Sciencemadness Discussion Board - Synthesis of malonic acid

Synthesis of malonic acid

Pages: 1 2 3

vulture - 1-3-2004 at 09:06

The usual synthesis route for malonic acid is chloroacetic acid + NaCN.

Now, chloroacetic acid isn't the problem, but NaCN is. A rather large quantity is required and ordering is not really an option. Synthesis of that much NaCN is rather cumbersome and the reagents need to be fairly pure.

Are there any other synthesis methods, maybe archaic ones?

Doesn't matter if the yield is low and the reagents are cheap and/or easily available.

Before you say anything, the malonic acid is required for an oscillating reaction.

If I'd were to use it for barbiturates, I'd just buy it (it's bloody expensive) because the profit would easily cover the expense.

Marvin - 1-3-2004 at 09:44

I dont have much and IVe been looking for a good way for a long time. I want to try the BZ reaction myself, though I think ferroin will be more of a problem.

Only method that looks like it might be doable at home, is oxidation of malic acid with chromic acid mixtures. Its possible nitric acid might work though it could go furthur to oxalic acid, etc.

Still searching for a good way. Decomposition of barbituric acid, oxidation of b-propylene glycol etc are all unlikley at home.

BromicAcid - 1-3-2004 at 10:11

Just inital speculation thumbing through my chemical encyclopedia (un complicated by essentials such as seperation of products and reaction conditions

)

Glycerol and magnesium sulfate react to form acrolein, this can be reacted to produce the 1,3-propylene glycol (I don't know excatly how it's done, under the entry for 1,3-propylene glycol it lists acrolein as what trimethylene glycol is derived from.) Then from here all you need is a good oxidizing agent, acidic dichromate, alkaline permanganate, etc. Thereby malonic acid is obtained.

Jeeze..... so much information is lacking from a complete picture, although glycerol, magnesium sulfate, and an oxidizing agent are easy starting materials to aquire so far...

vulture - 1-3-2004 at 11:41

Oi, it seems that propane-1,3-diol is even more expensive than malonic acid itself!

Geomancer - 1-3-2004 at 14:57

The glycerol->1,3 propanediol routes are what springs to mind. Unfortunately, glycerol seems to have a strong tendancy to dehydrate to acrolein, rather than making a nice long stop at the monodehydrated compound. Why is this? There might be a route there, though; trap the aldehyde as an EG acetal or something. Rehydrating acrolein seems roundabout.
Some bacteria are perfectly happy doing this conversion. It's being studied industrially apropos some new polyesters.

NaCN

Turel - 1-3-2004 at 16:03

Hoffman Isonitrile Synthesis would work.

Aqueous ammonia and sodium hydroxide will react with chloroform to produce aqueous sodium cyanide and sodium chloride. It may require the aid of phase transfer catalysis, as chloroform and water are immiscible.

-T

unionised - 1-3-2004 at 16:17

I think it would need the assistance of the Almighty to work very well. OTOH you can always see if there is a market for the formic acid that is produced by the hydrolysis of chloroform. If there is, then you can sell it and buy malonic acid.

Can you decarboxylate malic acid to give hydroxy propanoic acid then oxidise that to malonic acid?
I have a vague recollection that something like that is where malonic acid gets its name from. (Indirectly, from the latin name for apples)

funny you should mention that

Polverone - 1-3-2004 at 17:36

I tried using this method to produce NaCN some time ago, and tried again just recently. I didn't use a PTC, just added enough 95% ethanol so that everything went into solution. I heated an excess of NaOH with chloroform and clear household ammonia solution (all added to the EtOH) for an hour or so. It formed a light-yellow solution with a foul smell. The liquid failed to give a positive Prussian blue test. Did the ethanol interfere?

Vulture: why is homemade NaCN a problem? The high-temperature route that uses sodium cyanate as an intermediate, which has been discussed here before, is easy to perform. The main impurity in your NaCN made this way will be Na2CO3, and that shouldn't interfere with your synthesis.

further results

Polverone - 3-3-2004 at 22:15

I tried the isonitrile method again, two different ways. In the first, I tried using some swimming pool "winter algaecide" that listed "ACTIVE INGREDIENT: Alkyl (C14, 58%; C16, 28%; C12, 14%) dimethyl benzyl ammonium chloride ... 40%" as a phase transfer catalyst. Where the liquid contacted a bit of solid NaOH clinging to the neck of the flask, it instantly darkened and eventually became almost black. In the liquid mixture of NaOH, CHCl3, and aqueous ammonia (no EtOH this time), it rapidly turned brown with agitation. The brown-gunk was selectively concentrated in the blobs of chloroform sitting in the flask. After moderate heating for 4 hours or so, the chloroform blobs had disappeared. There was considerable brown-black gunk floating on the surface. There was an unpleasant scent like when I ran the experiment with EtOH cosolvent, but it wasn't as strong.

The unpurified liquid (still containing gunk) was mixed with acidified Fe(II)/(III), and gave a handsome Prussian blue.

I repeated the experiment, but used no cosolvent and no PTC. The chloroform took considerably longer to disappear (and I think it mostly disappeared by volatilization from the flask). Only the scents of ammonia and chloroform were noted in the final liquid. It grew slightly cloudy but there was no discoloration. I failed to obtain Prussian blue when the liquid was put to the test.

So, at least when using dilute (household) aqueous ammonia, it appears that a PTC is very helpful, EtOH cosolvent somehow interferes to give something else, and chloroform reacts very little without PTC or cosolvent. It also appears that my expedient PTC choice was not very good, because it gives rise to a bunch of degradation junk. Perhaps a simpler PTC would not degrade.

Actually, I've controlled the variables so loosely, it's impossible to determine for sure what the sources of my successes and failures were. Still, I didn't want the details of my crude and imperfect trials to fade from memory. This was also exciting because it's the first time I've been able to see the effects (or apparent effects) of a PTC.

Please Note

Turel - 3-3-2004 at 22:21

That PTCs only have effect during strong agitation or stirring, not while static. They aid homogenization, not facilitate it alone.

Very interesting results.

[Edited on 4-3-2004 by Turel]

Really?

Polverone - 3-3-2004 at 23:02

I was hand stirring/shaking it at frequent intervals for an hour or so (I don't have any stirbars!) But my understanding was that PTCs aren't used to make the solution homogeneous, but to shuttle back and forth between the aqueous and nonaqueous phases, carrying the desired anion into the nonaqueous phase. I can see that this shuttling would be enhanced by greater contact area between the two phases, but I would expect some effect even if they remain placid. Is this not the case?

True

Turel - 4-3-2004 at 22:12

No, your description of the operation of PTCs is correct. However, as catalysts for two-phase solvated reactions, they are largely ineffective unless used in conjunction with strong agitation or stirring.

I have always been taught that PTCs are worthless while static. The polar PTCs are themselves attracted to the polar solution, so net static diffusion is minimal, and never reaches anything close to the high degree of diffusion created under agitation.

Can you make a stirrer out of plastic or steel? Rapid uneven stirring works best for solvents that are not ruined by aeration, as the cavitation produced causes a very high degree of ion diffusion from the PTC into the nonpolar solvent.

I have never done an isonitrile synthesis myself, I have only read about them. But I have read about them in many places, and I got the drift that they were highly effective in organic synthesis. In this reaction, iso-hydrogen cyanide would be the product, and it is very unstable, most especially in polar solution, where H2O easily removes hydrogen from electropositive nitrogen. The result is a solvated cyanide ion.

Since only 1 in 8 people on average can detect the smell of cyanide, I would not recommend smelling it as a test for it's presence. I too find it disheartening to hear of the failure of Prussian blue. Ethanol should not interfere.

-T

Polverone - 4-3-2004 at 22:46

Thank you for the discussion of PTC use. My practical organic books are from before the common use of PTCs, and theory-oriented texts usually omit important practical details.

I will see what sort of stirrer I can improvise. I would hope magnetic stirring would work, if not be ideal, since I already have a magnetic stirrer and just need some stirbars for it... In this environment, I might be able to get away with using a paperclip, since it shouldn't be very corrosive to iron.

