Sciencemadness Discussion Board - The short questions thread (2)

The short questions thread (2)

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not_important - 28-5-2009 at 19:58

Iron stains want complexing and/or reducing conditions to remove. Oxalic acid usually does the job, citric may do it while a mix of oxalic and citric has always worked for me although fairly long treatment with a strong solution of them may be needed.

Oxidisers - chlorine bleach, peroxides - just make the stains work by pushing more iron into the red-brown Fe(III) state. Reducing agents drop it back to the pale green Fe(II), complexing agents and mild acids can remove that much more easily than the Fe(III).

Was the meat well cooked? If so you may experience digestive upsets from the decomposition productions - especially the diamines - but likely nothing worse. Et a lot of active culture yogurt over a number of hours, that can help.

Sedit - 28-5-2009 at 20:14

Thanks I never really considered reducing it I was always looking for a more brute force method. Ill give a mix of Oxalic and citric a try. Soak some smaller items overnight and see what I can get.

Yeh it was cooked very well and simmered over an 45 minutes after that while the water was boiling so bacteria should be no problem but I do feel slightly upset to the stomach nothing major though. Its my grandfather Im more concerned about because hes almost 90 and he loves my cooking so much he ate a plate and a half before we found it to not be so good and after we told him he just shrugged it off and finished eating

. Cooking is like easy organic chemistry but sadly I think I contaminated this reaction and the flask is more then likely gonna want to clean it self out tommorow

not_important - 28-5-2009 at 20:26

Again, yogurt and fresh ginger can help reduce the upset if there's wasn't too large amount of decay products.

Sedit - 28-5-2009 at 20:39

Fresh out of yogurt, with three kids that don't last more then a few hours when I get it although if it persist tommorow I will definitly go out and get some. I don't think it was two bad because I have a pretty good nose for it but I was just really busy with company over this time to pay close enough attention. Just a few minor knots in the belly but nothing major and its been a few hours now but like I said before its the 90 year old im concerned about since he ate alot and at the age of 90 the digestive system isn't the greatest as it is. If I find he has problems tonight I will run out no matter what time to get him something to help. I do have alot of different fresh mint growing which may aid in speeding the digestive system similar to ginger so that may be a quick option if really needed.

Thanks for the help N_I always a pleasure.

User - 29-5-2009 at 02:55

Let's see if i got this right.
I thought of concentrating my nitric acid by using the double volume of sulfuric acid.
I don't have access to a vacuum generator so I am doing without ( i know more decomposition etc).

Well as for my question : would it be useful to use a vigreux collom?

-I think an advantage would be that the product has a much higher concentration.

-The downside would be that there would be more decomposition because of the relative longer exposure to higher temp.

When one takes in account that without it a vigreux a second destilation would be needed.
Maybe even with a second time is needed.
I would go for a vigreux.

Maybe there are factors I am not considering..
Anyone?

[Edited on 29-5-2009 by User]

DJF90 - 29-5-2009 at 07:42

Distilling with a large excess of the sulfuric acid (98%) should give you nitric acid in the 90%+ range (so long as you start with conc. nitric acid - about 70%). Dont use a vigreux as this is more troublesome than helpful. If you need the acid to be more concentrated then use a vacuum. Be careful when you mix the acids (add the sulfuric to the nitric - this is more "water-like") There are plenty of threads about concentrating nitric acid over on E&W, but I'm sure you'll find some here too if you UTFSE.

pHzero - 30-5-2009 at 11:51

My quick question:
What's the [Ag(CN)2]- ion called? It's analagous to the dicyanaurate (I) ion ([Au(CN)2]-), but I don't know what would replace the "aur" part of the name?
dicyanargate (I) maybe?

entropy51 - 30-5-2009 at 13:22

[Ag(CN)2]- is called dicyanargentate(I) according to http://www2.ucdsb.on.ca/tiss/stretton/CHEM2/solubil7.htm

pHzero - 30-5-2009 at 15:44

Quote: Originally posted by entropy51

[Ag(CN)2]- is called dicyanargentate(I) according to http://www2.ucdsb.on.ca/tiss/stretton/CHEM2/solubil7.htm

Ah right, thanks

Just realised how stupid I was asking here though, rather than googling "Ag(CN)2- name" ah well..

manimal - 31-5-2009 at 14:50

I have noted an inability to "salt out" 91% isopropanol. Added salt does not induce formation of layers, even after a few days of standing. Is this because 91% is simply too concentrated to give up additional water by saturation w/ salt?

S.C. Wack - 31-5-2009 at 15:07

http://dx.doi.org/10.1021/ja01343a024

UnintentionalChaos - 31-5-2009 at 16:18

Alright, this is gonna sound pretty stupid coming from me, but I've been having issues recrystallizing sodium acetate trihydrate. It was made from boiled-down (3.8L) vinegar neutralized with baking soda, and by the time it was down to 600mL or so, it was a dark reddish-brown from condensed organic crap in the mix. I had a fine white precipitate seperate from the boiling solution which did not melt under reasonable heating and was somewhat picky about dissolving in more water. Upon further concentration and cooling, the dark reddish-brown supernatant formed masses of fragile clear (but stained with impurities) crystals which are without a doubt the trihydrate. I vacuum filtered the crystals and redissolved both forms in a little distilled water under heating, followed by filtering through cotton wool to remove any crud that drifted into the mix. The resultant solution is pale straw yellow colored. After sitting overnight, I got a compact white crystalline mass forming on the bottom. Agitating it brought a flocculent white precipitate of crystals. There is no sign of the glassy, brittle trihydrate crystals in the new solution. No more material will crystallize out.

Now for my question. Am I having anhydrous sodium acetate (or a lower hydrate of some kind) crystallize out of supercooled trihydrate which refuses to crystallize despite my best efforts? Usually I do very well with crystallizations and this is very frustrating. The reddish filthy swill that I harvested 2 crops from continues to lay down small amounts of trihydrate crystals as it evaporates, which makes it even more irritating.

