Sciencemadness Discussion Board - The short questions thread (1)

The short questions thread (1)

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DJF90 - 12-1-2009 at 22:16

I've worked through a thermodynamics question and cannot finish. The reaction is 1/2 N2 + 3/2 H2 <=> NH3. All I know is that 1 mole of ammonia was allowed to come into equilibrium at a given temperature and 1 atm pressure, and I have calculated the mole fractions for each component. So how do I go about finding the moles of each component?

Nicodem - 13-1-2009 at 01:47

You also need the equilibrium constant at the given temperature (http://en.wikipedia.org/wiki/Equilibrium_Constant). The equation for the equilibrium constant is derived from the reaction stoichiometry. Without this information we can not help you. Alternatively you could calculate the equilibrium constant from the free energy (deltaG=-RT*lnK), which can be calculated from the reaction enthalpy (deltaH), reaction entropy change (deltaS) and temperature (http://en.wikipedia.org/wiki/Gibbs_energy): deltaG=deltaH-TdeltaS

brew - 13-1-2009 at 04:52

It seems abit odd that it is a thermodynamic question and no temperature is given, no standard conditons etc. Perhaps it is just an equilibrium problem, and if at equilibrium Ke would have to equal 1.

Perhaps change the equation back to its origin N2 + 3H2 > 2NH3 and from there write an eqilibrium constant expression where K is one and NH3 is one mole, hence if able to, solve for the other reagants algebraically. It is late in the day, and perhaps I am wrong anyway.

DJF90 - 13-1-2009 at 08:25

Sorry I've worked through the problem, I'm just at the end of it and stuck with mole fractions that I am unable to convert into moles. Here is the full question:

Find the amounts of (i) N2, (ii) H2 and (iii) NH3 present when 1 mole of NH3 is allowed to come to equilibrium at the temperature where Kp=1 and 1 atm pressure.

Basically what I did was:

Kp = pi (x) [p(x)/p(standard)] raised to v(x), the stoichoimetric coefficent of x.

The partial pressures used were (1-X)p (where (1-X) is the mole fraction) for NH3, 0.25X for N2, and 3/4X for H2. after some algebraic manipulation, it was found that X = 0.7930..., which gives mole fractions of 0.2069, 0.1982, and 0.5947 for NH3, N2 and H2 respectively.

Now my problem is how to change these mole fractions into moles?

I know that:
n(NH3) / [n(N2) + n(H2) + n(NH3)] = 0.2069
n(N2) / [n(N2) + n(H2) + n(NH3)] = 0.1982
n(H2) / [n(N2) + n(H2) + n(NH3)] = 0.5947

But as I dont know what the sum of moles is at equilibrium (the denominator in the above fractions) I cannot work out the moles of the individual components. There is probably some logical thing I can do to work it out that is staring me in the face but I just can't see what to do. Sorry for the confusement earlier.

solo - 14-1-2009 at 17:08

I need to find the CAS number for Sodium Naphthalene, `i've found Sodium 1 naphthalenesulfonate and ` Sodium 2 naphthalenesulfonate in Aldrich I wonder if it's the same or even if these are good substitutes and the same sodium naphthalene save the different location of a sulfate ......mmmmm, is this what's called the complex of sodium naphthalene..........solo

sparkgap - 14-1-2009 at 17:53

Quote:

Originally posted by solo
I need to find the CAS number for Sodium Naphthalene, `i've found Sodium 1 naphthalenesulfonate and ` Sodium 2 naphthalenesulfonate in Aldrich I wonder if it's the same or even if these are good substitutes and the same sodium naphthalene save the different location of a sulfate ......mmmmm, is this what's called the complex of sodium naphthalene..........solo

The CAS number for sodium naphthalide (or naphthalenide, depending on who you ask) is 3481-12-7 .

sparky (~_~)

kclo4 - 15-1-2009 at 17:07

What is the most common way of determining alkaloid concentrations in plant material?
are the methods very different when using the wet material vs the dried material?

Nicodem - 16-1-2009 at 00:48

Most common method would be GC for volatile enough alkaloids (a FID detector is fine, but GC coupled with MS is better) and HPLC for all other UV active alkaloids. You need a pure sample for standard to determine the alkaloid both qualitatively and quantitatively. In some cases the alkaloid can be determined qualitatively already by making the TLC of the alkaloid fraction of the plant matter by comparison with a standard, but this is nearly not a 100% confirmation of its presence, neither does it tell anything about its concentration (but at least it is easily done by the average amateur).
Whether the material is wet or dry it does not make any difference for the determination itself, just the homogenization of the material and its extraction to give the alkaloid fraction is different.

hector2000 - 16-1-2009 at 02:09

does phenylacetone dissolve in acetone?what about ether?

bquirky - 16-1-2009 at 02:16

Hello,
I have a short question for the short question thread about salt solubility in acids,

It should be simple but its been puzzling me.

If you have say a %50 solution of sulfuric acid and then dissolve as much metal salt (say zinc sulfate) into it as possible. so you end up with a saturated zinc salt solution with a very low PH.

What is "supposed" to happen when a piece of zinc metal is then added to the solution ?

the sulfuric acid should want to react with the zinc but zinc sulfate is not as soluble in water as sulfuric acid.

