Sciencemadness Discussion Board

The Short Questions Thread (4)

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blogfast25 - 29-2-2016 at 16:55

Quote: Originally posted by luminouspath  
I'm planning on concentrating some 12% H2O2 to 30% by removing water under vacuum at ~30c, is this as safe as I think it is?


Assuming it's possible and you've got good control over final strength it should be.

I would give my glass ware a bit of a deep clean to avoid any oxidisables in there.

DraconicAcid - 29-2-2016 at 16:58

Quote: Originally posted by 1.6180339  
In areas with too much sulfur dioxide, could we[in a industrial process] convert it into sulfur trioxide, make it react with water to create sulfuric acid and then make it react with sucrose to make carbon which could be used for industrial processes?


There's far more cost-effective sources for carbon. Charring wood is much cheaper, and if you had to make it from sugar, it's easier to do it with heat than with sulphuric acid.

PHILOU Zrealone - 1-3-2016 at 06:57

Quote: Originally posted by 1.6180339  
In areas with too much sulfur dioxide, could we[in a industrial process] convert it into sulfur trioxide, make it react with water to create sulfuric acid and then make it react with sucrose to make carbon which could be used for industrial processes?

Not usefull except for desesperate people.
This is like killing a fly with an atomic bomb...way too much energy and money expenses for such a little result.
See differences between "efficacy", "efficiency", "effectuality", "effectiveness".

If you have that much SO2 better cath it by liquefaction or into NaOH solution (NaHSO3 and Na2SO3) and this would be more interesting as reducer.

The making of SO3 cost energy, the concentration of H2SO4 to a level able to dehydrate sucrose is another barrier.
The use of sucrose to make carbon is an energetical and financial crime.
The use of conc H2SO4 to dehydrate sucrose to C is another chemical crime.

karlos³ - 3-3-2016 at 05:44

Hello!
I need semipermeable membranes, the one I am talking about are these used to separate the electrodes in car batteries.
They are perfect for the experiments wih electrochemistry I want to start with, but I have no battery to wreck for them.

Can one buy these membranes somewhere separate, online or offline?
I guess that is possible but I don´t know how they are called, or where one would look after them.

It would be deeply appreciated if someone could help me out here.

[Edited on 3-3-2016 by karlos³]

[Edited on 3-3-2016 by karlos³]

Deathunter88 - 10-3-2016 at 08:44

Potassium Chlorate -> Barium Chlorate
I have some potassium chlorate that I made through electrolysis. Since I'm not really into pyrotechnics I want to convert it into another more soluble chlorate. At the top of my list are magnesium, calcium, strontium, and lithium chlorates. As an intermediate I plan to use the barium chlorate + sulphuric acid route to make chloric acid, and then react that with a suitable carbonate. But here lies the problem, I cannot find a way to convert potassium chlorate into barium chlorate. The 2 or 3 threads already on this forum about this issue all ended with a theoretical dead end so that is why I chose to post this question in the short questions thread.

This conversion does not seem doable with a metathesis reaction since the only chlorate salts that is less soluble than potassium chlorate are caesium and rubidium (both way too expensive to use). Perhaps a member here has some way?

Or maybe this is simply impossible and I will just toss my chlorate. :(

Metacelsus - 10-3-2016 at 08:52

Maybe this could be done with an ion-exchange resin of some sort, loaded with barium. However, this would likely be expensive.

Why couldn't you just put barium chloride in a chlorate cell (or for that matter, the chloride salt of your choice)?

Deathunter88 - 10-3-2016 at 09:27

Quote: Originally posted by Metacelsus  
Maybe this could be done with an ion-exchange resin of some sort, loaded with barium. However, this would likely be expensive.

Why couldn't you just put barium chloride in a chlorate cell (or for that matter, the chloride salt of your choice)?


I guess I could just electrolyse some calcium chloride...
But since I already have the KClO3 I don't want to throw it away, and I hate having to clean up the huge hexavalent chromium mess my cell makes after running which is why I was hoping I would never have to do electrolysis ever again.

karlos³

Sulaiman - 10-3-2016 at 14:12

I am a noob just starting to experiment with galvanic cells,
I think that for Lead-Acid cells it is more of a physical separator and to prevent dendrites shorting out the cell.
Almost any porous materials such as woven or wool fiberglass, plastic kitchen scourer etc...

In other situations such as the Daniel cell the membrane allows ions to squeeze through the pores
but prevents bulk mixing of electrolytes either side.

karlos³ - 11-3-2016 at 12:07

Quote: Originally posted by Sulaiman  
I am a noob just starting to experiment with galvanic cells,
I think that for Lead-Acid cells it is more of a physical separator and to prevent dendrites shorting out the cell.
Almost any porous materials such as woven or wool fiberglass, plastic kitchen scourer etc...

In other situations such as the Daniel cell the membrane allows ions to squeeze through the pores
but prevents bulk mixing of electrolytes either side.


