Sciencemadness Discussion Board

The short questions thread (2)

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Picric-A - 17-9-2009 at 22:31

Quote: Originally posted by entropy51  
There is an OTC product which consists mainly of sodium dithionite and carbonate. I don't know about the luminol synthesis, but it works for a number of other purposes.

Rit Color Remover http://www.ritdye.com/Fabric+Treatments.28.51.7.49.lasso

Dithionite can also be prepared by reducing SO2 with sodium formate in aqueous methanol. There is a reference posted around here somewhere on this method.

[Edited on 17-9-2009 by entropy51]


I have not found an OTC product containing Dithionite here in the UK.
As to the synthesis via reduction of methanol, there are lots of reperences and patents to that but i decided it only looked viable on an industrial scale.

Attachment: sodium dithionite from formate + SO2.pdf (603kB)
This file has been downloaded 598 times

entropy51 - 18-9-2009 at 07:15

Quote: Originally posted by Picric-A  

As to the synthesis via reduction of methanol, there are lots of reperences and patents to that but i decided it only looked viable on an industrial scale.
Oh did you now?

I guess an amateur like myself shouldn't attempt it then?:P

It works in the lab if one has source of SO2.

You can also prepare an aqueous solution by stirring up NaHSO3 with Zn dust.

Picric-A - 18-9-2009 at 07:24

Quote: Originally posted by entropy51  
Quote: Originally posted by Picric-A  

As to the synthesis via reduction of methanol, there are lots of reperences and patents to that but i decided it only looked viable on an industrial scale.
Oh did you now?

I guess an amateur like myself shouldn't attempt it then?:P

It works in the lab if one has source of SO2.

You can also prepare an aqueous solution by stirring up NaHSO3 with Zn dust.


I was not questioning your ability, i was meerly stating large reactors are needed, solutions containing mixtures of methyl formate, and lots of other chemicals are needed as a reaction base (see lit.) followed by a cheap SO2 source (my only source is H2SO4 + Cu- i am going to buy a cylinder soon though). Overall it seems like more trouble than its worth

I have given the NaHSO3 method a shot, but i didnt have powdered Zn, only granules. upon standing in a NaHSO3 solution for 24 hours, i got a colourless solution with the lump of Zn in and a lot of white ppt in the bottom (Zn(OH)2).
Could be a good method however powdered Zn if definitly needed

[Edited on 18-9-2009 by Picric-A]

1281371269 - 18-9-2009 at 07:51

Powdered Zn is dead easy to get. It was the first chemical I bought.

Jor - 18-9-2009 at 08:22

Garage chemist gas mentioned the anhydride of periodic acid (diiodine heptoxide) on this forum once. It is supposed to be an orange polymeric solid.

I added a spatula of periodic acid to 2mL of conc. H2SO4, and I then heated with a propane burner. The solid just dissolves, but no orange color is produced.
I added 3mL of DCM and shook vigorously, I then took the DCM layer, and boiled away the DCM. No residue remained. My guess is that I2O7 should be soluble in DCM, so no anhydride was formed.

Does anyone know how it can be synthesised? Without using things like oleum. Maybe heating periodic acid with phosphorus pentoxide?

entropy51 - 18-9-2009 at 09:29

Quote: Originally posted by Picric-A  
i was meerly stating large reactors are needed, solutions containing mixtures of methyl formate, and lots of other chemicals are needed as a reaction base (see lit.)

The patent you posted specifies a reactor because they're describing a large-scale industrial application of the reaction.

It works just fine in a stirred flask. You need only sodium formate, methanol and SO2.

The Zn reduction of NaHSO3 is even easier, but does require Zn dust.

bbartlog - 18-9-2009 at 14:54

What is produced at the anode if I electrolyze molten ammonium acetate (ammonia at the cathode, I expect)? Assume I use a PbO2 anode or the like, making the formation of a metal acetate or the like from the anode material unlikely. I'm having trouble seeing what the C2H3O2 anion would turn into, and the references I see online mostly talk about using ammonium acetate to assist in the electrolysis of something else (or turning it into nanocrystalline diamond via unorthodox voltage and electrodes).

ketel-one - 18-9-2009 at 19:15

^interesting question. The way I imagine it is that the AcO* radical would do a hunsdiecker type thing and make CH3* + CO2, two of the CH3* combining to make CH2CH2. Or it could make Ac-O-O-Ac, and decompose into something.

How do you know when an alcohol can be chlorinated with HCl? Is it correct that you can make tertiary halides from tertiary alcohols simply with HCl at RT?

UnintentionalChaos - 18-9-2009 at 19:45

Quote: Originally posted by bbartlog  
What is produced at the anode if I electrolyze molten ammonium acetate (ammonia at the cathode, I expect)? Assume I use a PbO2 anode or the like, making the formation of a metal acetate or the like from the anode material unlikely. I'm having trouble seeing what the C2H3O2 anion would turn into, and the references I see online mostly talk about using ammonium acetate to assist in the electrolysis of something else (or turning it into nanocrystalline diamond via unorthodox voltage and electrodes).


Quote: Originally posted by ketel-one  
^interesting question. The way I imagine it is that the AcO* radical would do a hunsdiecker type thing and make CH3* + CO2, two of the CH3* combining to make CH2CH2. Or it could make Ac-O-O-Ac, and decompose into something.

How do you know when an alcohol can be chlorinated with HCl? Is it correct that you can make tertiary halides from tertiary alcohols simply with HCl at RT?


Look up the Kolbe Electrolysis. I'm going to guess that you'd turn two NH4+ into 2 NH3 and a H2 at the cathode and perform the standard kolbe electrolysis at the anode, forming ethane and CO2.

ketel-one A mixture of conc. HCl and zinc chloride is called Lucas's reagent and may be of some assistance here. Usually the reaction is done at reflux to get good yields.

ketel-one - 18-9-2009 at 20:23

Many thanks Chaos, that explains everything.