Will the extra zinc sulfate precipitate onto the surface of the zinc and pasivate it ? or will the perciptate form somewhere else in the solution because of a temperature gradient ?

any ideas as to how these kind of processes might work ?

kclo4 - 16-1-2009 at 14:52

Thank you for your reply Nicodem. I guess if I wanted to determine a rough approximation of alkaloid content I'd probably just have to do an extraction and measure the end alkaloid mass relative to the begging plant material.

bquirky, I don't know if this would really help you out on answering your questions, but you might want to read about the common-ion effect.
http://en.wikipedia.org/wiki/Common-ion_effect

having the acid-water solution saturated with zinc sulfate would lower the levels of H+ in solution and that might play a roll in its reaction with zinc.

[Edited on 16-1-2009 by kclo4]

Sedit - 16-1-2009 at 20:46

Ok its about time i start beating down the short questions threed because i have a bunch and reading just isnt cutting it any more

For chlorination efforts would a fluoresent lightbulb with the Phosphor blown off to shine the UV/Blue light thru suffice for dissociation of chlorine?
It is generated by mercury vapor so im assuming that the output should be satisfactory for this application.
Is this correct?

BTW: i plan on clearing the Phosphor using a small tesla coil to clean the glass without breaking it for anyone woundering

not_important - 16-1-2009 at 20:57

Do you know that the Tesla discharge would do the job and still leave the lamp functional?

Light in the blue to UV range works, for moderate to large diameter reaction vessels the shorter wavelengths are strongly absorbed enough that they aren't as effective.

You want uncoated bulbs like these

http://www.hardwareandtools.com/invt/u252122?ref=gbase

http://www.prolighting.com/h33cd400.html

http://www.peclamp.com/buymvclear.html

[Edited on 17-1-2009 by not_important]

Sedit - 16-1-2009 at 22:45

"Do you know that the Tesla discharge would do the job and still leave the lamp functional?"

Yes sir iv been a mad scientist well before found this forum. Been building tesla coils for at lest 10 years now.

Would have to say i prefer the smaller ones that you can play with over the ones that one has to be a little cautious with.
Give me a little 2-3 incher diameter and i can clean a floresent light bulb in about 15 seconds leaving a pale blue glow that floresent objects shine in.

[Edited on 17-1-2009 by Sedit]

[Edited on 17-1-2009 by Sedit]

not_important - 16-1-2009 at 23:48

The high pressure in out and out mercury vapour lamps increases the output in the near UV and blue end of the spectrum, as compared to the low pressure of fluorescent lamps which radiate about half their light at 254 nm - too short for efficient driving of the chlorination reactions. Rough spectra and other information can be seen here

http://www.crystec.com/senlampe.htm

http://www.lamptech.co.uk/Documents/M3%20Spectra.htm

Sauron - 17-1-2009 at 04:52

This may or may not turn out to be a short question.

I want to prepare t-butyllithium. The lit. says I need a 2% alloy of sodium in lithium, preferably as dispersion in mineral oil.

I have not been able to find a supplier for this (98%Li 2% Na).)

I have hundreds of grams of Li wire in mineral oil. And a loy of Na.

Should I made my own?

Li mp 180 C. Na as I recall less than that. A melt of the two under a high boiling mineral oil stirred with a PTFE paddle ought to produce a fine sand dispersion of the alloy (2 g Na for every 98 g Li.)

Is this a plan? Or am I missing something?

Aldrich has a "high sodium" Li but it is 0.5% Na, while the lit. is quite specific about 2% Na.

[Edited on 17-1-2009 by Sauron]

appetsbud - 17-1-2009 at 05:05

can someone point me in the direction of a Paraformaldehyde synthesis? preferably from an aqueous formaldehyde solution

is it just as simple as heating with a little conc. acid and then removing the water via distillation?

thanks.

Sauron - 17-1-2009 at 06:44

There is a very recent US patent 7005083 on process for preparing alkyllithiums, that encompasses and describes preparing lithium-sodium alloy dispersions using an autoclave at a temperature above the mp of Li (180 C). The patent also states that such Li/Na alloy dispersions are commercially available from Postin Products Inc., in Faith, N.C., USA.

So I guess I can prepare it myself as I have a stirred heated autoclave.

[Edited on 17-1-2009 by Sauron]

Attachment: US7005083[1].pdf (360kB)
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hissingnoise - 17-1-2009 at 06:45

Paraformaldehyde is, I think, produced by simple evaporation of formalin without recourse to acid catalysts.
Distillation can be used to minimise losses.
Formalin normally contains small quantities of methanol to inhibit this polymerisation.
I have unused formalin which contains a fairly large precipitate of the polymer.

Sauron - 17-1-2009 at 06:49

There is an ACS monograph on Formaldehyde in the forum library for free download and it covers paraformaldehyde in detail.

panziandi - 17-1-2009 at 07:37

Sauron I was reading your post on t-BuLi and thought oooh "high sodium lithium" from Aldrich but when I got to the end of your post I realized it was no good! A quick google for "lithium sodium alloy" turned up a few springerlink journals but my proxy server sign in is playing up bastard thing! I think certainly making your own in an autoclave is the way forward.

