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[*] posted on 24-4-2016 at 03:12
Strange yellow precipitate in TiCl2 solution


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.)




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[*] posted on 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?




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[*] posted on 27-4-2016 at 00:17
question


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?




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[*] posted on 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.




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[*] posted on 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]




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[*] posted on 27-4-2016 at 11:53


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



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[*] posted on 28-4-2016 at 17:03
Figuring out molar equivelancies with various hydrates


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]
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[*] posted on 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...




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[*] posted on 29-4-2016 at 05:38


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



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[*] posted on 29-4-2016 at 10:42


Tin hydroxide dissolving in excess NaOH?
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[*] posted on 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




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[*] posted on 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.




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[*] posted on 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




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[*] posted on 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.



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[*] posted on 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]




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[*] posted on 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?




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[*] posted on 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.



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[*] posted on 30-4-2016 at 20:09
question


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

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[*] posted on 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.



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[*] posted on 1-5-2016 at 21:25
another question


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?




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[*] posted on 2-5-2016 at 05:50


Yes.



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shocked.gif posted on 3-5-2016 at 02:31
alternative of Sodium Borohydrate (NaBH4)


I want to reduce Mn+2 to produce Mn 0, can anybody suggest cheap alternative of NaBH4 ?:P
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[*] posted on 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.)




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[*] posted on 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]




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[*] posted on 4-5-2016 at 22:06
Hydroquinone, metol and millon's reagent


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?




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