Oh and that CH2CH2 should have been CH3CH3 obviously.

[Edited on 19-9-2009 by ketel-one]

not_important - 18-9-2009 at 20:34

Quote: Originally posted by Jor  
Garage chemist gas mentioned the anhydride of periodic acid (diiodine heptoxide) on this forum once. It is supposed to be an orange polymeric solid.


From Inorganic chemistry Egon Wiberg, Nils Wiberg, Arnold Frederick Holleman, 2001, comes the attached.

References from the 1990s talk about the heptoxide, mention it as orange, and state a very similar preparation to the hexoxide. Guess we need to find some journal references.

I2O6.png - 24kB

ketel-one - 18-9-2009 at 23:55

Alkylation reactions have to be performed in nonpolar solvent right? A 40% aqueous alkyl halide would have to be distilled out first, yes?

sonogashira - 18-9-2009 at 23:59

No not really. I think it is better in a polar solvent in fact. But likely it depends on what you are alkylating. Phenols are alkylated with dimethyl sulfate in water, alkyl halide in alcohols etc. (I have not heard of alkyl halide in water being used though - likely it is better to have it in alcohol but it may work.) As I say it depends what you have to alkylate and the size of the alkyl group though. :)

woelen - 19-9-2009 at 09:10

Quote: Originally posted by not_important  

From Inorganic chemistry Egon Wiberg, Nils Wiberg, Arnold Frederick Holleman, 2001, comes the attached.

References from the 1990s talk about the heptoxide, mention it as orange, and state a very similar preparation to the hexoxide. Guess we need to find some journal references.

I have the book "Chemistry of the Elements" from Earnshaw and this tells that I2O7 does not exist. There is no anhydride of periodic acid. The compound I2O6 certainly may exist, but I2O7 probably does not. The orange material, mentioned by Jor, probably is an impure compound, contaminated with iodine, which can explain the orange color.

[Edited on 19-9-09 by woelen]

Klute - 19-9-2009 at 14:23

40% aqueous alkyl halide solution? What alkyl halide is soluble to the extent of 40% in water? Are you not confusing with the amine? Or a solution in alcohol?

ketel-one - 19-9-2009 at 21:30

Lol, good point. I Must've had some unreacted alcohol.

Does anyone know... if you try to halogenate piperidine (or pyridine), would it halogenate at the 2 or the 4 position? Could iodine be used for that purpose with acetone as solvent?

Klute - 20-9-2009 at 07:06

Pyridine will be halogenated to the 3 and 5 positions with difficulty, requiring hot temperatures. See p. 58 in the book "Preparation of organic intermediates" that sonigasira posted in the aluminium alkoxide thread (org chem forum).

sonogashira - 20-9-2009 at 07:53

Quote: Originally posted by ketel-one  
Lol, good point. I Must've had some unreacted alcohol.

Does anyone know... if you try to halogenate piperidine (or pyridine), would it halogenate at the 2 or the 4 position? Could iodine be used for that purpose with acetone as solvent?


From Joule & Mills' Heterocyclic Chemistry:
"2-Bromo- and 2-chloropyridines can be made extremely efficiently by reaction of pyridine with the halogen, at 0-50C in the presence of palladium(II) chloride[25]"
And reference 25 is attached.

4-position is via the N-oxide


Attachment: s-1980-29025.pdf (129kB)
This file has been downloaded 772 times

Siddy - 21-9-2009 at 18:36

Sorry to butt in...

Say a clumsy chemist was to break a piece of glassware, what is the best adhesive they should use?

I fixed the outlet on the water jacket of a condenser with cyanoacrylate which works perfectly, but what if you were repairing something that may come into contact with organic solvents?

Cyanoacrylates are soluble in acetone, DCM etc.. before they set. (eg, directions on super glue says clean excess with acetone before dries).

Does anyone one if it will hold up once set, to at least acetone, hopefully DCM too?
Or is there a better adhesive to use?

JohnWW - 21-9-2009 at 20:56

Cyanoacrylate glue ("Superglue") is supposed not to work on materials containing alkalis, including soda glass. Two-pot epoxide glues - "Araldite" - are supposed to be the appropriate glue for such materials.

watson.fawkes - 21-9-2009 at 20:58

Quote: Originally posted by Siddy  
Say a clumsy chemist was to break a piece of glassware, what is the best adhesive they should use?
It's torch heat. Blow during and anneal afterward.

Really? If there were an adequate resin or polymer to repair glass, wouldn't there be labware made out it?—since it's cheaper to cast plastic than to blow glass.

Ephoton - 22-9-2009 at 01:30

dont bother :) you broke it hehe you buy another simple.

post whoredom is mine whaa ha ha ha

Siddy - 22-9-2009 at 02:34

Thanks JohnWW!

Jor - 22-9-2009 at 06:20

I'm currently doing the synthesis of 3-nitrophtalhydrazide. I added NaOH to a suspension of 2,11g 3-nitrophthalic acid until the solution turned pink (phenolphthalein). i then added a pinch of 3-nitrophthalic acid to make it yellow again. I then added 1,35g (slight excess) of hydrazine sulfate, and boiled the solution down to the point where a crystal crust formed. I then added about 15mL of glycerin, and heated the solution at 170-185C for 1 hour.

Now my first question is, at this point there was already quite some precitipate, while in the method on versuchschemie, a clear red liquid is obtained. How is this possible? Might it be possible that phtalhydrazide or 4-nitrophthalhydrazide precitipated, because of impurities in the 3-nitrophthalic acid (i made this myself).