Sauron - 17-1-2009 at 07:57

Making t-butyl chloride is easy. t-BuOH and conc HCl.

Per the patent, preparing the alloy dispersion is quite straightforward.

From there on calls for good technique with argon atmosphere and anhydrous conditions. But that is not a surprise.

Commercial solutions of t-Butyllithium are rather pricey, typically 800 mls of a 1.7 M solution in pentane works out to about $300 US by the time I get it. Too damned much for something I can prepare myself.

The US company in NC has no website and no published email address. Amazing in this day and age.

Here is the JOC paper from 1960 that is referenced from Paquette's book, this was the lit. I was referring to above. The prep of t-BuLi is in experimental section.

[Edited on 17-1-2009 by Sauron]

Attachment: jo01080a035.pdf (82kB)
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Sedit - 17-1-2009 at 10:21

Thank you not_important
So judging from the links you sent me A floresent light prepared that way would be effective since a small fraction is produced in the bandwidth area I would need but very far from efficient in terms of time and electrical input

How about those small hand held Black lights these seem to have there output in the early UV range. Just curiouse because I found on laying around last night.

Just found the frequency of the type of blacklight I have.
320-400 nm
I just dont know the strenght of it but that should be effeciant. Experimentation will tell

[Edited on 17-1-2009 by Sedit]

not_important - 17-1-2009 at 20:51

I suggest reading this http://www.faizkaskar.8k.com/light1.html particularly near the end when chlorination of toluene is being discussed. Then consider that a clear (no phosphor) high pressure mercury lamp, such as used for outdoor lighting, has peaks around UV (365nm), violet (406nm), blue (435nm), green (546nm), and yellow (578nm) while the low pressure fluorescent ones most of their energy in the 253 nm line. relying on the phosphor coating to convert that to the desired output.

By this http://www.standardpro.com/product/categories/fluorescentgui... blacklight bulbs have most of their output betweem 350 and 400 nm with the peak around 370 nm.

With fluorescent lamps the actinic bulbs used with aquariums or the bulbs designed for curing plastics and coatings are likely more effective than a blacklight bulb.

gnitseretni - 18-1-2009 at 11:00

I wanted to make some picric acid to make DDNP, so I got me some extra strength aspirin(500mg) and dissolved it in acetone. Unfortunately I ran out of acetone so I added some MEK. I'd say I had about 1 part acetone and 2 parts MEK.
Anyways, after boiling for a while and about one third of the original volume left, the solution started taking on a brown color. I took a syringe and sucked up about a ml or two and while swinging over to squirt it into a separate container, the stuff in the syringe solidified on me before I could do so.

What could this stuff be? It looks just like caramel.

I got another syringe and sucked up some more and quickly squirt it out onto the concrete. I thought it was completely solid, but found after hitting it with a hammer that it was a thick shell with fine brown powder inside. The powder was bone dry.

Anyways, I did this a week ago and don't have the stuff anymore so I can't do any tests to see what it is. I just thought I'd post it to see if anyone knows what I may have made.

Sedit - 18-1-2009 at 11:32

That was perfect Not_important thanks

Klute - 18-1-2009 at 12:51

appetsbud,

I have prepared paraformaldehdye from 37% formalin several times, by just vacuum distilling the solution until water distillate slowls down consoderably, and cooling the resulting thick limpid syrup, which slowly sets to waxy paraformaldehyde. tricky to get out of the flask at first, it cna be worth directly pouring it hot on a cool thick pyrex dish, and "crushing" it with a spatula tot acheive the granulometry you want. Once cold, it is very hard to further "grind" or cut up.

I finished by drying it over NaOh or P2O5 to get a very hard white material, which cna then be powdered with a pestle and mortar, but this isn't practical and it flies everywhere..

here is a picture of the produced paraformaldehdye:

sparkgap - 20-1-2009 at 07:35

Apparently my question was missed, so I'm gonna post it again:

Quote:

Alright, I will have to ask something after all. I am far away from my Merck Index, so could someone be so kind as to post the whole Merck Index entry for the antihyperlipidemic acipimox? I am especially looking for synthesis-related references.

I can't use my Merck Index on a Linux machine

so could someone be so kind?

sparky (~_~)

Acipimox; Merck Index 14th Edition

smuv - 20-1-2009 at 19:17

Monograph Number: 0000111
Title: Acipimox
CAS Registry Number: 51037-30-0
CAS Name: 5-Methylpyrazinecarboxylic acid 4-oxide
Additional Names: 2-carboxy-5-methylpyrazine 4-oxide
Manufacturers' Codes: K-9321
Trademarks: Olbemox (Pharmacia); Olbetam (Pharmacia)
Molecular Formula: C6H6N2O3
Molecular Weight: 154.12
Percent Composition: C 46.76%, H 3.92%, N 18.18%, O 31.14%
Literature References: Prepn: V. Ambrogi et al., DE 2319834; eidem, US 4002750 (1973, 1977 both to Carlo Erba). Prepn and toxicology: eidem, Eur. J. Med. Chem. 15, 157 (1980). Pharmacological profile: P. P. Lovisolo et al., Pharmacol. Res. Commun. 13, 151, 163 (1981). Pharmacokinetics: L. M. Fuccella et al., Clin. Pharmacol. Ther. 28, 790 (1980); L. Musatti et al., J. Int. Med. Res. 9, 381 (1981). Mechanism of action: K. Aktories et al., Arzneim.-Forsch. 33, 1525 (1983).
Properties: Crystals from water, mp 177-180°. LD50 orally in mice: 3500 mg/kg (Ambrogi).
Melting point: mp 177-180°
Toxicity data: LD50 orally in mice: 3500 mg/kg (Ambrogi)
Therap-Cat: Antilipemic.
Keywords: Antilipemic; Nicotinic Acid Derivatives.

