Sciencemadness Discussion Board - The Short Questions Thread (4)

The Short Questions Thread (4)

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Molecular Manipulations - 31-5-2015 at 22:55

Will the use of a steel file impart any significant amount of steel contamination to the powder?

Certainly some, but for most purposes it's not a big deal. Run a magnet over the shaving to be sure.
xfusion, I've heated potassium nitrate above that temp before. No bang to speak of (hard to speak with one's face covered in molten oxdizers and glass shards [JK]). In large amounts and/or confined, I'd stay away.

delta H/G for absorption of water by concentrated sulfuric

Insanus - 1-6-2015 at 07:12

Does anyone know where to find quantitative data for enthalpy/free energy changes from absorption of water by sulfuric acid? This would be useful for predicting reactions such as

CH3CHOHCH3 <----> CH2=CHCH3 + H2O

in acidic media.

Edit: NIST has reaction energetics for both sulfuric acid and water. Calculations would need to take into account concentration and the exact ionic species involved, but the data is there.

[Edited on 6/1/2015 by Insanus]

[Edited on 6/1/2015 by Insanus]

Funkerman23 - 3-6-2015 at 19:45

Does anyone know what diameter capillary tube I need to use for a theile tube/ non mel-temp machine melting point determination? I don't have a mel-temp machine so it has to be the older fashioned sample tied to a thermometer in a theile tube/ oil bath route..I'm stumped and I don't know if the diameter of the capillary tubing is important to the accuracy or not. I really appreciate any advice.

j_sum1 - 4-6-2015 at 02:06

Silica gel from cat litter
There are brands of cat litter claiming to be 100% silica gel (according to MSDS). However, not all of the crystals are clear -- there is perhaps 5% blue crystals in the mix. The label says that it is scented for some of the brands.

Does anyone have any experience with cat litter as a source for silica gel? It seems like it might be a useful thing to have on hand and rather cheap and simple to obtain.

diddi - 4-6-2015 at 03:24

ok as a spill kit perhaps...

Crowfjord - 4-6-2015 at 17:10

@Funkerman23: diameter shouldn't matter much, as long as it is fairly thin. I was unaware that melting point capillaries even came in different diameters. I think the ones I use (also with a Thiel tube and thermometer) have an inner diameter of 1.0 mm.

Zombie - 4-6-2015 at 17:19

Quote: Originally posted by j_sum1

Silica gel from cat litter
There are brands of cat litter claiming to be 100% silica gel (according to MSDS). However, not all of the crystals are clear -- there is perhaps 5% blue crystals in the mix. The label says that it is scented for some of the brands.

Does anyone have any experience with cat litter as a source for silica gel? It seems like it might be a useful thing to have on hand and rather cheap and simple to obtain.

Ali baba... $600.00 a ton.

The blue crystals you see may be the color indicator. WallMart sells it as a flower desiccant, and it turns blue as it absorbs moisture. It's about 5 bucks a pound there.

I never saw it in cat litter.

PHILOU Zrealone - 5-6-2015 at 05:52

Quote: Originally posted by blogfast25

Quote: Originally posted by DraconicAcid

It's chelated to make it sound more nutritious.

That won't work with chemophobes!

But wel with chemophages

...

Just tell chemophobes that:
Proteins are polyaminoacids, polypeptides very similar to polyamide plastic...they wont eat anything but starch...
--> Starch is a polymer of glucose similar to cellophane and cellulose.
--> They die starving

and Darwin wins

The Volatile Chemist - 10-6-2015 at 12:13

When chlorinating solutions of, say, dyes, how necessary is it to absorb excess chlorine? I don't have a fume hood.

Loptr - 11-6-2015 at 09:32

Since potassium fluoride had a melting point around 858*C, can I assume that I can easily dehydrate the dihydrate to the anhydrous form without it decomposing?

blogfast25 - 11-6-2015 at 10:09

Quote: Originally posted by Loptr

Since potassium fluoride had a melting point around 858*C, can I assume that I can easily dehydrate the dihydrate to the anhydrous form without it decomposing?

No. The dihydrate has an MP of only 41 C (Wiki). On melting it becomes a saturated solution of KF in water.

Heavily dissociated into K+ and F-, the fluoride ions reacts with water, because HF is a weak acid (pKa = 3.17, Wiki):

F-(aq) + H2O(l) < === > HF(aq) + OH-(aq). Water soluble fluorides do indeed yield alkaline solutions.

Since as HF is volatile, on heating it will partially 'boil off'. You end up with KOH, or at least KF heavily contaminated with KOH.

[Edited on 11-6-2015 by blogfast25]

Loptr - 11-6-2015 at 10:18

Quote: Originally posted by blogfast25

Quote: Originally posted by Loptr

Since potassium fluoride had a melting point around 858*C, can I assume that I can easily dehydrate the dihydrate to the anhydrous form without it decomposing?

