Sciencemadness Discussion Board

The Short Questions Thread (4)

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Panache - 21-4-2017 at 13:19

Quote: Originally posted by clearly_not_atara  
https://erowid.org/archive/rhodium/chemistry/redamin.dithion...

Quote:
General Procedures

Method A.

The oxime (20 mmol) was mixed with water (15 ml) containing sodium dithionite (28 mmol). The suspension was stirred overnight at room temperature. (Warming to 40°C reduced reaction times to several hours.) In some cases, a precipitate formed.This product was very high melting and, on treatment with 2 M hydrochloric acid, liberated the carbonyl compound and sulfur dioxide. It was therefore assumed to be the bisulfite addition compound of the carbonyl compound and was not isolated. A slight excess of 2 M hydrochloric acid was added to the reaction mixture and nitrogen was bubbled through the mixture to expel the sulfur dioxide. Solid sodium carbonate was added carefully to alkalinity; the aqueous mixture was allowed to stand for 30 min and was extracted with ether (2x10 ml) which was dried (MgSO4) and evaporated. The residue was essentially pure carbonyl compound (by t.l.c.)

Method B.

The reaction described under Method A was performed in the presence of sodium hydrogen carbonate (28 mmol). Cleavage by means of this modification appeared to proceed considerably faster. The usual workup gave the carbonyl compound in comparable yield.


Notably there are some claims that sodium dithionite reduces oximes, however, upon inspection, dithionite is only reported to reduce alpha-keto oximes, produced by nitrosation of the ketones. It's worth noting that alpha-keto oximes tend to be present as the nitroso compound rather than the oxime, unlike all other oximes, and the nitroso compound is more susceptible to reduction.


So unless my oxime is alpha keto dithionite likely will not work. Any other reactants you know of.
I've tried % sulphuric, 35% yield lots of oximetry back could be worked into a continuous method.
Problem is both the oxime and the ketone steam distill, so you end up with it dissolved in your ketone and that separation is tedious.
Any input is welcomed

S.C. Wack - 21-4-2017 at 14:12

The oxime, dissolved in 10 to 12 vol. of 50% aqueous ethanol,
was refluxed with 3.5 molar equiv of sodium bisulfite until
thin layer chromatography indicated complete reaction. After
removal of the ethanol by distillation, the residue was admixed
with chloroform and an excess of dilute hydrochloric acid, and
the ketone or aldehyde was extracted into the organic layer.
In the case of the aldehydes, hydrolysis of the bisulfite adduct
required stirring with acid for up to 30 min to obtain two clear
layers.

The extracts gave near-quantitative yields of “crude” product,
usually single spot by tlc.

JOC 31, 3446 (1966)

clearly_not_atara - 21-4-2017 at 16:28

Quote: Originally posted by Panache  

So unless my oxime is alpha keto dithionite likely will not work. Any other reactants you know of.
I've tried % sulphuric, 35% yield lots of oximetry back could be worked into a continuous method.
Problem is both the oxime and the ketone steam distill, so you end up with it dissolved in your ketone and that separation is tedious.
Any input is welcomed


I think you were confused by my wording. "Reducing" an oxime in almost all contexts refers to conversion to the primary amine, and I wanted to say that this would not happen. Conversion to a ketone or aldehyde is called "oxime cleavage" or "hydrolysis" because usually no redox occurs, although some reaction sequences cleave the N-O bond and hydrolyse the resulting imine.

Dithionite converts normal oximes to carbonyl compounds, and alpha-ketooximes to alpha-ketoamines, according to the literature.

However as noted, bisulfite alone is sufficient for cleavage of oximes into the respective carbonyl compounds. It's possible that in the dithionite methods, dithionite reacts with air and bisulfite is the actual cleaving reagent. Since NaHCO3 should accelerate the decomposition of dithionite (it's an alkali) this seems particularly likely.

[Edited on 22-4-2017 by clearly_not_atara]

[Edited on 22-4-2017 by clearly_not_atara]

RogueRose - 24-4-2017 at 17:12

Is it possible to safely heat glass media bottles made by Corning, Schott Duran, Pyrex or VWR? I have some of each and they all seem identical in size, shape and weight and looking at some data on them (from Corning), they seem to be made of borosilicate 7040 or 7070. They say they can be autoclaved and heated but I'm curious if anyone has ever tried heating them with a hot plate or flame.

I would think if it is the same glass as flasks and beakers, I don't see a reason it wouldn't be able to be heated - except maybe the plastic seal at the top - so cut it off or don't heat past XX degrees.

For the 7040 glass it showed a linear expansion with heat up to softening point of 800C - it was basically a 40-45 degree angle - so it seems like it may have the proper characteristics for heating.

Any ideas or experience with this?

Sulaiman - 24-4-2017 at 17:33

I suspect that the thickness and shapes of bottles would make thermal stresses much greater than in an rbf for example,
I'd have more confidence in a soda-glass rbf than a borosilicate-glass bottle.

Geocachmaster - 24-4-2017 at 17:42

I unfortunately don't have any media bottles, but I think that their glass is thicker than normal beakers and flasks. Thicker glass increases the chance of breaking while heating. Heating on a hotplate and autoclaving are two very different things. The heating in an autoclave is much more even and gradual than on a hotplate. A simple beaker and watch glass as a cover is going to be much safer on a hotplate, and cheaper. A water bath would probably be fine but I wouldn't risk direct heating.

JJay - 24-4-2017 at 18:50

I think a water bath would be fine for heating a media bottle too but wouldn't heat one directly on a hotplate and certainly wouldn't expose one to a flame. Several of my media bottles have plastic coatings that would be destroyed. While the others could probably withstand some uneven heating, I don't think they are really designed for it.