Summary of results, in case I expressed it unclearly before: Prussian blue test was positive with the batch using PTC, negative with ethanol cosolvent, negative with "plain" mix using no cosolvent and no PTC.

Magpie - 28-12-2006 at 18:55

This is an old thread concerning the synthesis of malonic acid but vulture's question was left unresolved.

Quote:

Are there any other synthesis methods, maybe archaic ones?

I also wanted to know the answer so tracked down an archaic reference, ie, that of von Dessaignes in Ann., 107, p.251, 1858, "Ueber eine durch Oxydation der Aepfelsaure erhaltene Saure." A friend was kind enough to provide me a summary in english.

The article provided no quantities which made my synthesis attempts difficult. Here's how it went on my 2nd attempt:

1. Dessaignes: oxidize the aqueous malic acid at room temperature with acidic potassium dichromate. Me: mix 50mL DI water, 1mL con H2SO4, 1.34g Na2Cr2O7*2H2O, and 0.91g malic acid in a 250 mL beaker. Temperature max = 16C. Noted some small amount of bubbling which I assumed was CO2. Color was green with hint of orange. pH = 1-2. Brought temperature up to 33C using a water bath.

2. Dessaignes: add an excess of milk of lime, filter, and save the filtrate . Me: added NaOH to pH = 9 as I was afraid of Ca malonate precipitation*; filtered, centrifuged filtrate, and saved final centrifugate.

3. Dessaignes: Add lead acetate to precipitate the malonic acid. Me: Added a total of 6.7g Pb(OAc)2*3H2O (which seems like too much) to the 80 ml of centrifugate. Color was pea soup green at first but eventually supernate became almost clear with a purple tint when all Pb(OAc)2 added. pH =6.

4. Dessaignes: Add dilute HNO3 but only enough to dissolve the malonic acid leaving the rest of the ppt. Me: Added 6N HNO3 to pH = 2 (about 9 mL).

5. Dessaignes: Filter, saving the filtrate. Me: filtered discarding the green ppt and keeping the dark blue tinted filtrate.

6. Dessaignes: Add ammonia to "3/4 saturation." Pb malonate crystals will reform. Me: added ammonium hydroxide to pH=8. Never got any crystals even after reducing the filtrate volume by at least half.

*(milk of lime [Ca(OH)2] was used on the 1st attempt with same results, ie, ultimately no crystals of Pb malonate.)

I'm sorry if this long dissertation has become tedious but I thought that only through providing all detail could I hope for any constructive suggestions. Although I have spent a lot of energy on this already I'd be willing to do more experimentation if I could get some encouraging suggestions about what to do differently. Your discussion and suggestions are welcomed.

guy - 28-12-2006 at 21:39

Quote:

Originally posted by Polverone
I tried using this method to produce NaCN some time ago, and tried again just recently. I didn't use a PTC, just added enough 95% ethanol so that everything went into solution. I heated an excess of NaOH with chloroform and clear household ammonia solution (all added to the EtOH) for an hour or so. It formed a light-yellow solution with a foul smell. The liquid failed to give a positive Prussian blue test. Did the ethanol interfere?

I got that several times and it really made me mad to waste a lot of my chemicals.

I got it to work once. Ethanol seems to be the culprit.

You probably know this, but water causes to formation of formate, and ethanol screws it up. What other solvents are good for this, preferbly OTC?

Sauron - 28-12-2006 at 22:12

I have some malononitrile on hand but no malonic acid at present.

Malonic acid $43 250 g from Acros or $18/100 g

Malononitrile same source $70 500g so it is a bit cheaper. Maybe not enough to bother with for the effort of the -CN to -COOH transformation.

1,3-propanediol same source $48 250 ml

Acros is a Belgian outfit. Connected with Fisher in USA. Sell worldwide, I buy from them here.

There's a series of rxns in Org Syn that will get you where you want to go but I don't remember if they start with glycerol, or with acetone. If halogenation of glycerol not so bad but if s,3-dihaloacetone, very nasty stuff, vesicant, crosslinks DNA strands I bet just like sulfur or nitrogen mustards, sym-dichlorodiethyl ether or lewisite. Carcinogens, mutagens, evil in molecular form, A direct assault on the genome.

Sauron - 29-12-2006 at 02:19

I checked in Org,Syn. and unfortunately no prep of trimethylene glycol. The prepsgiven in references in their malonic acid prep include hydrochloric acid oxidation of malononitrile and a carboxylation of carbon suboxide - which does not seem very practicable unless you have a cylinder of that on hand.

If diethyl malonate is cheaper than the acid you might saponify that,

There is a prep of 1,3 dichloroacetone by dichromate oxidation of 1,3-glyceroldichlorohydrin. Abd a prep for that too. But then you will have to figure out how to get rid of that 2-OH group, then hydrolyze to the glycol then oxidize to the acid. Sounds like 5 steps from glycerin and one of then undocumented.

How much malonic acid do you need?

Sauron - 29-12-2006 at 03:04

There's an alternate route but no advantage as you are still starting from chloroacetic acid or sodium chloroacetate. First prepare sodium cyanoacetate from NaCN; liberate the free acid in hod with HCl. Esterify with 95% ethanol. Use the ethyl cyanoacetate to prepare cyanoacetamide. Condense the cyanoacetamide to malononitrile. The Org.Syn. procedures use either PCl5 or POCl3 for this. Treat malononitrile with conc HCl, and you have your malonic acid. The only economy I can suggest is that CC can replace the phosphorus chlorides in the amide to nitrile step. (Cyanuric chloride),)

I have zipped up the Org.Syn.pdf's and attached the zip file below. But I think you will be better off buying the malonic acid or maybe malononitrile rather than doing all this laborios prep work.

YT2095 - 29-12-2006 at 03:10

Quote:

Originally posted by Polverone

I will see what sort of stirrer I can improvise. I would hope magnetic stirring would work, if not be ideal, since I already have a magnetic stirrer and just need some stirbars for it... In this environment, I might be able to get away with using a paperclip, since it shouldn't be very corrosive to iron.

try the plastic coated paperclips and put a drop of super-glue (cyano acrylate) on each end.
it should work as a temporary fix

Sauron - 29-12-2006 at 03:21

Zip file attached

[Edited on 29-12-2006 by Sauron]

Attachment: malonic acid.zip (577kB)
This file has been downloaded 1512 times

garage chemist - 29-12-2006 at 04:56

Quote:

Originally posted by Sauron
There's an alternate route but no advantage as you are still starting from chloroacetic acid or sodium chloroacetate. First prepare sodium cyanoacetate from NaCN; liberate the free acid in hod with HCl. Esterify with 95% ethanol. Use the ethyl cyanoacetate to prepare cyanoacetamide. Condense the cyanoacetamide to malononitrile. The Org.Syn. procedures use either PCl5 or POCl3 for this. Treat malononitrile with conc HCl, and you have your malonic acid. The only economy I can suggest is that CC can replace the phosphorus chlorides in the amide to nitrile step. (Cyanuric chloride),)

Why so damn difficult? Why don't you just treat the cyanoacetic acid with HCl to get your malonic acid?

Sauron - 29-12-2006 at 05:35

Damed if I know. Don't kill the messenger, I am only reporting what I found in Org.Syn. The same procedure is in Vogel - as I am sure many have noticed he cribbed a lot from Org.Syn. and just changed to scale, usually downward.

Magpie - 29-12-2006 at 11:52

@Sauron: I would be happy with 100g of malonic acid or diethyl malonate. That Acros price is reasonable but as an individual I cannot buy from the major supply houses. The acid is available over the internet but the seller requires a DEA form. I don't have anything to hide but I don't want a knock on the door either. If I want this bad enough I can synthesize some NaCN and use the method in Vogel. The Dessaignes method is at first glance so attractive as it uses relatively safe reagents readily available.

I have a gut feeling that the Dessaignes method would work in the hands of a skilled old-school chemist. That's what makes this such an intriguing challenge.

Sauron - 29-12-2006 at 14:42

I suppose the concern is about clandestine barbiturate manufacture, but I was blissfully aware that there is any such these days. The only one around here addicted to barbiturates is my dog, he's an epileptic. And his vet prescribes those.