Sedit - 31-5-2009 at 17:41

I find a wash of the ground sodium acetate with acetone takes great care of the organics left and after a filter im left with a pure white filter cake( very slightly off white until it drys). Have you tryed to seed the supercooled solution? I highly doubt you have anhydrous sodium acetate comming out of the solution.
Im sorry I can't be more help but my main problem when working with this is stoping the supersaturated solutions from crystalizing when Im trying to work with it and cloging filters and such.

Intergalactic_Captain - 6-6-2009 at 08:01

This ain't a chem question, but I don't currently have accounts any more appropriate forums...Feel free to delete if it's objectionable...

Can picric acid be detonated by way of a rifle round? Specifically, half a film canister full, pressed, and hit by a NATO (5.54x39 iirc) bullet? I have some amount that's been sitting around for quite some time and needs to be disposed of in a suitable fashion.

1281371269 - 6-6-2009 at 09:18

It should do, it's shock sensitive - a google search should give you all the information you need. However I worry that by shooting it you might detonate a bit and send the rest flying all over the place, especially if you did a bad shot. You could dilute down to a very very weak solution and flush it away.

[Edited on 6-6-2009 by Mossydie]

Lambda-Eyde - 6-6-2009 at 09:58

Why don't you just use a blasting cap?

[Edited on 6-6-2009 by Lambda-Eyde]

Intergalactic_Captain - 6-6-2009 at 10:12

Don't have any caps and I'm NOT putting it in my septic system... I'm going shooting with a buddy of mine today and figured it's as good a time as any to get rid of it - No loss on an incomplete detonation, I just don't want it on my shelf anymore.

hissingnoise - 6-6-2009 at 11:16

If its purity is high, why not hold on to it---you may kick yourself later if you get rid of it. . .
If you're unsure of the purity, recrystallise!

entropy51 - 6-6-2009 at 13:33

It sounds like you've found a non-chemical solution. Some hazardous materials handbooks recommend reduction to triaminophenol with tin and HCl, and then oxidation of the triaminophenol with acidified permanganate. These references recommend running the reaction on 8 grams or less at a time. Googling brought up some of these references if anybody else is interested.

The_Davster - 6-6-2009 at 14:03

I found a file saying picric dets 5/10 times with a .30 cal bullet.
Perhaps go something more powerful than your .221 like 7.62x39, .308 or 30-06 depending on what is available.

http://publications.drdo.gov.in/gsdl/collect/defences/index/...

If not, Sn/HCl reduction is the common way.

[Edited on 6-6-09 by The_Davster]

manimal - 7-6-2009 at 17:08

Do amino acids form ammonium salts in the presense of ammonia, e.g. ammonium glycinate?

Rich_Insane - 7-6-2009 at 18:15

If I have an Ahlinn flask that is slightly chipped on the narrow end, will it still work? Or should I sell it/throw it away

Formula409 - 7-6-2009 at 19:16

Can someone please write the reaction between a nitrate salt and aluminium powder which reduces the nitrate ion to the ammonium ion? I'm having an amazing mental blank.

Formula409.

sakshaug007 - 7-6-2009 at 19:55

Quote: Originally posted by Formula409

Can someone please write the reaction between a nitrate salt and aluminium powder which reduces the nitrate ion to the ammonium ion? I'm having an amazing mental blank.

Formula409.

I'm not sure exactly, my guess would be:

Al + KNO3 + H2O --> Al2O3 + KOH + NH3 (or NH4OH, in water)

I can't see how it would proceed without the presence of water.

Hope this helps.

UnintentionalChaos - 7-6-2009 at 20:29

Quote: Originally posted by Formula409

Can someone please write the reaction between a nitrate salt and aluminium powder which reduces the nitrate ion to the ammonium ion? I'm having an amazing mental blank.

Formula409.

Is this in solution, or do you mean upon acidification of the remains of a pyrotechnic mix?

If the latter,

3NaNO3 + 8Al -> 3NaAlO2 + Al2O3 + 3AlN

The aluminum nitride should hydrolyze in water, releasing ammonia, or in acid, being immediately trapped as ammonium.

Quote: Originally posted by Rich_Insane

If I have an Ahlinn flask that is slightly chipped on the narrow end, will it still work? Or should I sell it/throw it away

What in the hell, may I ask, is an Ahlinn flask?

An allihn condenser? Perhaps a Kjeldahl flask? Slightly chipped can mean a lot of things depending on who you ask. Take a picture of the damage.

[Edited on 6-8-09 by UnintentionalChaos]

Panache - 9-6-2009 at 18:30

Does anyone have any actual real experience with the shelf stability of whinchesters of analR conc nitric acid? I picked some up for a song but its circa late 90's (19 that is not 18, lol). The bottles are all unopened, i opened one and its essentially colourless with the density matching the range quoted on the label. Vogel 3rd edition suggests running N2 through nitric to flush out the NO, No2's etc but i imagine in that decade the nitric was far less pure than now.
I also got two whinchesters of analR n-hexane, anyone have suggestions for something interesting to do with this very pure paraffin?
Thirdly i also got a small amount (20L) of (manganese) parkerizing concenetrate. Looking into it it appears as this may be the perfect solution to my iron retort bases constantly rusting up. Anyone parkerized anything before and subsequently used it in the laboratory? Tips? I'm fairly sure its managanese base, but the MSDS says it contains nickel salts which is odd?

UnintentionalChaos - 9-6-2009 at 21:07

Why would nitric acid be that much less pure just 12 or so years ago? People have been making it for centuries, and lets face it, no matter what, it's better than anything you distill yourself. The biggest difference between then and now is probably that nitric acid for trace metals analysis is now available in PFA and FEP bottles which leach next to nothing (even compared to glass) into the acid.

if it's colorless and the bottle was sealed, it's as good as brand new. It has been properly stored for it's lifetime.