Yeah, that's what I was afraid of, and I don't feel like working with HF at this time.

Thanks blogfast!

[Edited on 11-6-2015 by Loptr]

greenlight - 21-6-2015 at 00:37

Quick question, just concentrated some 70% Nitric acid via distilling with 98% Sulphuric acid and continued until no more nitric came over. The sulphuric acid obviously has a large quantity of water in it now, is it fine to boil the water back off on heat so its reconcentrated for re-use again.

[Edited on 21-6-2015 by greenlight]

The Volatile Chemist - 21-6-2015 at 19:03

At the very least, do it in a fume cub., as the dense vapours of the sulfuric acid will form on heating, and are rather unpleasent.

learningChem - 25-6-2015 at 17:29

closely related question :

How long (approx.) should it take for ~94% sulfuric acid to reach 98%? I think I tried once, and yes got nasty clouds of acid, but the concentration didn't change much. Seems like a lot of acid evaporates, not just the water?

[Edited on 26-6-2015 by learningChem]

greenlight - 25-6-2015 at 20:39

Try what I have been doing by just heating at 110-120.C in a beaker or other glass container on a hotplate instead of boiling the acid as well by heating strongly. This will boil the water which will evaporate off but leave the acid so no acid clouds.
You can put a watch glass or beaker bottom or anything that has a glass surface over the acid heating container to tell when all the water is gone (I use a glass saucepan with glass lid). If the glass surface becomes fogged up with condensation, there is still water in the acid. It will also tend to lightly fume white fumes when all the acid is nearly gone but these fumes will not fog up the glass.

learningChem - 26-6-2015 at 11:52

Thanks greenlight! I'll give your method a try. I was under the impression that, as the acid became more concentrated, more energy was needed to get rid of the remaining water, but experiment should clear this up =)

greenlight - 26-6-2015 at 20:40

I have used this method many times usually starting with clear 60% Sulphuric acid and afterwards it is much darker yellow and thicker and will char paper towel black on contact.
Just periodically check for water condensation and when none is visible let it heat for an additional 10 minutes to ensure good concentration

lenner - 27-6-2015 at 03:53

Whats the best way to produce AlBr3 in a solvent? Mixing Al with bromine will form AlBr3 but the reaction is very exhotermic. In a solvent water or oxidation could be avoided as much as possible. I have doubts about what type of solvent would be adequate and what type of precaution should I take since the reaction is so exhothermic (which is what keeps me from just trying it).

Oscilllator - 27-6-2015 at 21:13

Quote: Originally posted by lenner

Whats the best way to produce AlBr3 in a solvent? Mixing Al with bromine will form AlBr3 but the reaction is very exhotermic. In a solvent water or oxidation could be avoided as much as possible. I have doubts about what type of solvent would be adequate and what type of precaution should I take since the reaction is so exhothermic (which is what keeps me from just trying it).

I think Dichloromethane would be ideal here as it is non-flammable, bromine is miscible in it and does not react with bromine. Just use an ice bath and/or a reflux condensor to catch the inevitable DCM fumes.
Provided the reaction is spontaneous (It might not be) this should be a very safe way to do it, as the bromine will always be quite dilute. The only issue you might have is adding a whole bunch of bromine at the start and then having the reaction take off and boiling away all your DCM.

xfusion44 - 28-6-2015 at 00:30

Can mineral spirits (white spirit) go bad? I have (probably) pretty old can of it and it should be colourless, but mine seems to be a little bit brown/orange (maybe there's some rust in can, which causes that?) and it's also imposible to ignite it, but it should burn... I'm confused.

Thanks for answer

xfusion44 - 28-6-2015 at 12:50

Quote: Originally posted by Molecular Manipulations

Quote: Originally posted by Brain&Force

Will the use of a steel file impart any significant amount of steel contamination to the powder?

Thanks

byko3y - 29-6-2015 at 11:07

lenner, see US patent 2975214 for preparation of AlBr3 in different solvents.

learningChem, concentrated sulfuric absorbs water from air as crazy. Some moisture guard is a must, the highest concentration I've obtained without using the protection is something like 92% in a receiver and 96% in a source pot.

greenlight, 98% sulfuric acid can be dehydrated above 300°C, 90% needs more than 270°C, 80% needs at least 210°C.
You need to make at least 50% sulfuric acid content in a resulting mixture (120°C dehydration temperature for H2SO4-H2O system) to leave no way for the nitric acid to come dilluted from the mixture. Of course, I'm supposing you have azeotropic or higher initial concentration of nitric acid.
http://www.qvf.com/qvf-process-systems/mineral-acids/concent...

[Edited on 29-6-2015 by byko3y]

learningChem - 6-7-2015 at 23:42

Thanks byko3y.