Thank you for your reply!
I am also just beginning with electrochemistry but I already know that a lot of alternatives for the cell divider can be used.
But in my case I have to use a different strenght of eletrolyte in each part of the cell.
These membranes look mostly attractive to me because they don´t have a high resistance so the solution does not heat up very much.
I could also take something from unglazed porcellain but then I have to cool my cell because of that high resistance.

I have to research if a daniel cell would be a suitable alternative for my experiment but it seems unlikely.

Maker - 15-3-2016 at 10:33

What advantage does a RBF offer over a conical flask for distillation/reflux?

I'm considering a jointed conical flask or two for distillation as I don't have a heating mantle but there must be a reason why RBFs are used far more often.

Sulaiman - 15-3-2016 at 11:08

For (partial) vacuum distillation an Erlenmeyer would have to be made of thicker glass than an rbf
especially the flat base.
Thick glass cannot support as large a temperature shock or difference as thin
http://www.kavalier.cz/en/section/32-simax-glass-mass.html
so maximum operating temperatures of an Erlenmeyer must be less than an rbf

Zephyr - 15-3-2016 at 11:13

I believe the main reason RBF's are preferred is because they are better suited to high temperatures. When an RBF reaches high temperatures and the glass expands slightly, it expands outwards evenly because of its spherical shape. However, when the bottom of a erlenmeyer heats up and expands, it cannot evenly expand outwards because it is flat, and if the heating is extreme enough or the glass low quality enough, it will often fracture.

That being said conical flasks are usually acceptable at low temperatures, but for higher temp operations I recommend looking into a sand, air, oil, or salt bath if you do not have access to a heating mantle.

Edit: Looks like Sulaiman beat me to the punch!

[Edited on 3-15-2016 by Zephyr ]

Maker - 16-3-2016 at 09:15

That makes sense. :D Oil bath sounds like the way forward.

I've seen a couple of other types of flask at school too, one is pear shaped which is apparently to allow one to scrape thick/solid residues off the walls and another shaped like an inverted pear (Like an RBF but with a pointed bottom). What are they both called and what is the latter one for?

Arthur Dent - 20-3-2016 at 05:54

Hi all,

Quick question about rust. Just acquired a piece of equipment with rusty bolts and nuts... They are custom bolts and can't be replaced, so I removed all of them and dipped them in rust remover solution (diluted phosphoric acid) which successfully removed all the rust... rinsed the nuts in hot water to remove all traces of phosphoric acid, dried them thoroughly with hot air, and then dropped the nuts in a jar filled with WD40... a day later, i patted dry all the hardware.

Before reinstalling the nuts, is there something I should do to prevent oxidation from coming back since they are now down to bare metal with no protective coating on them? I would rather not paint them since the nuts and bolts will be screwed on/off quite often.

I remember a long time ago, I had a spray can lubricant called Moly-kote which protected quite well bare metal parts but on the down side, was a stain hazard for fabrics...

Any suggestions for non-staining metal protection? Thanks!

Arthur Dent - 20-3-2016 at 06:00

Here's an image of the parts in question, they're rack-mount brackets, which originally had a golden iridescent coating (cadmium?) but it had started rusting at the edges and close to the threads. The phosphoric acid bath brought it down to the bare gray metal.

IMG_2198.jpg - 61kB

Maker - 20-3-2016 at 06:26

I would put a bolt through them (To protect the threads) and give them a going over with some black rattlecan.

alexleyenda - 21-3-2016 at 14:52

In class, I have prepared a Nickel complex where I change 2 CH3COO ligands for 2 NO2 ligands. The procedure asks us to mix NaNO2 with NH4CH3COO and then add it to the complexe. I don't understand why we don't just use directly NaNO2 instead of making NH4NO2 ??

What can I do with Rosemary?

j_sum1 - 28-3-2016 at 03:09

I have a large rosemary bush at my place. I like to keep it trimmed down to six foot but it is again reaching the eaves of my house and so needs 3-4 foot chopped off (yet again).

It tastes great by the way -- but obviously much more than can be used for culinary purposes.

I recall reading recently of some interesting compounds that might be extracted from the plant, but I cannot remember what they are. In the absence of any better information I thought I'd begin with a steam distillation of the "essential oil". (I haven't done a steam distill before so it will be a good exercise.)

Are there any good syths or extractions I can do from the oil or from the leaves directly?

Morgan - 28-3-2016 at 06:16

Why do the Chinese use this ubiquitous shape to display so many products?
http://www.alibaba.com/product-detail/Original-Taste-Pure-In...

PHILOU Zrealone - 28-3-2016 at 06:52

Quote: Originally posted by j_sum1  
I have a large rosemary bush at my place. I like to keep it trimmed down to six foot but it is again reaching the eaves of my house and so needs 3-4 foot chopped off (yet again).

It tastes great by the way -- but obviously much more than can be used for culinary purposes.

I recall reading recently of some interesting compounds that might be extracted from the plant, but I cannot remember what they are. In the absence of any better information I thought I'd begin with a steam distillation of the "essential oil". (I haven't done a steam distill before so it will be a good exercise.)