[Edited on 6-10-09 by UnintentionalChaos]

Rich_Insane - 9-6-2009 at 21:33

Quote:

What in the hell, may I ask, is an Ahlinn flask?

I'm sorry, I meant to say condenser (long night

). Just a bit on the narrow end is chipped.

I'm wondering if it can be used , because that's where you would either put a desiccator, a thermometer or something, and there's usually adapters for those.

Unfortunately the camera's broken (it's all red and fuzzy, probably the lens). By Chipped I meant rather small, perhaps a 1.5 cm2 piece of glass, or 2. The end isn't completely off, and the damage only goes for about 1 cm.

[Edited on 10-6-2009 by Rich_Insane]

[Edited on 10-6-2009 by Rich_Insane]

UnintentionalChaos - 9-6-2009 at 21:52

Narrow end? Is this a condenser with ground glass joints and the male cone is broken? If you can get a strip of grease all the way around the remainder of the joint, it can still be used, but it may eventually break from strain on the preexisting damage.

kclo4 - 9-6-2009 at 22:43

How much do you think this fume hood is worth?
http://www.thevespiary.org/0601091238.jpg

I'll try to get measurements later if its needed, but it is obviously only meant to be used by one person at a time. Maybe 40" wide?
It has an attachment for water, a small cone shaped "sink" below the water inlet. The cone sink my be 250ml? Also has an attachment for some sort of gas such as natural gas or propane.
The fume hood also has a 24" fluorescent light above the glass on the entrance side.

It has very little rust, works perfectly, the hose has a hole in it, but will be fixed soon I'd think. The glass on the back comes up, and the shield in front opens up.

Oh also, wtf does this go to, or how does it work? I'm sure it is a water distillation set up for the lab that is incomplete. It plugs in, but it doesn't seem to do anything since there is no place to heat, etc. I haven't plugged it in yet though.
http://www.thevespiary.org/0601091523.jpg

hissingnoise - 10-6-2009 at 04:29

Panache, UC's reply to your query contained no insult, real, veiled or implied. . .
You, though, do seem to have some kind of problem.
And you owe UC an apology. . .

Panache - 10-6-2009 at 04:52

hmmmm,

perhaps i over-reacted.

UC, i apologise whole-heartedly if your response was meant in an inquisitive manner, if it was meant in a snide manner, i stand by what i accuse you of. Please respond as to whether it was the former or the latter.

Thank-you hissinghoise for bringing this up.

[Edited on 10-6-2009 by Panache]

hissingnoise - 10-6-2009 at 05:01

We all have "our moments", Panache, and you've done the right thing. . .

einstein(not) - 11-6-2009 at 09:38

I have approx. 5lbs. of calcium sulfate (perhaps calcium hydrogen sulfate) left over from the production of nitric acid via calcium nitrate (fertilizer grade) and sulfuric acid (drain opener). What do I actually have and what would be the best way to purify it for later use?

User - 11-6-2009 at 09:58

Well if calcium(hydrogen)sulfate is strongly heated one can obtain SO3.
First calciumpyrosulfate will appear due to the loss of water.
When heated further ->500deg pure SO3 will come over.
Heating it very strongly would be a start cleaning it.

Also , when a solid nitrate is added and heated nitric vapors appear. They would be heavily decomposed under influence of the heat required.

[Edited on 11-6-2009 by User]

kclo4 - 11-6-2009 at 10:04

Solubility, and pH will determine if it is calcium hydrogen sulfate or calcium sulfate.

I personally think if it is calcium hydrogen sulfate you can add it to water to get dilute sulfuric acid as the calcium sulfate precipitates. If you add calcium sulfate to sulfuric acid/water it doesn't dissolve into calcium hydrogen sulfate the last time I checked. If it does, I am missing out!

So what I would do to purify it is just rinse it with water until it has a neutral pH. Calcium sulfate is hardly soluble in water but it can be very annoying and remain in the solution as small particles for a long time.

If I were doing this and were serious about yield I would boil/reflux the solution + calcium sulfate to get larger CaSO4 crystals. It would help break the suspension and make it easier to filter. Attempting to filter CaSO4 is often a mistake, and I would decant off the solution above it then rinse with distilled water to get rid of any acids or soluble salts, etc.

einstein(not) - 11-6-2009 at 13:33

I'm having trouble with the reaction equation. Ca(NO3)2 + 2H2SO4 -----> 2CaSO4 + 2HNO3 + H I don't beleive any hydrogen was evolved so obviously I've got something wrong. Right?!?

[Edited on 11-6-2009 by einstein(not)]

hissingnoise - 11-6-2009 at 14:19

That should be: Ca(NO3)2 + H2SO4--->CaSO4 + 2HNO3. . .
Yay! Your H2SO4 is going twice as far!

einstein(not) - 11-6-2009 at 15:27

Quote: Originally posted by hissingnoise

That should be: Ca(NO3)2 + H2SO4--->CaSO4 + 2HNO3. . .
Yay! Your H2SO4 is going twice as far!

Thanks!

UnintentionalChaos - 11-6-2009 at 20:19

Calcium sulfate, as it turns out, is appreciably (as appreciable as a typically "completely insoluble" powder can be) soluble in nitric acid solutions. I learned this the hard way in one of those "identify this white powder" labs in inorganic chemistry. "Oh, a little spatula dissolved in nitric acid, so it can't be CaSO4" This of course was before I knew what Ksp was. Distillation is in order to get clean acid.

[Edited on 6-12-09 by UnintentionalChaos]

Formula409 - 13-6-2009 at 03:53

Can Methyl Tosylate be made by simple Fischer Esterification of p-Toluenesulfonic acid?