Sauron - 20-1-2009 at 21:48

Recently someone posted about a chemical-equation writing utility software or applet,

I have tried UTFSE without avail.

Can some one point me to the proper place?

UnintentionalChaos - 21-1-2009 at 00:37

Sauron, do you mean for drawing out organic reactions? I personally use this freeware: http://www.acdlabs.com/download/chemsk.html

Sauron - 21-1-2009 at 01:47

No, not a 2D structure drawing app. I use MDL ISISDraw for that.

This is something that automatically handles chemical EQUATIONS:

NaOH + HCl -> NaCl + H20

It puts these in bold and subscripts the numbers automatically

Apparently it is a TSR applet.

DJF90 - 21-1-2009 at 13:02

I am currently stuck on a question I have come up with. If I were to dehydrate ethanol to ethene using conc. sulphuric acid then what kind of elimination would this be (E1/E2)?

I dont think it would be E1 because the primary carbocation isn't stable enough, and I dont think E2 can occur because that only happens in basic conditions (and we need the acid so that the hydroxyl group becomes a good leaving group; water). The only base that could be present is the HSO4- ion, which I understand works for the case of t-BuOH (its a poor base but an even worse nucleophile), but I E2 needs a strong base to deprotonate a carbon atom.

After attempting a google search I have come across pages saying the mechaism is E2 (ie. via the carbocation) but these are GCSE help guides and I dont think that the carbocation is stable enough to be formed.

sparkgap - 21-1-2009 at 14:32

Well, you have to take into account that your concentrated sulfuric acid actually transforms your ethanol first into ethyl hydrogen sulfate, CH3CH2-O-SO2-OH . So your idea of transforming the -OH group into a better leaving group is correct.

sparky (~_~)

DJF90 - 21-1-2009 at 14:56

Right, you got a mechanism for the production of the bisulphate ester? I tried to draw it starting with nucleophilic attack of the acid by the alcohol (like in a fischer esterification) and I had little luck creating a mechanism that seemed likely. From the bisulphite I assume E2 elimination to occur, with a H being removed off of the methyl group causing creation of the double bond and departure of the bisulphate ion? What is the base that removes the proton?

Is this the reaction mechanism solely because its a primary alcohol (the sterics of a sec/tert alcohol prevent ester formation?)? I know elimination from t-BuOH isn't via the bisulphate ester because protonation of OH allows it to leave as water and create a carbocation, which then is deprotonated by the bisulphite ion thus producing isobutene.

[Edited on 21-1-2009 by DJF90]

raiden - 21-1-2009 at 22:16

Is it possible to dehydrate Aluminium Chloride Hexahydrate by resubliming it?

Sauron - 21-1-2009 at 22:34

Maybe I should move my post to the Short Unanswered Questions thread?

DJF90 - 22-1-2009 at 00:06

Not something I've seen Sauron. Although it would be a really useful app, especially when you have to write a document with 10^n equations in it.

not_important - 22-1-2009 at 00:10

Quote:

Originally posted by raiden
Is it possible to dehydrate Aluminium Chloride Hexahydrate by resubliming it?

No, you can't even just plain sublime it; it loses HCl to form basic chlorides and then Al(OH)3, Al(O)OH, and Al2O3 as heating is continued.

Sauron - 22-1-2009 at 01:40

I saw this on the forum just 3-4 days ago and it appeared to be a fresh post. But I have tried and tried to find it via FSE and nothing turns up.

I search on

writing chemical equations
chemical equation writing
equation writing
writing equations
and each of those appended by "software" or "program"

No results.

DJF90 - 22-1-2009 at 09:23

http://www.sciencemadness.org/talk/viewthread.php?tid=4268&a...

Post by ayush, 31-12-2008 at 19:16,

Quote:

This multifunctional periodic table is an excellent tool for both students and serious researchers. It has over 20 types of data on each element and all known isotopes in customizable, user-expandable tables. All sixteen numeric data types (plus any user-added data) can be corelated in particle or line graphs. Includes a powerful chemical equation balancer that can solve the most complex organic reactions and calculate molecular weights and amounts of reagents.

Not exactly what you describe but along the right lines?

Sauron - 22-1-2009 at 12:07

Sorry, nope, that's not it.

