worried about losses of RP or potentially dangerous redox reactions occurring.

You should really read more carefully the posts Zyrk, it's the seconth day in a row I answer you with a quote :p

Mailinmypocket - 16-1-2014 at 18:11

Read up on nitrosyl chloride (formed in aqua regia) and it's chemistry and reactions, with phosphorus if possible. It helps to know what sorts of reactions it participates in with other materials as well, the forum library also has some good inorganic chemistry books that will be helpful. Don't forget about the glue and other stuff in the striker that lead to a mess when combined with concentrated strong acids.

Side note: you are willing to waste precious aqua regia to get a tiny amount of phosphorus?

[Edited on 17-1-2014 by Mailinmypocket]

Zyklon-A - 16-1-2014 at 18:53

I did read it carefully, he said that he didn't try it, not that he didn't know weather it reacts or not, also you didn't really answer my question, which I asked him.
I don't know why you call me Zyrk, there's no 'r' in Zyklonb, nor is it so long that you have to shorten it at all.

Mailinmypocket - 16-1-2014 at 19:16

I should have used the quote function to quote the OP. I was not necessarily responding to you zyclonb, sorry for the misunderstanding.

Zyklon-A - 16-1-2014 at 20:02

Quote: Originally posted by Mailinmypocket

I should have used the quote function to quote the OP. I was not necessarily responding to you zyclonb, sorry for the misunderstanding.

Lol, I was responding to Alex, not you, I should have used the quote, this is confusing, anyway no problem.

[Edited on 17-1-2014 by Zyklonb]

thebean - 17-1-2014 at 11:29

Quote: Originally posted by Mailinmypocket

Side note: you are willing to waste precious aqua regia to get a tiny amount of phosphorus?

[Edited on 17-1-2014 by Mailinmypocket]

I would be using my poor man's aqua regia made from HCl and nitrate salts so it wouldn't be a big deal to me.

DubaiAmateurRocketry - 17-1-2014 at 13:38

Does Acetyl Isocyanate(AcNCO) exist ? or possible to synthesize ? such as (AcCl + NaNCO > NaCl + AcNCO) ?

Brain&Force - 17-1-2014 at 17:43

Acetyl isocyanate does exist, and there's a supplier for it.
http://www.chemspider.com/Chemical-Structure.10489205.html

DubaiAmateurRocketry - 18-1-2014 at 02:21

Where is the supplier ? my search engine is filled with trichloroacetyl isocyanate..

alexleyenda - 18-1-2014 at 21:18

I know that black rubber stoppers do not resist to nitric acid and I was wondering what else I should avoid using with them, I could not find informations about it. Same question for cork stoppers, what chemicals should be avoided with cork stoppers? Thank you

[Edited on 19-1-2014 by alexleyenda]

Organikum - 22-1-2014 at 06:42

Quote: Originally posted by alexleyenda

I know that black rubber stoppers do not resist to nitric acid and I was wondering what else I should avoid using with them, I could not find informations about it. Same question for cork stoppers, what chemicals should be avoided with cork stoppers? Thank you

[Edited on 19-1-2014 by alexleyenda]

Acids burn cork stoppers.

Cork stoppers wrapped well in PTFE tape, not the cheapest thinnest kind but 0,2mm min. thickness and wider is better too, those corks withstand about everything and can in case just be re-wrapped.
If using stopppers with holes the tape has to be wound through this to make it perfect.

My preferred connector actually, the elasticity is often advantageous too.