So I did a quick test. Heated a few mls of ~94% sulfuric acid in a small beaker. At ~120C white fumes can already be seen, but I don't think there's any water in them. I tried to condense the vapor on a piece of glass and didn't see any kind of droplets. I did see condensation/refluxing on the inside walls of the beaker, a few milimeters above the acid surface, but I suppose that's just acid (though it's a bit strange that the drops 'climb' the wall at that low temperature?).

Anyway, just as I originally thought, removing water from already concentrated acid does require a fair amount of heat/time? (apart from some sort of moisture guard)

The Volatile Chemist - 10-7-2015 at 11:05

Quote: Originally posted by xfusion44

Can mineral spirits (white spirit) go bad? I have (probably) pretty old can of it and it should be colourless, but mine seems to be a little bit brown/orange (maybe there's some rust in can, which causes that?) and it's also imposible to ignite it, but it should burn... I'm confused.

Thanks for answer

Red-orange stuff is probably rust, and it may not burn because it has collected a lot of water from the air over the years. Try distilling it and see how much water's in it. The distillate should definitely a-flame.

Gooferking Science - 13-7-2015 at 09:44

When setting up a vacuum distillation, are there any precautions I should take to avoid damaging my vacuum pump? I'm afraid that vapor from the distillation will get in the pump and damage it.

gdflp - 13-7-2015 at 10:32

That depends on what type of pump you have as well as what you're planning on distilling. I'm assuming that you're using a rotary pump, in which case, and solvents or corrosive vapors will need to be trapped. The standard way of doing this is to use a liquid nitrogen trap, but LN is difficult to come by for most. An alternative, is to use a dry ice/acetone trap, this will still condense most gases. If what you're distilling an aqueous or other reasonably noncorrosive compound, the vacuum pump can be used without a trap, or with only a water trap, but I would change the pump oil quite frequently, every 1-3 months depending on usage.

If you have a diaphragm pump, then you don't need to worry. Very few things can attack a diaphragm pump, I have never run into a situation where I need to use a trap.

kecskesajt - 15-7-2015 at 02:31

What happens if I nitrate ascorbic acid?

learningChem - 18-7-2015 at 23:15

At what temperature does hydrated sodium bisulfate become anhydrous sodium bisulfate?

I heated the hydrate for a while and it boils and turns into a syrup-like liquid. Then at some point it mostly stops boiling and white fumes start to appear. I'm wondering if at that moment the substance is anhydrous bisulfate, or rather pyrosulfate?

According to wikipedia the mp of anhydrous bisulfate is 315C but it also lists 315C as the decomposition point so I'm not sure if heating to 315C will dry the bisulfate or decompose it.

[Edited on 19-7-2015 by learningChem]

Zephyr - 19-7-2015 at 17:55

I recently got an old Gast Rotary Vane vacuum pump, model 0522-V3-G180X.
Its not in very good shape, and hums without spinning when switched on.
Does anyone here know much about this pump or have any ideas on how to fix it?
Thanks!

Texium - 20-7-2015 at 15:31

Does anyone know what compound(s) are responsible for the characteristic smell of cork? I've always really liked the smell, which always seems to be pretty much the same whether it's in the form of a cork stopper, a bulletin board, or the backing of my ruler. I wasn't able to find anything helpful upon Googling it. Some results about "cork taint" in wine came up, although apparently that is not always caused by cork, and if it is, it is irregular and unrelated to compounds found commonly in cork.

Zephyr - 22-7-2015 at 17:32

Finally got the pump working!
Huge thanks to Steve_hi for all his help.

Oscilllator - 23-7-2015 at 17:04

Quote: Originally posted by kecskesajt

What happens if I nitrate ascorbic acid?

I'm no expert, but ascorbic acid is a decent reducing agent (it can reduce Cu(II) to Cu(0)) and Nitric acid is a strong oxidiser. So I don't think you will actually get an ascorbic acid nitrate, more likely horrible mess.

chemrox - 23-7-2015 at 17:39

Back to your question: I would use an NaOH filter between the glassware and the pump. I don't have a cylinder large enough to make one and I stripped an intermediate that had some organic acid and anhydride vapors in my vacuum oven. I changed my pump oil afterward. None of my tubing is big enough and filter flask though thick walled are flat. Not sure about pulling 28 psi on one of these for 12 hours. It sure helped seal the door gasket ;^)

j_sum1 - 25-7-2015 at 02:57

Is there a simple qualitative test for Al3+ ions in solution?

kecskesajt - 25-7-2015 at 04:24

Quote: Originally posted by j_sum1

Is there a simple qualitative test for Al3+ ions in solution?