Next I made a very stupid mistake. I added water about 1,5x the volume, to the still 150C solution (I was impatient) and there was some boiling.
Now what i forgot is possible hydrolysis. Is it possible that at those temperatures a lot of my 3-nitrophthalhydrazide hydrolysed? How prone is a hydrazid group to hydrolysis at slightly acidic pH?

ketel-one - 22-9-2009 at 12:48

Quote: Originally posted by not_important  

You can use iodine, water, and H2S, which give aqueous HI and elemental sulfur; iodine, water, and SO2, which gives you aqueous HI and H2SO4; or I2, H2O, and red phosphorous, the favorite of meth makers.

I've thought of an easier way, (no sulfur or phosphorous required); iodine is extracted very easily from medicinal iodine, and it reacts very violently with aluminum (you don't even need aluminum powder, just small squares cut with scissors out of aluminum foil); reaction starts spontaneously, produces lots of heat, and should be contained or lots of iodine will vaporize. That's to get Al-I3.

Then what I'm thinking is:
Al-I3 + 33%HCl = AlCl3 + 63%HI

sonogashira - 22-9-2009 at 13:42

Quote: Originally posted by Jor  
Now what i forgot is possible hydrolysis. Is it possible that at those temperatures a lot of my 3-nitrophthalhydrazide hydrolysed?


I would imagine it is more difficult to hydrolyse than a normal amide as the adjacent nitrogen causes more delacalization into amide bond - so I would think it is not a big problem - but that is only a guess.:)

[Edited on 22-9-2009 by sonogashira]

ketel-one - 23-9-2009 at 18:07

Does styrene react with aqueous HCl to make 1-phenyl,1-chloroethane? Trying to make styrene from styrofoam worked horribly bad, but I think I can perfect it.

UnintentionalChaos - 23-9-2009 at 18:16

Quote: Originally posted by ketel-one  
Does styrene react with aqueous HCl to make 1-phenyl,1-chloroethane? Trying to make styrene from styrofoam worked horribly bad, but I think I can perfect it.


No.

Dry HCl or HBr is usually used in an organic solvent or by passing directly into the compound to be hydrohalogenated, assuming it is a liquid.

ketel-one - 23-9-2009 at 18:42

Ok, that's logical. No reason HCl would ever want to leave a polar solution.

Sedit - 23-9-2009 at 18:52

Perhaps a phase transfer catalyst and hard stirring could work I would think.

To make styrene from styrofoam sounds like a mess due to the huge bulk. I have had fine success on the Pyrolysis of polystyrene forks one time. Its so cheep and easy to get I would never go that way if I really wanted it as it would be much easier to just distill from over the counter products.

Panache - 23-9-2009 at 19:06

Firstly perhaps this thread should be restarted as No.3, its getting somewhat long.

When nichrome wire is used in/as a heating element it is typically coiled in a small diameter coil (~3mm), then this coil is coiled around some ceramic mounting. Is small coiling simply to maximise the length and hence surface area of nichrome over a minimum volume or is it neccessary for some inductive effect and hence critical to the nichrome's resistance and heating ability.
Why is thicker nichrome wire seemingly rarely used?

not_important - 23-9-2009 at 21:04

Quote: Originally posted by Panache  
...
Why is thicker nichrome wire seemingly rarely used?


Resistance inversely is proportional to cross-sectional area, and thus to the square of the diameter. Double to the diameter and thus quarter the resistance.

As most heaters run from the mains, a fixed voltage level, cutting the resistance to 1/4 means pulling 4 times as much current, and generating 4 times as much heat per given length of wire. Continue to increase the diameter and you reach a point where the mains can't supply the current, or the heating becomes excessive and the resistance wire fails.

This is made worse because the radiative surface is proportional to the diameter (pi*2r) and so the area able to radiate heat increases linearly with increasing diameter while the heat generated increases as the square of diameter.

You can use four times the length of the twice-as-think wire, and draw the same amount of current and generate the same amount of heat, but now you've used 16 times as much nichrome - 4 times the cross section area x 4 times the length. And the heat generated per unit length decreases to 1/4 that of the thinner wire, giving the opposite problem of above - now the heating element may be too cool.

You could run the wire at 1/2 the voltage, giving 1/2 the current and thus the same amount of power dissipation. But transformers are heavy and expensive, so there is little call for that unless you need a short section of wire running hot.


12AX7 - 24-9-2009 at 18:52

I have a 1 ohm resistor that's wound open-frame with fairly heavy nichrome, probably around 12 gauge. It's probably good for about 300W, being about 8" long, 3" across, 1" thick (the winding is on a hollow rectangular ceramic form).

Tim

UnintentionalChaos - 25-9-2009 at 09:50

Does anyone know of a fairly simple synthesis of diacetyl? (butane-2,3-dione) Vogel's has precisely diddle on the matter and a benzoin condensation is worthless on enolizable substrates.

Klute - 25-9-2009 at 10:35

I think benzyl chloride posted a easy synthesis of diacetyl, from butanone oxime, after nitrosatation of MEK IIRC, or something similar.. It was a text or pdf paper, part of a multistep synthesis..

EDIT: Gotcha:

http://www.sciencemadness.org/talk/viewthread.php?tid=11966#...

Ask benzylchloride1 for details regarding the preparation of diacetyl, looks pretty easy to do.

[Edited on 25-9-2009 by Klute]

Inoxia - 25-9-2009 at 12:45

Concerning the ariel oxidation of benzaldehyde. If i left a small dish of around 2ml of 'Almond flavouring' (which uses i guess benzaldehyde) out in the open for 12 hour then boiled off the alcohol/water the next morning would i be left with benzoic acid (presuming it contained benzaldehyde) crystals?

