Sciencemadness Discussion Board

chloral nitrate?

woelen - 8-7-2007 at 12:57

I now have a small quantity of chloral hydrate. This is a compound, with two hydroxy groups on a single carbon atom. In some sense, it can be regarded as an alcohol:

Cl3C-CH(OH)2

Many alcohols can be nitrated, and as such, I tried with chlorale hydrate. I mixed concentrated nitric acid with an equal volume of concentrated sulphuric acid, and carefully added just a few crystals of chloral hydrate.
When this is done, then there is no visible reaction. I heated a very little bit (under a running tap. with hot water, appr. 60 C). Still no visible reaction. I let it cool down again. After that, nothing interesting was visible and I loosely stoppered the test tube, and left the whole mix in a place at room temperature.

Several hours later, I came back, and to my surprise I found this:



There are beautiful long needle-like transparent crystals at the inside of the test tube. These crystals, however, quickly disappear, when water is allowed to enter the test tube. A single drop of water makes them disappear quickly.
There was no strong smell, just as with chloral hydrate itself, which also only has a faint smell.

Could this be chloral nitrate, CCl3CH(ONO2)2?

Eclectic - 8-7-2007 at 13:10

That looks like the sort of crystals you get when s-trioxane (formaldehyde trimer) sublimes. Maybe this is the trimer of anhydrous chloral?

You could try the same experiment without the nitric acid...

garage chemist - 8-7-2007 at 13:18

I dont think so, chloral nitrate would be a very unstable compound. The crystals are likely simply chloral hydrate.

From chloral hydrate and four times its volume of concentrated sulfuric acid, you can obtain chloral (CCl3-CHO) by gentle heating, which is a liquid and separates as an immiscible layer from the H2SO4. By adding water to the chloral, the hydrate is formed again.

By the action of NaOH solution upon chloral hydrate, it is cleaved to chloroform and formate, which was once used as a method to obtain extremely pure chloroform completely free from CCl4, which is always an impurity in chloroform prepared by conventional means like haloform reaction or methane chlorination.

The action of nitric acid on chloral hydrate is used to prepare trichloroacetic acid by oxidation. Red nitric oxides are produced in this reaction, which is initiated by heating. Try heating some chloral hydrate with 65% HNO3 until red fumes are produced.
Trichloroacetic acid is also cleaved to chloroform by an excess of NaOH.

How do you always make such good photos of test tubes and other near objects? Do you zoom in on the object, and what distance of the camera to the object do you use?
I find it difficult to make good photos of things like crystals in flasks and reactions.

Nicodem - 8-7-2007 at 23:13

Chloral hydrate can not be nitrated in acidic ambient since it eliminates water to form chloral. Also, the geminal substitution tend to be most stabile with the increasing nucleophilicity of the substitents even if stabilized with an electron withdrawing group like in chloral hydrate. Since the nitrate is so many magnitudes less nucleophilic than hydroxide there is little chance that chloral hydrate dinitrate could be an existing compound at any given condition. Furthermore if it existed, it could only be at extremely low temperature, since otherwise the elimination of HNO2 would occur and thus you would end up with trichloroacetic acid. Those crystals might just be trichloroacetic acid after all. If you want a confirmation, try to recover some, recrystallize from petroleum ether (or whatever alkane mixture you have access to) and measure its mp.

PS: Nice picture!

woelen - 9-7-2007 at 04:19

Good to read these responses. I'll certainly retry with different experimental conditions (H2SO4 only, HNO3 only). I already had some feeling that the nitrate or dinitrate would not be formed (otherwise I could have found references on Internet), but I just wanted to know for sure.

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For this kind of pictures, I use a Pentax Optio S 3.2 MPixel digital camera. It is a very small camera, which I keep in one hand, and the test tube I keep in the other hand. The small size of this camera and ease of operating it, combined with its strong macro properties, allow me to make these pictures. The distance of the test tube from the camera is quite low, for this picture it is over 10 cm, but for some of the more detailed pictures, it can be as low as 6 cm.

http://www.dpreview.com/news/0301/03010801pentaxoptios.asp

For general photography, this certainly is not the best camera, but for macro photography it remains unbeaten (except maybe by some newer models of the Pentax Optio series).

I also have a Canon Powershot A620. This is a great camera for general photography and it makes nice movies, but for macro photography it is not nearly as good as the Pentax Optio S. It is much more bulky and I hardly manage to control this camera with one hand, while keeping the test tube with dangerous fuming hot stuff :D in the other hand. For the same macro level, you also need to go closer to the object you want to make a picture of. You need to go as near as 1.5 cm, but the perspective distortion then is terrible for objects, which are not perfectly flat, and the depth of sharpness width then almost drops to 0, making good pictures of the round test tubes very hard to make.

Most important, however, I found, is the very small size and weight of the Pentax Optio S, combined with the ease of controlling this camera with one hand. The other hand really is 100% free to do other things.

YT2095 - 9-7-2007 at 04:26

it`s a Perfect pic, I thought I recognised that test tube ;)

woelen - 11-7-2007 at 14:11

I did the experiment with only H2SO4 (96%, lab grade, no drain cleaner). This results in formation of an oily layer on top of the sulphuric acid layer, just as described by garage chemist. No special crystals are formed. I let this material settle for two days. After these two days, I had a lot of white crystals along the glass wall of the test tube, and the oily layer had become somewhat waxy. Apparently, the material has picked up water, despite the fact, that I had stoppered the test tube with a rubber stopper. Most likely, the water was taken out of the H2SO4 again and made the chloral into the chloral hydrate again.
All liquid material still was perfectly colorless and all solid material still was perfectly white after these days. So, no charring or condensation into brown crap occurred.