How about malononitrile? Same dilemna or no?

This really is a ridiculous situation. Do you have to put up a billboard and take out a newspaper ad that says THIS IS NOT A CLANDESTINE LAB?

Ever try talking to the concerned authorities (local, state, federal) and simply explaining that you are an amatuer experimental chemist and would like advice on how to proceed without sending out false alarms. Don't wait for a knock on the door. Invite them. If they see you aren't making drugs (or othings of concern like explosives or chemical weapons, then they will lose all interest. Worth a try? Assuming you an an upstanding citizen with roots in your community and no criminal record or radical affiliations. The DEA employs a lot of chemists, you know. There will be sympathetic souls around.

I dunno. It depends on where you live and what your situation is. I haven't lived in the US for two decades and am 10,000 miles away geographically. I can't judge.

garage chemist - 29-12-2006 at 15:09

I have a feeling that the Dessaignes method is not a laboratory method for the preparation of malonic acid.
If I had gotten this german article if I was in need of malonic acid I would never have attempted this synthesis. It just screams out "This will end up in a mess and wasted chemicals".
I know other preparations from which I can say by looking at them that I'll never get them to work, even if I have all the chemicals and apparatus.

Tell me, what is this synthesis for which you need diethyl malonate? Maybe we can find another way to make the target compound without use of malonic acid.

It indeed is a stupid situation to be deprived of chemicals because of nonsensical laws (who makes clandestine barbiturates anyway?).
In those situations you have to remind yourself that you have nothing to hide and therefore can fill out any forms necessary for the procurement of the substances.
When i bought acetic anhydride I also had to fill out forms about my lab and especially what I intended to do with the chemical. I simply stated the truth (cinnamic acid, cumarin, acetylcellulose), and that worked.

If malonic acid is still your target, I'd investigate the manufacture of NaCN from commercial ferrocyanides. There are ways to do this without dealing with gaseous HCN:

Simple heating of ferro- or ferricanides with charcoal powder under exclusion of air (loosely stoppered test tube) furnishes a mixture whose only water- soluble compound is KCN.
However, most of the cyano groups in the ferrocyanide are lost as N2 in the process.
The best method is still the one with gaseous HCN.

jimmyboy - 29-12-2006 at 15:09

if you can make malonic by oxidizing malic with dichromate then it should be possible to use permanganate as well - just mix it in a small quantity and keep at a very low temp (close to freezing) for a day or so - permanganate is much stronger - malonic decomposes to acetic acid with heat so it should be easy to figure out if you had any success

DeAdFX - 29-12-2006 at 16:13

In response to the one mentioning bZ oscillating reaction.

You know you do not need malonic acid for the BZ reaction. Other polycarboxylic acids will work too such as citric acid[this is pretty damn easy to come by]. However oxalic acid does not work at all. The color is not nearly as good as the malonic acid but as long as you can see the color change that is good enough.

pantone159 - 29-12-2006 at 16:31

I'm also interested in malonic acid, I want it for the BZ oscillating reaction. (Which is very sensitive to Cl-, btw. I think you need fairly pure malonic acid, the 'pure enough for most purposes' in the Org Syn prep might not be pure enough for the BZ, don't know for sure.)

It seems stupid that they worry about barbituates, I also thought that nobody made these clandestinely. The legal issues are even more complicated for me, legally (Texas) I'd have to apply for a permit just to have malonic acid, and I don't like the idea of introducing my lab to the police. As someone said, that might not actually be a bad idea, but it still makes me uncomfortable.

Sauron - 29-12-2006 at 17:46

I second jimmyboy's suggestion. The oxidation of malic acid (by whatever means) is producing a keto-dicarboxylic acid as intermediate; in this C4 instance the carbonyl group is alpha to one end and beta to the other. So voila, decarboxylation, and you have your C3 dicarboxylic acid (malonic) instead of C4 dicarboxylic (malic is a hydroxysuccinic acid).

Where the old prep gets murky is in lack of clarity and in the (unnecessary IMO) complexity of the workup. It looks like all those steps and reagents are merely to quench the hexavalent chromium and isolate the acid. Why not look for a way to oxidize the malic acid in a nonaqueous solvent instead?

I realize I am not suggesting any specific solvent/oxidizer system but that is because one is impoverished by the wealth of choices. After all the malonic acid synthesis (from diethyl malonate and sodium ethoxide, to higher molecular weight compounds) also depends on just such a decarboxylation of a beta-carboxtlic acid. so why not review the workups normally used there and look for a method that eschews all that calcium hydroxide and lead acetate malarky?

Or, rather than making any salt of the malonic acid at all, why not just extract the malonic acid into an organic solvent that is impervious to the chromic acid or permanganate or whatever your choice of oxidizer is?

Then rotavap off the solvent.

Wash the product till it is completely free of traces of the chromic acid)

Sauron - 29-12-2006 at 17:47

Of course if you need a DEA form to buy malic acid too then you are back in the box.

Magpie - 29-12-2006 at 18:51

No, in the US malic acid is readily and cheaply available at health food stores and from some wine making suppliers.

Sauron I believe you have stated the correct mechanism and this is confirmed in my text by Fieser & Fieser. But I don't see where KMnO4 is an advantage over Na2Cr2O7, especially if you say that the dichromate is the milder oxidant. (Dessaignes specified room temperature and I never exceeded 33C. I didn't see as much CO2 evolution as I expected.)

Making the diethyl malonate per Vogel requires chloroacetic acid. Making this requires red P. So I'm screwed on that route.

Malonic acid esters are shown to be very useful for synthetic chemistry due to acidity at C2. I have an experimental procedure exploiting this. That is my only reason for wanting the diethyl malonate which I intended to make from malonic acid.

In reference to your understandable feelings about the authorities attitude toward home labs, well, that is just a fact of life in the US today. Twenty, or even ten years ago that would not have been the case. Early in my time on this forum I proposed the very thing you mentioned, ie, go directly to the local police and state your case. One forum member promptly said that that was a completely idiotic idea. I decided that he might very well have a good point. I have not heard of anyone in the US that has actually done this. Even if someone did get a favorable reception I would think that this is something entirely local and immediately subject to change. For that reason I have chosen to keep a low profile. I feel that even though home chemistry is not illegal in the US they would find a way to close me down if they chose.

Sauron - 29-12-2006 at 19:53

All right then. Form a small business, incorporate, get an EIDN with the IRS and a business license from your city and state as a research laboratory. Depending on your zoning situation you might not be able to do this in your home. But then again you might. I can't overemphasize that this is not a formula for setting up a drug lab as the DEA and the local police etc do check up on commercial labs; they can audit your lab and if you are making a scheduled substance or have made one previously chances are they can find enough trace of it to be a problem. Don't buy drug precursors (especially direct ones like phenylacetone) although sometimes depending on what LEGITIMATE work you are doing this is unavoidable. Explosives are a sticky issue. ATF will have to license you and approve storage and local and state authorities may well have their own requirements, Doing this in a residence is very likely to be incompatible under the law.

In short to be a legitimate lab you must behave as a legitimate lab.

If you behave like a clandestine lab then even if you aren't making drugs, or explosives, or worse, you will certainly be seen as a clandestine lab and you will always have problems with suppliers of equipment and chemicals. The govt has put the burden legally on lab supply houses to know who they are dealing with and to report any and all suspicious activity. In short every supplier is now an informer. If the feds find that a supplier sold to a drug lab and did not rat them out, the seller goes to jail. In some locations it gets worse as there are vigilante groups like CAMP going after drug operations.

So keep your nose clean.

If you can't sly under the radar, make sure you have a transponder with IFF (that's Identify Friend or Foe.)

Sauron - 29-12-2006 at 20:08

BTW red P is not mandatory for chlorination of acetic acid, it merely accelerates the rate of the chlorination. Direct chlorination is UV mediated so you need reliable sunlight or preferably a UV reactor. OR chlorinate with CuCl2 (see thread) or N-chlorosuccinimide, both of which monochlorinate the alpha position of carboxylic acids I think. NBS does for sure and bromoacetic acid would work just as well,

Furthermore chloroacetic acid is cheao and as far as I know unrestricted. (Am I wrong?)