Add NaOH solution.If Al ions are present,a goo will form.

j_sum1 - 25-7-2015 at 04:40

Well, that's a nice idea. But it might not work in practice. I have some element wire that I have dissolved in acid. It appears that it is not nichrome as I originally thought. I am wondering if it is Kanthal (Cr-Fe-Al). NaOH will also precipitate Fe (which I know is present) as well as the Cr. Something a bit more selective than NaOH would be preferable.

gdflp - 25-7-2015 at 06:43

You could use the alizarin test as described here if you have alizarin.

Alternatively, you may be able to add a large excess of sodium hydroxide, then centrifuge the solution. Fe(OH)₃ and Cr(OH)₃ are insoluble in this solution and will precipitate, but Al(OH)₃ dissolves in an excess of sodium hydroxide to give sodium aluminate. After centrifuging, decant the supernatant and slowly add a dilute solution of hydrochloric acid(or another acid). If a precipitate forms, then this should be reasonable confirmation of the presence of aluminum, though this precipitate will redissolve if too much acid is added.

j_sum1 - 25-7-2015 at 07:12

Thanks for that.
Alizarin looks to be reay useful and versitile as an indicator. This video mentions Cr, Fe, W, and group II as well as Al.
https://m.youtube.com/watch?v=IFlIi_Xk3FE
The Wikipedia article mentions Ca as well as the Mg mentioned in the vid. It looks like it takes some care and a bit of knowledge to use. I'd love to know more.

I'm inspired by aga's current dissolving and analysis of rocks. This could be a good tool.

DraconicAcid - 25-7-2015 at 12:23

Quote: Originally posted by gdflp

You could use the alizarin test as described here if you have alizarin.

Alternatively, you may be able to add a large excess of sodium hydroxide, then centrifuge the solution. Fe(OH)₃ and Cr(OH)₃ are insoluble in this solution and will precipitate, but Al(OH)₃ dissolves in an excess of sodium hydroxide to give sodium aluminate. After centrifuging, decant the supernatant and slowly add a dilute solution of hydrochloric acid(or another acid). If a precipitate forms, then this should be reasonable confirmation of the presence of aluminum, though this precipitate will redissolve if too much acid is added.

Chromium(III) is also amphoteric, and will dissolve in excess sodium hydroxide. If your solution is not colourless, add hydrogen peroxide and heat it to convert the chromium anion into chromate, then acidify.

[Edited on 25-7-2015 by DraconicAcid]

j_sum1 - 6-8-2015 at 22:16

I bought some oxalic acid from the hardware store. Is it likely to be anhydrous or dihydrate? It might make a difference if I want anything in a stiochiometric ratio.

gdflp - 7-8-2015 at 09:06

It is most likely the dihydrate. In my experience, most MSDS's specify whether it is the dihydrate or anhydrous, and all I have seen have specified the dihydrate. Anhydrous oxalic acid is typically unavailable unless you go to a lab supplier.

Argentum - 16-8-2015 at 15:14

I will be working with fuming nitric acid soon and I need some help with choosing my gloves. Nitrile and latex are not an option, but what about vinyl gloves? I heard that (butyl?) Gloves work fine, but they are pretty difficult to find, and not cheap at all. Any other choice?

Safety first, always!

Question

Volanschemia - 19-8-2015 at 02:08

Does anyone know the solubility of Tin(IV) Nitrate and Tin(II) Nitrate?

omnilytic - 19-8-2015 at 04:53

Hi guys,

I am trying to conduct capillary electrophoresis on salt water, and was wondering if you had any ideas for getting rid of the sodium and chloride?

So far viable options include using an ion exchange resin for chloride and applying complexing agent, such as a crown ether to the sodium.

Cheers.

Gooferking Science - 20-8-2015 at 06:33

I have a math course question. I am trying to decide whether I should take precalculus or trigonometry. I plan on doing calculus 1, 2, and 3 in the future. Also, the trig class offered is not at my high school. It is at a local junior college and is called "Plane Trigonometry". The precalculus class at my high school is an entire year. The trig class at the junior college is only two months, but is said to be a very rigorous course. My question is: Would I be just fine taking trig then calculus, or would I be missing important concepts covered in precalc? Just FYI I took college algebra this summer, so I have a good knowledge of college level algebra already. I have heard that precalc is considered to be trig + college algebra, is this true? Sorry if this is a confusing post, but I'm kind of freaking out since school starts next week and my schedule is all messed up still!

solo - 20-8-2015 at 06:40

.....take pre-calculus, there will be time for trig and plane geometry down the line....solo

The Volatile Chemist - 21-8-2015 at 12:25

Any opinions as to how easy it is to skip from AP Calc AB to BC? I know my derivatives...

j_sum1 - 5-9-2015 at 21:57

Made something pretty by mistake.
Nice pink shade of crystalline solid. Totally unexpected in what I was doing. I was wondering if anyone could cast some light onto what it is.