My question is thouh:

Will a hybrid combination of neodym and electromagnet prevent loss of magnetisation in the neaodym part at higher temperatures and still have the combined power? f course some iron is to be included in this, say and iron pot with the neodym in it and a coil around the pot.

thx
/ORG

alexleyenda - 22-1-2014 at 11:36

Thank you, and interesting question I cannot answer

thebean - 22-1-2014 at 13:21

How do you guys purify your vitamins and what not from store bought pills? In specific I want to know about inositol, PABA, and nicotinic acid (niacin).

confused - 22-1-2014 at 16:18

Sublimnation of PABA (340°C)

liquid-liquid extraction of nicotinic acid using DMSO?
(http://lxsrv7.oru.edu/~alang/onsc/solubility/allsolvents.php...)

inositol could also be extracted using liquid-liquid extraction
(not a real reference, but http://www.researchgate.net/post/Does_anyone_know_about_solu...)

the purification technique really depends on what you're using it for

thebean - 22-1-2014 at 17:04

The PABA sublimation will come in handy but liquid-liquid extraction isn't useful seeing as I'm removing insoluble binders from the pills. Recrystallization from water solution would work fine but whenever I do it with inositol I seem to destroy the inositol, so I'm nervous about destroying what little inositol and nicotinic acid I have left.

Question

Sublimatus - 23-1-2014 at 12:56

Does anyone know of any preparations of solid hydronium salts or a discussion of them in literature?

Wikipedia mentions they do indeed exist, but only gives a few odd and exotic examples.

[Edited on 1/23/2014 by Sublimatus]

UnintentionalChaos - 23-1-2014 at 13:31

Quote: Originally posted by Sublimatus

Does anyone know of any preparations of solid hydronium salts or a discussion of them in literature?

Wikipedia mentions they do indeed exist, but only gives a few odd and exotic examples.

There are quite a few solid strong acid monohydrates that are in fact solid hydronium salts. For example p-toluenesulfonic acid monohydrate is hydronium p-toluenesulfonate. Indicative of it being a salt is that the monohydrate's melting point is considerably higher than the anhydrous acid. Perchloric acid monohydrate is similar.

Sublimatus - 23-1-2014 at 13:39

Thanks, UC.

I've actually made p-toluenesulfonic acid monohydrate, but never considered it from that perspective.

If anyone can point me to preparations for similar salts, I'd be excited to read them.

DraconicAcid - 23-1-2014 at 13:40

Most hydronium salts are probably thought of as acid monohydrates, whatever their structure. If you look at the phase diagram of nitric acid/water mixtures:

You will see the bump marked "NAM" (nitric acid monohydrate). This is a solid that melts at about 238 K (-35 oC), and I would assume that there is extensive hydrogen bonding between the hydronium ion and the nitrate ion (I expect that it'd be closer to being hydronium nitrate than a water molecule hydrogen bonded to nitric acid).

Actually, this paper http://peer.ccsd.cnrs.fr/docs/00/71/66/36/PDF/PEER_stage2_10... references an X-ray crystal structure, and says that it is indeed made of hydronium and nitrate ions.

ETA: Ninja'd by UC!

[Edited on 23-1-2014 by DraconicAcid]

Brain&Force - 23-1-2014 at 15:28

If I mix Cl^- and Cu²⁺, the tetrachlorocuprate complex ion is formed. Am I producing tetrachlorocupric acid when I mix HCl and CuCl₂, or is it just a mixture? I'm guessing it depends on the complexation constant for the ion.

Similarly, are there isolable salts of ABC_x, where A is a alkali/alkaline earth metal, B is a first-row transition metal and C is a halide? (a hypothetical example being potassium tetrachloronickelate)

Continuing the nitric acid questions, does orthonitric acid (H₃NO₄) exist?

DraconicAcid - 23-1-2014 at 16:40

Quote: Originally posted by Brain&Force

If I mix Cl^- and Cu²⁺, the tetrachlorocuprate complex ion is formed. Am I producing tetrachlorocupric acid when I mix HCl and CuCl₂, or is it just a mixture? I'm guessing it depends on the complexation constant for the ion.

Yes. It is an aqueous mixture of hydronium ions, and ions such as tetrachlorocuprate, so it's referred to as tetrachlorocupric acid.