Where I am neither malonic acid not diethyl malonate are restricted. I haven't bought any but have obtained quotes on the latter. I did buy malononitrile without the slightest hassle. So you might want to look into malononitrile and then you can treat it with conc hydrochloric acid to obtain malonic acid.

I thought you could

Ephoton - 30-12-2006 at 04:56

could also use sulfur for this chlorination as it generated
chlorides of sulfur which then chlorinated the acid.

polymer - 30-12-2006 at 07:33

You could use sodium carbonate and Nitrogen. Sodium carbonate can come from even a chemistry set and nitrogen can come from fertilizer. To form the compound, they obviously use covalent bonds, so you could probably use a method similar to electrolysis.

JohnWW - 30-12-2006 at 15:41

Quote:

Originally posted by Sauron
(cut) Where I am neither malonic acid not diethyl malonate are restricted. I haven't bought any but have obtained quotes on the latter. I did buy malononitrile without the slightest hassle. So you might want to look into malononitrile and then you can treat it with conc hydrochloric acid to obtain malonic acid.

If you treated malononitrile (the nitrile of malonic = propanedicarboxylic acid, NC-CH2-CN ) with concentrated HCl to hydrolyse it to malonic acid, HOOC-CH2-COOH, the byproduct would be mostly NH4Cl which is left in solution; but there is the possibility of side-reactions involving some decomposition to form HCN and ClCN, which are volatile and highly poisonous.

jimmyboy - 30-12-2006 at 16:09

i just suggested permanganate because it is so much easier to acquire and should work ..

our freedoms to experiment in science in the US are all but gone in chemistry -- either you hide what you have and risk getting harrassed or arrested for your desire of knowledge or you annually pay the government a few grand that you dont have and face the same bs.. no wonder our backwards country is so far behind everyone else - we punish innovation with our laws - but i am getting off subject

has anyone actually tried to the dichromate route and seen the results?

Sauron - 30-12-2006 at 18:42

Malonic acid by rxn of malononitrile with HCl (Conc) is an Org.Syn. prep. Not something I dreamed up. Classic procedure. Checked and verified, annotated and with safety notes and references.

I think anyone competent would recognize that adding a nitrile to mineral acid might release HCN and therefore, this is a procedure to be done in a GOOD HOOD. ClCN less likely but same caution. GOOD HOOD. If you were to attempt this on a larger scale I would say GOOD HOOD, SCBA, and CAUSTIC SCRUBBER, the oversize scrubber interposed between the reaction flask and the hood exhaust.

jimmyboy - 31-12-2006 at 00:12

hmm yet another interesting route -- lets see - cyanoacetic acid + ammonia -> cyanoacetamide + phosphorus oxychloride -> malononitrile + HCL -> malonic acid + HCN

hmm i think thats right anyway

cyanoacetates can be found in superglues - solvents

Sauron - 31-12-2006 at 02:40

Superglue is an ester of cyanoacrylic acid not cyanoacetic acid. Methyl cyanoacrylate it I recall.

garage chemist - 31-12-2006 at 03:45

Adding HCl to nitriles does not release any HCN at all.
I have done the hydrolysis of benzyl cyanide with H2SO4 and trust me, there was not the slightest bit of HCN smell. I even did it outside my fume hood.
What happens is that the C-N triple bond gets cleaved completely (hydrolysis) and the nitrogen ends up as ammonium ion. Neither the C nor the N atom does change oxidation state here.
Liberation of HCN would require C-C bond fission which will simply not happen because there is no oxidiser present.

joeflsts - 31-12-2006 at 07:55

Quote:

Originally posted by Sauron
I second jimmyboy's suggestion. The oxidation of malic acid (by whatever means) is producing a keto-dicarboxylic acid as intermediate; in this C4 instance the carbonyl group is alpha to one end and beta to the other. So voila, decarboxylation, and you have your C3 dicarboxylic acid (malonic) instead of C4 dicarboxylic (malic is a hydroxysuccinic acid).

Where the old prep gets murky is in lack of clarity and in the (unnecessary IMO) complexity of the workup. It looks like all those steps and reagents are merely to quench the hexavalent chromium and isolate the acid. Why not look for a way to oxidize the malic acid in a nonaqueous solvent instead?

I realize I am not suggesting any specific solvent/oxidizer system but that is because one is impoverished by the wealth of choices. After all the malonic acid synthesis (from diethyl malonate and sodium ethoxide, to higher molecular weight compounds) also depends on just such a decarboxylation of a beta-carboxtlic acid. so why not review the workups normally used there and look for a method that eschews all that calcium hydroxide and lead acetate malarky?

Or, rather than making any salt of the malonic acid at all, why not just extract the malonic acid into an organic solvent that is impervious to the chromic acid or permanganate or whatever your choice of oxidizer is?

Then rotavap off the solvent.

Wash the product till it is completely free of traces of the chromic acid)

Check this out:
http://cgi.ebay.com/The-Fascinating-Oscillating-Kit-Kits_W0Q...

pantone159 - 31-12-2006 at 13:01

Quote:

Originally posted by joeflsts
Check this out:
http://cgi.ebay.com/The-Fascinating-Oscillating-Kit-Kits_W0Q...

Tempting. I'm trying to decide how to deal with the fact that shipping would nearly double the cost ($10 to $18). Even 5g of malonic acid is a lot better than wat I got now.

BTW - I am thinking that permanganate oxidation of malic acid will yield only oxalic acid and not malonic acid.

My reasoning... (Forgive the fact that I don't know organic chem very well so I may be totally off.)

Malic acid structure: HOOC-CH2-CHOH-COOH
I think the alcohol group will first oxidize to ketone. (Not sure though.) My books describe the mechanism of KMnO4 oxidation as going through the enol form of the ketone, i.e.

Malic acid oxidized to ketone: HOOC-CH2-C(=O)-COOH
This might form the enols as:
A: HOOC-CH=C(-O)-COOH, or
B: HOOC-CH2-C(-O)=COOH
But B can't happen because the last (rightmost) carbon already has all four hands full with the acid group, and can't spare one to make the enol, so only A can happen.

KMnO4 then grabs onto the double-bonded enol carbons, and further mechanism breaks the chain there, turning both of the enol carbons into carboxylic acid groups, thus yielding:
HOOC-COOH and HOOC-COOH, i.e. two molecules of oxalic acid.

However, as soon as the new year passes, I hope to stop by my local brewing supplier, which I hope will have malic acid, and I will try oxidizing it, to see what happens.

Questions:
1- How can one distinguish oxalic from malonic acid? I think malonic will decarboxyate under heating, forming acetic acid (someone mentioned this already), while oxalic won't. That is one test, but I'd like more.

2 - In Magpie's prep, if malonic acid was formed in the first oxidation step, why wouldn't it precipitate with 'an excess of milk of lime' as calcium malonate?

[Edited on 31-12-2006 by pantone159]

Sauron - 31-12-2006 at 13:29

Better have a look at the behavior of beta-kto acids.

Furthermore, it is quite simple to distinguish between oxalic and malong acid.

Melting points. ALl you need is a Thiele tube, some melting-point capillaries, and mineral oil for medium. And a heat source. And patience. First you need to crystallize the acid and keep recrystallizing till you obtain a constant m.p.. Also you can obtain an authentic sample of each acid. If you mix small equal amounts of the product and the acid you want, if they are same acid the m.p. remains same, if they are different the m.p. will be lowered.

Derivatives: esters, amides, chlorides etc. Diethyl oxalate and diethyl malonate have different b.p.'s. The amides are usually solids. Oxalic acid forms oxalyl chloride with cyanuric chloride, this reaction fails for malonic acid.

This is first year organic practical lab stuff. See Vogel's book.