UnintentionalChaos - 25-9-2009 at 12:53

Quote: Originally posted by Klute  
I think benzyl chloride posted a easy synthesis of diacetyl, from butanone oxime, after nitrosatation of MEK IIRC, or something similar.. It was a text or pdf paper, part of a multistep synthesis..

EDIT: Gotcha:

http://www.sciencemadness.org/talk/viewthread.php?tid=11966#...

Ask benzylchloride1 for details regarding the preparation of diacetyl, looks pretty easy to do.

[Edited on 25-9-2009 by Klute]


Thanks! I remember that thread, but not well enough to have gone looking for it.

[Edited on 9-25-09 by UnintentionalChaos]

Nicodem - 26-9-2009 at 00:08

Quote: Originally posted by ketel-one  
Does styrene react with aqueous HCl to make 1-phenyl,1-chloroethane? Trying to make styrene from styrofoam worked horribly bad, but I think I can perfect it.

Yes, you can prepare 1-chloro-1-phenylethane by refluxing styrene with conc. HCl(aq) and then steam distilling. You will need a few milligrams of 2,6-di-t-butyl-4-methylphenol or some other hydrophobic polymerisation inhibitor (hydroquinone might not work well or give only low yields).
See also JOC, 45, 3527-3529 (I think the paper was already posted on this forum so UTFSE).

Jor - 26-9-2009 at 10:28

I have some 3-nitrophthalhydrazid now, and as I have not recieved my sodium dithionite yet, I'm looking for another way to make luminol. I will only use about 20% of my current batch, and reduce the rest with sodium dithionite.

What would be my best bet? I was thinking of Sn/HCl, so add 250mg 3-nitrophthalhydrazid to 1mL conc. HCl, and add 300-400mg of tin. Then reflux (test-tube). The problem is that the starting material is insoluble, and tin is a solid as well.
How long should I reflux it?
Do you think this works, or do you have alternative methods?

entropy51 - 26-9-2009 at 11:00

Jor, you could probably make enough dithionite for a 250 mg reduction by stirring some NaHSO3 up with Zn dust.

Alternatively I would expect SnCl2 or sodium sulfide to reduce the nitro group.

If you want to use Sn metal, you might check to see if the hydrazide is soluble in AcOH. If so, you could use Sn/AcOH I believe.

Try it on an even smaller scale than 250 mg, if possible.

Picric-A - 27-9-2009 at 01:04

I have tryed Sn/HCl method- failure.
I also tryed the folowing method taken from ORganic Chemistry by P.J.Durrant. one gram of the solid was dissolved in 10ml of water and 8g of NaOH was dissolved in the water as well. A steady stream of H2S was then bubbled in until the NaOH stopped absorbing the gas then the mix was refluxed for 30 mins. Yield- none. No ppt was formed after addition of GAA.

densest - 27-9-2009 at 15:03

A (hopefully) easy question: how to destroy nitrate (NO3-) ions in an aqua regia solution after dissolving metals presumed to contain precious ones. The protocol which I have read says to boil the liquid down until it's syrupy, add concentrated HCl, repeat twice more. This seems to me both wasteful and toxic. What I propose to do is to add urea (CO(NH2)2) and heat, assuming that the reaction 6 HNO3 + 5 CO(NH2)2 -> 5 CO2 +8 N2 + 13 H2O is possible. I haven't found anything explicit either way: what I do find is that urea is used to remove excess NO2 from nitric acid solutions.... is the analogy likely? Thanks for any pointers!

entropy51 - 27-9-2009 at 15:45

The Energetic Materials crowd can offer better advice, but might you not form urea nitrate, which I believe can be explosive under certain cirumstances? Just a thought.

EDA from EDTA

Sedit - 27-9-2009 at 16:46

Can anyone offer some suggestions or point me in the right direction on how I would cleave an acetate function from an amine?

My goal is to synthesis Ethylenediamine from Ethylenediamine tetraacetate or NaEDTA respectively. I have found some information on using a dilute (aq)solution of Ammonia but this all takes place at high temperatures and pressures which I would like to avoid if possible. I remember seeing a suggestion in another threed sometime back on how to handle a situation like this but I can not for the life of me seem to find it anymore.

Thanks in advance for anyhelp,
~Sedit

Jor - 27-9-2009 at 16:57

I can obtain some a-Benzoin oxime for a acceptable price. Google doesn't give much useful hits, only that it can be used to detect metals and seperate some metals from eachother.

Are there any interesting complexes wich can be isolated, with this compounds as the ligand, or other interesting uses?

What's good for dissolving HDPE?

Formatik - 30-9-2009 at 14:35

I've come up with a list of some compounds based on some listed incompatabilities (just below), but am asking also anyone with experience in dissolving the plastic.

Thionyl chloride, chlorosulfuric acid, aqua regia, nitric acid (95% plus), oleum, H2CrO4/H2SO4, bromine, ethylene chlorohydrin, halogenated hydrocarbons (tetrachloroethylene, perchloroethylene, chlorobenzene, chloroform, dichloromethane, tetrachloromethane, etc.), hydrocarbons and others (cyclohexane, ethylbenzene, kerosene, turpentine, diethyl ether, petroleum ether, and the like).

I remember using Et2O a while back to attack HDPE allowing the plastic to soak for some days, but it didn't do a good job. The same with dichloromethane. Neither dissolved significant amount of the plastic.

not_important - 30-9-2009 at 15:20

HDPE is in effect an alkane, mostly CH2 groups with a small amount of branching giving tertiary hydrogens, and a tiny amount of unsaturation. Anything immiscible with petroleum ether and that doesn't readily chemically attack the same is not going to do much to HDPE.