Are there any good syths or extractions I can do from the oil or from the leaves directly?

Essential oils of the terpenic familly found in Rosemary:
Borneol
Cineol
Eucalyptol
Camphene
Pinene

UC235 - 28-3-2016 at 07:05

Quote: Originally posted by alexleyenda  
In class, I have prepared a Nickel complex where I change 2 CH3COO ligands for 2 NO2 ligands. The procedure asks us to mix NaNO2 with NH4CH3COO and then add it to the complexe. I don't understand why we don't just use directly NaNO2 instead of making NH4NO2 ??


Probably to buffer the reaction pH. Ammonium acetate is a pH=7 buffer
Quote: Originally posted by j_sum1  
I recall reading recently of some interesting compounds that might be extracted from the plant, but I cannot remember what they are. In the absence of any better information I thought I'd begin with a steam distillation of the "essential oil". (I haven't done a steam distill before so it will be a good exercise.)

Are there any good syths or extractions I can do from the oil or from the leaves directly?


Possibly not a great paper, but the composition of most rosemary essential oil is probably pretty similar. Additionally, rosemary leaf is about 1.5-2.5% essential oil when fresh, so a very large amount of material must be steam distilled to produce a reasonable yield. It's certainly better than some other materials used for essential oils, though.

http://idosi.org/aejaes/jaes5(1)/13.pdf

You'll note that the main components are alpha-pinene, camphene, 1,8-cineole, and camphor. It's a pretty complex mixture without a dominant component and the components would all be much more easily available from other essential oils. Camphor can just be purchased pure, pinene makes up the majority of turpentine, eucalyptus globulus oil is almost all 1,8-cineole, and nutmeg essential oil is mostly camphene.

arkoma - 30-3-2016 at 09:34

damned ferlilizer I picked up the other days is giving me fits--the prills have a green dye on them and if I boil a Urea solution do I risk getting a bunch of biuret contamination?

chemrox - 30-3-2016 at 14:37

This thread could have been easier to find...now here's the dumbass question: when would I resort to "computational chemistry?" Is it a biotech thing? Do any of you use it and what sorts of problems can you address with it?

[Edited on 30-3-2016 by chemrox]

gluon47 - 3-4-2016 at 00:30

does this compound exist?

no matches for pubchem and chemspider.

any help would be appreciated.

furan.jpg - 571kB

HeYBrO - 3-4-2016 at 01:15

no i doubt it, the doubly bonded carbon ( called a cumulated diene or allenes) is sp2 so would be linear and hence there would be significant ring strain and the bond to N and C being stretched would make it very unstable. any one else want to add?

[Edited on 3-4-2016 by HeYBrO]

gluon47 - 3-4-2016 at 11:05

Ok, thanks for your help

CharlieA - 3-4-2016 at 18:10

Quote: Originally posted by HeYBrO  
no i doubt it, the doubly bonded carbon ( called a cumulated diene or allenes) is sp2 so would be linear and hence there would be significant ring strain and the bond to N and C being stretched would make it very unstable. any one else want to add?

I have to go with you on this one. With my framework Darling models I couldn't get the ring to close with a sledgehammer... and I didn't want to break my models.

Ramium - 3-4-2016 at 20:45

Would diacetone alcohol work as a catalyst for the production of potassium from potassium hydroxide and magnesium? Its a tertiary alcohol just like other common catalysts such as tert-butanol.

Any ideas?

Darkstar - 3-4-2016 at 23:49

Quote: Originally posted by Ramium  
Would diacetone alcohol work as a catalyst for the production of potassium from potassium hydroxide and magnesium? Its a tertiary alcohol just like other common catalysts such as tert-butanol.


There are two problems here. First, diacetone alcohol has a carbonyl group that would get reduced to a ketyl by both magnesium as well as potassium. Second, the highly basic conditions would just deprotonate the carbon between the carbonyl and hydroxyl group, dehydrating the diacetone alcohol to mesityl oxide, which now has a conjugated pi-system that is even more prone to reduction.

Ramium - 4-4-2016 at 00:29

Oh ok. Thats unfortunate. Thanks anyway

CharlieA - 4-4-2016 at 05:44

Quote: Originally posted by CharlieA  
Quote: Originally posted by HeYBrO  
no i doubt it, the doubly bonded carbon ( called a cumulated diene or allenes) is sp2 so would be linear and hence there would be significant ring strain and the bond to N and C being stretched would make it very unstable. any one else want to add?

I have to go with you on this one. With my framework Darling models I couldn't get the ring to close with a sledgehammer... and I didn't want to break my models.




I tried to connect the allene moiety with models this morning. The best I could do was an 8-membered ring containing the allene oiety and 1 tetrahedral N and 4 tetrahedral carbons, and this structure also seemed to have a little strain.

One more question

Ramium - 4-4-2016 at 20:58

Would chlorbutol (trichloro-2-methyl-2-propanol) work as a catalyst for potassium production from KOH and Mg?

Its a tertiary. Maybe the three chlorine atoms would be a problem?