RogueRose - 24-4-2017 at 22:25

Quote: Originally posted by Geocachmaster  
I unfortunately don't have any media bottles, but I think that their glass is thicker than normal beakers and flasks. Thicker glass increases the chance of breaking while heating. Heating on a hotplate and autoclaving are two very different things. The heating in an autoclave is much more even and gradual than on a hotplate. A simple beaker and watch glass as a cover is going to be much safer on a hotplate, and cheaper. A water bath would probably be fine but I wouldn't risk direct heating.


I agree with the thicker part being more prone to cracking. I've seen some thick wall flasks (ideal for filtering) and they seem about the same thickness as the bottles.

There are some small dots on the bottom of the bottle so they don't stick to counters or form suction if set in water - these wouldn't work well with a hot plate I would think, which is why I suggested a flame (and not a torch or bunsen burner) like an alcohol lamp.

Maybe I'll try a water bath, then oil bath followed by an alcohol burner (slow steady heat) and see how it handles it. I'd try a sand bath but I don't have that ATM..

Jimmymajesty - 26-4-2017 at 13:55

Hi folks! I bought a memory foam mattress but after a week I had to get rid of it, I could not bear the smell. It constantly oozed some sweet odour that first seemed the "new smell" but later it became ubearable.
It is funny for me to be chemofobic considering all the crazy stuff I inhaled in the last decades still I could not bear the thought that I sleep on cancer...
Do you have anything on this? I pinched out small pieces from the foam then analyzed it with HSGC-MS, nothing could be detected which is funny because I could smell it on the matress during the sampling.
Anyway thanks for your help in advance. And please dont buy this cancerous memory crap especially if its from china, so basically never ever buy these kind of things, you have been warned!

j_sum1 - 26-4-2017 at 16:34

Smell is one thing. Cancer-causing is quite another.
Both are reasons to get rid of a mattress but do not conflate the two.

Not sure why you could not detect it though.

Jimmymajesty - 27-4-2017 at 13:45

I did not add one thing to the story which is important to me, but for you it might be just a subjective observation, which actually is.
So after the 4th day I started to develop -apart from the constant headache- funny illnesses like dry eyes, twitching arm muscles, ear ringing etc. These were all mysteriously disappeared 3 days after I dumped the mattress.
I read some stuff on the internet:) based on the writings these mattress manufacturers are worse than Hitler and Stalin combined, however polybrominated diether type flame retardans may indeed cause something unwanted...

yobbo II - 27-4-2017 at 15:25


The auld story. Bad memories from a memory mattress encounter.

Panache - 30-4-2017 at 16:47

Quote: Originally posted by RogueRose  
Is it possible to safely heat glass media bottles made by Corning, Schott Duran, Pyrex or VWR? I have some of each and they all seem identical in size, shape and weight and looking at some data on them (from Corning), they seem to be made of borosilicate 7040 or 7070. They say they can be autoclaved and heated but I'm curious if anyone has ever tried heating them with a hot plate or flame.

I would think if it is the same glass as flasks and beakers, I don't see a reason it wouldn't be able to be heated - except maybe the plastic seal at the top - so cut it off or don't heat past XX degrees.

For the 7040 glass it showed a linear expansion with heat up to softening point of 800C - it was basically a 40-45 degree angle - so it seems like it may have the proper characteristics for heating.

Any ideas or experience with this?


If they are from schott my experienc is they are absolutely fine on a hot plate or direct flame heating (as in Meker type Bunsen burner used for sterilising or drying glass.
They are fairly useless on a hot plate as far as heating rate goes so I scrunch some al foil wetted with silicon oil between them and the plate.
Thousands of times I have heated these bottles fairly vigourously on the flame and the only time i experienced a failure was with new generically branded bottles I Was testing. Schott never failed

ficolas - 5-5-2017 at 14:06

Today I boiled water and iron hydroxide, to form iron oxide, but a penetrating kinda sulfur like smell scared me. Is that normal? Something like welding fumes?

Panache - 6-5-2017 at 20:39

Brauer calls for using diluted conc nitric to wash cro3 crystals, but I just ran out (well I ran out weeks ago but haven't replenished. Anyone know of or can think of an alternative minerL acid or salt solution that's suitable?

JJay - 7-5-2017 at 16:02

That's a tough one... it would have to be something extremely acidic and not reducing that will evaporate easily. I think you could probably use phosphoric acid, but it doesn't evaporate as easily. I would like to know the answer to this also.

clearly_not_atara - 7-5-2017 at 19:34

M/ethanesulfonic acid are my first guesses. Maybe ring-chlorinated benzoic acids.

Question

TheNerdyFarmer - 12-5-2017 at 18:36

I have a question regarding making solutions. I know that this is a super basic knowledge but the trouble I am having with this regards making specific percentages (I.e 13%) in a really weird amounts (I.e. 127ml). I get the concept sort of. I know that 30 grams of something dissolved 100 grams of water is a 30% solution of whatever.
The part I really choke up at is when I need to make an odd solution in an odd amount. Especially with like and acid. (I.e 36% HCl to a 12% HCl solution) and pretty much the same with other acids. I always end up guesstimating instead of doing exact calculation.
I guess what I am asking is if there are any formulas that I need to know or if anyone could shed some light on this. It would also be great if someone has like a link to a helpful video or page that explains how to precisely calculate a solution percentage. (Molar and by weight). Hope this isn't too simple of a question. Thanks in advance.

JJay - 12-5-2017 at 18:57

30 grams of something dissolved in 100 grams of water is a 23% solution. 30 grams / (30 grams + 100 mL * s.g. of water).

JJay - 12-5-2017 at 19:00

molarity: weight of solute / mol.wt of solute / (wt of solute + wt of solvent) / s.g. of solution / 1000

j_sum1 - 12-5-2017 at 19:05

Very simple. But you need some algebra.