I was distilling some paint stripper for dichloromethane. The MSDS says that it is methylene choride, toluene and methanol. The label is more specific: 870g/L methylene chloride, 130g/L methanol and "contains toluene". The actual product evidently contains other things as well -- it is a viscous green gelatinous goop. I assume there is some kind of dye in there as well.
Anyway, I distilled. Distillate was cloudy but uncoloured. I washed with sodium chloride solution, separated with a sep funnel and dropped it in a flask containing a small quantity of anhydrous calcium chloride to remove the remaining methanol and water. Distillate came over at just under 40°C as expected. Just at the last stage of distillation a pink/magenta coloration appeared on the calcium chloride. Like I said, not expected.

Later when I was cleaning up, I added a reasonable quantity of water to wash the flask. It occurred to me that the colour was reminiscent of phenophthalein in basic solution. On a whim I added some to a test tube and added a drop of HCl. The pink colour disappeared. (The test tube also filed with a white cloudy vapour.) I also tried adding some NaOH and (unsurprisingly) got a cloudy precipitate of Ca(OH)2. The pink colour remained.

Thinking that this might be worth investigating further I decided to boil my solution dry. The solution at this stage was pretty weak and the colour was not vivid. On applying heat the colour disappeared completely from the solution. In the final stage of recrystallisation, a little pink reappeared but not nearly as bright as my original crystals.

Unless I hit upon an enormous fluke I am guessing that I did not make phenolphthalein. But what I did produce has some similar properties. Any idea what it might be?

mayko - 6-9-2015 at 20:49

Why does acetaldehyde seem to be so hard to come by? It don't know of it being a precursor, and I don't think it's as toxic as the widely-available formaldehyde

PHILOU Zrealone - 7-9-2015 at 08:25

Quote: Originally posted by mayko

Why does acetaldehyde seem to be so hard to come by? It don't know of it being a precursor, and I don't think it's as toxic as the widely-available formaldehyde

It has some use as cyclic tetramer (metaldehyd) pesticide for mollusces but toxic for childrens, cats, dogs,...neurotoxic.
In Europa, those pesticides have lowered their active % to diminish the risk of poisoning from 80% to less than 5%.

The cyclic trimer (paraldehyd) is used as hypnotic/sedative, also neurotoxic.

This might explain the difficulty to obtain it.

The Volatile Chemist - 11-9-2015 at 12:29

Yes...indeed. Is it still common in the scholarly lab, or is it one of the many one has to make? I don't recall seeing it in Vogel's.

Eddygp - 17-9-2015 at 02:41

Is there any feasible way for a home chemist to substitute a methyl group attached to a nitrogen (N-methyl) to a 2-propynyl group (N-(2-propynyl))?

[Edited on 17-9-2015 by Eddygp]

The Volatile Chemist - 17-9-2015 at 14:12

Do you know of any feasible routes that the amateur 'can't' do? I can't think of anything from Vogel's or 'E-Z Organic Chemistry', but I have very little education in such.
Also, this thread is getting rather long...

[Edited on 9-17-2015 by The Volatile Chemist]

Crowfjord - 17-9-2015 at 14:26

@Eddygp: I think it depends on the type of amine. I know that a methyl group can be removed from a tertiary amine, with the Von Braun reaction being the classic example. The demethylated amine could probably then be reacted with 2-propynyl iodide or bromide. I don't think I have ever heard of a way to demethylate a secondary amine.

[Edited on 17-9-2015 by Crowfjord]

Eddygp - 18-9-2015 at 08:04

Quote: Originally posted by Crowfjord

@Eddygp: I think it depends on the type of amine. I know that a methyl group can be removed from a tertiary amine, with the Von Braun reaction being the classic example. The demethylated amine could probably then be reacted with 2-propynyl iodide or bromide. I don't think I have ever heard of a way to demethylate a secondary amine.

[Edited on 17-9-2015 by Crowfjord]

I was actually thinking about the ones in cyclic compounds, either aromatic or not. So yes, mostly tertiary amines. In fact, there are a few secondary ones in guanine that, if their H were to be substituted for an R group, they would be like the ones I mentioned. I was thinking about methylated uric acid, actually.

[Edited on 18-9-2015 by Eddygp]

Crowfjord - 18-9-2015 at 11:35

Sounds interesting. I think that the Von Braun reaction would work on such a substrate, but cyanogen bromide is pretty nasty. There is also a demethylation reaction that employs alkaline aqueous potassium ferricyanide, which works on some tertiary amines (see JOC, 1951, 16(8), pp 1303-1307). Perhaps it could could work on your substrate.

MeshPL - 19-9-2015 at 14:01

Can amides be N-alkylated by haloalkane or DMSO4?

Can I make nitromethanes by adding lighter, but liquid hydrocarbons to an excess fuming nitric acid and maybe even bubbling some NO2 as well?