Formula409.

hissingnoise - 13-6-2009 at 12:22

So. . .Thinking I hadda way of making hexamine dinitrate sans HNO3, I dissolved a small quantity of dry NH4NO3 in a little fresh formalin.
I expected the formalin smell to fade as the NH4NO3 reacted but after adding more than enough to bind up the CH2O the choking smell was undiminished. . .
I decided to evaporate the solution to dryness; the smell of hot aldehyde reached my kitchen door and my ball-and-cha. . .er, my good wife picked up on it.
At around 100*C the CH2O smell faded somewhat, being replaced largely, by the smell of NO2.
The residue which formed was white at first, slowly becoming orange-coloured and flatly refused to dry, so it was heated strongly again and the NO2 smell began to take on a fishy character.
Now I've got a wet, orangey-amber gloop which smells alternately of HNO3 and CH3NH2.
I dunno, the methanol stabiliser might have complicated things.
The original idea was to nitrolyse the dried HDN to RDX, thereby saving on HNO3.
Has anyone else tried doing this?
Or have any thoughts on the reaction. . .

UnintentionalChaos - 13-6-2009 at 16:33

Does anyone know a cheap supplier or happen to have a spare PTFE stopcock I can buy off of them?

I need a 1:5 taper 15.2/30 with a single 3 or 4mm straight bore. I could also use a single PTFE washer for a 11/25 stopcock, although I can deal without it. Thanks

hissingnoise-

the reaction of formaldehyde with ammonia to produce hexamine probably proceeds via a methylimine intermediate, which requires nucleophilic ammonia to react with the formaldehyde. Ammonium ion is not a nucleophile and ammonium nitrate solution is mildly acidic, keeping the concentration of free ammonia very low.

The gunk left over probably contains a lot of paraformaldehyde and ammonium nitrate. Small amounts of condensation probably occured despite the acidity, which is where the methylamine smell is coming from. perhaps methylimine copolymerized with the formaldehyde? I'm guessing wildly here.

[Edited on 6-14-09 by UnintentionalChaos]

not_important - 13-6-2009 at 17:49

Remember the old standard method for making the methylamines is to heat a mixture of formaldehyde and ammonium chloride, I suspect that the same reaction is occurring.

Do consider that if secondary amines are formed, then it is likely that nitrosoamines will also be formed.

hissingnoise - 14-6-2009 at 06:29

Thanks guys. Hexamine dinitrate is likely to be thermally unstable so heating the solution was a mistake, it seems. . .
But the ease with which hexamine is obtained from 4NH3 + 6CH2O threw me off.
Evaporation under (greatly) reduced pressure should give the dinitrate but I'm having some work done on my house and the plumbing is in a state of limbo, right now.
In theory, though, solutions of NH4NO3 in formalin should contain hexamine and enough HNO3 to form RDX.
Adding excess acetic anhydride should precipitate RDX since that is the synthesis that replaced direct nitration. . .
Yes, I know paraform is used in practise but the principle is the same.

If I tried dessicating the solution over H2SO4 would H2O be removed without affecting the HNO3 also present, does anyone know?

[Edited on 14-6-2009 by hissingnoise]

User - 14-6-2009 at 07:56

For the record , why not make HDN out of HCl, AN and hex. ?
It does't give good yields, but kilo's could be made for just a couple a bucks.
It seems cheap enough to me.

hissingnoise - 14-6-2009 at 08:20

Two reasons User,---I'm low on HCl and hexamine and I have a surplus of AN and formalin, and the formalin route has a hint of novelty about it. . .
And the reaction of formalin with ammonium salts is supposed to release the acid!
It's also possible that at low pressure, (dilute) HNO3 could be distilled from the solution.
But it's all guesswork. . .

Rich_Insane - 17-6-2009 at 12:05

Would you dissolve menthol with a polar or nonpolar solvent?

I have figured out that the hydroxyl group in menthol makes it polar on that side. The solubility in water is listed as "slightly soluble".

[Edited on 17-6-2009 by Rich_Insane]

497 - 17-6-2009 at 22:46

Two questions that shouldn't be too hard to answer:

1. In general, will methods of reductive amination for ketones also apply to aldehydes with similar yields? Are there any general differences?

2. Is there an appreciable difference between oxidation of benzyl alcohols and 2-phenethyl alcohols? For example, a catalytic H2O2 oxidation of benzyl alcohols yields 95% benzaldehyde, will it be similar for 2-phenethyl alcohol to phenylacetaldehyde?

Edit: I think I've answered #1 for myself, so cancel that one..

[Edited on 19-6-2009 by 497]

UnintentionalChaos - 18-6-2009 at 20:28

Quote: Originally posted by 497

2. Is there an appreciable difference between oxidation of benzyl alcohols and 2-phenethyl alcohols? For example, a catalytic H2O2 oxidation of benzyl alcohols yields 95% benzaldehyde, will it be similar for 2-phenethyl alcohol to phenylacetaldehyde?

The oxidation of 2-phenylethanol will proceed much like the oxidation of ethanol. The benzylic position confers special reactivity to benzyl alcohols, halides, etc due to the benzyl carbon's relation with the pi-cloud.

[Edited on 6-19-09 by UnintentionalChaos]

Formula409 - 19-6-2009 at 17:14

Imines form water-soluble hydrochlorides, don't they?

Formula409.
Edit: Yes they can : http://www.journalarchive.jst.go.jp/english/jnlabstract_en.p...

[Edited on 20-6-2009 by Formula409]

Sedit - 21-6-2009 at 18:09

Organic Hypochlorites like EtOCl are effective chlorinating agents for the conversion of toluene to Benzyl Chloride.

They are synthesised in a cold dim enviroment to prevent explosion risk from R-OH + NaOH + Cl2 = R-OCl + NaCl + H2O

My question is could this possibly be used insitu to chlorinate Toluene or is purification a must? My logic is the faster it is used for chlorination and turned back into EtOH the less one risk the explosion from it. The main problem I see here is the water formed interfering with the conversion of BnCl by Hydrolysis to HCl + BnOH. Drying with MgSO4 I guess would be possible but would also take a large amount of drying agent I would think.