Concentration is expressed in amount per litre. This could be

(a) a molar amount. 0.25 mol per litre means 0.25 moles of the substance has been dissolved in the solvent and the n the volume increased to 1 litre. (The 1 litre is the total amount of the solution and not the total amount of solvent used. They can be quite different for concentrated solutions. You can for example make a litre of sucrose solution using only 500mL of water.)
(b) percentage solutions. This is based on mass. I generally do this on a scale. For a 13% solution you would weigh out 130 of your solute and then make up to 1kg with the solvent. (Note that your description is not quite correct here. For 30% you do not use 30 grams dissolved in 100g of water. Instead you use 70g of water.)

If you wish to make up a different amount than 1 litre or one kilogram then you simply scale up or down appropriately. Thus for 225 mL of a molar solution (in (a) above) you would simply calculate the amount that you need and then multiply by 225/1000.
It is pretty normal to make up a stock solution of the concentration that you need and then measure out the amount that you need. So, make up 500mL and then measure out your 127mL leaving 373mL for later. Just label everything.



Diluting a solution is pretty easy if you are using molar amounts. There is a tidy little formula to use:
c1v1=c2v2

Concentration times volume is constant. Add some more solvent and the concentration decreases while the volume increases. But the product remains the same. Rearranging this formula allows you to calculate what your new volume needs to be for your new concentration and hence how much you need to dilute your solution.

You can do exactly the same for percentage solutions (b). A little care is required though. I always restate the concentration in grams per kg rather than as a percentage. This reinforces to me that I need to be calculating and measuring mass and not volume.

I don't know of any videos off the top of my head but I suspect that Khan Academy would have something pretty explanatory.


[edit] Pipped by Jjay. Also corrected some formatting.



[Edited on 13-5-2017 by j_sum1]

j_sum1 - 12-5-2017 at 19:15

I should add that weight/volume solutions are a bit more tricky. This is because the density of your solution changes upon dilution.
If your solution is sufficiently dilute then it probably does not matter. For example say you are mixing up a solution of phenolphthalein, a 0.5 gram per litre solution will be pretty much the same as a 0.5 gram per kilogram solution. The same could not be said for a 50% solution of sulfuric acid. If you want to pull out your density tables you could probably work it out.

TheNerdyFarmer - 13-5-2017 at 06:53

Thank you all very much! This has been quite helpful. I see that I have a lot of practicing to do lol.

mesanaw - 25-5-2017 at 08:44

For the extraction and isolation of piperine using IPA, can I substitute KOH with NaOH? KOH is difficult to procure in my area.
Thanks.

S.C. Wack - 25-5-2017 at 15:23

If someone has KOH flake and NaOH prill and needs to dissolve one or the other in an alcohol, the KOH is often chosen because it's easier.

chemistry

rka - 26-5-2017 at 23:20

How to, chemically and not electrolytically, get Ag from AgSCN powder?

PirateDocBrown - 26-5-2017 at 23:57

Heat.

AJKOER - 27-5-2017 at 04:20

Boiling AgCl in aqueous sugar is frequently performed (I have done it successfully a few years ago) to retrieve metallic Silver. This may provide a path here, but be wary of products produced and disposal issues thereof.

Please note, not all sugars are equally effective in this reaction! See discussion at https://www.finishing.com/195/29.shtml .

Eddygp - 31-5-2017 at 07:44

Can anyone point me in the direction to find some literature about glutamate dehydrogenase inhibitors? I haven't been able to find anything so far.

Alice - 31-5-2017 at 08:38

google scholar: glutamate dehydrogenase inhibitor / inhibition. Gives various hits. Or are you looking for something specific?

yobbo II - 1-6-2017 at 11:55

There was a thread/discussion on the board somewhere or other regarding seperation of sodium chlorate from sodium perchlorate using acetone. It was mostly figured out by (I think) Hennig Brand.
Can someone link me to the thread. I simply CANNOT find the thing.

Thanks,
Yob

Element mirrors

j_sum1 - 3-6-2017 at 03:00

The silver mirror (Tollen's reagent) is very well known.
I am interested in attempting the same thing with different elements. I have heard that something similar can be done with copper using hydrazine as a reducing agent.

I was wondering if anyone had any experience with mirroring other elements. Specifically bismuth.

J.

Question

Plunkett - 5-6-2017 at 04:07

Around what concentration does sulfuric acid cease to be an effective drying agent? More specifically, can my drain cleaner sulfuric acid (~93%) be used to dry bromine?

Question.

TheNerdyFarmer - 5-6-2017 at 11:47

Today was making a solution of copper (II) chloride via mixing a 1:1 ratio of HCl and 3% H2O2 and adding that to copper. I have done this before but I had a little bit of a different result this time. I left it for around two hours and it soon became this dark brown liquid. I tested it by adding a few drops to aluminum and it made the famous vigorous reaction between CuCl2 an Aluminum. I also dilluted a small sample and it turned into an olive green color.
My guess on what it is is a saturated solution of copper (ii) chloride or that it's just really contaminated with copper (i) chloride. What chemical do I have here???

cant think of a username - 5-6-2017 at 13:13

question

hi,so i just rewatched nurdrage's video for making luminol and i was wondering what exctly happens when he mixes aluminium and sodium metabisulfite?does it form sodium dithionite in situ? or is it something else.

DraconicAcid - 5-6-2017 at 13:45

Quote: Originally posted by TheNerdyFarmer  
Today was making a solution of copper (II) chloride via mixing a 1:1 ratio of HCl and 3% H2O2 and adding that to copper. I have done this before but I had a little bit of a different result this time. I left it for around two hours and it soon became this dark brown liquid. I tested it by adding a few drops to aluminum and it made the famous vigorous reaction between CuCl2 an Aluminum. I also dilluted a small sample and it turned into an olive green color.
My guess on what it is is a saturated solution of copper (ii) chloride or that it's just really contaminated with copper (i) chloride. What chemical do I have here???


I think the dark brown is either the tetrachlorocuprate(II) ion, or a complex ion containing both copper(I) and copper(II).

TheNerdyFarmer - 5-6-2017 at 13:54

Is there any way to know the percent of copper ii chloride so I can try and extract it?.

DraconicAcid - 5-6-2017 at 14:22

Quote: Originally posted by TheNerdyFarmer  
Is there any way to know the percent of copper ii chloride so I can try and extract it?.

If you dilute it sufficiently, any copper(I) chloride should precipitate.

Determinig KCl vs NaCl solution - easy way to differentiate?

RogueRose - 7-6-2017 at 10:00

Is there any way to test a solution to see if it is sodium or potassium chloride

Texium - 7-6-2017 at 10:31

A flame test is a good way to distinguish, if it is pure. A wooden splint dipped in KCl solution and held in a flame will give a light purple color, while an NaCl solution will give a bright yellow-orange color. The presence of sodium will cover up the color from potassium though, so it isn't perfect.

Another way that comes to mind is that some potassium salts are much less soluble than their sodium analogues, such as with chlorate and perchlorate. So if you happened to have some sodium chlorate, you could add some of your solution to a reasonably concentrated solution of it and see if you get a precipitate.

RogueRose - 7-6-2017 at 12:02

Quote: Originally posted by zts16  
A flame test is a good way to distinguish, if it is pure. A wooden splint dipped in KCl solution and held in a flame will give a light purple color, while an NaCl solution will give a bright yellow-orange color. The presence of sodium will cover up the color from potassium though, so it isn't perfect.

Another way that comes to mind is that some potassium salts are much less soluble than their sodium analogues, such as with chlorate and perchlorate. So if you happened to have some sodium chlorate, you could add some of your solution to a reasonably concentrated solution of it and see if you get a precipitate.


Thank you, the flame test should be sufficient to verify what the solution is as it should be a pure solution (99%+ purity of original salt mixed with distilled water). Thanks again!

RogueRose - 7-6-2017 at 12:32

Quote: Originally posted by zts16  
A flame test is a good way to distinguish, if it is pure. A wooden splint dipped in KCl solution and held in a flame will give a light purple color, while an NaCl solution will give a bright yellow-orange color. The presence of sodium will cover up the color from potassium though, so it isn't perfect.

Another way that comes to mind is that some potassium salts are much less soluble than their sodium analogues, such as with chlorate and perchlorate. So if you happened to have some sodium chlorate, you could add some of your solution to a reasonably concentrated solution of it and see if you get a precipitate.


well I took a long wooden match stick, dipped it into the solution and put it into the blowtorch flame. I had 2 known (NaCl and KCl) solutions and an unknown. the NaCl gave a nice yellow/orange flame almost as if I was adding an accellerant to the flame. If put at the base of the flame, the length of the flame turned yellow/orange. When I did the KCl I still got a similar flame, but it didn't seem as "energetic" - if that makes sense, it woulnd't propagate the flame the same way when placed at the base (all solutions are super-saturated at room temp).

The unknown does look more similar to the KCl but it is not a purple flame except maybe a couple mm immediately off the match stick.

The purple wasn't easily noticable in the known KCl solution as they all gave a yellow orange flame but the NaCl never produced the small purple near the wood.

Is this the results I should expect? I was concerned that maybe the wood was a little different to give the purpleish color.

The sodium flame was impressive though!

DraconicAcid - 7-6-2017 at 12:45

The trouble with the flame test is that soooo many things are contaminated with just enough sodium to overpower the other colours. A common trick in the old days (when people relied on the test) is to view the flame through didymium welder's glass, which blocks out the yellow quite selectively. Or using a platinum wire instead of a wooden stick.

Sulaiman - 7-6-2017 at 18:17

first find something to hold your sample in the flame which does not itself give a yellow flame ... surprisingly difficult.
Nichrome resistance wire works ok, I think nickel spatulas but I'm not sure,
carbon often gives a similar yellow.

even then, if the 1% impurity is sodium (quite possible) then it may be enough to show as sodium.

I have tried borax bead tests, if I have nothing else to do I may re-visit .... useful but tedious.
https://en.wikipedia.org/wiki/Bead_test

Texium - 7-6-2017 at 19:14

Even so, if his sample was sodium chloride, or contained a substantial amount of it, it would have been very obvious. I think it's safe to conclude that the unknown is KCl, assuming that it could either be NaCl or KCl, and not anything else.

Sulaiman - 7-6-2017 at 19:42

true,

I just thought that the borax bead test was worth mentioning

(plus, I was showing off the entire extent of my knowledge in this area :P

ninhydric1 - 10-6-2017 at 19:49

@RogueRose

If you are really desperate, make a sodium cobaltinitrite solution. Potassium cobaltinitrite is insoluble in water and will precipitate out, while sodium cobaltinitrite won't. But it may contaminate the potassium chloride solution, so perform this test in a separate reaction vessel.

[Edited on 6-11-2017 by ninhydric1]

The Aussie - 10-6-2017 at 20:51

Question

For a distillation, is it possible to set up a Syphoning system if you don't have a water pump on hand, would this work? I was thinking something like this and just lifting the water at the bottom to the top by hand.

thanks.

PS this is my first time here so yea, sorry if this has already been discussed, although i did look , I could not find anything.

syphoning.png - 6kB

j_sum1 - 10-6-2017 at 21:08

I'd be more inclined to use a regular tap.
A siphon is not going to be a recirculating system unless you babysit it and tip the water back into the top bucket.
Also, the flow-rate will diminish as the had decreases.
It just all seems awkward. Either spend $20 bucks on a pond pump or plan on flushing some water down the drain.

Sigmatropic - 11-6-2017 at 02:13

A syphon is certainly a way to cool a condensor. The flowrate can be easily adjusted by raising the upper basin. This allows you to control the temperature of the output quite precisely. As has been said above you need to baby-sit this setup and be prepared to reestablish the syphon several times per run. Aquarium pumps are the way to go.

Acetic Acid from Sodium acetate/diacetate

RogueRose - 13-6-2017 at 17:09

I was looking at the procedure of mixing H2SO4 and CH3CO2Na (sodium acetate) and was wondering if sodium diacetate could be used NaH(CH3CO2)2. The equations I've come up with are as follows:

Sodium acetate + sulfuric acid
2(CH3CO2Na) + H2SO4 = 2(CH3COOH) +Na2SO4

Sodium acetate + sulfuric acid
NaH(CH3CO2)2 + H2SO4 = 2(CH3CO2H) + NaHSO4
So with the diacetate the result would be sodium bisulfate which may be more useful for some than sodium sulfate. Now I'm not sure how soluble NaHSO4 would be in acetic acid or if it would push a reaction back towards sulfuric acid - both would be good reasons this isn't used but I figured it would be worth knowing.

clearly_not_atara - 13-6-2017 at 17:25

I think the product will in all cases depend on the stoichiometry of H2SO4 vs Na+. I am also pretty sure you want to distill acetic acid off.

I think you can convert it to the calcium salt somehow? Not sure.

EDIT: Combining hot concentrated aqueous solutions of calcium chloride and sodium acetate precipitates calcium acetate in about 50% yield, judging by solubity tables.

[Edited on 14-6-2017 by clearly_not_atara]

JJay - 13-6-2017 at 18:02

Can you store acetone in HDPE wash bottles, or will it evaporate?

tsathoggua1 - 14-6-2017 at 01:10

Out of methyl carbonate, ethylene carbonate and propylene carbonate, which is preferable for the purpose of production of small quantities of alkali metals via electrolysis?

Thing is, I have a reasonable sized quantity of Na, but its in the form of one large brick, hermetically vacuum-packed and heat-sealed in an inner, thick plastic liner filled with mineral oil which itself is inside a hermetically sealed bag full of inert gas. I'd sooner make a little and not have to open it, due to shelf-life considerations, I don't need the entire 100g brick right this moment, only maybe 15g or so, so it makes more sense to try electrolysis in one of the carbonate esters of hydroxide, and in any case its something I've wanted to add to my solvent arsenal for a while.

If that doesn't produce the desired results, I'll go with Nurdrage's excellent looking method of thermiting NaOH or KOH using magnesium dust, grinding up the slag in a blender (only variation, backfilling this with an argon purge first and grinding it under positive pressure of Ar. Might have to tweak the bender a little bit but its only ever been me that uses the thing so no loss there), and then separating the sodium out using 1,4-dioxane prepared by the sulfuric acid dehydration of ethylene glycol, which I have plenty of, as I do of concentrated H2SO4. I just don't want to open the Na brick I already have since it won't have an indefinite shelf life, and there is perhaps 15g or so of a ketoxime awaiting preparation so as to be reduced to the corresponding amine with Na/alcohol.

Sulaiman - 14-6-2017 at 07:17

Quote: Originally posted by JJay  
Can you store acetone in HDPE wash bottles, or will it evaporate?


Acetone, b.p. 56oC
so you can store it in an hdpe wash bottle, until it evaporates away :)
I find hdpe bottles with hdpe:hdpe sealed bottle caps ok for acetone starage with negligible loss due to evaporation.

tsathoggua1 - 14-6-2017 at 07:51

I'd expect the result of destructive distillation performed on the Ca salt to be acetone rather than CH3COOH, that is a known route to P2P, albeit not one optimized for yields, pyrolysis of either the Ca or Pb salt of phenylacetic acid results in dehydration to the ketone, not liberation of the acid. H2SO4 is a fairly strong dehydrating agent itself when concentrated, so I could see either formation of acetone, or perhaps dehydration and elimination to ethene (or worse, ketene, although granted that would make acetic acid if put to use, but its dreadfully toxic stuff)

[Edited on 14-6-2017 by tsathoggua1]

Texium - 14-6-2017 at 08:00

Ignoring tsathoggua's usual cruddy, off-topic, drug chem laced post, yes, you should definitely be able to distill a mixture of sulfuric acid and sodium diacetate to get acetic acid. There would be no reason to convert it to the calcium salt. You do have to distill though. Don't count on getting it to work otherwise.

JJay - 14-6-2017 at 09:27

Quote: Originally posted by Sulaiman  
Quote: Originally posted by JJay  
Can you store acetone in HDPE wash bottles, or will it evaporate?


Acetone, b.p. 56oC
so you can store it in an hdpe wash bottle, until it evaporates away :)
I find hdpe bottles with hdpe:hdpe sealed bottle caps ok for acetone starage with negligible loss due to evaporation.


I don't have any sealed caps for my HDPE wash bottles. I wonder if it is still OK to store acetone in them....

Geocachmaster - 14-6-2017 at 09:42

What kind of wash bottle is it? I have two, one has the spout through the lid, the other through the side. The one which comes out the side has a cap that lets air in (not sure why, they didn't want air bubbling through I guess). In the former case I think a bunch of PTFE tape would seal it well! As for the latter option I'm not sure.

Edit: the surface area of acetone that is exposed though the "straw" is probably not significant. That would be the only way for acetone to escape.

[Edited on 6/14/2017 by Geocachmaster]

JJay - 14-6-2017 at 09:56

The straw on my wash bottles actually goes through the cap. They are about 1/10 the cost of the wash bottles with the straw at the side, and I see very little difference in function....

I guess I'll try filling one with some acetone and see what happens.

[Edited on 14-6-2017 by JJay]

Question

TheNerdyFarmer - 18-6-2017 at 19:43

I know this is probably a simple question but how exactly does one discern the peroxide ion from the dioxide ion? Is it something to do with oxygens electronegativity?

JJay - 18-6-2017 at 20:16

Dioxide and peroxide aren't ions... dioxide means it contains two oxygens.

Peroxide means it contains at least two oxygens with a single bond to each other.