Quote:

Similarly, are there isolable salts of ABC_x, where A is a alkali/alkaline earth metal, B is a first-row transition metal and C is a halide? (a hypothetical example being potassium tetrachloronickelate)

Yes- I've made K₃FeF₆, and I've seen preps for KCuCl₃ ( http://infohost.nmt.edu/~jaltig/CuCmpds.pdf ). I've also made (NEt₄)₂NiCl₄; the potassium salt would probably also be isolable, though not as stable.

Quote:

Continuing the nitric acid questions, does orthonitric acid (H₃NO₄) exist?

Mmm...don't think so.

[Edited on 24-1-2014 by DraconicAcid]

DraconicAcid - 24-1-2014 at 11:14

Quote: Originally posted by DraconicAcid

Yes- I've made K₃FeF₆, and I've seen preps for KCuCl₃ ( http://infohost.nmt.edu/~jaltig/CuCmpds.pdf ).

I just tried this prep. I did not get "red needles" or "ruby-red crystals", as the compound is described as. I got a deep green solution, and a lighter green precipitate. Bah.

TheChemiKid - 24-1-2014 at 14:49

I have tried to make nitrocellulose many times, but it has failed every time. I used this method, but substituted sodium nitrate instead of the potassium nitrate. I correctly scaled down stoichiometrically. Every time I have tries this the cellulose gets attacked by the sulfuric acid and becomes a mush. Any pieces I could salvage were not at all flammable.
Can someone please help?

thebean - 24-1-2014 at 16:23

Does anyone know the products of paracetamol and sodium hydroxide other than 4 aminophenol? If I write the reaction out it is missing only sodium atoms in the products. I think it might be a salt of a carboxylic acid or possibly sodium nitrate.

DraconicAcid - 24-1-2014 at 17:00

Quote: Originally posted by thebean

Does anyone know the products of paracetamol and sodium hydroxide other than 4 aminophenol? If I write the reaction out it is missing only sodium atoms in the products. I think it might be a salt of a carboxylic acid or possibly sodium nitrate.

4-aminophenol and sodium acetate.

HOC₆H₄NHC(O)CH₃ + NaOH -> HOC₆H₄NH₂ + NaO₂CCH₃

[Edited on 25-1-2014 by DraconicAcid]

TheChemiKid - 24-1-2014 at 17:44

Can a more detailed writeup of this equation be posted? What solvent should be used?

UnintentionalChaos - 24-1-2014 at 19:19

Quote: Originally posted by TheChemiKid

Can a more detailed writeup of this equation be posted? What solvent should be used?

For hydrolysis of paracetamol? Use HCl instead. The reaction mixture turns dark from oxidized crap from the free aniline and deprotonated phenol.

If it's anything like acetanilide, it needs very little heating to split the amide group in ~20% HCl. Bring it up to reflux and hold for maybe 10 minutes. There should be no solids by the time you reach reflux.

Cool, and neutralize with base to precipitate p-aminophenol It can probably be recrystallized easily from hot, distilled water.

alexleyenda - 24-1-2014 at 20:57

The names he gave you before the equation is the more detailed way to describe an organic molecule as they are too complex to be accurately described in a simple linear equation. First, in case you don't have much knowledge in o-chem, (Something)-R is a shortcut to write organic molecules, the R is any carbon chain. I'll try to explain to you what happened. First, take a look at the molecule's structure http://commons.wikimedia.org/wiki/File:Paracetamol-skeletal.... . According to Draconic, NaOH will attack the bond between the amine (R-NH-R) and the Carbon that does a double bond with the Oxygen. I have not seen amide reactions yet in my o-chem classes, but according to the current knowledge I have, I should be able to describe what happened :

The carbon double bonded with the oxygen have a partial positive charge as the oxygen is very electronegative and attracts the electrons much stronger to itself. A partial positive charge is a perfect weak spot for the negatives OH- ions from NaOH to attack.