Even if BnOH was the main product this still would not pose a problem since BnO is the final target compound to begin with. I know of all the other ways to get there so lets not make it a discussion on that please. Just curious about this reaction in general.

UnintentionalChaos - 21-6-2009 at 19:49

Quote: Originally posted by Sedit

Organic Hypochlorites like EtOCl are effective chlorinating agents for the conversion of toluene to Benzyl Chloride.

They are synthesised in a cold dim enviroment to prevent explosion risk from R-OH + NaOH + Cl2 = R-OCl + NaCl + H2O

My question is could this possibly be used insitu to chlorinate Toluene or is purification a must? My logic is the faster it is used for chlorination and turned back into EtOH the less one risk the explosion from it. The main problem I see here is the water formed interfering with the conversion of BnCl by Hydrolysis to HCl + BnOH. Drying with MgSO4 I guess would be possible but would also take a large amount of drying agent I would think.

Even if BnOH was the main product this still would not pose a problem since BnO is the final target compound to begin with. I know of all the other ways to get there so lets not make it a discussion on that please. Just curious about this reaction in general.

consider perhaps tert-butyl hypochlorite which is more stable and somewhat less suicidal to work with.

do you have a reference for the conversion of toluene to benzyl chloride?

Sedit - 21-6-2009 at 20:38

Yes, It was posted by SC Wack some time back in the organic hypochlorite threed. Its the first post on the second page.

http://www.sciencemadness.org/talk/viewthread.php?tid=1896&a...

Link to the file: http://www.sciencemadness.org/talk/files.php?pid=47308&a...

Suicidal don't even appear to be the word for these compounds and its mostly just for academia at the moment.

Im curious of BnCl or more then likely BnOH could be formed directly after the EtOCl or which ever organic hypochlorite thats used does. Excess NaOH could be used to neutralize the formed HCl I would think and more then lead to Benzyl Alcohol. But since BnOH in concentrated acid is suppose to give BnCl also im a little lost as to the product of performing this in one pot if its at all possible.

[Edited on 22-6-2009 by Sedit]

The_Davster - 21-6-2009 at 20:39

I looked for a hypochlorite chlorination of toluene (only google) and my cursory results only turned up ring chlorination (attachment)
EDIT: Ah, excellent I knew I had seen the sidechain chlorination reaction somewhere before.

However ethyl hypochlorite is not too bad, at least Axt and myself have made it without incident. Smells horrible though, perhaps tBu would stink less with its lower vapour pressure. Of course, tBu also offsets the explosion risk

Attachment: ref1 chloro.pdf (599kB)
This file has been downloaded 816 times

[Edited on 22-6-09 by The_Davster]

Sedit - 21-6-2009 at 20:42

Quote from the paper

Quote:

Results and Discussion
Reaction with Toluene.
-Toluene was chosen as
an initial substrate for our study since the occurrence
of side chain or ring substitution of halogen
is in general diagnostic for radical or polar processes.
Preliminary experiments at 40' using excew
(3: 1) toluene and azobisisobutyronitrile (AIBN)
or light as initiator showed a rapid reaction after a
variable induction period, complete consumption
of hypochlorite, and the formation of benzyl chloride
and t-butyl alcohol as major products. A more
detailed study of the effect of reaction variables on
rate gave the results appearing in Table I, while

sparkgap - 22-6-2009 at 06:07

I may just not be looking hard enough, but are there any sites like SDBS that have two-dimensional spectra (e.g. COSY, NOESY, HSQC, HMBC)?

sparky (~_~)

jokull - 22-6-2009 at 12:58

Dear sparky,

I don't know such a kind of database, but I do know there is a 2D NMR predictor from ACDLabs, it also have the capability of being tuned to improve the predictions.

I don't rememeber if this module is included within the acdlabs 11 suite posted somewhere in the forum, but I hope so.

Your question pulled out my curiosity...I found this site:

http://www.ebi.ac.uk/nmrshiftdb/

maybe it would be helpful.

[Edited on 22-6-2009 by jokull]

Sedit - 23-6-2009 at 09:47

Can an enamine reduction using formic acid take place in a highly acidic enviroment of 60% H2SO4?

I see nothing stoping it but then again if I was a good chemist I wouldn't be here asking you folks

Attachment: Formic Reduction.pdf (331kB)
This file has been downloaded 792 times

Jor - 23-6-2009 at 14:45

IIRC sulfuric acid of such concentration would decompose the formic acid to carbon monoxide.

Sedit - 23-6-2009 at 14:47

Ahh yes good point.

Saber - 24-6-2009 at 03:40

What is the best type of UV light for radical chlroinations?
I understand these work with sunlight however is the UV from blacklights or from those water sterilising tubes better?
thanks,

not_important - 24-6-2009 at 07:53

This has been covered in the threads that discuss making the alpha-chloro-tolulenes. Light in the blue to near UV is the best, for a variety of reasons. A non-coated (clear bulb) mercury vapor lamp for outdoor lighting applications gives both good wavelengths and high intensity, and is much more cost effective than the two choices you offered.

Aubrey - 24-6-2009 at 12:13

Has anyone synthed benzonitrile from benzoic acid and sulphamic acid before? I've not seen this preparation described anywhere and am interested in learning a little more about it

kclo4 - 25-6-2009 at 07:15

What is wrong with sciencemadness? I haven't been able to post anything lately and it takes me too the main screen.. has anyone else had that problem?

[Edited on 25-6-2009 by kclo4]

not_important - 25-6-2009 at 08:06

KCLO4 - I had that problem when I tried to post a reply which included a equation in which I'd made a double headed arrow <=&GT; (lesser than)(equals)(greater than)

I'm pretty sure the < > combination is behind the problem, I had to use the HTML code "& LT ;" "& GT ;" (without the internal spaces) to get them in this message.

kclo4 - 25-6-2009 at 08:16

That must be something with HTML or something, I was posting a link to a google search, and a link to another sciencemadness thread.