[Edited on 31-12-2006 by Sauron]

JohnWW - 31-12-2006 at 14:29

Quote:

Originally posted by garage chemist
Adding HCl to nitriles does not release any HCN at all.
I have done the hydrolysis of benzyl cyanide with H2SO4 and trust me, there was not the slightest bit of HCN smell. I even did it outside my fume hood. What happens is that the C-N triple bond gets cleaved completely (hydrolysis) and the nitrogen ends up as ammonium ion. Neither the C nor the N atom does change oxidation state here. Liberation of HCN would require C-C bond fission which will simply not happen because there is no oxidiser present.

I would not be so sure, in view of a particular case in point. The compound with one less C than malonic acid, namely oxalic acid, (COOH)2, is fairly easily decomposable by heating to CO2 and H20, by C-C bond fission (which is why rhubarb, which contains it, has to be cooked before eating), because of the repulsion due to the two vicinal highly polar groups. The same would apply to a slightly lesser extent to the corresponding nitrile, which in fact is cyanogen, (CN)2, the thermal decomposition of which in the presence of HCl would certainly result in formation of some HCN (prussic acid) and ClCN (cyanogen chloride). So, although the repulsion of the polar groups would be significantly less in malononitrile, it is likely to be still sufficient for decomposition involving C-C bond rupture if the hydrolysis by HCl was done at a sufficiently high temperature.

pantone159 - 31-12-2006 at 14:44

Quote:

Originally posted by Sauron
Melting points. ALl you need is a Thiele tube, some melting-point capillaries, and mineral oil for medium. And a heat source. And patience. First you need to crystallize the acid and keep recrystallizing till you obtain a constant m.p.. Also you can obtain an authentic sample of each acid. If you mix small equal amounts of the product and the acid you want, if they are same acid the m.p. remains same, if they are different the m.p. will be lowered.
[Edited on 31-12-2006 by Sauron]

All I currently lack is the Thiele tube, the capillaries, and the authentic malonic acid sample. (And possibly the patience

)

However, I did decide to get one of those kits, and the seller also had a Thiele tube, so I added that, so I soon should have 2 of 3.

Another test, btw, for malonic acid vs oxalic acid ought to be if the oscillating reaction works or not.

P.S. Based on the ingredients list, the kit is probably the Briggs-Rauscher reaction (see Shakhashiri, Chemical Demonstrations, Vol 2, 7.1) with sulfamic acid substituted for sulfuric acid.

[Edited on 31-12-2006 by pantone159]

Magpie - 31-12-2006 at 15:38

Malonate and oxalate would both be precipitated with Ca++ and that is why I tried NaOH vs Dessaignes' milk of lime [Ca(OH)2]. Oxalate is especially insoluble. Dessaignes does say that oxalates are byproducts. This would be the advantage of using Ca++, ie, it would take them out. If I try this again I would use milk of lime but have enough volume to keep the slightly soluble malonate in solution.

Fieser & Fieser in Introduction to Organic Chemistry, 1957, p. 251 do say that the oxidation proceeds through the ketone and then elimination of the C=0 as CO2 to give malonic acid.

As a side note sugar beets (or beets in general) are supposed to be rich in Ca malonate, with the total organic acid content being around 2.6%. It sounds like another interesting (and time consuming

) challenge to try to isolate malonic acid from that source.

Edit:

I wanted mention that a really slick way to help identify an organic acid is through titration. By titrating you determine the molecular weight.

[Edited on 1-1-2007 by Magpie]

Sauron - 31-12-2006 at 19:17

Malonic acid is soluble in Et2O

Oxalic acid (anhydrous) is only slightly soluble in Et2O

Malonic acid is soluble in water @20 C 1400 g/L

Oxalic acid (anhydrous) is soluble in water only 90 g/ @20 C.

Malonic acid melts at 130-135 C

Oxalic acid (anhydrous) sublimes at 101-157 C and melts at 185-198 C

There are other differences in physical and chemical characteristics.

(Acros tech data sheets q.v.)

See also Aldrich catalog, Merck Index, Vogel's chapter on dicarboxylic acids. Many other reference works. CRC Handbook.

Obviously these two are not at all difficult to distinguish from each other, without the need to resort to exotic solvents or derivatization.

I bet the crystalline structures are also distinctly different. I know what oxalic acid crystals look like. Cubic and translucent. Haven't seen any for about 40 y\s but it's a clear memory. Don't recall what malonic acid crystals look like but it will be in the Merck Index monograph.

On points like this it is best to forget Kantian pure reason and go for the facts.

Sauron - 31-12-2006 at 19:31

Merck Index 12th Ed:

Malonic acid: small crystals mp c.135 C with decomposition; sublimes in vacuo. 1 g dissolves in 0.6 ml water; in 2 ml alcohol; in 1.1 ml methanol; in 3 ml propanol; in 13 ml ether; in 7 ml pyridine.

Oxalic acid: crystallized from glacial acetic acid forms orthorhombic crystaks, the shape being pyramidal or elongated octahedra. Hygroscopic, mp 189.5 C (dec); sublimes best at 157 C.

pantone159 - 1-1-2007 at 11:21

Quote:

Originally posted by Sauron
Malonic acid is soluble in water @20 C 1400 g/L

Oxalic acid (anhydrous) is soluble in water only 90 g/ @20 C.

Are you sure about these numbers? (I.e. a ratio of 15.6, not 1.56.) If so, then oxalic and malonic ought to be easily separable by dissolving in a limited amount of water, the malonic would dissolve but the oxalic not.

Sauron - 1-1-2007 at 15:36

You can easily check the same sources I did.

Acros tech data sheets - shall I attach them?

Merck Index 12th Edition, monographs for oxalic acid and malonic acid.

Malonic acid 1400 g/L H2O 20 C
Oxalic acid (anhydrous) 90 g/ L H2O 20 C.

That's c. 15X more solubility in the commonest solvent we have at room temperature.

So, identification and seperation are EASY. It doesn't get any easier.

Merck merely restates the identical difference in solubility in a slightly different way.

1 g malonic acid dissolves in 0.6 ml water. In other words, water dossolves 1.4X its own weight of malonic acid. How much does a liter of water weigh? 1000 g by definition. So, how much malonic acid will dissolve in that liter at 20 C? 1400 g.

I believe Merck Index is on the Reference forum in digital form. I am using the hardback I bought in mid 90s. For a long time that and an elderly Aldrich were the only chemical references I had here.

pantone159 - 1-1-2007 at 15:54

Quote:

Originally posted by Sauron
Acros tech data sheets - shall I attach them?

Sure, if it is convenient. Can I find these online? I wasn't disbelieving you, btw, those numbers just sounded too good to be true. My CRC merely describes both as soluble in water. Separation from unreacted malic acid may be an issue as well.

Quote:

I believe Merck Index is on the Reference forum in digital form.

I didn't know that, I'll look for it. Otherwise I'll head to the library to read the dead-tree version. Thanks.

Sauron - 1-1-2007 at 16:25

http://www.acros.be

is website of Acros Organics. In the Catalong Search field just enter "amlonic acid" and the page will come up. In upper right there's an adove acrobat PDF icon, click on that and you get the page in PDF rather than HTML, right click on that and save it to your hard disk. Repeat for exalic acid.

I did not look up malic acid.

As mentioned I think you can download the entire Merck Index from rapidshare, there are links on the Reference forum. I think so. I did not look at them because I have the book already on my shelf.

Sauron - 1-1-2007 at 16:26

NOTE that the hydrated form of oxalic acid may have different solubility than the anhydrous acid.

pantone159 - 1-1-2007 at 16:55

Quote:

Originally posted by Sauron
NOTE that the hydrated form of oxalic acid may have different solubility than the anhydrous acid.

The dihydrate has solubility 138 g/L at 20 C, which is still plenty different to separate them.

Sauron - 1-1-2007 at 17:19

Yes, lower than malonic acid by an order of magnitude.

This happy situation does not always occur, you know? But in this case, it's very very simple and easy.

jimmyboy - 1-1-2007 at 19:22

heck - get the oscillating kit - 10 bucks for 5 grams of malonic.. easy.. haha of course that would be cheating and you wouldnt learn anything

Decarboxylation of oxaloacetic acid

pantone159 - 15-2-2007 at 17:59

Note: The decarboxylation of the beta-keto acid formed by oxidizing does NOT yield malonic acid, but rather pyruvic acid.