Acetic acid

Romix - 5-4-2016 at 04:34

Can acetic acid be made from recrystallized copper acetate?

By displacing copper with lead, and adding HCl to it?

Last time I tried, it didn't crystallize out in cold. Because of low concentration?

Acetate of copper seems to be more stable then chloride.


[Edited on 5-4-2016 by Romix]

Metacelsus - 5-4-2016 at 07:41

Yes, the chlorine atoms would be a problem, as you noticed.

Gold anodes

Jstuyfzand - 5-4-2016 at 11:18

Can gold be used as an anode in brine electrolysis or things like copper sulphate?
I can get gold plated things from electronics, its much more common than platinum.
I tried a small gold plated wire, it turned black!
Thoughts?

Thanks!

Darkstar - 5-4-2016 at 21:53

Quote: Originally posted by Ramium  
Would chlorbutol (trichloro-2-methyl-2-propanol) work as a catalyst for potassium production from KOH and Mg?

Its a tertiary. Maybe the three chlorine atoms would be a problem?


Yes, they would be a problem, as already mentioned. Keep in mind that magnesium and potassium are both extremely powerful reducing agents. This means you're going to need an alcohol that doesn't have any easily reducible functional groups or they will react with both your initial magnesium metal, as well as any potassium metal that does form. The reaction conditions are also going to be highly basic as well, so your alcohol can't have a bunch of acidic protons that will get removed by the hydroxide and alkoxide ions, or by reduction from the magnesium or potassium metal. (although, in this case, there aren't any)

Fischer Esterification

Bean - 8-4-2016 at 09:03

I was wondering if sulfamic acid could be used as a catalyst for Fischer esterification or even making ethers such as diethyl ether. Does anyone have experience using this. According to Wikipedia it can be used for esterification, but I am unsure if this applies to all reactions for example between ethanol and ethanoic acid etc. It also has a pKa of 1.0, is this suitable. Also, apparently when it reacts with alcohols organosulfates are formed- could this allow it to be used to produce ethers?

Thanks.

[Edited on 8-4-2016 by Bean]

Calcium Nitrate

IceDahl - 8-4-2016 at 09:31

I found something that I find quite strange. Calcium nitrate is rated "3" on the Instability/reactivity part of the Fire Diamond. While potassium nitrate and sodium nitrate is 0 on the yellow.

The "3" on the yellow is desctibed as: "Capable of detonation or explosive decomposition but requires a strong initiating source, must be heated under confinement before initiation, reacts explosively with water, or will detonate if severely shocked".

Is calcium nitrate really so unstable compared to sodium or potassim nitrate?
I have heated small amounts of calcium nitrate with a butane torch before but nothing happened.

Mabus - 8-4-2016 at 09:52

Is that Fire Diamond used for anhydrous Ca(NO3)2 or for the tetrahydrate form?
Because I can see reason for the anhydrous form, but not for the hydrate.

UC235 - 8-4-2016 at 10:15

No, sulfamic acid undergoes solvolysis in alcohols to yield ammonium alkylsulfates which are only a slightly acidic salt. It's an interesting prep but not too useful.

Even if it didn't undergo solvolysis, for Fischer esterification, you really need a strong acid to have a good reaction rate. The acid needs to protonate the alcohol. Alkyloxoniums have a pKa of ~-1.7 so you ideally want an acid with a lower pKa than that to achieve significant quantities of the intermediate.

Bean - 8-4-2016 at 11:44

Thanks for the reply.

The Volatile Chemist - 9-4-2016 at 17:30

I recently dissolved some 99% Sn solder in HCl. All that's left out of solution is the resin core (the rest was 1% Pb). Is it feasible to crystallize out a chloride of tin, or should I just precipitate it out as a hydroxide or basic carbonate or something...?

Bean - 10-4-2016 at 03:13

Sodium hydrogen sulfate (bisuslfate) as Fischer esterification catalyst ?

Thanks in advance,
Bean.

Deathunter88 - 10-4-2016 at 04:10

Quote: Originally posted by The Volatile Chemist  
I recently dissolved some 99% Sn solder in HCl. All that's left out of solution is the resin core (the rest was 1% Pb). Is it feasible to crystallize out a chloride of tin, or should I just precipitate it out as a hydroxide or basic carbonate or something...?


Yes, it is very possible to crystallise out tin(II) chloride if you leave some excess tin in solution and wait for the evaporation. (no heating!)

S.C. Wack - 10-4-2016 at 06:57

Heating isn't necessarily an issue. Blanchard gives directions that depend on heating (among other things), as the dihydrate crystallizes out of the saturated solution on cooling. Vaporization by air will provide acidic air, dealt with in whichever way.

The Volatile Chemist - 12-4-2016 at 14:01

Quote: Originally posted by S.C. Wack  
Heating isn't necessarily an issue. Blanchard gives directions that depend on heating (among other things), as the dihydrate crystallizes out of the saturated solution on cooling. Vaporization by air will provide acidic air, dealt with in whichever way.