Solvents based on aromatics and halocarbons are the best bet, and generally need to be hot to boiling. You still don't get a good solution, often a gel is the result. General hydrocarbons, including kerosene, and other compounds that have long carbon chains (biodiesel) will also act as solvents when rather hot.

I believe that ethylene under very high pressure dissolves PE, I remember reading a paper on MW fractionating polyethylene during production by controlling the pressure of the remaining ethylene with higher MW PE dropping out first as the pressure was reduced.

Solvents based on phenols or hydrogen donors such as tetralin and decalin 'dissolve' HDPE by enhancing thermal degradation to random mixtures of mostly linear chain alkanes over a wide range of carbon numbers.





[Edited on 30-9-2009 by not_important]

ketel-one - 30-9-2009 at 20:51

Phenolic hydroxy groups don't count as alcohols and don't react with HCl to replace the hydroxy with a chlorine, right? Also I don't suppose there's any easy way to turn a phenolic hydroxy group to a phenolic methoxy group, right?

ketel-one - 30-9-2009 at 20:52

Phenolic hydroxy groups don't count as alcohols and don't react with HCl to replace the hydroxy with a chlorine, right? Also I don't suppose there's any easy way to turn a phenolic hydroxy group to a phenolic methoxy group, right?

mr.crow - 30-9-2009 at 21:03

Quote: Originally posted by ketel-one  
Phenolic hydroxy groups don't count as alcohols and don't react with HCl to replace the hydroxy with a chlorine, right? Also I don't suppose there's any easy way to turn a phenolic hydroxy group to a phenolic methoxy group, right?


As far as I can tell you can't replace phenolic OH with something else

Making methoxy is easy, there are several threads about methylating agents. Phenolic OH easily turns into a phenolate salt, which can undergo Williamson ether synthesis.

ketel-one - 30-9-2009 at 23:06

Thanks, I've found it NaOH + alkyl halide, in retrospect that should have been really easy to realize but alright I got it, thanks.

dann2 - 1-10-2009 at 06:00

Just wondering (as one does)
What is the reaction product between Titanium Tetrachloride and Acetone?


http://cgi.ebay.co.uk/THE-REACTION-PRODUCT-OF-ACETONE-TITANI...

Dann2

Formatik - 1-10-2009 at 10:41

Quote: Originally posted by not_important  
HDPE is in effect an alkane, mostly CH2 groups with a small amount of branching giving tertiary hydrogens, and a tiny amount of unsaturation. Anything immiscible with petroleum ether and that doesn't readily chemically attack the same is not going to do much to HDPE. ...


Thanks not_important, I will be trying the hot hydrocarbon when I get the chance to.

Formatik - 1-10-2009 at 10:43

Quote: Originally posted by densest  
A (hopefully) easy question: how to destroy nitrate (NO3-) ions in an aqua regia solution after dissolving metals presumed to contain precious ones. The protocol which I have read says to boil the liquid down until it's syrupy, add concentrated HCl, repeat twice more. This seems to me both wasteful and toxic. What I propose to do is to add urea (CO(NH2)2) and heat, assuming that the reaction 6 HNO3 + 5 CO(NH2)2 -> 5 CO2 +8 N2 + 13 H2O is possible. I haven't found anything explicit either way: what I do find is that urea is used to remove excess NO2 from nitric acid solutions.... is the analogy likely? Thanks for any pointers!


Urea will destroy nitrogen oxides from HNO3, but with aq. HNO3, you get urea nitrate. Aqua regia and urea forms something like near quantitative amount of nitrous oxide.

Reduction Stannic to Stannous

smuv - 1-10-2009 at 19:28

I am looking to recycle Sn(II) used for a reduction. I was wondering if I could reduce Sn(IV) to Sn(II) in an acidic solution by means of SO2.

I have looked at the reduction potential tables and it seems that SO2 is a strong enough reducing agent; I wonder If I am missing something which may make this fail in practice.

I do know that Ferrous salts have been used for this reduction, but I would like to achieve this reduction without adding any metal salts that would precipitate under basic conditions (I plan to isolate the tin from solution by basifying it to collect the hydrated oxide).

Thanks

Nicodem - 1-10-2009 at 23:21

Quote: Originally posted by dann2  
Just wondering (as one does)
What is the reaction product between Titanium Tetrachloride and Acetone?


http://cgi.ebay.co.uk/THE-REACTION-PRODUCT-OF-ACETONE-TITANI...

Dann2

Titanium tetrachloride is a relatively strong acid so its reaction with enolisable ketones should give normal acid catalysed self condensation products (in amounts and ratios depending on temperature and other conditions). The acid catalysed self condensation products of acetone are numerous, but the most famous ones are mesityl oxide, phorone and mesitylene.

ketel-one - 2-10-2009 at 21:13

Does anyone know any common high BP boiling solvents? What I'm looking at (decarboxylation of amino acids to corresponding amines) asks for cyclohexanol (BP 160C) but I'm wondering if something common has similar or higher boiling point.

crazyboy - 2-10-2009 at 21:19

Quote: Originally posted by ketel-one  
Does anyone know any common high BP boiling solvents? What I'm looking at (decarboxylation of amino acids to corresponding amines) asks for cyclohexanol (BP 160C) but I'm wondering if something common has similar or higher boiling point.


Xylene is probably the most common solvent with the highest BP.

ketel-one - 4-10-2009 at 21:30

It's alright, I've found that solvent is not totally necessary, strongly heating amino acid with burner makes it fume a lot of white CO2 (fumes don't smell bad or anything). If you heat it too strongly it will start to turn brownish and melt and will easily ignite even if it never directly touches flame. I think the main disadvantage over doing it with solvent is that it might get heated unevenly, some parts burning and some parts undecarboxylated.