Can I make 2,5-dimethylcyclohex-2-enone from 4-methyl-5-oxohepthanal via aldol elimination? (As a note other byproducts would likely have to contain a 4-membered ring so they wouldn'be favoured. Also not sure if I named those compounds properly)

The Volatile Chemist - 19-9-2015 at 14:04

Interesting. If the ferricyanide works, the procedure is feasable for even I. I wonder if it would work on a phase transfer catalyst, like Aliquat 336.

MeshPL - 19-9-2015 at 14:15

What concerns turning methylamines into isopropylamines... maybe (if wikipedia says so...) you could turn tertiary amine into amine oxide and make it undergo Meisenheimer rearrangement or Polonovski reaction. Maybe you could attempt to hydrolyse resulting hydroxylamine or acetamide/iminium. Look those reactons up on the net.

Crowfjord - 20-9-2015 at 09:47

Concerning N-demethylation, it seems that that one paper I sited earlier is the only reference that uses ferricyanide. I, at least, could not find any more references on the subject. No follow-ups or anything. This could be a ripe area for amateur research, as the study in that old paper was not exactly thorough. I'll post the text in the references section, in Recent Journal Articles of Interest.

Also, I found this paper, for another N-demethylation method.

[Edited on 20-9-2015 by Crowfjord]

Morkva - 20-9-2015 at 10:16

A method of preparing arylsulfonic acid chlorides is to oxidize the thiol with chlorine in glacial acetic acid. If the oxygen is coming from the acetic acid then must not the byproduct be acetyl chloride??

Either that or the CO2 elimination product methyl chloride, or am I missing something?

The Volatile Chemist - 22-9-2015 at 18:20

Are athere any common modifications to citric acid known which ar feasable to the amateur? Oxidation seems to just leadto decomposition. Obviously I'm asking about organic reactions, not making salts of it. I don't suppose one can chlorinate the hydroxyl group and collect the product?

Zephyr - 22-9-2015 at 18:29

An easy conversion is to citrazinic acid by melting with urea, although the yields aren't great:
https://www.youtube.com/watch?v=jbw9dt12qsI

Solubilities

AngelEyes - 23-9-2015 at 01:15

So...short('ish) question...hopefully there'll be a quick answer.

When I was at school we were taught many things, some of which I even learned, but one of them was that all sodium salts are soluble and all nitrates are soluble.

Now, it's been a while since I was in a place of learning but are those two 'facts' still correct? If so, then why sodium and nitrates...what's so special about them? Are there any others that fall into the 'always soluble' category? And by that I mean appreciably soluble, not <1g / L
And if it's not true any more, then what salt was found that violated the rule?

Cheers

Angel

DraconicAcid - 23-9-2015 at 07:02

Quote: Originally posted by AngelEyes

So...short('ish) question...hopefully there'll be a quick answer.

When I was at school we were taught many things, some of which I even learned, but one of them was that all sodium salts are soluble and all nitrates are soluble.

Now, it's been a while since I was in a place of learning but are those two 'facts' still correct? If so, then why sodium and nitrates...what's so special about them? Are there any others that fall into the 'always soluble' category? And by that I mean appreciably soluble, not <1g / L
And if it's not true any more, then what salt was found that violated the rule?

It's not that any new salt was found that violated the rule, it's just that the exceptions are unimportant for first-year chemists (or high school students). I teach my students that all alkali metal salts are soluble, despite knowing that there are insoluble potassium salts (perchlorate and chloroplatinate) and insoluble sodium salts (sodium zinc uranyl acetate hexahydrate, according to my 1961 copy of Vogel), and that all nitrates are soluble, despite knowing that the green isomer of [Co(en)2Cl2]NO3 is insoluble.

Brain&Force - 23-9-2015 at 10:16

I have a sizable chunk of dysprosium ( ‎http://www.imgur.com/BD86qse ) but I need a way to break it into manageable chunks, preferably without losing any small pieces. Wikipedia said it should be possible to cut it with a knife but that hasn't worked at all. Got any ideas on how to efficiently break this into ~5 more manageable chunks?

[Edited on 23.9.2015 by Brain&Force]

PHILOU Zrealone - 25-9-2015 at 11:06

Quote: Originally posted by The Volatile Chemist

Are athere any common modifications to citric acid known which ar feasable to the amateur? Oxidation seems to just leadto decomposition. Obviously I'm asking about organic reactions, not making salts of it. I don't suppose one can chlorinate the hydroxyl group and collect the product?

Look at:
-citraconic acid (HO2C-CH=C(CH3)-CO2H)
-aconitic acid (HO2C-CH=C(CO2H)-CH2-CO2H)
-acetonedicarboxylic acid (HO2C-CH2-CO-CH2-CO2H)
All made from citric acid and common chems or processes.

The Volatile Chemist - 5-10-2015 at 12:50

Quote: Originally posted by Pinkhippo11

An easy conversion is to citrazinic acid by melting with urea, although the yields aren't great:
https://www.youtube.com/watch?v=jbw9dt12qsI

I saw that, and tried it... It all turned to tar... And I couldn't get anything to precipitate...