Hmm, makes sense I guess, but I've seen many other metal solutions hydrolyzed (I think that's canon usage), so I'll try it with a little product first.

yobbo II - 15-4-2016 at 14:39

(18) If f (−1) = −7 and f (x) = g(−6 · x), what point must satisfy y = g(x)?

Can someone do this?

TIA

aga - 15-4-2016 at 14:55

Twelvety, obviously.

OK. This has been happening a lot recently.

It's time the wizard responsible owns up to synthesising a Hormunculus, again.

Yes, the Globe Bottles and any dried Mouse Blood will be taken away if you do, but that is all.

If not, and we Will find you, your bodkin shall be severely shriven.

yobbo II - 16-4-2016 at 13:59


This dude says sixthedy.
http://openstudy.com/updates/53b95693e4b09f140ca68b23

Thanks googness shakespeareannn is my first language.

Yob (licenced to diff.)

Dr.Q - 17-4-2016 at 23:08

What is the solubility of formaldehyde in isopropanol and ethyl acetate

Strange yellow precipitate in TiCl2 solution

j_sum1 - 24-4-2016 at 03:12

I have made some TiCl2 for a series of experiments in Ti chemistry. I added some Ti powder to HCl and let it stew for a few days. Reaction was vigorous at first and I overflowed the flask I was using. I was left with a deep purple acidic solution of titanium chloride with a sediment of unreacted Ti powder.

The thing that has me puzzled is a bright yellow sediment that stuck to the flask and also appeared in the filter paper. It looks kind of reminiscent of sulfur but maybe a shade less brilliant yellow. It stuck to the flask and was difficult to clean off. Mechanical scrubbing did almost nothing. I tried an oxidising environment and a strongly reducing environment and it failed to budge. Eventually I got it to dissolve in a concentrated NaOH solution with a bit of warming.

Obviously I have some impurities. The HCl used was from the hardware store -- crystal clear with no detectable iron. I have never had issues with it before. The Ti powder was from a chemical supplier but no assay given. I am going to guess that whatever impurities present came in the Ti.

So, the next question is what is this stuff?
If I was to guess I would say that I have a vanadium compound in the V oxidation state. But which vanadium compounds are likely? I can't really see that it is pentoxide since it precipitated in strongly acidic conditions. Or maybe it is another transition element compound. Tungsten, molybdenum and niobium all have some yellow compounds. Any suggestions? Any tests that I could do? (I have a kg of the Ti powder and so it won't be too hard to get more of the precipitate.)

j_sum1 - 26-4-2016 at 17:24

Bump to my previous question. Paging woelen if he is around.


New (unrelated) question.
I recall reading a nice little pdf on this site on identification of polymers using burn tests. I thought I had saved it but can't locate it. My google-fu is failing me too. Does anyone know of a good little manual for polymer identification?

question

Ramium - 27-4-2016 at 00:17

Does anyone know if copper forms a complex with isopropylamine?

if so, how would one prepare this complex. Maybe addition of freebase isopropylamine to copper sulphate solution? would the complex from as a precipitate?

DraconicAcid - 27-4-2016 at 07:40

Quote: Originally posted by Ramium  
Does anyone know if copper forms a complex with isopropylamine?

if so, how would one prepare this complex. Maybe addition of freebase isopropylamine to copper sulphate solution? would the complex from as a precipitate?


It should. Add isopropylamine to a concentrated solution of copper(II) sulphate- you will get an precipitate of copper(II) hydroxide at first, and then excess amine will dissolve the precipitate (if it indeed forms a complex). If it stays in solution (and I expect it would be soluble), add alcohol to precipitate the complex.

Metacelsus - 27-4-2016 at 08:03

I know that it forms a complex with ammonia and with butylamine (I've made both, with chloride as counterion), so I don't see why it wouldn't form one with isopropylamine. I don't think the complex would precipitate after it formed. You might be able to add a countersolvent to get it to precipitate (as DraconicAcid suggested).

[Edited on 4-28-2016 by Metacelsus]

Ramium - 27-4-2016 at 11:53

Ok, thanks guys. I'll definitely try making it

Figuring out molar equivelancies with various hydrates

RogueRose - 28-4-2016 at 17:03

I need some help to make sure I am calculating molar equivalencies correctly with hydrates.

Sodium carbonate has an anhydrous density of 2.54g/cm^3 and a monohydrate of 2.25g/cm^3

molar mass of anhydrous = 106g/mole

anhydrous
Na2CO3 -> 2(23) + 12 + 3(16) = 106

Monohydrate
Na2CO3 + H2O -> (2(23) + 12 + 3(16)) + ((2)1 + 16) = 124


5 mole = 530g anhydrous
5 mole = 620g monohydrate

Is this correct?

Does the volume of liquid added factor in as then solubility and density will come into play.



[Edited on 29-4-2016 by RogueRose]

blogfast25 - 28-4-2016 at 17:29

Quote: Originally posted by RogueRose  


molar mass of anhydrous = 106g/mole

anhydrous
Na2CO3 -> 2(23) + 12 + 3(16) = 106

Monohydrate
Na2CO3 + H2O -> (2(23) + 12 + 3(16)) + ((2)1 + 16) = 124


5 mole = 530g anhydrous
5 mole = 620g monohydrate

Is this correct?