Hope the problem can be fixed soon, but seems like it might be hard to diagnose. I noticed it late last night/early morning.

Saber - 25-6-2009 at 09:39

Has anybody here attempted the monoethanolamine synthesis?
I understand it is made by reacting ethylene oxide (ethene + O2/Ag) with aqueous ammonia, producing a mix of the amines.
I have also read however monoethanolamine reacts with aqueous ammonia to produce diethylamine… Is there any way I can stop this reaction, or how long does it take?
I understand to produce mainly monoethylamine I must use an excess of ethene oxide…

As for the intended usage, I plan on making some taurine

12AX7 - 25-6-2009 at 14:18

The classic monoalkylamine synthesis involves phthalimide as a protecting group. Because the two carbonyls polarize it, the N-H is fairly acidic and the H can be abstracted with potassium carbonate or hydroxide. An electrophile, like ethyl chloride, can substitute easily. The phthalimide is then hydrolyzed in acid, leaving the monoalkylamine salt.

Not sure if this will work with your particular alkyl. I would suppose the hydroxyl of 2-chloroethanol would sit still, it's less electrophilic. Not sure how you get it either, add HCl to ethylene oxide maybe? Also sounds hazardous or noxious...

Tim

sparkgap - 25-6-2009 at 17:16

There's also the variation that replaces phthalimide in the Gabriel synthesis with saccharin; go look for the refs yourself, but I do recall that the saccharin has to be oven-dried for the reaction to work. Otherwise, there's also Delepine (which uses hexamine).

sparky (~_~)

Saber - 27-6-2009 at 08:38

Sorry just figured out my post would not work
I will do a bit more research

[Edited on 27-6-2009 by Saber]

DJF90 - 27-6-2009 at 10:39

Wikipedia says that the synthesis of taurine starts with 2-hydroxyethanesulfonic acid, obtained from the reaction of ethylene oxide with sodium bisulfite. From the intermediate compound you might try alcohol => halide (appel reaction? or Sn2 with HI or possibly HCl/ZnCl...) followed by gabriel or delepine reaction, or you could go alcohol => aldehyde (PCC, oppenauer oxidation, TEMPO, DMP etc...) and then reductive amination (leuckart, NH4+ HCO2-?) Or alcohol => acid (jones reagent) and then a schmidt reaction (H2SO4/NaN3). Personally the first idea looks most attractive (I dont think the appel reaction (Ph3P/CCl4) or gabriel synthesis will cause problems with the sulfonic acid group), although you might have trouble obtaining the reagents.

Sedit - 1-7-2009 at 18:17

Once again I need the wisdom of you folks here.

I was distilling a small amount of bromoethane maybe about 25-35ml and I allowed it to distill to dryness at which point there was a pop and went over to find the bottom of the flask went out.

I smelled no smell but within couple seconds (10 tops)I could taste a funny taste in my mouth simular to when a funky perfume or something gets sprayed to close possibly bug spray. It is semi indoors without a fume hood but its well ventilated(few feet from a big wide open garadge door) so I turned it off and am letting it air out and got the hell out of there.

I know Bromoethane is no good for you...

How concerened should I be long term...short term.... over a small exposure.

I expected this to have a smell of some kind and yet there was nothing but the taste is obvious and not placebo effect.

crazyboy - 1-7-2009 at 18:29

"8.1 Single exposure
8.1.1 Inhalation

In a single exposure study, groups of five mice and five rats of each sex were exposed to bromoethane at concentrations of 0, 2800, 5700, 11 000, 23 000, or 45 000 mg/m3 (0, 625, 1250, 2500, 5000, or 10 000 ppm) for 4 h (Roycroft, 1989). Mortalities were observed at 5700 mg/m3 (1250 ppm) or more in mice and only at 45 000 mg/m3 (10 000 ppm) in rats. At 45 000 mg/m3 (10 000 ppm), clinical signs of toxicity prior to death included dyspnoea, hyperactivity, and incoordination. No further information was available. LC50 values determined from these data were 21 200 mg/m3 (4681 ppm) for rats and 12 300 mg/m3 (2723 ppm) for mice.

As a comparison, 1-h LC50 values were calculated as 122 000 mg/m3 (26 986 ppm) for rats and 73 500 mg/m3 (16 230 ppm) for mice (MacEwen & Vernot, 1972). The primary response seen in both rats and mice was central nervous system (CNS) depression, although no further details were given. The authors reported seeing diarrhoea in rats exposed to the highest concentration, 180 000 mg/m3 (40 000 ppm), and in mice exposed to concentrations ranging from the highest of 90 000 mg/m3 (20 000 ppm) down to 57 000 mg/m3 (12 600 ppm). (The lowest concentrations used were 90 000 mg/m3 [20 000 ppm] for rats and 45 000 mg/m3 [10 000 ppm] for mice.)

Other LC50 values quoted are 53 000 mg/m3 (11 700 ppm) for rats and 36 000 mg/m3 (7950 ppm) for mice following 2-h exposures, with "general damage to the nervous system" occurring (Izmerov et al., 1977).

Early, generally poorly reported, studies in mice, guinea-pigs, and cats confirm the toxicity of bromoethane by inhalation (Müller, 1925; Bachem, 1927; Glaser & Frisch, 1929; Sayers & Yant, 1929). In guinea-pigs, for example, toxic signs were observed within 1 min of exposure to 45 000 mg/m3 (10 000 ppm) (Leuze, 1922; Sayers & Yant, 1929; Abreu & Emerson, 1940).
8.1.2 Oral

An oral LD50 of 1350 mg/kg body weight in rats has been listed with no further details (Izmerov et al., 1977), although in another study, Miller & Haggard (1943) observed "no serious ill effects" in rats from an oral dose of 1200 mg/kg body weight in olive oil.
8.1.3 Dermal

Schwander (1936) reported no signs of toxicity following a study in which an unspecified quantity of liquid bromoethane was placed in direct contact with the skin of one rabbit for 6 h. "

manimal - 2-7-2009 at 14:39

Since concentrated ammonia solution is so useful but rather scarce, I was thinking that a promising way to prepare it would be to heat intimately mixed ammonium sulfate and calcium hydroxide in a metal can and pipe the fumes into water.