Intergalactic_Captain - 22-1-2009 at 18:26

This may have been answered before, but it's been nagging at me today so I figured I'd ask. How are superconducting magnets, like those in NMR machines, magnetized? Once they're cooled to a superconducting temperature, is an external magnetic field applied to induce current flow, or is it induced by some sort of cryogenic voodoo magic?

not_important - 22-1-2009 at 20:27

They are DC electromagnets - http://en.wikipedia.org/wiki/Superconducting_magnet#Operatio...

497 - 25-1-2009 at 00:31

Sauron, I don't know if this is exactly what you mean, but it will balance equations and has some other nice features. I use it quite a bit.
http://home.c2i.net/astandne/help_htm/dwnload1.htm

Oh, I found it, here near the top: https://sciencemadness.org/talk/viewthread.php?tid=4268&...

Now for a question of mine: Can you make nitrite esters of long chain alkanols like dodecanol or cetyl alcohol the same way you make nitrites like isopropyl nitrite, NaNO2 + HCl? I wonder how well it would work because they're not soluble in water... Maybe with the addition of another solvent to increase their solubility?

After looking around some more I've found that lauryl alcohol is like 0.0003% soluble in plain water, but when there is some sodium lauryl sulfate present, its solubility is greatly increased, up to maybe 1 or 2%, which should be plenty for it to react, considering that is how soluble amyl alcohol is, and it is known to work with the one-pot NaNO2 + acid route.. Here's the ref. The lauryl sulfate should be stable in the conditions right? Hopefully the foaming doesn't become too problematic... fatty alcohols are used specifically as defoamers after all..

[Edited on 25-1-2009 by 497]

Sedit - 25-1-2009 at 11:19

Does Hydrogen have a disassociation frequency such as chlorine does?

panziandi - 26-1-2009 at 16:08

What the F**K has been up these past 24hrs or so? Sciencemadness had been down, files deleted from server or something?! I felt like a junkie without smack! Totally distraught and lost amongst the sea of Roguesci! Thank god this is back up now!

not_important - 26-1-2009 at 17:47

Quote:

Originally posted by Sedit
Does Hydrogen have a disassociation frequency such as chlorine does?

Yes, way down in the deep ultraviolet, disassociation happening at 840 A or shorter.

http://www.nature.com/nature/journal/v118/n2973/abs/118592c0...

Sedit - 26-1-2009 at 19:02

Thank you.
Once this was accomplished would this cause the hydrogen to behave as it does in the presents of catylist and be useful for hydronation efforts?
A radical hydrogen seems like it could be very useful if it could be produced from the use of light in a reaction vessle.

not_important - 26-1-2009 at 22:20

This - 840 A, the absorption starts at about 1200 A IIRC - is in the extreme or vacuum UV, although a pure nitrogen atmosphere is transparent in the range. It's not easy to generate much light in this range, and the photons carry a lot of energy - above 10 eV. I suspect that you'd find that most organics fragment under the required conditions.

Sedit - 27-1-2009 at 00:03

Yeh i didnt even take that into consideration never mind, may be useful for a few things but like you said it will pretty much cook anything at those levels

Alumina TLC Plates

smuv - 27-1-2009 at 15:08

Can alumina TLC plates be used for efficient chromatography of polar molecules?

I could get a good deal on some alumina TLC plates, but have no experience w/ alumina. All of the chemistry I do is with fairly polar species (carbonyls, alcohols, amines etc.); would I get bad separation of polar compounds and/or smearing?

I would appreciate if anyone could share any TLC tips they have picked up using alumina plates.

Panache - 27-1-2009 at 22:19

Quote:

Originally posted by smuv
Can alumina TLC plates be used for efficient chromatography of polar molecules?

I could get a good deal on some alumina TLC plates, but have no experience w/ alumina. All of the chemistry I do is with fairly polar species (carbonyls, alcohols, amines etc.); would I get bad separation of polar compounds and/or smearing?

I would appreciate if anyone could share any TLC tips they have picked up using alumina plates.

In addition to this question, does anyone have any experience on whether prepared tlc plates have a shelf life if stored appropriately (cool dry), i uncovered a box of various types and the like but most are dated in the early 80's. They appear fine.

jokull - 28-1-2009 at 06:32

Several years ago I had to work with some old stock of alumina and silica TLC plates. The results I obtained were the same when carried out on brandnew plates. So, I think you can employ old plates with confidence. By the way, I was dealing with carotenoids

Aubrey - 28-1-2009 at 08:28

i've just been making phthalimide from phthalic anhydride and ammonia. Even with googles, breather and holding it at arms length ammonia is pretty nasty stuff (dont have fume hood).
Next time i was thinking replacing ammonia with urea. Is this any less smelly?

Klute - 28-1-2009 at 11:27

Alumina plates work very well for polar compounds, I use this routinely at work as most complexes I work on are sensible to silica gel.. Doesn't change much from silica plates from a practical point of view, the Rf are different but still similar. I guess as long as they are well stored they can be kept years with no problem!

raiden - 28-1-2009 at 13:43

Would combining 32 percent hydrochloric acid and DCM give 2 layers, one of anhydrous HCl and one of water?

Jor - 28-1-2009 at 13:57

Quote:

Originally posted by Aubrey
i've just been making phthalimide from phthalic anhydride and ammonia. Even with googles, breather and holding it at arms length ammonia is pretty nasty stuff (dont have fume hood).
Next time i was thinking replacing ammonia with urea. Is this any less smelly?