Malic acid, HOOC-CH2-CHOH-COOH, is first oxidized to the keto-acid oxaloacetic acid, HOOC-CH2-C(=O)-COOH.

This will decarboxylate easily. BUT, it doesn't lose the carbon we want. The leftmost carbon is the one that is at the right position for decarboxylation, not the rightmost one. So, decarboxylation instead yields pyruvic acid, CH3-C(=O)-COOH.
(FWIW, Wikipedia states the same thing.)

If we want to get from oxaloacetic acid to malonic acid, we need to get that other end carbon to leave. I think this really may require KMnO4, which can chop up the carbon chain, rather than other oxidants which will go no further than the keto-acid, and all we can get from there is pyruvic acid.

PS - I have made some tries oxidizing malic acid with KMnO4. The major product seems to be a colorless liquid, with an odor I would describe as 'soapy'. Possibly this could be pyruvic acid, but I don't know. I do have some solid crystals in the product, however. The liquid is evaporating slowly (while standing at RT) and I might even get some reasonable crystals I can try to analyze. I intend to post details, but I haven't got my write-up into shape yet. My workup needs improvement as well.

PPS - I finally got to the dead-tree library to look up the Fieser reference that Magpie gave. It does say that the process goes through the keto-acid, but doesn't discuss how one gets from there to malonic acid.

[Edited on 16-2-2007 by pantone159]

1,3-propanediol, aka "Bio-PDO"

pantone159 - 18-2-2007 at 11:13

It occurs to me that 1,3-propanediol, HO-CH2-CH2-CH2-OH, aka "Trimethylene glycol" ought to be a very suitable precursor for malonic acid. Oxidize the two primary alcohol groups to acids by any of the well-developed methods, and one should have malonic acid.

So, where to get 1,3-propanediol? It currently does not seem to be common (a search of my usual sources did not list it for sale anywhere. Aldrich does list it, $60 for 500 g, but I can't get it there.) "Propylene glycol", 1,2-propanediol OTOH is widely available, cheap, and not suitable for malonic acid synthesis. (The names of these compounds seem to be used somewhat inconsistently, beware.)

This may soon change. DuPont apparently has just built a plant to produce 1,3-propanediol from corn, with the brand name "Bio-PDO". The first commercial shipments was just under 3 months ago. The following press release is full of platitudes about 'non-petroleum, bio, renewable' but I personally just care about moving that pesky OH group to the end.
http://news.thomasnet.com/companystory/500722

The material will be branded "Zemea" for personal-care uses, e.g. cosmetics, which are likely to be a good OTC source, once the stuff becomes available.

Also see the Wikipedia page:
http://en.wikipedia.org/wiki/1,3-Propanediol

EDIT: I just re-read this thread from the start, and noticed the previous talk about the 1,3-diol, which I had forgot about.

[Edited on 19-2-2007 by pantone159]

unionised - 18-2-2007 at 11:19

Potentialy interesting stuff, though I had to smile at this bit of their site.

"Tate & Lyle uses innovative technology to transform corn, wheat and sugar into value-added ingredients ....

In the last decade, Tate & Lyle established fermentation as a core technological competency and is now one of the world's major fermentation producers with 17 fermentation plants on four continents. In addition to Bio- PDO(TM), its portfolio includes citric acid, biogums ... and both fuel grade and potable ethanol."
"Inovative technology" seems to include making rum.

[Edited on 18-2-2007 by unionised]

not_important - 18-2-2007 at 11:59

As it is a new year, I will suggest a 'new' method.

The polyunsaturated fatty acids linoleic and linolenic have respectively one and two =CH-CH2-CH= groupings in them. Treatment with ozone, followed by workup undex mild oxidising conditions, or just treatment with oxidisers under controlled conditions, will yield mixtures of mono- and di- carboxylic acids.

These acids are propionic, caproic, malonic, and azelaic - HO3C(CH2)7CO2H
Oleic acid yields perargonic - CH3(CH2)7CO2H - and azelaic.

That is somewhat of a mess, but still not too difficult to separate out. Raw (unboiled) linseed oil is roughly 5% oleic, 60% linoleic, and 25% linolenic. Note that is the unboiled oil, not the more common boiled; at a higher price the edible form can be purchased as flax or flaxseed oil.

The weight percentage yield is terrible, but in some location linseed oil can be had quite cheaply. Using ozone to accomplish the first stage of oxidation reduces the cost in oxidisers.

Sandmeyer - 18-2-2007 at 14:38

Quote:

Originally posted by Sauron

I think anyone competent would recognize that adding a nitrile to mineral acid might release HCN and therefore, this is a procedure to be done in a GOOD HOOD. ClCN less likely but same caution. GOOD HOOD. If you were to attempt this on a larger scale I would say GOOD HOOD, SCBA, and CAUSTIC SCRUBBER, the oversize scrubber interposed between the reaction flask and the hood exhaust.

Nonsense. Hydrolysis of the nitrile could happen, then if heating is continued the CO2 and not HCN is evolved as a result of decarboxylation.

Quote:

Originally posted by not_important
As it is a new year, I will suggest a 'new' method.

The polyunsaturated fatty acids linoleic and linolenic have respectively one and two =CH-CH2-CH= groupings in them. Treatment with ozone, followed by workup undex mild oxidising conditions, or just treatment with oxidisers under controlled conditions, will yield mixtures of mono- and di- carboxylic acids.

That sounds like an impractical idea.

[Edited on 18-2-2007 by Sandmeyer]

[Edited on 18-2-2007 by Sandmeyer]

tr41414 - 29-1-2008 at 12:54

The oxidation of double bonds could be done with acidic CuO, which would make the process quite cheap...

Another thing that comes to my mind for synthesising malonic acid is oxidation of citric acid.
Using hot H2O2(?) http://jxb.oxfordjournals.org/cgi/reprint/4/2/129.pdf Could anyone please fetch the ref?

Also would it be possible to react trichloroethylene with cyanide and then hydrolyse it (or maybe using HgSO4 for hydrolysis to make it more practical)?

tr41414 - 1-2-2008 at 02:22

The article uses 100% H2O2, so that is not really useful, and it also seems that malonic is readily decomposed in rxn

I might give the CuO oxidation of oil a try

I guess the resulting mess could be "purified" by forming esters and distilling them...

Aqua-regia - 1-2-2008 at 03:46

The synthesis of malonic acid:

http://www.versuchschemie.de/htopic,10447,.html

pantone159 - 1-2-2008 at 07:35

Quote:

Originally posted by tr41414
The article uses 100% H2O2, so that is not really useful, and it also seems that malonic is readily decomposed in rxn

Actually it said '100 vol.' which I think is different from 100% (H2O2 has some special concentration units which are often used and I am not really up on.) It did say:

Quote:

very low final yield of malonic acid

and didn't give much detail for the citric-to-malonic oxidation, so I also concluded that this procedure was not useful, too bad.

garage chemist - 1-2-2008 at 09:09

May I suggest buying some ethyl cyanoacetate and hydrolysing that with NaOH (free malonic acid decarboxylates upon boiling of the aqueous solution, so alkaline hydrolysis has to be used)?
Ethyl cyanoacetate is seriously cheap, it would actually cost less to make malonic acid this way than to buy it.
And ethyl cyanoacetate isnt a regulated chemical, I think.

tr41414 - 2-2-2008 at 11:04

I wrongly assumed that 100 vol. = 100 vol % (it actually corresponds to volume of O2 released by volume of peroxide), so it is around the usually available 30%.

Also the article mentions decomposition of malonic to acetic acid, which might be affected by adding some base...

benzylchloride1 - 3-2-2008 at 10:46

Has any one tried the haloform reaction on 2,4-pentadione? This reaction should produce sodium malonate and chloroform. What would be the best method to isolate the malonic acid out of the reaction mixture?

Nicodem - 3-2-2008 at 13:10

A haloform reaction on 2,4-pentadione (acetylacetone) would result in 3,3-dihalogenation which upon cleavage with hydroxide gives acetic acid (and chloroform). See the mechanism of the haloform reaction:
http://www.organic-chemistry.org/namedreactions/haloform-rea...
http://en.wikipedia.org/wiki/Haloform_reaction
http://www.chem.ucalgary.ca/courses/351/Carey5th/Ch18/ch18-3...