Does the volume of liquid added factor in as then solubility and density will come into play.


Yes, correct.

Don't understand your last point/question though...

a nitrogen rich explosive - 29-4-2016 at 05:38

Would IPA form a complex with copper perchlorate much like HMTD?

Romix - 29-4-2016 at 10:42

Tin hydroxide dissolving in excess NaOH?

xfusion44 - 29-4-2016 at 13:08

What would be the best option to separate a mixture of isobutanol, methanol, 2-butoxyethanol, acetone and methyl acetate? I don't have fractionating column, but even if i had it, it would probably be impossible to separate acetone and methyl acetate with only 0.4degC difference in boiling point. What about extractive distillation? How do I know which of those compounds will form azeotrope and how to get rid of that?

Thanks

DraconicAcid - 29-4-2016 at 13:12

Quote: Originally posted by xfusion44  
What would be the best option to separate a mixture of isobutanol, methanol, 2-butoxyethanol, acetone and methyl acetate? I don't have fractionating column, but even if i had it, it would probably be impossible to separate acetone and methyl acetate with only 0.4degC difference in boiling point. What about extractive distillation? How do I know which of those compounds will form azeotrope and how to get rid of that?

Thanks


You could get rid of the methyl acetate by boiling the mixture with aqueous sodium hydroxide.

xfusion44 - 29-4-2016 at 18:26

Quote: Originally posted by DraconicAcid  
Quote: Originally posted by xfusion44  
What would be the best option to separate a mixture of isobutanol, methanol, 2-butoxyethanol, acetone and methyl acetate? I don't have fractionating column, but even if i had it, it would probably be impossible to separate acetone and methyl acetate with only 0.4degC difference in boiling point. What about extractive distillation? How do I know which of those compounds will form azeotrope and how to get rid of that?

Thanks


You could get rid of the methyl acetate by boiling the mixture with aqueous sodium hydroxide.


Does that mean I'd destroy it or that I'd be able to distill it after that? I forgot to mention that I want to keep all of these solvents if possible - at least methanol and methyl acetate, because I don't have those two and the other ones are all pretty much in lower quantities. BTW, the mix of those solvents is actually a paint stripper - I've already separated the toluene out of it, but I don't know how to proceed with others.

Also, wouldn't NaOH also react with other solvents? At least partially?

I was thinking about adding another solvent or salt, to make some of those solvents in the paint stripper less soluble or to make them boil at higher temperatures and thus increasing the difference between boiling points, but I don't know much about that.

Thanks

DraconicAcid - 29-4-2016 at 19:20

That would destroy the methyl acetate, or at least convert it into methanol and nonvolatile sodium acetate. Sodium hydroxide won't react with the other solvents.

gdflp - 29-4-2016 at 19:30

Quote: Originally posted by DraconicAcid  
Sodium hydroxide won't react with the other solvents.

Hot sodium hydroxide will cause the acetone to undergo an aldol condensation with itself.

[Edited on 4-30-2016 by gdflp]

xfusion44 - 29-4-2016 at 23:02

So, would the resulting diacetone alcohol dehydrate to form mesityl oxide? And; would diacetone alcohol or mesityl oxide be beneficial in terms of distilling that mixture of solvents or would it just make it harder?

Both diacetone alcohol and mesityl oxide have much higher b.p. than methanol, so I guess that'd be fine?

How concentrated should NaOH be? (there is already a lot of water in that mixture, I forgot to say) And how long should it boil? If I understand correctly, this should be done with reflux setup?

Mabus - 30-4-2016 at 11:44

Just noticed the largest urea crystals I grew are hollow. Anyone can tell me if it's normal? Cause I don't know anything about how urea crystallizes.

question

Ba(ClO3)2 - 30-4-2016 at 20:09

we've been thinking about making a bit of ethyl bromide.

does it have to be stored in the lab freezer?

why is it usually stored this way?

is it ok to store it for any length of time without a freezer in a bottle?

thanks in advance

Metacelsus - 1-5-2016 at 07:05

Ethyl bromide is quite volatile (vapor pressure 52 kPa at 20 C, boiling point 38 C), and definitely not something you want to breathe. Keeping it cold means less will evaporate. However, as long as you have a good seal on your storage container, it's probably fine to store at room temperature.

another question

Ramium - 1-5-2016 at 21:25

I'm trying to make semicarbazide hydrochloride from hydrazine sulphate.

I found one procedure of the synthesis using hydrazine sulphate and potassium cyanate.

http://www.prepchem.com/synthesis-of-semicarbazide-hydrochlo...

But I don't have potassium cyanate. Would sodium cyanate work if I recalculated the stoichiometry?

Metacelsus - 2-5-2016 at 05:50

Yes.

alternative of Sodium Borohydrate (NaBH4)

veerenyadav - 3-5-2016 at 02:31

I want to reduce Mn+2 to produce Mn 0, can anybody suggest cheap alternative of NaBH4 ?:P

j_sum1 - 3-5-2016 at 02:42

Why?