Any thoughts on this?

Apparently it is possible, but as for yields, I haven't an idea. An interesting Popular Science article recommends it for experimental preparation. http://books.google.com/books?id=mCYDAAAAMBAJ&pg=RA1-PA1...

Reading that, I can see that Popular Science has gone downhill since 1944.

[Edited on 3-7-2009 by manimal]

[Edited on 3-7-2009 by manimal]

woelen - 3-7-2009 at 02:20

This could indeed be done. Mix the two chemicals (crushing to a powder may be necessary, before mixing) and add a few drops of water to make a slurry of the chemicals. carefully heat the mix, but do not heat very strongly, the mix should remain wet and you should not boil off the water in the mix.

Even easier is the making of ammonia from solid crushed ammonium sulfate and solid crushed sodium hydroxide. You must add a few drops of water to the mix, such that the material is just wet. Initially the reaction sets of, but at the end you may need some heating.

Simply leading the gas into water may be somewhat risky. There is a chance of suckback into the flask in which the ammonia is made (remember: ammonia dissolves in water exceptionally well). Better is to use an inverted small funnel and put that inverted funnel just on the water surface.
Another good thing to do is start the preparation with household ammonia instead of plain water. The first 5% of ammonia you then already have in the liquid.

User - 3-7-2009 at 03:00

Short question:
I might have the chance to buy a aspirator pump.
The seller says that it sucks 200 Mbar ( i think he means milli

)
So how must is look at this, is it strong enough to use for normal lab apps?

hissingnoise - 3-7-2009 at 03:45

Quote: Originally posted by woelen

Better is to use an inverted small funnel and put that inverted funnel just on the water surface.

The inverted funnel doesn't work well with NH3 as NH4OH is less dense than water.
The gas dissolves in the top few mm and is lost from the funnel when those few mm are saturated. . .
Bubbling the gas into the bottom of the container of water works well but a container or plastic bulb fitted inline to the hose to accommodate suckback is necessary.

entropy51 - 3-7-2009 at 11:13

Put a small magnetic stir bar ("flea") in the beaker under the funnel and spin it just enough to mix the NH4OH collecting in the funnel with the water beneath it. NH3 sucks back so easily that the funnel is a good idea. Using a setup like this I can dissolve so much NH3 in the water that it requires external cooling; it's quite exothermic as more and more NH3 dissolves.

I agree that NaOH seems to work better than Ca(OH)2, if the NH4SO4 is finely ground.

Manimal, aren't those old Popular Science articles great? The author of that one, Kenneth Swezey, later collected the articles in a great book called Chemistry Magic. Sometimes you can find it for sale on abebooks.com.

[Edited on 3-7-2009 by entropy51]

Paddywhacker - 3-7-2009 at 14:06

You should be able to get ammonia by heating almost any ammonium salt with almost any non-volatile basic compound. As a schoolboy I had my first experience of ammonia when I heated red lead with ammonium chloride. Got a huge toot of ammonia, and probably some lead volatiles too, when I gave it a good sniff.

1281371269 - 3-7-2009 at 18:02

I will try this on a small scale with about 5g ammonium chloride and sodium hydroxide. I'm guessing the equation looks like:
NH4Cl + NaOH -> NH3 + NaCl + H2O

An issue with this is that I will not end up with particularly conc. Ammonia solution, because I will also be boiling off water...

entropy51 - 3-7-2009 at 18:11

As Woelen said "add a few drops of water to make a slurry of the chemicals. carefully heat the mix, but do not heat very strongly, the mix should remain wet and you should not boil off the water in the mix."

That is there is little water to boil off and the NH3 comes off without strong heating.

You will not end up with a particularly concentrated Ammonia solution because 5 g of NH4Cl will yield very little NH3. If you scale it correctly, you can get quite strong ammonia solution, especially if you repeat it a few times. But keep the absorbing water cool. Otherwise it becomes quite hot and will dissolve no more NH3. Cold water absorbs much more NH3 than warm water.

[Edited on 4-7-2009 by entropy51]

1281371269 - 3-7-2009 at 18:33

I see, so about 60 degrees as opposed to a couple of hundred?
And yes I'm not really doing it to make Ammonia solution to use, more to try it out (on a test tube scale)

entropy51 - 3-7-2009 at 18:39

Quote: Originally posted by User

Short question:
I might have the chance to buy a aspirator pump.
The seller says that it sucks 200 Mbar ( i think he means milli

)
So how must is look at this, is it strong enough to use for normal lab apps?

1 bar = 750 mTorr = 750 mm Hg

So 200 mbar = 1/5 bar = 150 mm Hg

Normally aspirators connected to the sink produce 10 or 20 mm Hg depending on the water temperature, and that is often a good pressure for vacuum distillation. Many materials boil about 100C less than their atmospheric BP at 10 mm Hg. It depends on the boiling point of the material you are distilling as to whether a higher pressure is acceptable. 150 mm is on the high side for many distillations.

I have a pump that pulls about 150 mm Hg and I sometimes use it for vacuum filtrations, but it's not that good and I mostly use the sink aspirator.