I have tried this method, using urea, and it works like a sharm. You are done with the synthesis in a few minutes, because you just have to melt it, and after some time the mixture suddenly froths up, indicating you are done.
Then I think you will have to purify from any excess urea, but I skipped this as it did not interfere with my next step.
And yes, carbon dioxide and ammonia are evolved. You will get all the ammonia in a short period of time, so capture it with a wash bottle. However, just work outside. Good thing about ammonia is that it will very quickly rise up (it's lighter than air), so it won't just hang around like bromine/chlorine/NO2 do.

Urea method is cheap, very easy, and uses not expensive reagents.

acetylation of amine with acetyl chloride

chemrox - 28-1-2009 at 17:51

Acetyl chloride seems much more reactive than acetic anhydride and so I'm planning to use it for an acteylation I plan. Question: do I expect the product amide to be an HCl salt or free amide?

gsd - 28-1-2009 at 18:12

Combining 32 % HCl with DCM (or with any other organic solvent for that matter) will give 2 layers. Depending on the Distribution Coefficient (also called as partition coefficient) of the system, each layer will contain different proportions of HCl, water and the organic solvent.

HCl has extremely strong affinity towards water. it is difficult to strip HCl out of water short of consuming it by reaction with alkali.

gsd

sparkgap - 28-1-2009 at 18:27

Quote:

Originally posted by chemrox
Acetyl chloride seems much more reactive than acetic anhydride and so I'm planning to use it for an acteylation I plan. Question: do I expect the product amide to be an HCl salt or free amide?

The amide group itself is not particularly basic (electron delocalization). Only if you have more basic groups present in your compound will you wind up with the salt.

sparky (~_~)

chemrox - 28-1-2009 at 19:04

Ahh .. yeah .. I forgot to mention the pyrrolidine moiety and that was the crux of the question.. woops

sparkgap - 28-1-2009 at 19:23

Quote:

Originally posted by chemrox
Ahh .. yeah .. I forgot to mention the pyrrolidine moiety and that was the crux of the question.. woops

If so, yes, it comes out as the salt; more so if the sp3 N is not tertiary.

sparky (~_~)

Oboe - 31-1-2009 at 08:35

Hello, I have perhaps very simple question of whether it is possible for this reaction

Glycine
H2N-CH2-CO2Et convert to Cl-CH2-CO2Et

with use of NaNO2 and HCl or if there is problem with forming this not wanted compound. May it be best to form alpha acid first?
Thank you if you can help me

[Edited on 31-1-2009 by Oboe]

sparkgap - 31-1-2009 at 09:12

Quote:

Originally posted by Oboe
Hello, I have perhaps very simple question of whether it is possible for this reaction

Glycine
H2N-CH2-CO2Et convert to Cl-CH2-CO2Et

with use of NaNO2 and HCl or if there is problem with forming this not wanted compound. May it be best to form alpha acid first?
Thank you if you can help me

Apparently.

sparky (~_~)

Oboe - 31-1-2009 at 14:13

Yes that is alternative Mr sparkgap. Ester is target though but it seems to stabilize diazo compound to get stable compound I linked. Perhaps someone knows if is possible to go from this stable compound to chloride?

smuv - 31-1-2009 at 14:34

you could first make the chloroacetic acid and then form the ester; It is easier to form esters with chloroacetic acid than acetic acid because it is more electrophilic.

Picric-A - 1-2-2009 at 08:33

My school has recently given be two very old bottles (1985) of Phthalic acid and phthalic anhydride.
Both are in very good condition, and i see no reason why the phthalic acid is affected in any way by its age.
I am not sure about the anhydride however, could it have absorbed water to hydrate itself to form phthalic acid?
I am not sure how readily phthalic anhydride combines with water, i hear some anhydrides have to be refluxed with water for days to hydrate.
thanks,

chemrox - 1-2-2009 at 11:05

both are solids at STP... I don't see why there'd be a lot of interconversion if they've been kept in bottles in cabinets... you're not in Kentucky or Vietnam are you?

crazyboy - 1-2-2009 at 14:19

I recently made acetic acid by distillation of sodium acetate and sulfuric acid. The sodium acetate had been dried in the oven at 500F and the sulfuric acid was technical grade. I took 150ml of this and added two spoonfuls of anhydrous MgSO4 to dry it then I filtered it. The resulting solution was distilled with another spoonful of MgSO4 and the first 125ml that came over were collected.

I have placed the acid in the refrigerator and although it is 7C it is not totally frozen. What did I do wrong or what can I do to make it anhydrous?

Jor - 1-2-2009 at 14:35

Are you sure it did not 'supercool'. My GAA was still liquid at 5C, but when I shaked the bottle, it all solidified, a beautiful process to see.

crazyboy - 1-2-2009 at 16:42

I shook it and it didn't freeze. also the acetic acid before purification partially froze in the refrigerator and froze solid in the freezer but after the dessication step it wont even freeze below 0C.

This is the result of several combined bathes in the later batches I used a slightly dirty sulfuric acid but the product looked the same maybe this has something to do with it?