I also did the experiment with only HNO3 (52%). I did not manage to get any reaction. I boiled the liquid for some time, but no brown vapor/gas is produced. The solid chloral hydrate simply dissolves in the HNO3, and that is all. So, there is no vigorous (runaway) oxidation of the chloral hydrate. Could it be that chloroacetic acid is formed? But if that is the case, then why did I not see the brown NO2?
After two days, I still have a colorless solution, no separation of crystals, nothing but colorless liquid.

Altogether, I can conclude for 100% certainty that no chloral nitrate is formed. Most likely there was no oxidation at all.

[Edited on 11-7-07 by woelen]

garage chemist - 13-7-2007 at 04:00

I looked up the synthesis of trichloroacetic acid from chloral hydrate, fuming nitric acid (density 1,5) is required for the oxidation. So your 52% HNO3 probably did not react.

The fuming HNO3 is slowly added to molten chloral hydrate as it reacts. From the crude reaction mixture, trichloroacetic acid is isolated by direct distillation in vacuum. Its melting point is 57°C.

Trichlororacetic acid is used against warts and for skin peelings because it is fat-soluble while being a very strong acid at the same time, allowing its deep penetration through the skin. It is not very toxic apart from its strong acidity.

fractional - 13-7-2007 at 05:33

Quote:
Originally posted by woelen
I let this material settle for two days. After these two days, I had a lot of white crystals along the glass wall of the test tube, and the oily layer had become somewhat waxy. Apparently, the material has picked up water, despite the fact, that I had stoppered the test tube with a rubber stopper. Most likely, the water was taken out of the H2SO4 again and made the chloral into the chloral hydrate again.


The solid formed in the oily layer was most likely the chloral-trimer known as metachloral.
From Merck-Index: "...[chloral] polymerizes under the influence of light and in presence of sulfuric acid forming a white solid trimer called metachloral."

This reaction is reversible. By dry heating the metachloral to about 180 degC you can distill off the chloral.

woelen - 13-7-2007 at 13:39

Fractional, I think you are 100% right. I added a lot of water to the test tube with the H2SO4 and the white solid. The H2SO4 of course reacted with the water, giving off heat and it dissolved/mixed quickly, but the white solid did not dissolve at all. Chloral hydrate, on the other hand, very quickly dissolves in water, so this white solid I have now must be the trimer.

It was also quite funny to see that now, again a few days later, the oily layer had formed a nice disk with a thickness of 1.5 mm or so. From a mechanical point of view, this disk is remarkably stable. Only with very hard/fast shaking of the test tube with water, and the disk in it, it broke in two parts.

fractional - 14-7-2007 at 01:01

Quote:
Originally posted by woelen
It was also quite funny to see that now, again a few days later, the oily layer had formed a nice disk with a thickness of 1.5 mm or so. From a mechanical point of view, this disk is remarkably stable. Only with very hard/fast shaking of the test tube with water, and the disk in it, it broke in two parts.

Yes, that's exactly how I know it.
In my experience this effect can be used for purifying chloral. The standard procedure starts from chloralhydrate, re-crystalling it from chloroform and petrolether. But for this to work the chloralhydrate must be absolutely dry, otherwise two phases will be formed. But you may have noticed it is not that easy to dry chloralhydrate without significant losses.

Metachloral can be washed in H2O will little loss. During the dry distillation step all the fractions with b.p. below 180 degC can be discarded, leaving practically pure chloral at decomposition temperatures between about 180...190 degC.

A very useful synthetic reagent, chloral, but I think your main interests are more in the anorganic line, aren't they?.....;)

woelen - 14-7-2007 at 09:06

I also have been thinking of this as a method to make chloral from chloral hydrate, but only after I read your post. I did not know of the existence of metachloral. But now you tell me, it does not really surprise me, because low-number aldehydes (formaldehyde, ethanal) also polymerize very easily and the ethanal trimer also is known as "meta".

My main line of interest indeed is in inorganic chemistry. This little excursion into chloral is because I recently received a small quantity of chloral hydrate. I now know that nitrating this compound is not possible, but I may try to make small amounts of chloroacetic acid, and see how this reacts with metal ions to make complexes. E.g. chloroacetato-copper(II) or something like that, and how it compares to normal acetato-copper complexes. But from previous posts, I understand that chloroacetic acid tends to decompose, giving chloroform at higher pH. Is it possible at all to make chloroacetates?

The reason why I only do little experimenting in organic chemistry, is that such things usually require non-aqueous solvents (with all vapor and health issues involved + more waste and more cost). Also, organic reactions frequently are slower and more difficult to visualize, and more glassware and apparatus is needed for interesting organic chemistry. At the moment I simply do not have the room in my house, where I can do that kind of things.

fractional - 14-7-2007 at 13:10

@woelen
It's definitely possible to make chloroacetates. Here are some references I could find:
a.) In an old edition of the Merck-index: "Sodium salt, C2Cl3NaO2, sodium trichloroacetate...Yellow deliquese powder, mp > 300deg. Solubility in water at 25deg: 1.2 kg/l, sol in ethanol. Caution: Very corrosive!...Use: As a decalcifier and fixative in microscopy, also as a precipitant of proteine. As herbicide."
b.) Beilstein, 4th ed., Vol.2, p.208, gives a listing of metal salts of TCA, but nothing specific on complexes. I have attached the relevant page ( in German)

Attachment: 0218_0001.djvu (86kB)
This file has been downloaded 461 times

leu - 15-7-2007 at 04:20

Gaatbaug disclosed an addition compound of chloral and nitric acid in Norwegian Patent 54,575 in 1934 :P For more information about the alpha and beta forms of the polymers of chloral see: Chattaway & Kellett, JCS 2709-14 (1928) :cool: