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PHILOU Zrealone
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Quote: Originally posted by mayko  | | Why does acetaldehyde seem to be so hard to come by? It don't know of it being a precursor, and I don't think it's as toxic as the widely-available
formaldehyde |
It has some use as cyclic tetramer (metaldehyd) pesticide for mollusces but toxic for childrens, cats, dogs,...neurotoxic.
In Europa, those pesticides have lowered their active % to diminish the risk of poisoning from 80% to less than 5%.
The cyclic trimer (paraldehyd) is used as hypnotic/sedative, also neurotoxic.
This might explain the difficulty to obtain it.
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
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The Volatile Chemist
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Yes...indeed. Is it still common in the scholarly lab, or is it one of the many one has to make? I don't recall seeing it in Vogel's.
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Eddygp
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Is there any feasible way for a home chemist to substitute a methyl group attached to a nitrogen (N-methyl) to a 2-propynyl group (N-(2-propynyl))?
[Edited on 17-9-2015 by Eddygp]
there may be bugs in gfind
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The Volatile Chemist
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Do you know of any feasible routes that the amateur 'can't' do? I can't think of anything from Vogel's or 'E-Z Organic Chemistry', but I have very
little education in such.
Also, this thread is getting rather long...
[Edited on 9-17-2015 by The Volatile Chemist]
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Crowfjord
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@Eddygp: I think it depends on the type of amine. I know that a methyl group can be removed from a tertiary amine, with the Von Braun reaction being
the classic example. The demethylated amine could probably then be reacted with 2-propynyl iodide or bromide. I don't think I have ever heard of a way
to demethylate a secondary amine.
[Edited on 17-9-2015 by Crowfjord]
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Eddygp
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| Quote: Originally posted by Crowfjord | @Eddygp: I think it depends on the type of amine. I know that a methyl group can be removed from a tertiary amine, with the Von Braun reaction being
the classic example. The demethylated amine could probably then be reacted with 2-propynyl iodide or bromide. I don't think I have ever heard of a way
to demethylate a secondary amine.
[Edited on 17-9-2015 by Crowfjord] |
I was actually thinking about the ones in cyclic compounds, either aromatic or not. So yes, mostly tertiary amines. In fact, there are a few secondary
ones in guanine that, if their H were to be substituted for an R group, they would be like the ones I mentioned. I was thinking about methylated uric
acid, actually.
[Edited on 18-9-2015 by Eddygp]
there may be bugs in gfind
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Crowfjord
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Sounds interesting. I think that the Von Braun reaction would work on such a substrate, but cyanogen bromide is pretty nasty. There is also a
demethylation reaction that employs alkaline aqueous potassium ferricyanide, which works on some tertiary amines (see JOC, 1951, 16(8), pp 1303-1307). Perhaps it could could work on your substrate.
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MeshPL
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Can amides be N-alkylated by haloalkane or DMSO4?
Can I make nitromethanes by adding lighter, but liquid hydrocarbons to an excess fuming nitric acid and maybe even bubbling some NO2 as well?
Can I make 2,5-dimethylcyclohex-2-enone from 4-methyl-5-oxohepthanal via aldol elimination? (As a note other byproducts would likely have to contain a
4-membered ring so they wouldn'be favoured. Also not sure if I named those compounds properly)
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The Volatile Chemist
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Interesting. If the ferricyanide works, the procedure is feasable for even I. I wonder if it would work on a phase transfer catalyst, like Aliquat
336.
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MeshPL
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What concerns turning methylamines into isopropylamines... maybe (if wikipedia says so...) you could turn tertiary amine into amine oxide and make it
undergo Meisenheimer rearrangement or Polonovski reaction. Maybe you could attempt to hydrolyse resulting hydroxylamine or acetamide/iminium. Look
those reactons up on the net.
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Crowfjord
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Concerning N-demethylation, it seems that that one paper I sited earlier is the only reference that uses ferricyanide. I, at least, could not find
any more references on the subject. No follow-ups or anything. This could be a ripe area for amateur research, as the study in that old paper was
not exactly thorough. I'll post the text in the references section, in Recent Journal Articles of Interest.
Also, I found this paper, for another N-demethylation method.
[Edited on 20-9-2015 by Crowfjord]
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Morkva
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A method of preparing arylsulfonic acid chlorides is to oxidize the thiol with chlorine in glacial acetic acid. If the oxygen is coming from the
acetic acid then must not the byproduct be acetyl chloride??
Either that or the CO2 elimination product methyl chloride, or am I missing something?
http://www.carrotmuseum.co.uk/falcarinol.html
“Science is the ultimate pornography, analytic activity whose main aim is to isolate objects or events from their contexts in time and space. This
obsession with the specific activity of quantified functions is what science shares with pornography.”
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The Volatile Chemist
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Are athere any common modifications to citric acid known which ar feasable to the amateur? Oxidation seems to just leadto decomposition. Obviously I'm
asking about organic reactions, not making salts of it. I don't suppose one can chlorinate the hydroxyl group and collect the product?
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Zephyr
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An easy conversion is to citrazinic acid by melting with urea, although the yields aren't great:
https://www.youtube.com/watch?v=jbw9dt12qsI
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AngelEyes
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Solubilities
So...short('ish) question...hopefully there'll be a quick answer.
When I was at school we were taught many things, some of which I even learned, but one of them was that all sodium salts are soluble and all nitrates
are soluble.
Now, it's been a while since I was in a place of learning but are those two 'facts' still correct? If so, then why sodium and nitrates...what's so
special about them? Are there any others that fall into the 'always soluble' category? And by that I mean appreciably soluble, not <1g / L
And if it's not true any more, then what salt was found that violated the rule?