I have no idea why it won't freeze it doesn't make sense. :mad:

sparkgap - 1-2-2009 at 17:03

crazyboy, I have a feeling this is freezing-point depression at work. Your acid is not as pure as you think it is, and the "slightly dirty sulfuric acid" is a very probable suspect.

sparky (~_~)

crazyboy - 1-2-2009 at 17:10

That is what I suspected. As I said the first batches froze partially at 7C and totally at 0C but after I combined them with my latest one nothing freezes it.

The first few times I was using drain cleaner sulfuric acid which had a yellow tinge. I ran out of that so I switched to my gallon bottle of the same brand but the color was much darker.

I'm assuming it is some sort of buffer in the acid. Any idea what to do or should I just toss the acetic acid?

Oboe - 1-2-2009 at 17:28

Thank you for mr sparkgap and mr smuv for help with my question

mr crazyboy. perhaps to add clean sulfuric and distill again?

chemrox - 2-2-2009 at 14:27

You could try some hardware store zeolites as molecular sieves to pull the last bit of water out. Have you checked the boiling pt? It should be elevated a bit by the water if there.

Sedit - 2-2-2009 at 16:42

What would the zeolites be marketed as in the hardware store? Just zeolite or is there a brand to look for?

chemrox - 3-2-2009 at 00:03

Zeolite.. they will be with the water purification materials. Also do a google for zeolite..
http://www.bearriverzeolite.com/brz-specifications.htm

[Edited on 3-2-2009 by chemrox]

Panache - 3-2-2009 at 01:04

dryer cartridges for refrigeration fluids are an easily sourced source of zeolites. Danfoss is a large supplier manufacturing out of mexico i believe. Almost every suburb in the western world will have a refrigeration mechanics supplier as it is such an essential industry for food safety.

Picric-A - 3-2-2009 at 11:20

Has anybody got any information of the synthesis of Hydrazine from chlorourea and sodium hydroxide?
I understand chlorourea can be made by passing chlorine through a soloution of urea with a suspention of ZnO (can CaCO3 be used instead of ZnO?)
I guess i could just acidify with H2SO4 to precipitate the sulphate from the reaction mix..
thanks in advance.

Panache - 3-2-2009 at 16:06

Does silver(II) form any volatile salt or complex at stp?
Are the ammoniated silver(II) complexes, that are reputed to be explosive, so sensitive and/or easily formed, that one should avoid any contact between ammonium solutions and silver(II) solutions or are the salts/complexes only explosive when dry?

Jor - 3-2-2009 at 16:25

Yes, it forms a pyridine-complex, and it can be isolated as the peroxidisulfate salt:

http://www.versuchschemie.de/topic,11731,-Tetrapyridinsilber(II)-peroxodisulfat.html

If you can't read it:
-dissolve 15g K2S2O8 or 13,2g Na2S2O8 in 400ml water.
-Dissolve 4,5g AgNO3 and 5ml pyridine in 100mL water.

Add the solutions to eachother, a orange color is produced and precitipate is formed. Leave the solution to stand for 30min.

woelen - 4-2-2009 at 00:18

Interesting thing about silver(II)-pyridine complex. Nice to try on a reagent tube scale.

@Panache: You are talking about ammoniated silver(II) complexes, but these are plain silver(I) complexes. If you mix ammonia and a solution of a silver(I) salt and allow to stand for a few hours, then a solid silver(I) compound (some amminie/nitride/amide complicated badly specified complex) is formed, which is extremely sensitive to mechanical agitation. Simply decanting the wet solid from a container may lead to an explosion already.

stoichiometric_steve - 4-2-2009 at 14:17

Are Swagelok threads a proprietary thing or is there a standard naming for it? I have a manometer with a 1/8" (or so i think, the small one) outer threading and i need to adapt it to a tube nipple. I haven't found anything so for that could do this.

Panache - 6-2-2009 at 05:15

Quote:

Originally posted by stoichiometric_steve
Are Swagelok threads a proprietary thing or is there a standard naming for it? I have a manometer with a 1/8" (or so i think, the small one) outer threading and i need to adapt it to a tube nipple. I haven't found anything so for that could do this.

I haven't found anything that fits other than Swegelok stuff, however i have not asked the Swagelok rep directly, not that he would know neccessarily being a sales rep, lol. I try to avoid the system given the joins are not guaranteed once loosened and retightened (as compared to the Prochem system which is). Swagelok however do offer a full range of thread to nipple adaptors which could be ordered online easily enough.

Alternatively buy 20 old GC's for $20 and salvage the nipples.

hector2000 - 12-2-2009 at 10:19

what componet will produce if water add to phenylacetyl chloride?

sonogashira - 12-2-2009 at 17:12

^ Carboxylic acid will be formed. Phenylacetic acid.

Intergalactic_Captain - 13-2-2009 at 09:50

http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=CV1...

Is there any reason why the above reaction is not used more often in the preparation of nitrostyrenes? Granted, it is for unsubstituted benzaldehyde, but I can't see any reason why it isn't used in any, for example, 2c-h syntheses.