Mn is not too tricky to get.
With care, it is possible to reduce MnO2 to the metal using Al in a thermite reaction.

(And I think it is sodium borohydride that you were referring to.)

a nitrogen rich explosive - 3-5-2016 at 13:57

MnO2 is inside batteries.
A good thermite composition here:
30% MnO2
30% Al powder
20% acetone peroxide
10% azidotetrazole
10% neat hydrazine

Get what I mean? :D :D :D

[Edited on 3-5-2016 by a nitrogen rich explosive]

Hydroquinone, metol and millon's reagent

j_sum1 - 4-5-2016 at 22:06

I have the opportunity to save these three reagents from the throw-out pile and am wondering if it is worth doing so.

I look at hydroquinone and metol and see a doorway to the world of phenols but have insufficient OC knowledge to step through that door. I guess a bit of reading is needed. Suggestions? Possible simple synths for an OC noob?

Millon's reagent is simply Hg dissolved in nitric acid and diluted. I have no idea of the concentration of the bottle I have -- only that there is more than half a litre of the stuff. On reading it seems that it is used as an indicator for proteins and that it interacts with phenol groups. I am not going there. It looks to me like its intended use lands the experimenter right in the middle of a mess of organomercury compounds. But if I stick to inorganic chemistry, is there something that I might do with this? I guess I could reduce back to metallic mercury but if that is my goal, I might as well break a thermometer and avoid the hassle. Are there any other sensible possibilities for this reagent?

Solubility of various chemicals

a nitrogen rich explosive - 5-5-2016 at 00:52

Solubility for the following, please:
Maltodextrin
Silicon dioxide (the type found in pills)
Aminoguanidine hydrochloride

I need aminoguanidine salts for synthesising tetrazoles, but the usual aminoguanidium bicarbonate synthesis is a very long process and I am challenging myself to synthesis azidoazide azide (1,1-azobis(tetrazole)) from OTC materials.

j_sum1 - 5-5-2016 at 01:31

Silicon dioxide is insoluble by any standards.

a nitrogen rich explosive - 5-5-2016 at 02:29

Thanks. So is magnesium stearate. And from what I know, maltodextrin is only capable of forming a suspension with water. Is aminoguanidine hydrochloride also insoluble (hopefully, it isn't.)

Loptr - 5-5-2016 at 08:26

Are you getting this from a pill? What pill contains aminoguanidine HCl?

a nitrogen rich explosive - 5-5-2016 at 08:28

These ones:

http://www.supersmart.com/en--Blood-Sugar-Glycation--Aminogu...

Chem Rage - 5-5-2016 at 10:23

Quote: Originally posted by a nitrogen rich explosive  
These ones:

http://www.supersmart.com/en--Blood-Sugar-Glycation--Aminogu...


Why not just open up an account with Sigma Aldrich and order from them? Saves the hassle of dismantling gelatine capsules, lol!

Here you go: http://www.sigmaaldrich.com/catalog/product/aldrich/396494?l...

Chem Rage - 5-5-2016 at 10:24

Quote: Originally posted by a nitrogen rich explosive  
These ones:

http://www.supersmart.com/en--Blood-Sugar-Glycation--Aminogu...


Why not just open up an account with Sigma Aldrich and order from them? Saves the hassle of dismantling gelatine capsules, lol!

Here you go: http://www.sigmaaldrich.com/catalog/product/aldrich/396494?l...

a nitrogen rich explosive - 5-5-2016 at 11:31

Aldrich doesn't sell to hobbyists, no matter what.

aga - 5-5-2016 at 11:48

Sigma Aldrich Sells according to a System/Procedure.

Fundamentally they are a sales outfit, just with some restrictions on Who they sell to, What they sell, and in what quantitity.

As with all Systems, figure out how that system works, then approach it along the correct line that will not cause that System any concern, and it will operate according to it's design.

Ebay is a lot easier.

[Edited on 5-5-2016 by aga]

PHILOU Zrealone - 6-5-2016 at 09:16

Quote: Originally posted by a nitrogen rich explosive  
Would IPA form a complex with copper perchlorate much like HMTD?

IPAmine with Cu(II) perchlorate --> of course.

Did you meant HMTD (hexamethylene triperoxyde diamine) or HMTA (hexamethylenetetramine)? I think that once again you did confuse the two!

What complex with Cu(II) perchlorate?
The one from Laboratory of Liptakov? Then it is HMTA!

--> Think twice³ before posting!

[Edited on 6-5-2016 by PHILOU Zrealone]

PHILOU Zrealone - 6-5-2016 at 09:21

Quote: Originally posted by a nitrogen rich explosive  
MnO2 is inside batteries.
A good thermite composition here:
30% MnO2
30% Al powder
20% acetone peroxide
10% azidotetrazole
10% neat hydrazine

Get what I mean? :D :D :D

[Edited on 3-5-2016 by a nitrogen rich explosive]

Stop doing this!:mad:

The two first ingredients are indeed a thermite...
But the rest is wrong info (misinformation) to make someone innocent hurt himself...and set the surrounding in fire because of the gaseous blast of incandescent molten metallic drops...