A new question: This doesn't work for acetic anhydride, but can nicotinic anhydride be made by nicotinic acid + conc. H2SO4? I'm basing this on the fact that nicotinic acid is a solid while glacial acetic acid is a liquid, only the former I would believe is capable of dissolving sulfuric acid without dehydrating.

chloric1 - 5-10-2009 at 05:48

Quote: Originally posted by smuv  
I am looking to recycle Sn(II) used for a reduction. I was wondering if I could reduce Sn(IV) to Sn(II) in an acidic solution by means of SO2.

I have looked at the reduction potential tables and it seems that SO2 is a strong enough reducing agent; I wonder If I am missing something which may make this fail in practice.

I do know that Ferrous salts have been used for this reduction, but I would like to achieve this reduction without adding any metal salts that would precipitate under basic conditions (I plan to isolate the tin from solution by basifying it to collect the hydrated oxide).
Thanks


Crystallize your tin chloride and then dissolve in warm concentrated hydrochloric acid. Add tin metal in small pieces and keep warm until dissolved. Leave a single piece of tin at the bottom to keep the stannous ions "fresh".

manimal - 5-10-2009 at 13:16

Has anyone ever made sodium pyrosulfate from sodium bisulfate by heating? I am having trouble. I can't tell if the molten yellow liquid is pyrosulfate or still bisulfate. Also, there is much smoking.

Picric-A - 5-10-2009 at 13:53

Manimal- check out the oleum from NaHSO4 thread.
The smoking is most probably the decomp of the sodium pyrosulphate to SO3 which immedietly reacts with water vapour to produce fumes of H2SO4.
To make the pyrosulphate slowly heat the bisulphate to +-10deg C 315 degrees, till it melts and starts to release water. when the water is removed (may take a few hours) then SO3 (lots of fumes) should form. This is the point you want to stop at to make the pyrosulphate.

[Edited on 5-10-2009 by Picric-A]

manimal - 5-10-2009 at 15:19

Yes, I suppose you're right. I didn't think it was H2SO4 mist, because I thought H2SO4 mist would be intensely irritating, whereas with this, it only had an odor that I would describe as like hot ashes or rock dust.

Swede - 6-10-2009 at 07:12

I am interested in the use of sulfur dioxide gas to reduce chlorate ion contamination in electrolytically-prepared perchlorate. In the past, I have used (with success) Potassium Metabisulfite added to hot, acidic aqueous solutions, but the contact of the evolved SO2 gas with the solution is very brief, and the amount of potassium metabisulfite necessary to "clean" the perchlorate is higher than I would like.

SO2 gas has an inverse solubility, with higher SO2 solubilities at lower temperatures. This is in conflict with the salt of interest, potassium perchlorate, which has a lousy solubility except at higher temperatures.

Essentially, the challenge is to maximize contact of the SO2 gas and use the least amount of metabisulfite per unit of chlorate contamination, which is an estimated 0.5%.

I'm thinking perhaps the best way is to start cold, introduce the SO2, then slowly heat, which will dissolve the contaminated perchlorate, with hopefully enough SO2 remaining to reduce the contamination.

Any thoughts on obtaining the most economical use of the metabisulfite?

Picric-A - 6-10-2009 at 07:19

Quote: Originally posted by manimal  
Yes, I suppose you're right. I didn't think it was H2SO4 mist, because I thought H2SO4 mist would be intensely irritating, whereas with this, it only had an odor that I would describe as like hot ashes or rock dust.


Maybe becuase SO3 evoloutiuon had just started so it was a vry dilute acid mist. Either way i dont reccomend smelling it to detect H2SO4 presence. use indicator paper suspended above the molten mix. When all its red stop heating becuase you have the pyrosulphate

watson.fawkes - 6-10-2009 at 07:42

Quote: Originally posted by Swede  
Any thoughts on obtaining the most economical use of the metabisulfite?
I'd modify a fractional distillation column, using counter-current circulation of the gas (up) and the liquid (down). Really, this just means a pump and a way of introducing the gas. You could continue to use the metabisulfite, or alternately use a sulfur candle. The goal in intimate gas-liquid contact, reusing rather than venting the gas.

MJ_ - 6-10-2009 at 23:17

I know this might be a stupid question, but is it possible to drink lab reagent grade ethanol that is diluted to 40%? This is no different to vodka yes? I already have ethanol however i dont want to be buying vodka if what i have is perfectly safe when diluted appropriately.

Picric-A - 7-10-2009 at 10:36

Depends if it is denatured or not. Reagent grade ethanol means nothing.
Rectified spirit (~98% ethanol) can be drunk if diluted however it would be cheaper to buy vodka than use that.
Most 100% ethanol has been made by azeotropic distillation with benzene to remove water and as such the ethanol may contain traces of benzene with its accociated hazards.
Just buy vodka and dotn try to use laboratory chems :) safer and cheaper

MJ_ - 7-10-2009 at 12:43

Thanks for your reply, it is definately not denatured and i believe it is 100% ethanol but the thought of trace amounts of benzene is enough to put me off.

Jor - 16-10-2009 at 11:32

I can currently obtain a lot of silica gel, however, it is a powder used for chromotography. Can this also be used for used a dessicant in my dessicator, instead of the granules?

Picric-A - 16-10-2009 at 12:30

Yes i use a powder however everytime you change the thing you are drying i reccomend you give the powder a bit of a stir up.
Granules are of course better but the powder works nonetheless

sparkgap - 16-10-2009 at 20:07

Quickie: I tried checking in Google Scholar but can't seem to sort it out; has epichlorhydrin ever been used in a Friedel-Crafts alkylation?

sparky (~_~)

smuv - 17-10-2009 at 08:11

The problem with epichlorohydrin is that it has two labile positions. The epoxide moiety will react under FC conditions to yield an aryl-substituted alcohol (the ratio of primary/secondary depends on the nucleophilicity of the aromatic you are reacting it with). Additionally epichlorohydrin is chloro-substituted which will as well lead to FC products. One could even make the argument that the alcohol formed from the ring opening of the epoxide would as well react under some conditions.