Quote: Originally posted by PHILOU Zrealone

Quote: Originally posted by The Volatile Chemist

Look at:
-citraconic acid (HO2C-CH=C(CH3)-CO2H)
-aconitic acid (HO2C-CH=C(CO2H)-CH2-CO2H)
-acetonedicarboxylic acid (HO2C-CH2-CO-CH2-CO2H)
All made from citric acid and common chems or processes.

Thanks, I might try them.

gluon47 - 7-10-2015 at 22:25

Can nitric acid be feasibly produced via the reaction of potassium nitrate and hydrochloric acid? I've seen many vids of people making it with sulphuric acid and potassium nitrate but none using hydrochloric acid.

j_sum1 - 7-10-2015 at 23:40

Add some copper and your answer is yes.
Lower yield. Lower concentration possible. And in many parts of the world, sulfuric acid is more readily available than hydrochloric (but not where I live). That's why it is used.

j_sum1 - 10-10-2015 at 15:58

The pic below shows a nice gas discharge tube.
But what exactly is causing it to glow? How might I construct something like this?

[Edited on 10-10-2015 by j_sum1]

The Volatile Chemist - 11-10-2015 at 12:02

That's cool! I assume it's argon filled. I totally forget how that happens, though I'm sure you could do something similar with a fluorescent bulb tube placed next to a source of charge.

j_sum1 - 26-10-2015 at 03:23

Bump.

And another question...

I was demonstrating a precipitation reaction to my class today -- as a means of separating. I mixed up some copper sulfate pentahydrate and sodium chloride and then dissolved it. I challenged the students to select something that would cause a copper compound to precipitate. Sodium sulfide was chosen -- which I thought was a good idea. Copper (II) sulfide is nice and insoluble and I thought would filter out well.

What I was not expecting is that the filtrate was a lemon yellow colour. My first thought was that the sulfide had reduced the copper ions to Cu(I). I added a little peroxide to see if Cu would oxidise back to a blue colour. Instead, the solution went clear. I am confident that there is no Cu in the solution. But I am unsure what might be giving the yellow colour. Maybe some Cl-S-O anion that I am unfamiliar with?

The Volatile Chemist - 27-10-2015 at 17:03

Quote: Originally posted by Brain&Force

I have a sizable chunk of dysprosium ( ‎http://www.imgur.com/BD86qse ) but I need a way to break it into manageable chunks, preferably without losing any small pieces. Wikipedia said it should be possible to cut it with a knife but that hasn't worked at all. Got any ideas on how to efficiently break this into ~5 more manageable chunks?

[Edited on 23.9.2015 by Brain&Force]

heat it up and trow it in an ice bath?

Sorry, my ideas always manage to be terrible, but it seems it might work if Dy has a high thermal expansion coefficient. For $42 here's your answer...(http://www.sciencedirect.com/science/article/pii/00219614849...)
This article has expansion data for the rare earth oxides (and is thus interesting, worth posting) but n/a (http://www.osti.gov/scitech/servlets/purl/4840970/).
The coefficient is here (https://books.google.com/books?id=SFD30BvPBhoC&pg=PA160&...) and it looks rather small, probably not enough to shatter it, but who knows?

hissingnoise - 1-11-2015 at 06:49

DEA throwing in the towel ─ here's acetic anhydride for sale on ebay . . . ???

Pricy, too!

battoussai114 - 1-11-2015 at 17:51

I was checking the mechanism for Citric Acid / Ethylene Glycol gel formation (studying sol-gel deposition for an upcoming project) and based on what I saw it seems that any diol would work for the purpose of polymerizing and getting the gel part of sol-gel... since the Citric Acid is responsible for locking the metal cations in place, the diol characteristics seem minor as it's only there to form a polyester.
Am I missing something? Why do people specifically use EG?

[Edited on 2-11-2015 by battoussai114]

Tabun - 2-11-2015 at 07:34

Does anybody know a book on drugs synthesis?Not recipes nor "easy to make drugs"...I want to understand exactly what and how it's happening.I can come up with some ideas but only in theory and I don't really know where to look for more info so a book on this subject would be the best.I don't want to cook or anything similar.I mean...there's a book with war gases in this site's library and explosives are a part of this site so you should understand what I'm trying to say...I just like them and I'm interested in them.

aga - 2-11-2015 at 07:46

Alexander Shulgin's TiHKAL and PiHKAL seem to be The books on drug synths.

mayko - 8-11-2015 at 22:01

I've tried to make iron sulfide from the elements, once by stoichiometric proportions and once by teh 7:4 mass ratio prescribed by Nile Red. I can't seem to get it to ignite! Touching with hot glass or using a magnesium ribbon fuse just creates burning blue sulfur, and a direct blowtorch flame kicks up a bunch of golden burning iron sparks. Is there something I am missing?