Cheers
Angel
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DraconicAcid
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| Quote: Originally posted by AngelEyes | So...short('ish) question...hopefully there'll be a quick answer.
When I was at school we were taught many things, some of which I even learned, but one of them was that all sodium salts are soluble and all nitrates
are soluble.
Now, it's been a while since I was in a place of learning but are those two 'facts' still correct? If so, then why sodium and nitrates...what's so
special about them? Are there any others that fall into the 'always soluble' category? And by that I mean appreciably soluble, not <1g / L
And if it's not true any more, then what salt was found that violated the rule? |
It's not that any new salt was found that violated the rule, it's just that the exceptions are unimportant for first-year chemists (or high school
students). I teach my students that all alkali metal salts are soluble, despite knowing that there are insoluble potassium salts (perchlorate and
chloroplatinate) and insoluble sodium salts (sodium zinc uranyl acetate hexahydrate, according to my 1961 copy of Vogel), and that all nitrates are
soluble, despite knowing that the green isomer of [Co(en)2Cl2]NO3 is insoluble.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Brain&Force
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I have a sizable chunk of dysprosium ( http://www.imgur.com/BD86qse ) but I need a way to break it into manageable chunks, preferably without
losing any small pieces. Wikipedia said it should be possible to cut it with a knife but that hasn't worked at all. Got any ideas on how to
efficiently break this into ~5 more manageable chunks?
[Edited on 23.9.2015 by Brain&Force]
At the end of the day, simulating atoms doesn't beat working with the real things...
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PHILOU Zrealone
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| Quote: Originally posted by The Volatile Chemist | | Are athere any common modifications to citric acid known which ar feasable to the amateur? Oxidation seems to just leadto decomposition. Obviously I'm
asking about organic reactions, not making salts of it. I don't suppose one can chlorinate the hydroxyl group and collect the product?
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Look at:
-citraconic acid (HO2C-CH=C(CH3)-CO2H)
-aconitic acid (HO2C-CH=C(CO2H)-CH2-CO2H)
-acetonedicarboxylic acid (HO2C-CH2-CO-CH2-CO2H)
All made from citric acid and common chems or processes.
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
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The Volatile Chemist
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I saw that, and tried it... It all turned to tar... And I couldn't get anything to precipitate...
| Quote: Originally posted by PHILOU Zrealone | | Quote: Originally posted by The Volatile Chemist | | Are athere any common modifications to citric acid known which ar feasable to the amateur? Oxidation seems to just leadto decomposition. Obviously I'm
asking about organic reactions, not making salts of it. I don't suppose one can chlorinate the hydroxyl group and collect the product?
|
Look at:
-citraconic acid (HO2C-CH=C(CH3)-CO2H)
-aconitic acid (HO2C-CH=C(CO2H)-CH2-CO2H)
-acetonedicarboxylic acid (HO2C-CH2-CO-CH2-CO2H)
All made from citric acid and common chems or processes. |
Thanks, I might try them.
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gluon47
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Can nitric acid be feasibly produced via the reaction of potassium nitrate and hydrochloric acid? I've seen many vids of people making it with
sulphuric acid and potassium nitrate but none using hydrochloric acid.
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j_sum1
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Add some copper and your answer is yes.
Lower yield. Lower concentration possible. And in many parts of the world, sulfuric acid is more readily available than hydrochloric (but not where I
live). That's why it is used.
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j_sum1
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The pic below shows a nice gas discharge tube.
But what exactly is causing it to glow? How might I construct something like this?
[Edited on 10-10-2015 by j_sum1]
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The Volatile Chemist
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That's cool! I assume it's argon filled. I totally forget how that happens, though I'm sure you could do something similar with a fluorescent bulb
tube placed next to a source of charge.
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j_sum1
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Bump.
And another question...
I was demonstrating a precipitation reaction to my class today -- as a means of separating. I mixed up some copper sulfate pentahydrate and sodium
chloride and then dissolved it. I challenged the students to select something that would cause a copper compound to precipitate. Sodium sulfide was
chosen -- which I thought was a good idea. Copper (II) sulfide is nice and insoluble and I thought would filter out well.
What I was not expecting is that the filtrate was a lemon yellow colour. My first thought was that the sulfide had reduced the copper ions to Cu(I).
I added a little peroxide to see if Cu would oxidise back to a blue colour. Instead, the solution went clear. I am confident that there is no Cu in
the solution. But I am unsure what might be giving the yellow colour. Maybe some Cl-S-O anion that I am unfamiliar with?
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The Volatile Chemist
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| Quote: Originally posted by Brain&Force | I have a sizable chunk of dysprosium ( http://www.imgur.com/BD86qse ) but I need a way to break it into manageable chunks, preferably without
losing any small pieces. Wikipedia said it should be possible to cut it with a knife but that hasn't worked at all. Got any ideas on how to
efficiently break this into ~5 more manageable chunks?
[Edited on 23.9.2015 by Brain&Force] |
heat it up and trow it in an ice bath?  Sorry, my ideas always manage to be
terrible, but it seems it might work if Dy has a high thermal expansion coefficient. For $42 here's your answer...(http://www.sciencedirect.com/science/article/pii/00219614849...)
This article has expansion data for the rare earth oxides (and is thus interesting, worth posting) but n/a (http://www.osti.gov/scitech/servlets/purl/4840970/).
The coefficient is here (https://books.google.com/books?id=SFD30BvPBhoC&pg=PA160&...) and it looks rather small, probably not enough to shatter it, but who knows?
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