Overall you have 2 or 3 possible reactions, based on this, I think it would be very difficult to do any FC chemistry on this compound without observing a mixture of many products (including polymers).

[Edited on 10-17-2009 by smuv]

sparkgap - 18-10-2009 at 07:35

Hmm, I forgot to take into account that alkyl groups activate aromatic rings towards FC alkylation. :( I was hoping 1,3-diphenyl-2-propanol was as easy as epichlorhydrin + benzene, but I guess I'll look for something else.

Thanks smuv.

sparky (~_~)

[Edited on 18-10-2009 by sparkgap]

DJF90 - 18-10-2009 at 08:03

You could try using 2eq. Phenylmagnesium bromide? It should attack at the least hindered position (unless you add a strong lewis acid) of the epoxide, yielding a halohydrin, which would then react with the second equivalent to displace chlorine. Shouldnt be too hard?

smuv - 18-10-2009 at 12:28

DJF90: phenylmagnesium bromide might work well in this case; however the mechanism probably would be a little different than you proposed. The first eq of phenylmagnesium bromide would open the epoxide to form the corresponding chlorophenylpropanoate, this should immediately cyclicize, displacing chloride, to form 2-(phenylmethyl)-oxirane. This epoxide will react with the second eq of PhMgBr to form 1,3-diphenyl-propan-2-oate.

DJF90 - 18-10-2009 at 13:10

Yes sorry wasnt thinking straight. I've drawn the mechanism for that reaction (grignard on epichlorohydrin) before as well! Its been a long day. But yea, I suggest that would be the appropriate way to go, so long as PhMgBr isnt a problem.

Nicodem - 19-10-2009 at 00:47

Quote: Originally posted by sparkgap  
Hmm, I forgot to take into account that alkyl groups activate aromatic rings towards FC alkylation. :( I was hoping 1,3-diphenyl-2-propanol was as easy as epichlorhydrin + benzene, but I guess I'll look for something else.

There are a couple of examples of Friedel-Crafts alkylation of benzene derivatives with epichlorihydrin, however the products are Ar-CH2CH(OH)CH2Cl or bisalkylated products. See: European Journal of Medicinal Chemistry, 43 (2008) 300-308 (on mesitylene); Khimiya Geterotsiklicheskikh Soedinenii (1985) 1319-1321 (bisalkylation of p-xylene). I could find no examples of ArCH2CH(OH)CH2Ar formation, but you could treat the Ar-CH2CH(OH)CH2Cl with NaOH to form the 3-arylpropene oxide and then do another FC alkylation with it.

sparkgap - 19-10-2009 at 06:29

Thanks to Nicodem, smuv, and DJF90. :)

@Nicodem:

"There are a couple of examples of Friedel-Crafts alkylation of benzene derivatives with epichlorihydrin, however the products are Ar-CH2CH(OH)CH2Cl or bisalkylated products."

- As you said, the fact that it stops at the halohydrin isn't too much trouble. :) I'll check those refs now... hopefully an excess of benzene should minimize the polyalkylation products in my case.

@smuv and DJF:

I already thought of that idea, but was asking about FC because I was trying to avoid Schlenk-ware use (long story). :) But I'll fall back on that if FC doesn't pan out.

sparky (~_~)

DJF90 - 19-10-2009 at 11:14

You shouldnt need schlenk-ware, just normal ground joint glassware. Make the grignard and react it with the epichlorohydrin - the order of addition shouldnt matter - you should get the product regardless of whether you add the epichlorohydrin to the grignard or vice versa.

Neo-6 - 21-10-2009 at 08:40

How much lithium is in cr2032 battery or any of the cr batteries?
This can be calculated and if someone could find the mg of lithium per mAh it would be greatly appreciated:)

densest - 21-10-2009 at 09:01

Quote: Originally posted by Neo-6  
How much lithium is in cr2032 battery or any of the cr batteries?
This can be calculated and if someone could find the mg of lithium per mAh it would be greatly appreciated:)


These pieces of information are easily available from common references:

The rest is arithmetic and/or algebra

:D

Neo-6 - 21-10-2009 at 20:57

is this correct?

cr3032 has 500mAh = 0,5Ah
0,5Ah * (60min*60sek) = 1800 coulombs
1800 / 96500 = 0,018653 moles
witch is in lithiums case ~ 130mg

NaHSO4 substituting KHSO4

Picric-A - 22-10-2009 at 10:45

Can sodium bisulpate replace potassium bisulphate in dehydration reactions, specifically glycerol--> acrolein.
thanks,

chemrox - 22-10-2009 at 11:30

Quote: Originally posted by Neo-6  
is this correct?

cr3032 has 500mAh = 0,5Ah
0,5Ah * (60min*60sek) = 1800 coulombs
1800 / 96500 = 0,018653 moles
witch is in lithiums case ~ 130mg


what is "cr3032" ??

12AX7 - 22-10-2009 at 14:51

Quote: Originally posted by chemrox  

what is "cr3032" ??


An odd case, the newbie knows something a forum regular does not:
http://www.google.com/#q=cr2032


[Edited on 10-22-2009 by 12AX7]

watson.fawkes - 22-10-2009 at 14:56

Quote: Originally posted by chemrox  
what is "cr3032" ??
It's the designation of a flat coin cell.

Jor - 27-10-2009 at 10:21

Does anyone know if phosphorus pentoxide attacks glass t temepratures of 250C? I am planning to make a small amount (10mL) of the chemical, by heating P4O10 with NaCl, where POCl3 will be formed at 250C (Gmelin).