What about the 3,4,5 substitution pattern? It seems like it would be easier to use NaOH as the "catalyst" (though not finding a mechanism, I'd imagine that it proceeds through sodium methyl-nitronate) than trying to find/prepare some odd amine.

Nicodem - 13-2-2009 at 09:56

Quote:

Originally posted by Intergalactic_Captain
http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=CV1...

Is there any reason why the above reaction is not used more often in the preparation of nitrostyrenes?

Maybe you just did not look hard enough. I have seen a few examples at the Hyperlab forum.

Intergalactic_Captain - 13-2-2009 at 10:04

The Hyperlab? Not speaking russian, I've never checked it out. Would you be able to give any examples of reactions tried and what the yeilds were?

...When I said it was not used often, what I meant was that in my beilstein searches I've not seen it more than a couple of times, and in googling and searching the old rhodium archives it's never come up (everyone seems to use EDDA).

Ebao-lu - 14-2-2009 at 02:56

Quote:

...When I said it was not used often, what I meant was that in my beilstein searches I've not seen it more than a couple of times, and in googling and searching the old rhodium archives it's never come up (everyone seems to use EDDA).

So you are concerning the catalyst? Oh dear, there are many catalysts except EDDA. For example, ammonium acetate, cyclohexylammonium acetate. Alkalines will also do. Just google smth like " Henry reaction catalysts"
I thought first you did not find at rholdium any examples of henry reaction at all..

chemrox - 14-2-2009 at 17:00

I'm sorry am I reading this correctly? You didn't find examples of the Henry reax at Rhodium archives? No that can't be right. My question is in the Henry reaction between benzaldhyde and nitroethane have better catalysts been found since the Org Syn article posted by Nicodem? I seem to recall this being done with ammonium acetate but the mixture was kept in the dark for two weeks until yellow crystals of nitropropene formed.

Formula409 - 14-2-2009 at 21:16

Is it possible to isolate the Schiff Base of a ketone such as P2P/MVK/MDP2P in order to get 100% imine formation, aiming to increase the yields of an Al/Hg reduction?

Formula409.

solo - 15-2-2009 at 08:27

HOw can ifigure out the size of a reflux column when refluxing in various size flasks and volume in flask,,,,,solo

chemrox - 15-2-2009 at 10:01

Quote:

Originally posted by Formula409
Is it possible to isolate the Schiff Base of a ketone such as P2P/MVK/MDP2P in order to get 100% imine formation, aiming to increase the yields of an Al/Hg reduction?

Formula409.

I see two questions here. Yes Schiff bases are isolated all the time before the reduction is done. With ketones you need a catalyst like p-toluenesulfonic acid. A few crystals.

Does that get you 100% yield? Usually not. The reduction doesn't always get you 100% and various methods require different workups that all have yield implications. You're planning on an aluminum amalgam redux? You will have some washing up to do after that.

[Edited on 15-2-2009 by chemrox]

Sedit - 15-2-2009 at 14:20

Is stainless steel resistant to oxalic acid?

Im evaporating a solution of oxalic acid in a stainless steel vessle and the solution is taking on a green color.

Is this from contamination or it the stainless under attack?

Ok I just changed to a glass evaporation plate and on a good note the solution wasnt what was green it just appeared that way because the stainless steel had turned green.

Have I lost some oxalic due to this or am I still good to go?

I am extracting out of a dilute rust remover solution and I was suprized to find oxalic crystalized nicely on the bottom as soon as I opened it but Im still awaiting the final yeild from this solution

[Edited on 15-2-2009 by Sedit]

DJF90 - 15-2-2009 at 20:23

I would expect the green colour is from the nickel in the stainless reacting with the oxalic acid. Some yield is therefore likely to be compromised. http://physchem.ox.ac.uk/msds/NI/nickel_II_oxalate_dihydrate...

Sedit - 15-2-2009 at 22:09

Chances are your right, the color seems to fit and it dont appear that the steel is being heavyly corroded just discolored with a very fine sediment, after removel of the slight heat there is crystal formation so tommorow Ill be able to tell whats going on.

Always fun to know when you accidently create a cancer causing toxic substance

Formula409 - 15-2-2009 at 23:46

Quote:

Originally posted by chemrox

I see two questions here. Yes Schiff bases are isolated all the time before the reduction is done. With ketones you need a catalyst like p-toluenesulfonic acid. A few crystals.

Are they isolated just like amines (ie. extract, acidify, risnse/repeat)?

Quote:

You're planning on an aluminum amalgam redux? You will have some washing up to do after that.

Actually, I was planning on exploring whether thiourea dioxide (http://www.sciencemadness.org/talk/viewthread.php?tid=11785) is suitable for performing the reduction, naturally I will need some sort of way of quantitatively measuring yield, so the reduction procedures detailed on Rhodium are unsuitable for research.

Formula409.

[Edited on 16-2-2009 by Formula409]

Ebao-lu - 16-2-2009 at 00:34

As i know, the shiff base can be easily extracted by organic solvent. I've seen a procedure, where to P2P(probably, in some org solvent) was added some 25% ammonia solution, and this mixture was stirred to get a solution of P2P-imine in organic layer, which was further separated and reduced with NaBH4. Now i don't remember where i've seen it, but it seems to be a basic procedure for imines

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