As you can see the haloform reaction is based on base catalyzed enolization and subsequent halogenation. Therefore, there is no way for the terminal methyls of acetylacetone to get involved in the halogenation steps.

PHILOU Zrealone - 5-2-2008 at 05:51

Hydrolyse of speudocarboxylic groups....in basic or acidic media.
Thus nitriles, amides, esters, anhydrides, acid chlorides, alfa trichloromethylcetones, trichloromethyl hydrocarbons can be precursors
-CN --> -CO-NH2 --> -CO2NH4 --> -CO2H + NH4(+)
-CO2-R --> -CO2H + R-OH
R-CO-O-CO-R --> 2R-CO2H
-CO-CCl3 --> -CO2H + HCCl3
-CCl3 --> -CO-Cl +2HCl –> -CO2H +3 HCl
(trichloromethyl hydrocarbons are related to orthoformiates and can be seen as a transient trihydroxymethyl group...-CCl3 --> -C(OH)3 --> -CO2H + H2O)

As a side note the following known reaction on refluxing:
Cl2C=CHCl + H2SO4 (H2O) --> ClCH2-CO2H + 2 HCl
This goes maybe via H2O addition and then HCl elimination and ceto-enol equilibration…and finally hydrolysis of the resulting acid chloride…
Cl2C=CHCl + H2O --> Cl2CH-CHCl(OH) --> ClHC=CCl(OH) + HCl <----> ClCH2-CO-Cl --> ClCH2-CO2H + HCl

As a conclusion: the following compounds could be used
CCl3-CH2-CCl3
CCl2=CH-CCl3
CCl3-CH2-CN
CCl3-CO-CH2-CO-CCl3
CCl3-CO-CH2-CCl3
…

[Edited on 5-2-2008 by PHILOU Zrealone]

[Edited on 5-2-2008 by PHILOU Zrealone]

[Edited on 5-2-2008 by PHILOU Zrealone]

Jamjar - 5-2-2008 at 14:09

A haloform reaction on ethyl acetoacetate should produce disodium malonate and chloroform?

http://en.wikipedia.org/wiki/Ethyl_acetoacetate
"Ethyl acetoacetate may be prepared via the Claisen condensation of ethyl acetate. Two moles of ethyl acetate condense to form one mole each of ethyl acetoacetate and ethanol."

Regarding Acetoacetate instability
http://en.wikipedia.org/wiki/Acetoacetate
"The acid form has a half-life of 140 minutes at 37º C in water, whereas the basic form (the anion) has a half-life of 130 hours."

I'm a beginner so I could be talking out of my ass.

Nicodem - 5-2-2008 at 14:38

Two posts above I explained why the haloform reaction on acetylacetone yields acetate and chloroform.
For the same reason ethyl acetoacetate can't yield the malonate, but acetate, carbonate (AcOH and CO2 after acidification), ethanol and CHCl3.
Once again, the haloform reaction is based on base catalysed enolisation, meaning that the most acidic hydrogens will be substituted with the halogen.

PHILOU Zrealone - 6-2-2008 at 05:08

Nicodem and Jamjar,
CH3-CO-CH2-CO-CH3 and CH3-CO-CH2-CO-O-CH2-CH3 will react, as mentionned Nicodem, in basic media and halogen:
1°) First on the two H of the CH2 between the two C=O
2°) Once they were used, under further treatment with exces halogen it will jump on the H of external CH3.

This would mean that in the end products, one will get:
For pentandione: HCCl3, and HO2C-CCl2-CO2H (dichloromalonic acid) (maybe giving HO2C-CHCl2 + H2CO3 or doubtfully CH2Cl2 + 2H2CO3)
For ethylacetylacetate: HCCl3, H2CO3, CH3-CH2-OH and CHCl2-CO2H (or maybe further CH2Cl2+ H2CO3?).

I'm not sure CCl2 will be able to split off like CCl3 partly because dichloacetic (dichloroethanoic) acid is stable acid and I never have heard or read about such splitting.

I'm not sure about the stability of 2.2-dichloromalonic acid but it is mentionned in a few (7 or so) internet references aside with dichlormethane

[Edited on 6-2-2008 by PHILOU Zrealone]

Nicodem - 13-2-2008 at 12:53

Unfortunately it is not that easy. Me-CO-CCl2-CO-Me in the presence of hydroxide is just as easily cleaved to acetate and Me-CO-CHCl2 as CCl3-CO-R type of ketones are in the haloform reaction of methyl alkyl ketones. So, you can not perchlorinate acetylacetone and similar beta-dicarbonyl compounds under such conditions - they cleave faster.

PHILOU Zrealone - 14-2-2008 at 09:11

Quote:

Originally posted by Nicodem
Unfortunately it is not that easy. Me-CO-CCl2-CO-Me in the presence of hydroxide is just as easily cleaved to acetate and Me-CO-CHCl2 as CCl3-CO-R type of ketones are in the haloform reaction of methyl alkyl ketones. So, you can not perchlorinate acetylacetone and similar beta-dicarbonyl compounds under such conditions - they cleave faster.

So final products of perchlorination are only CHCl3 and H2CO3?

theobromacacao - 14-2-2008 at 10:15

Whilst it may not be strictly related to this thread, does anyone know how I can synthesise 1,10 phenathroline monohydrate (Ferroin) for use in a BZ reaction? Alternatively, does anyone know where an individual (who is based in the UK) would be able to obtain any? Thank you to all in advance!

Pixicious - 28-2-2008 at 07:33

I thought malonic acid was derived from apple acid?

http://216.239.59.104/search?q=cache:k82cozVDy8wJ:www.ital.sp.gov.br/bj_old/brazilianjournal/ed_especial/10.pdf+apples+%22malonic+acid%22&hl=en&am p;ct=clnk&cd=18&gl=uk

"...Dr Nitschke said the basic ingredients of the barbiturate were urea and malonic acid, found in apples. (me for a peaceful water-bed and some zzzzs) ..."

Something I found.

[Edited on 28-2-2008 by Pixicious]

[Edited on 28-2-2008 by Pixicious]

PHILOU Zrealone - 28-2-2008 at 08:25

Quote:

Originally posted by Pixicious
I thought malonic acid was derived from apple acid?

"...Dr Nitschke said the basic ingredients of the barbiturate were urea and malonic acid, found in apples. (me for a peaceful water-bed and some zzzzs) ..."

Something I found.

[Edited on 28-2-2008 by Pixicious]

[Edited on 28-2-2008 by Pixicious]

A bit confusing with those old grocery names:
-Malonic acid is propandioïc acid
-Maleic acid is cis-butenedioïc acid
-Malic acid is 2-hydroxy butandioïc acid
-Fumaric is trans-butendioïc acid

Malic refers to "Malus" in latin what means..."Apple" and for a not too obscure reason, without entering too much into religion aspects

... "Evil".
Maybe on a metaphoric/poetic view because Eve took the Apple from the knowledge tree to give it to Adam...it was the original sin, cause of all troubles on Earth

--> Evil.
I suppose Evilic acid or Metaphoric acid would have had less succes than Malic and doomed by the Church

Pixicious - 1-3-2008 at 15:03

lol.

I was just wondering.

I know I am a bit of a newbie but I was wondering, is it true beetroot contains a calicum salt of malonic acid? If so why isn't this synthesis.

Pix

[Edited on 2-3-2008 by Pixicious]

Formatik - 21-4-2008 at 22:13

Malonic acid was discovered by the oxidation of malic acid in 1858. Scheele first isolated and described malic acid from apple juice in 1785 (it’s not a good yield though

, apple juice contains 7g or more of malic acid per liter). Nevertheless, if you want to be adventurous and see some ancient methods then it’s mentioned in Beilstein Vol. 1, p. 648 -649, and from supplement 1, p. 280: in addition to some others, well known or usual, malonic acid preparation:

By the oxidation of malic acid (Dessaignes, A. 107, 251) or raw L-(+)-lactic acid (Dossios, Z. 1866, 449) with K2Cr2O7.