JJay - 8-11-2015 at 22:20

What size particles of iron are you using?

Eddygp - 10-11-2015 at 08:19

What can I do with 2.7g of bis(tetraethylammonium) tetrachlorocuprate(II)?

DraconicAcid - 10-11-2015 at 08:36

You can complain about the fact that you didn't make it in time for the copper carnival.

Mesa - 10-11-2015 at 11:50

Is it possible to efficiently recover an amine from a silicotungstic or phosphotungstic acid complex?
Is there a significant variation in properties of these if the complexed amine is primary, secondary, or tertiary?

mayko - 10-11-2015 at 13:07

Quote: Originally posted by JJay

What size particles of iron are you using?

Iron filings, not otherwise specified

The Volatile Chemist - 10-11-2015 at 15:30

Quote: Originally posted by Eddygp

What can I do with 2.7g of bis(tetraethylammonium) tetrachlorocuprate(II)?

Wow...where'd you get that one...?
You could 'still out the tetraethylammonia if you don't own any, though that's a bit wasteful...

karlos³ - 12-11-2015 at 11:05

Can somebody tell my if its possible to prepare mandelic acid from styrene?
I need some mandelic acid, I´ve ordered some but my supplier is taking too long to deliver it.
Styrene is on hand, lots of it.

I know it is possible to make mandelic acid from styrene via enzymatic synthesis, but I don´t know if it can be chemically made.
Eating some styrene and extracting my urine to get mandelic acid is not an option!

Is there an easy and chemical way to prepare this compound from styrene?
Otherwise I will just wait till it gets delivered.

I´ve searched far and wide but did not get any references for its chemical synthesis, its only mentioned in the measurement of exposure to styrene via urine analysis.

Texium - 12-11-2015 at 14:16

Does anyone know how to use these stemless fritted funnels? I've always used the kind with stems, but I found one of these at school and want to use it for a project I'm working on.

User123 - 12-11-2015 at 14:21

It's a Gooch crucible.

aga - 12-11-2015 at 14:26

Have you ever used a Gooch crucible ?

If so, please explain how one is used (i already read the Google results).

What is not clear in the photo is whether the frit is open at the bottom, or if that is a sealed bottom with a stand, so it might be something different.

Texium - 12-11-2015 at 14:36

It is open at the bottom, it does appear to be a gooch crucible. But I second aga: how do you use it? It can be used as a normal fritted funnel still, can't it?

User123 - 12-11-2015 at 14:48

You need a Buchner flask and a ring of rubber.

CharlieA - 12-11-2015 at 17:54

You have to first filter a suspension of some type of medium, by vacuum, to put down a filter "mat" over the perforations in the bottom. Then you filter your product. Back in the day (I won't say when), we used a suspension of asbestos fiber to make the mat.

CharlieA - 12-11-2015 at 18:02

What is the general opinion of using a tank of propane (like the kind used in a BBQ grill) to fuel a Bunsen burner? hBarSci doesn't recommend it, but then proceeds to give a detailed list of the parts needed, various brass fittings, mainly. All brass to brass joints are sealed with a Teflon tape similar to plumber's joint tape, but I think it is a more particular kind of tape, for this application.

Texium - 12-11-2015 at 18:27

I use a barbecue propane tank with my Bunsen burner and it works great, but it requires a special type of regulator. Fortunately for me, I was able to get one from someone who used to use a similar setup but didn't need it anymore. I'm not sure exactly what type of regulator it is, but I should be able to post a picture of it later.

CharlieA - 12-11-2015 at 18:36

zts16: thanks for the quick response. I would very much like to see a photo, when it is convenient for you to post one. Perhaps your valve has some kind of model number/company name on it That would be good to know too.

The Volatile Chemist - 15-11-2015 at 13:55

Does anyone use the typical lab-type/school-type Bunsen burners at all in their home labs? They seem impractical for the amateur...

DraconicAcid - 15-11-2015 at 14:10

Quote: Originally posted by CharlieA

You have to first filter a suspension of some type of medium, by vacuum, to put down a filter "mat" over the perforations in the bottom. Then you filter your product. Back in the day (I won't say when), we used a suspension of asbestos fiber to make the mat.

No, you don't. We use just a straight frit to collect precipitates all the time.

CharlieA - 15-11-2015 at 18:45

I believe that, strictly speaking, a Gooch crucible has a perforated bottom, not a frit. If some kind of mat is not laid down over the perforations ("holes"), about the only thing you can filter will be BB's or marbles.

gdflp - 15-11-2015 at 18:57

Depends on the crucible. Glass Gooch crucibles typically have fritted bottoms, while porcelain Gooch crucibles need some kind of additional filter aid such as an asbestos mat.

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