I know polyphosphoric acids will attack it, but does P4O10 (with ofcourse traces of polyphosphoris acids present , because the oxide is so incredibly hygroscopic.


JohnWW - 27-10-2009 at 13:24

According to Perry's Chemical Engineers' Handbook (see References section under Chemical Engineering books for links to access the 7th and 8th editions), chapter 23, glass is slowly attacked by H3PO4, at a concentration of about 35%, from about 300ºF (149ºC) upwards; and attacked more rapidly by concentrations greater than about 35% at about 400ºF (204ºC). I presume that laboratory borosilicate glass, which is somewhat alkaline, is meant; ordinary soda bottle glass would be even more rapidly attacked.

woelen - 27-10-2009 at 23:55

Oh yes, I expect the P4O10 to react with the alkaline glass to form pyrophosphates of sodium. The orthosilicate in the glass easily can give off an oxygen which is taken up by P4O10. Four of such oxygens then can lead to formation of P2O7(4-) ions.

The reaction will not be so fast that your glass is eaten away in a single run, but it will become opaque. You could try the synth and keep that etched piece of glassware for other such etching reactions (reactions involving fluoride or hydrazine-based reactions at elevated temperatures).

Jor - 28-10-2009 at 01:26

In that case I will use cheap glassware. Maybe even a test tube, wich is bent, just like in the synthesis of phosphorus. I think that if I heat slowly, most POCl3 will condense in the other side of the test tube. This is because POCl3 boils quite high, and even if you heat one part of a DURAn test tube very strongly, the other part remains cool. But heating should be slow, so the vapour cannot greatly heat that side of the test tube.
Now i really need some glasswool, to close the test tube. Because i think POCl3 will just chlorinate the OH-groups on cotton wool.
Should be good to preapre about 10mL in a few runs, I hope, wich is enough for my experiments.
Whemn I try, i will post the results.

Hydroflouric acid

Paddywhacker - 29-10-2009 at 01:57

A safety inspector that was looking over my lab told me that hydroflouric acid was a watched chemical, but he couldn't say why... thought it was some sort of explosive.

Anybody know what that was about? I'm not in the 'cooking' scene, but I've read parts of the Rhodium archive ... there's a lot of interesting chemistry there ... but never come across requirements for large amounts of HF.

Jor - 29-10-2009 at 07:28

Quote: Originally posted by Paddywhacker  
A safety inspector that was looking over my lab told me that hydroflouric acid was a watched chemical, but he couldn't say why... thought it was some sort of explosive.

Anybody know what that was about? I'm not in the 'cooking' scene, but I've read parts of the Rhodium archive ... there's a lot of interesting chemistry there ... but never come across requirements for large amounts of HF.

Because it is used in the synthesis of nerve agents like sarin. Here sodium fluoride and potassium fluoride, are also watched, although small quantities (less than 250g) is no problem.

Soxhlet Extractor

Picric-A - 29-10-2009 at 08:27

Thimbles are expencive for Soxhlet extractors.
Will they still work if you bung up the hole with cotton wool instead of a thimble?
I am extracting amygdalin from peach kernels btw :)

Sedit - 29-10-2009 at 12:15

I have a couple related questions that should not be to hard.


Question one what temperature does an Ethyl alcohol flame burn at

Question two: How would I beef up the energy and temperature of something like an alcohol lamp. I have considered using a mix of Nitromethane/MeOH like that used in RC cars but better judgement says not to. Any suggestions on a fuel source other then denatured alcohol which is only getting a steady reading with my Pyrometer of 900 degrees F which seems quite low to me(its a bit chilly outside as well though).

Picric-A - 29-10-2009 at 14:01

hmm maybe adding ~3% conc (~50%) H2O2 to dry meths?

Sedit - 29-10-2009 at 14:13

35% H2O2 is all I have and I don't know if it would be enough to balance the lose of heat associated with the H2O being present in the mixture. Im still thinking about the Nitromethane more now and just diluting it a bit more from the 30% solution down to about 5-10% and seeing how that compairs.

I still have yet to locate the temperature at which EtOH burns and all I keep getting in web searchs is how well alcoholic peoples fat burns in spontaneous human combustion:mad:. God I swear the internet has gone down hill in the last 8 years.

UnintentionalChaos - 29-10-2009 at 15:40

The answer is to use a larger flame or stop using alcohol and graduate to a bunsen or meker burner. A flame temperature is limited by the enthalpy of combustion, gas volume produced, how well it is being mixed with oxygen, and flame size (limits heat transfer to the surrounding air).

Sedit - 29-10-2009 at 16:28

Its not really a factor of needing to graduate or anything of that nature its just that I used something simular as This to melt lead a few days ago for the reduction of Pottasium nitrate and it got me to thinking of the range of temperatures that could be achived using different fuels in a simular device. It worked very well and burned for 45 minutes or so on a single charge which kind of impressed me to say the least.

I was under the impression that there is an upper limit to the temperature at which a specific fuel can burn at assuming complete combustion and is independent of flame size.

gsd - 29-10-2009 at 16:51

Quote: Originally posted by Sedit  

I was under the impression that there is an upper limit to the temperature at which a specific fuel can burn at assuming complete combustion and is independent of flame size.


That is correct. It is called as "Adiabatic Flame Temperature".
For a flame (produced by any fuel) that temperature will be way above 900 F or even 900 C.

It seems your problem is not the temperature of the flame but concentrating the flame on to your target.

Remember radiation losses are to the 4th power of absolute temperature of the flame. So if you have a tall and slender flame or multiple small flames then only a vary small portion of the heat will be delivered to your area of application.

gsd
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