PHILOU Zrealone - 6-5-2016 at 09:28

Quote: Originally posted by a nitrogen rich explosive  
Solubility for the following, please:
Maltodextrin
Silicon dioxide (the type found in pills)
Aminoguanidine hydrochloride

I need aminoguanidine salts for synthesising tetrazoles, but the usual aminoguanidium bicarbonate synthesis is a very long process and I am challenging myself to synthesis azidoazide azide (1,1-azobis(tetrazole)) from OTC materials.

If you don't say in what solvent and under what conditions, hard for someone to answer!

Silicon dioxyde is soluble in HF, in NaOH, in N2H4, in concentrated Na2CO3 and some form (Si(OH)4) into water.

SiO2 is soluble in molten Al2O3...

DraconicAcid - 6-5-2016 at 09:47

Quote: Originally posted by PHILOU Zrealone  


Silicon dioxyde is soluble in HF, in NaOH, in N2H4, in concentrated Na2CO3 and some form (Si(OH)4) into water.

SiO2 is soluble in molten Al2O3...


Silicon dioxide is soluble in hydrazine? Sez who?

NaOH + H2O2 - smell emitted - WTF is it!?

RogueRose - 7-5-2016 at 12:43

I mixed some ~30% H202 with a 50/50 NaOH/H2O solution. A frothy white foam formed in a vigorous exothermic reaction. The smell that was emitted is pretty vulgar and triggered my gag reflex, pretty severely. Any idea what the cause of the smell may be? I have no problems with mixing strong NaOH solutions (gives off smell for sure) nor do I have a problem with sodium salts in general. Ideas?

Metacelsus - 7-5-2016 at 13:10

Aerosol of sodium hydroxide. (This has been the discussion topic of numerous other threads.)

How safe would Th(ClO3)4 flash powder be?

Ba(ClO3)2 - 10-5-2016 at 19:44

We are thinking of making a bit of thorium chlorate.

Would setting off a small amount of thorium chlorate-magnesium flash powder be extremely dangerous?

We were concerned about the radioactive smoke that would likely be produced.

Any suggestions wellcome

j_sum1 - 10-5-2016 at 21:13

Seriously, why?

Thorium is difficult enough to get hold of -- I have no idea why you would want to perform a reaction where you could not recover it. As for releasing it into the environment: particularly in a finely divided form that could be breathed in -- I would not even consider it. I am not sure the exact consequence of breathing in radioactive thorium, but it is not something that is on my wish list.


If it is more than a curiosity -- that is, you want to do some actual science...
set off your flash powder in a closed vessel under vacuum with a filter to catch all of the product. Perform a before and after with a geiger counter to make sure all your thorium is where you think it is. Plan how you will reprocess it and calculate the yield / losses at each stage.

Ba(ClO3)2 - 10-5-2016 at 22:08

Sounds like more trouble than its worth. We were just curious what it would be like.
We dont think we'll try it now. You make a good point, the thorium is hard to come by.
It would be a wast.

But Purerly theoretically, do you think the flash powder would work?

We gather it would based on the last section of your post.

j_sum1 - 10-5-2016 at 22:20

Dunno.
I would want to think about the reactants and products carefully and work out if it is thermodynamically viable. I do know that thorium-anything converting to metal is a serious uphill battle. But I don't really have any idea about chlorates in this context.

Paging blogfast25. He would know: and probably without looking anything up.

Ba(ClO3)2 - 10-5-2016 at 22:39

ok thanks

Eosin Y - 11-5-2016 at 10:18

What is the pH of periodic acid? I have found some for sale at OnyxMet, and I want to try some simple reactions with it.
NaAcetate + metaperiodic acid = sodium periodate + acetic acid. Rubbish or not?

DraconicAcid - 11-5-2016 at 10:45

Quote: Originally posted by Eosin Y  
What is the pH of periodic acid? I have found some for sale at OnyxMet, and I want to try some simple reactions with it.
NaAcetate + metaperiodic acid = sodium periodate + acetic acid. Rubbish or not?


The CRC gives its Ka as 2.3e-2, which is fairly strong. Acetate ion should deprotonate it readily.

Ramium - 13-5-2016 at 21:07

Does manganese (iii) oxide (Mn2O3) react with hydrochloric acid?

I can't find any information on this.

DraconicAcid - 14-5-2016 at 00:29

Quote: Originally posted by Ramium  
Does manganese (iii) oxide (Mn2O3) react with hydrochloric acid?

I can't find any information on this.


Depends on how well-dried it is. If it's fresh, it will react easily. If it's calcined completely, then it's a rock.

glymes - 16-5-2016 at 12:59

Why is it that there are a lot of organic nitrates (nitrotoluene, nitrobenzene etc.) but virtually no chlorates of the same ilk? This interests me.
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