Campa Chem MSDS?

ninhydric1 - 18-6-2017 at 20:35

I plan to buy some Campa-Chem original holding tank fluid, but it seems that I can't access the MSDS to determine the concentration of formaldehyde and methanol. Does anyone have access to such records and/or does anyone have the concentrations of the two chemicals above?

DraconicAcid - 18-6-2017 at 20:37

Quote: Originally posted by JJay  
Dioxide and peroxide aren't ions... dioxide means it contains two oxygens.

Peroxide means it contains at least two oxygens with a single bond to each other.


The peroxide ion is O2 with a charge of -2, as is found in barium peroxide and potassium peroxide.

sparkgap - 18-6-2017 at 21:33

Quote: Originally posted by ninhydric1  
I plan to buy some Campa-Chem original holding tank fluid, but it seems that I can't access the MSDS to determine the concentration of formaldehyde and methanol. Does anyone have access to such records and/or does anyone have the concentrations of the two chemicals above?


I found the MSDS in the Wayback Machine (which I've attached), but there doesn't seem to be any info on the concentration.

sparky (~_~)

Attachment: Campa Chem.pdf (47kB)
This file has been downloaded 1698 times

JJay - 18-6-2017 at 22:25

Ok - to be fair, "dioxide" is not an ion.

"Peroxide" can be an ion if both of the oxygens are ionically bonded. But it is not a very stable one, and while ionic peroxides certainly should not be ignored, covalently bonded peroxides are more common. Technically, an "oxide" ion also exists, as does a "superoxide" ion, although like peroxide, these ions usually don't last very long as independent species.


[Edited on 19-6-2017 by JJay]

Darkstar - 19-6-2017 at 01:19

@JJay:

Dioxide and peroxide are both ions, and are formed through single-electron transfer to molecular oxygen. The dioxide ion (also called "superoxide") is O2− and is the result of one electron transfer, while the peroxide ion is O22− and the result of two electron transfers:

O=O + e− → [•O−O:]− (dioxide)

[•O−O:]− + e− → [:O−O:]2− (peroxide)

So to answer TheNerdyFarmer, the difference is that the dioxide ion is two oxygen atoms connected through a single bond where one has an unpaired electron and the other a full octet, while the peroxide ion is two oxygen atoms connected through a single bond but with both oxygen atoms having full octets. Potassium dioxide (KO2) is an example of the former, and potassium peroxide K2 O2 ) is an example of the latter.

JJay - 19-6-2017 at 01:26

Superoxide and dioxide are not the same thing. Carbon dioxide is not carbon superoxide, for example.

If you disagree, feel free to cite a reputable reference that describes the dioxide ion.

[Edited on 19-6-2017 by JJay]

Darkstar - 19-6-2017 at 01:31

"Dioxide" is the correct IUPAC nomenclature for the superoxide anion. For example, if you check the Wikipedia page for potassium superoxide, you will notice that it lists "potassium dioxide" as the IUPAC name. This is also mentioned on the Wikipedia page for the superoxide ion as well.

JJay - 19-6-2017 at 01:39

Quote: Originally posted by Darkstar  
"Dioxide" is the correct IUPAC nomenclature for the superoxide anion. For example, if you check the Wikipedia page for potassium superoxide, you will notice that it lists "potassium dioxide" as the IUPAC name. This is also mentioned on the Wikipedia page for the superoxide ion as well.


The IUPAC name for the superoxide ion is dioxide (1-), and the IUPAC states that superoxide is an acceptable alternative name. The IUPAC name for the peroxide ion is dioxide (2-). While it is permissible to omit the charge notation where there is no ambiguity under IUPAC nomenclature, that is not the case with dioxygen. Again, "dioxide" is not an ion.

Source: https://www.iupac.org/fileadmin/user_upload/databases/Red_Bo...




[Edited on 19-6-2017 by JJay]

Darkstar - 19-6-2017 at 02:11

Right, but I am assuming that the context of his question was the difference between the anions in salts like "potassium dioxide" and "potassium peroxide." We can argue nomenclature technicalities all day, but unless the anion is specifically designated as dioxide(2-) or the cation is prefixed with di- (like "dipotassium dioxide"), it's probably a safe bet to assume that it is the superoxide anion in simple salts like those.

JJay - 19-6-2017 at 02:19

See page 80 for a discussion of this issue. "Potassium dioxide" is a simple stoichiometric name, *NOT* a compositional name.

[Edited on 19-6-2017 by JJay]

Darkstar - 19-6-2017 at 03:29

Nice edit, but I am not "confused." I'm agreeing with you that "dioxide" by itself can be somewhat ambiguous. But the poster asked for the difference between "dioxide" and "peroxide" as if they were two different anions. Can they mean the same thing or be confused for one another? Yes. But in my experience you usually see "peroxide" for dioxide(2-), and either "superoxide" or "dioxide" for dioxide(1-). That's all I'm saying. It may not be 100% correct, but that's what I see.

I agree that using "dioxide" by itself can create ambiguity depending on the cation, I'm not denying that.

Edit: And just so it's clear, my original response was to your reply to Draconic before you edited it to include superoxide ions.

[Edited on 6-19-2017 by Darkstar]

gluon47 - 19-6-2017 at 11:40

My school has given me the opportunity to do a few experiments in there lab. I noticed they had a bottle of anhydrous propionic acid, so I thought I might make a propionate ester.

I would like to try isopropyl propionate. Fischer esterification should work for making this ester right?

Would adding anhydrous magnesium sulphate to the reaction mixture during reflux help to increase yield?

Any advice would be much appreciated.

ninhydric1 - 19-6-2017 at 13:06

So does anyone know the concentration of formaldehyde and methanol in Campa-Chem, or is there a method (such as titration) to figure out the concentration?

sparkgap - 19-6-2017 at 13:19

Quote: Originally posted by gluon47  
My school has given me the opportunity to do a few experiments in there lab. I noticed they had a bottle of anhydrous propionic acid, so I thought I might make a propionate ester.

I would like to try isopropyl propionate. Fischer esterification should work for making this ester right?

Would adding anhydrous magnesium sulphate to the reaction mixture during reflux help to increase yield?

Any advice would be much appreciated.


Fischer should work for both primary and secondary alcohols. Isopropyl alcohol is cheaper than propionic acid, so use that in excess. If you can use a Dean-Stark trap, do so; magnesium sulfate won't do much in those reaction conditions.

sparky (~_~)

mesanaw - 19-6-2017 at 13:25

Quote: Originally posted by gluon47  


I would like to try isopropyl propionate. Fischer esterification should work for making this ester right?

Would adding anhydrous magnesium sulphate to the reaction mixture during reflux help to increase yield?



Yes, a Fischer esterification would occur. Considering you are already using anhydrous propionic acid and presumably 99% isopropanol, the use of magnesium sulfate during reflux would not materially affect yield. I have seen desiccants used after reflux and separation in order to remove the residual water, but not during reflux.

[Edited on 19-6-2017 by mesanaw]

JJay - 19-6-2017 at 14:11

I don't think I'd put any desiccant in the reaction mixture itself, but you might want to put some in the reflux path: https://www.youtube.com/watch?v=Ah5ds_3s5BI

Magnesium sulfate might clump up if you put it in a Soxhlet and run wet solvents through it, so molecular sieves are probably better.

gluon47 - 20-6-2017 at 01:17

Awesome, thanks!:). I wont use MgSO4

gluon47 - 20-6-2017 at 21:34

Performed the reaction today on a 1/6 mole scale with 20% excess of isopropanol. It was interesting to smell the propionic acid. A lot like glacial acetic acid, but more rancid, I like it:).

Isopropyl propionate seems to have the characteristic sweet odour of most simple esters, but quite reminiscent of glue and slightly less sweet.

I'm going back to the lab next week for workup and purification. Can't wait!

Saccharin

Geocachmaster - 29-6-2017 at 16:49

The first photo is me trying to make sense of the pathway from methyl anthranilate to saccharin found on Wikipedia, shown in picture two. I don't have access to the reference given or any other papers I find about the synthesis of saccharin.

The diazotization in step one seems pretty straightforward. In step two the diazonium compound is reacted with sulfur dioxide to produce the sulfonyl chloride at number three. The Wikipedia page on sulfonyl halides says that phenyldiazonium chloride reacts with sulfur dioxide and HCl, so I'm assuming setp two needs an acid catalyst. In step three the sulfonyl halide is reduced by SO2 to the sulfinic acid. Steps two and three would be carried out together and at low temperatures because the sulfonyl chloride will react with water. Extra SO2 will exclude air and prevent oxidation of the sulfinic acid. What I'm confused most about is the two next steps. Wikipedia just says Cl2 and NH3. I read that sulfonamides can be prepared by reaction of an NHR2 with a sulfonyl chloride. For this reason I assumed that chlorine would react at step four to remake the sulfonyl chloride seen at step three. This would then react with ammonia to produce the sulfonamide. After this I'm thiking that the H+ on wiki is a hydrolysis which is followed by a ring forming step which must happen automatically.

Does anyone with more experiance than me think this is plausible/makes sense? I think making an artificial sweetener that is 300x sweeter than sugar would be really cool! It's something I want to do in the future. I'll be purchasing 250g (or maybe 1000g!) of phthalic anhydride and that would be my starting point.

Any input is greatly appreciated :)


SaccharinPathway - Copy.jpeg - 189kB WikiSac.jpeg - 260kB
Sorry for the bad quality :(, you have to open the pic to actually see anything...

UC235 - 29-6-2017 at 18:34

Quote: Originally posted by Geocachmaster  
The first photo is me trying to make sense of the pathway from methyl anthranilate to saccharin found on Wikipedia, shown in picture two. I don't have access to the reference given or any other papers I find about the synthesis of saccharin.

The diazotization in step one seems pretty straightforward. In step two the diazonium compound is reacted with sulfur dioxide to produce the sulfonyl chloride at number three. The Wikipedia page on sulfonyl halides says that phenyldiazonium chloride reacts with sulfur dioxide and HCl, so I'm assuming setp two needs an acid catalyst. In step three the sulfonyl halide is reduced by SO2 to the sulfinic acid. Steps two and three would be carried out together and at low temperatures because the sulfonyl chloride will react with water. Extra SO2 will exclude air and prevent oxidation of the sulfinic acid. What I'm confused most about is the two next steps. Wikipedia just says Cl2 and NH3. I read that sulfonamides can be prepared by reaction of an NHR2 with a sulfonyl chloride. For this reason I assumed that chlorine would react at step four to remake the sulfonyl chloride seen at step three. This would then react with ammonia to produce the sulfonamide. After this I'm thiking that the H+ on wiki is a hydrolysis which is followed by a ring forming step which must happen automatically.

Does anyone with more experiance than me think this is plausible/makes sense? I think making an artificial sweetener that is 300x sweeter than sugar would be really cool! It's something I want to do in the future. I'll be purchasing 250g (or maybe 1000g!) of phthalic anhydride and that would be my starting point.

Any input is greatly appreciated :)



Sorry for the bad quality :(, you have to open the pic to actually see anything...


Sulfinates can be chlorinated to sulfonyl chlorides. A similar reaction occurs during chlorination of bunte salts (see: http://www.sciencemadness.org/talk/viewthread.php?tid=9921&a...). Hydrolysis of the intermediate sulfur chlorides ends with a modestly stable sulfonyl chloride.

Treatment with ammonia gives the sulfonamide.

If you're only after a few grams, saccharin is readily isolated from US Sweet-n-low (or equivalent off-brands) where it is present as its sodium salt at roughly 3.5% by weight. The solid sweetner is dissolved in roughly twice it's weight in water, filtered to remove anticaking agents, and acidified with HCl. The saccharin free acid precipitates as a white powder (I recommend chilling it in an ice bath or fridge for a while). Vacuum filter off the liquid and rinse with ice water to remove contaminating glucose.

JJay - 29-6-2017 at 20:48

Does anyone know offhand where to get a few grams of n-phenylanthranilic acid?

Question

TheNerdyFarmer - 1-7-2017 at 12:01

Does anyone know of a relatively cheap source of absolute ethanol? I am having trouble finding good affordable sources of it. It would also be nice if there was an affordable source of lower concentration that I could dry myself.

Morgan - 1-7-2017 at 17:19

This was $12.99 for 750ml, a couple of dollars cheaper than the other 190 proof Everclear next to it in the store.
http://wongdrinks.blogspot.com/2012/04/i-drank-that-clear-sp...

Oxygen-containing cations

j_sum1 - 12-7-2017 at 20:16

Just a curiosity mostly.
Oxyanions abound. But are there any polyatomic cations that contain the element oxygen?
I could not think of any but then I have not had much dealing with polyatomic cations except for ammonium.

ninhydric1 - 12-7-2017 at 20:23

j_sum1 do complexes count? Because [Co(H2O)6]2+ is a 'cation' that contains water, which in turn contains oxygen.

True polyatomic cations containing oxygen are probably impossible due oxygen's high electronegativity.

EDIT: Never mind, Harristotle has some.

[Edited on 7-13-2017 by ninhydric1]

[Edited on 7-13-2017 by ninhydric1]

Harristotle - 12-7-2017 at 20:24

Vanadyl and Uranyl are the ones that spring to mind VO2+ and UO22+.

Other than that, dunno.

Strange purple chromium compound

j_sum1 - 12-7-2017 at 20:38

I am always cautious with hexavalent chromium and like to reduce it to Cr(III) before adding to the Cr waste bucket. Today I was cleaning up some solutions and observed a transient purple colour that I have not seen before. Details follow:

This was a demonstration I did for my students on oxidation states. We began with solutions of trivalent and hexavalent chromium, oxidised one and reduced the other to have their colours switch.

I had one of those "what was in there?" moments as I hadn't seen that purple before and was certainly not expecting it.
On reflection I guess it was either a peroxide complex or a transient Cr(IV) compound. Can anyone enlighten me?

j_sum1 - 12-7-2017 at 20:40

Quote: Originally posted by Harristotle  
Vanadyl and Uranyl are the ones that spring to mind VO2+ and UO22+.

Other than that, dunno.


Cool. Thanks. I should have thought of those. But since polyatomic cations are a rarity for me my mind just did not gravitate in that direction.

Tdep - 12-7-2017 at 21:09

Quote: Originally posted by j_sum1  

On reflection I guess it was either a peroxide complex or a transient Cr(IV) compound. Can anyone enlighten me?


Hello!
Blue/purple chromium in the presence of acidified peroxide, that sounds like Chromium(VI) oxide peroxide! I have a gif of me making some on twitter, hopefully you can see it... https://twitter.com/Explosions_Fire/status/87898935016739226...
I'll hopefully have a video up in the next two weeks where I give better footage of it. It's quite unstable in water so only appears transiently, so it seems to match what you're describing pretty well

j_sum1 - 13-7-2017 at 00:18

That looks very much like it. Mine was a lot lighter (concentration) and a bit more towards the purple and was cloudy with the precipitate at the time.

But it is cool learning new stuff like this. Thanks for that link.


edit: spelling.

[Edited on 13-7-2017 by j_sum1]

JJay - 13-7-2017 at 11:01

It can actually be extracted with ether and used as a high-yielding chromium oxidizer in anhydrous conditions (so it can turn primary alcohols to aldehydes without proceeding to a carboxylic acid, for example). It forms a somewhat stable adduct with pyridine and a highly stable adduct with bipy. The bipy adduct can be stored on the shelf for months.

ficolas - 16-7-2017 at 14:17

In the combustión of an inorgánicos ester, like triethyl borate, what product does the acid part form? The acid again? So, boric acid in the case of thriethyl borate?

DraconicAcid - 16-7-2017 at 14:30

Quote: Originally posted by ficolas  
In the combustión of an inorgánicos ester, like triethyl borate, what product does the acid part form? The acid again? So, boric acid in the case of thriethyl borate?

Or just the oxide.

ficolas - 16-7-2017 at 15:01

So it would be both, depending on the acid decomposition temperature or something?
And if its the oxide that forms, I guess it would be the oxide the acid decomposes to, right?

ninhydric1 - 18-7-2017 at 15:06

I plan to get one of these:

http://www.ebay.com/itm/192238825824

in the near future with the tight budget I have.

I'm planning to somehow adapt it to my current non-magnetic stirrer hotplate, so I was wondering how far above the magnetic stirrer would the magnetic stir bar still spin adequately?

TheNerdyFarmer - 19-7-2017 at 03:48

I know we have science madness patches for sale on this forum, but do we have things like bumper stickers for sale??

TheNerdyFarmer - 21-7-2017 at 18:51

Okay, this is something that has confused me for some time. I hear from some people that the Mica window on a Geiger Muller tube enables it to be able to detect alpha radiation. But now I go on eBay and see some seemingly high quality pancake probes but then I will find out that they can only detect beta and GT gamma radiation. Does or doesn't the Mica window enable the probe to detect alpha particles??

quenching a grignard/NH4Cl substitute

karlos³ - 29-7-2017 at 11:08

I need to quench a grignard reaction, and in this case the reaction is usually quenched using a saturated ammonium chloride solution.
However, I used all up and need to quench the reaction anyway.
What could be a possible good substitute here? I thought about using ice and diluted HCl, like 5%, but I am not sure, the substrate could be sensitive to it(an indolylketone, by the way).
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