A good aspirator connected to the faucet doesn't use all that much water, especially for distillations in a system free of leaks, unless you are using it 24/7.

entropy51 - 3-7-2009 at 18:45

Quote: Originally posted by Mossydie

I see, so about 60 degrees as opposed to a couple of hundred?
And yes I'm not really doing it to make Ammonia solution to use, more to try it out (on a test tube scale)

The test tube will contain a very strong ammonia solution and very little heat should be needed to release it, certainly less than 200 C. Normally by the time an aqueous NH3 solution reaches the boiling point of water all the NH3 has escaped as NH3 gas.

1281371269 - 3-7-2009 at 19:10

So one could effectively concentrate dilute ammonia solution by heating it gently and bubbling the resulting gas through water - would this not be easier than the ammonium salt / basic compound method?

querjek - 3-7-2009 at 19:33

Can aluminum chloride hexahydrate be dehydrated by heating somewhere below the sublimation point (178 C)?

not_important - 3-7-2009 at 20:06

No, it gives off HCl leaving various chlorohydrates, and I believe finally AlO(OH). SFAIK there is no way to dehydrate the hydrate chloride that is less challenging than directly making the anhydrous chloride from the metal.

Sedit - 3-7-2009 at 20:11

I have seen a patent that performed the dehydration of the hydrate but N_I's words still ring true, your better off just starting anhydrous and going that route.

entropy51 - 4-7-2009 at 07:21

Quote: Originally posted by Mossydie

So one could effectively concentrate dilute ammonia solution by heating it gently and bubbling the resulting gas through water - would this not be easier than the ammonium salt / basic compound method?

Well you can do that, but it's a long and slow process. Dilute ammonia is maybe 6% around here but if you can find the stronger "janitorial" strength the effort would be more worthwhile. Most ammonia I see nowadays also contains detergents which may foam too much when heating. You can make tons of NH3 with a bag of (NH4)2SO4 fertilizer and NaOH or Ca(OH)2.

UnintentionalChaos - 4-7-2009 at 09:42

I attempted to make some 2-bromopropane yesterday in the same way that smuv did in an ollllllllllld thread. This was basically H2SO4 + HBr + iPrOH. Does anyone know what specifically is that horrible acrid smell when the distillation begins?

As an aside, I'm pretty sure I got mostly diisopropyl ether and propene, followed by some actual isopropyl bromide containing (possibly, as I couldn't smell it, which is odd) dissolved bromine and a vast quantity of HBr fumes. I'm going to try a different way.

On the other, other hand, the reaction mix turned this incredible indigo-purple color halfway through, which hydrolyzed when I added some water. Is it possible that I had some polyatomic bromine cations floating around?

entropy51 - 4-7-2009 at 11:13

UC, that might go better without the H2SO4, which I don't think you need. Could the acrid smell be SO2? HBr, some of which probably distills?

I've made 2-Br propane by distilling constant boiling HBr and iPrOH per Vogel, page 387 of the 4th edition (probably also in the older editions).

UnintentionalChaos - 4-7-2009 at 11:26

Quote: Originally posted by entropy51

UC, that might go better without the H2SO4, which I don't think you need. Could the acrid smell be SO2? HBr, some of which probably distills?

I've made 2-Br propane by distilling constant boiling HBr and iPrOH per Vogel, page 387 of the 4th edition (probably also in the older editions).

Well, the thread claimed 90% yield, so I gave it a shot (also because I don't have any HBr) It was more like a rancid nasty organic smell. It came over before any distillate and passed through my NaOH scrubber, so I don't know.

I think I'll just make some HBr next time and go that route.

I do seem to have squeezed a roughly 30% yield out of it, although I am not happy with this.

[Edited on 7-4-09 by UnintentionalChaos]

entropy51 - 4-7-2009 at 11:58

UC, here's a nice HBr prep in case you don't already have a favorite method. I hope it's useful.

Attachment: HBr.pdf (1.2MB)
This file has been downloaded 3234 times

Maja - 4-7-2009 at 15:06

what would one use large amounts of KBrO3 ?

manimal - 5-7-2009 at 19:35

Does anyone know the percentages of xylene isomers contained in commercial xylene solvent?

And: what the fuck happened to roguesci.org?

[Edited on 6-7-2009 by manimal]

Paddywhacker - 5-7-2009 at 20:52

it varies. Look at the wiki page for xylene.

JohnWW - 5-7-2009 at 22:00

The technical-grade stuff contains ortho-, meta-, and para-xylenes, C6H4(CH3)2, of the same composition. There may be minor amounts of ethylbenzene (which has the same composition and molecular weight as xylenes); and possibly of other substances such as ethyl-substituted homologs, toluene, trimethylbenzenes, and durene.

UnintentionalChaos - 5-7-2009 at 23:01

Quote: Originally posted by JohnWW

The technical-grade stuff contains ortho-, meta-, and para-xylenes, C6H4(CH3)2, of the same composition. There may be minor amounts of ethylbenzene (which has the same composition and molecular weight as xylenes); and possibly of other substances such as ethyl-substituted homologs, toluene, trimethylbenzenes, and durene.

You're just about guaranteed a few percent of ethylbenzene. This can interfere, depending on what you want the xylene for. I recrystallized sulfur from hardware store xylene a while back and when I dried it, found clumps of gummy gray material clinging to my sulfur. The reason?

ethylbenzene + sulfur -(heat)-> H2S + styrene

And some of the styrene probably polymerized on me. Come to think of it, the little globs looked an awful lot what you get when you splash acetone on expanded polystyrene foam. I've wondered if this is actually a good way to remove ethylbenzene entirely and make a more useful product. Perhaps a KMnO4 wash to cleave the styrene's double bond before distillation would be in order.

Unrelated question:

Does anyone know where to get tiny hard polystyrene prills? About the consistency of ion exchange resin beads is needed, because that is more or less the project.

[Edited on 7-6-09 by UnintentionalChaos]

manimal - 6-7-2009 at 01:31

Ethylbenzene can be present in quite large amounts. According to it's MSDS, Startex brand xylene contains 20% ethylbenzene.

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