The negative charge from OH- forms a bond with the partial positive carbon, and the bond between N and the carbon breaks and the electron goes to the N. In simple, OH takes the place of R-NH, it is a substitution reaction. At this point, you have HOC6H4NH with a negative charge on N (extremely basic but stabilised by resonance with the cycle), you have Na + ions in solution and you have CH3COOH aka acetic acid http://commons.wikimedia.org/wiki/File:Acetic-acid-2D-skelet... . The N- from HOC6H4NH being extremely basic, it will steal the H from the OH on the acetic acid (simple acid-base neutralisation). You get a now neutral HOC6H4NH2, the acetic acid lost a proton and has become CH3COO- and you still have the Na+ ions in solution so of course you get the sodium acetate salt http://commons.wikimedia.org/wiki/File:Sodium-acetate-2D-ske... .

For the solvant, paracetamol is a bit soluble in water, its solubility gets much better at high temp and of course NaOH is very soluble. It should be a good solvant that does not produce undesired side reactions. I can't think of any other solvant. The use of other solvants could cause not wanted side reactions because of the negative R-NH created during the process that attacks pretty much everything : double bonds, triple bonds, alcools, partial positive spots created by links with halogens and ketones.... It also reacts with water, but it does not change anything as it will create OH- that steal the hydrogen from the acetic acid anyways.

That was a good revision of my o-chem class :p I hope it helped.

Edit: Chaos, the formation of a sodium salt with the "phenol part" should not be a problem as if this happen, it means that there will be acetic acid left (as long as stoechiometric amounts are used) that will protonate the basic salt back in "phenol".

[Edited on 25-1-2014 by alexleyenda]

DJF90 - 25-1-2014 at 09:53

Its good that you're trying to apply your knowledge Alexleyenda, but you'll need an excess of sodium hydroxide as one equivalent will deprotonate the phenol first; a second equivalent is needed (though it would be best to use an excess, maybe 5 eq. in total) for the amide hydrolysis. Phenoxide anions tend to be sensitive to oxidation, hence the dark tarry crap formed. For paracetamol at least, hydrochloric acid is an easier alternative.

alexleyenda - 25-1-2014 at 11:18

@DJF90
Right, I did not even take time to think about this, in my head I decided for no reason that the hydrolysis would occure before the deprotonation of the acidic phenol, that was quite stupid! However, I don't understand how the phenoxide would oxidize, could you explain it?

At least I guess I correctly understood the mechanism as you did not say anything about it

[Edited on 25-1-2014 by alexleyenda]

DJF90 - 26-1-2014 at 07:20

Your description of the mechanism was so long I didn't bother to read it. In hindsight I should have because it is wrong.

What you describe is Sn2 displacement of the R-NH(-) species. Due to stereoelectronic effects (essentially overlap of the reacting orbitals) Sn2 cannot occur at an sp2 center (e.g. amide carbonyl carbon, as we are discussing here). What happens is a stepwise mechanism, first attack of the nucleophile at the carbonyl carbon and breaking of the C=O double bond to form a C-O(-) structure, called a "tetrahedral intermediate". The negative charge on the oxygen can then close down to reform the C=O double bond, and the RNH(-) group is displaced.

I beleive this is not quite accurate though, as it doesn't account for the energetic favourability for the displacement of hydroxide vs the amine anion (technically an "amide" but lets not go there...) as difference in pKa and leaving group ability between OH(-) and RNH(-) is significantly large. It is more likely that the tetrahedral intermediate is deprotonated also to form a double anion (both residing on the oxygens) and then elimination of the amine anion occurs, leaving the carboxylate salt. This mechanism would provide more of a driving force for elimination of the RNH(-) vs the OH(-).

[Edited on 26-1-2014 by DJF90]

DubaiAmateurRocketry - 26-1-2014 at 12:18

Does Aluminum amide exist ? Al(NH2)3 ?

HeYBrO - 26-1-2014 at 13:14

Just to clarify, would adding HCl to potassium hydrogen phthalate produce the free acid and KCl?

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