Through the oxidation of allylene or propylene with a cold KMnO4 solution (Berthelot, A. Spl. 5, 97).

From hexabromomethylethylketone and fuming HNO3 (Demole, B. 11, 1714). CBr3.CO.CH2.CBr3 + O + 4 H2O = C3H4O4 + CO2 + 6 HBr.

By boiling barbituric acid with potash (Kali) (Bayer, A. 130, 143).

Boiling cyanoacetic acid with alkalis or HCl acid (Kolbe, A. 131, 349; H. Mueller, A. 131, 352).

From b-dichloroacrylic acid ester and silver oxide at 125 deg. (Wallach, Hunaeus, A. 193, 25). CCl2:CH.CO2.C2H5 + Ag2O + H2O = CO2H.CH2.CO2.C2H5 + 2 AgCl.

By boiling mucobromic acid C4H2Br2O3 with Ba(OH)2 (no reference given).

Preparation process from chloroacetic acid ethyl ester or chloroacetic acid with KCN is briefly described also (p. 648-9).

By the oxidation of d-quercitol (acorn sugar) with KMnO4 (Killiani, Schaefer, B. 29, 1763).

Through the addition of alkalis on acetonetricarboxylic ester (Willstaetter, B. 32, 1284).

If I can find the other triple in Beilstein, maybe there is more about its preparation.

A. = Liebig's Annalen der Chemie; A. Spl. (supplement).
B. = Berichte der Deutschen Chemischen Gesellschaft.
Z. = Zeitschrift für Chemie.

There is also supposed to be a significant amount of calcium malonate in red beet juice. I have no idea on how to extract it. The salts of malonic acid are hardly soluble in water, except the alkali salts. Ca.C3H2O4 + 1 ¾ H2O is almost insoluble in cold water. 100 parts water solubilize at t ° = 0.2897 + 0,0.sub.2.423 544 1.t – 0,0.sub.4.233 510 3.t2 parts of the anhydrous salt. There are various hydrates of the calcium salt.

[Edited on 21-4-2008 by Schockwave]

Magpie - 22-4-2008 at 10:13

Malonate salts are indeed present in the juice of beets. Precipitation of Ca malonate is a problem for industrial refiners of sugar beets. It is present at something around 2% in sugar beets IIRC. It would be an interesting challenge to extract malonate from beets IMO.

garage chemist - 22-4-2008 at 17:50

Here, it would be significantly cheaper to prepare malonic acid from (commercial) ethyl cyanoacetate than to buy it.
Crystalline malonic acid is expensive and does not find too much use outside the laboratory.
Even diethyl malonate is much cheaper than the free acid, and cyanoacetates are even cheaper.

benzylchloride1 - 11-3-2009 at 21:47

Cohens Practical Organic Chemistry available in the Sciencemadness Library mentions that malonic acid can be prepared by oxidizing malic acid. Malic acid is a common chemical that can be obtained cheaply from suppliers catering to home brewers. What type of oxidizing agent could oxidize malic acid to malonic acid and carbon dioxide. Cohen's book dose not give any references to a procedure. This could be a useful procedure if it procedes in a decent yield. Formatik upthread gives references to paper from the 1850's and 60's relating to this oxidation. My university does not have access to many foreign chemistry journals even through inter library loan. If any one has copies of these articles it would be much appreciated to post them on this thread. I have recently aquired some malic acid and would be interested in making malonic acid.

[Edited on 11-3-2009 by benzylchloride1]

Sauron - 11-3-2009 at 22:01

The usual suspects

Aqueous permanganate
Chromic-sulfuric acid
Nitric acid

I urge you to consult the literature, directly (Beilstein, the ACS search engine, Houben-Weyl, etc. Merck Index, Kirk-Othmer, Ullmann's.

Almost all the non-heterocyclic volumes of Beilstein are available in References. Much of Houben-Weyl as well. You really need to work on acquiring your own library this way. It is time well spent.

benzylchloride1 - 11-3-2009 at 22:18

Sauron thank you for the advice. I have been slowly downloading material from the reference forum, but my internet connection is extremly slow, so it has been a very tedious process. I have some permangante, dichromates and nitric acid in decent quantities. If you want to make carboxylic acids via the malonic ester synthesis, 2,4-pentadione may work. Alkylate with your chosen halide. Dialkylated 2,4 pentadione could then be turned into the disubstitued malonic acid via the haloform reaction. No enolate ion can form at the 3rd carbon because it is disubstituted. After acidfing the reaction mixture, the dialkylated malonic acid could be isolated by extraction with an organic solvent. The malonic acid could decarboxylate during the work up. I wonder if a Grignard reagent could be prepared from an active methylene halide, methylene chloride is out of the question because it is used as a solvent in some Grignard reactions. The Grignard reagent could then be reacted with dry ice to make malonic acid. I do not know if dibromomethane or diiodomethane can form grignard reagents. If sodium is availiable, an organosodium reagent could be eaisly prepared and reated with dry ice to form malonic acid.

Sauron - 11-3-2009 at 22:40

2,4-pentanedione aka acetylacetone, is indeed a very useful reagent. But if I want to make anything via malonic ester synthesis I go buy diethyl malonate. (I know this is a headache in the US, I am happily not in the US.)

Acetylacetone is better purchased than made, as it is rather cheap while the prep requires a LOT of expensive BF3-etherate as I recall. See Org.Syn. for why DIY fails on economic grounds in this instance. Then look up the price of the end product versus the catalyst.

Per Acros

2,4-pentanedione $46 Kg MW 100 so $4.50/mol

Diethyl malonate $74 2.5 Kg MW 160 so $5 mol

I would not make acetylacetone into a malonic ester unless I had no other choice. We have not yet factored in cost of alkylating agent and the time and effort.

[Edited on 12-3-2009 by Sauron]

benzylchloride1 - 11-3-2009 at 22:51

I planning on obtaining some diethylmalonate soon. I get my 2,4-pentadione for free in small quantites from a friends company, so the cost is not an issue, also I use small quantites, under 25 grams per synthesis usually. I have some BF3 etherate, but would rather not waste it. I mainly use acetylacetone for making coordination complexes, but I plan on alkylating some to run a malonic ester type synthesis of a some a carboxylic acid, I have not decided yet which acid I want to make. 2,4-pentadione is 2 times more expensive then diethyl malonate.

[Edited on 11-3-2009 by benzylchloride1]

Sauron - 11-3-2009 at 23:35

From what supplier is acetylacetone 2x as expensive as malonic ester?

On a molar basis they are 10% apart in price, and both $5 or less.

If you are using 25 g of the dione you are working on a 250 mmol scale as the molar mass is 100. So even if you were not getting it for free the cost would be about $1.12. It's $45/Kg and a Kg is 10 mols.

Malonic ester (diethyl malonate) is $48/Kg or $73/2.5 Kg, As you can see it does not pay to buy the smaller packing. That is $30/Kg and a Kg is 6 mols so $5/mol. $1.25 on your scale.

If you crunched the numbers for the 1 Kg bottle then yes it would be close to 2X. But that is not how to buy it. And no there is no similar economy in the 2.5 Kg pack of the dione.

Magpie - 12-3-2009 at 08:25

re: synthesis from malic acid. See page 1 of this thread. I didn't have much luck with the Dessaignes synthesis from the 1800's. Here's wishing you better success!

Globey - 12-3-2009 at 13:30

Why not just fractionate the diethylmalonate from synthetic black cherry aroma oil (~25%), and de-esterify with NaOH?

garage chemist - 12-3-2009 at 14:25

Acetylacetone doesn't give malonate with NaOCl, not the slightest bit! Read the rest of this thread to find out why!
The methylene group of the acetylacetone gets chlorinated first, and then the molecule is cleaved into two C2 fragments and CHCl3.

benzylchloride1 - 12-3-2009 at 20:32

The experiment that I am planning to conduct with the acetylacetone does not involve synthesizing malonic acid directly, but produces an alkylated 2,4-pentadione which then will be converted to an aliphatic acid using the haloform reaction. The haloform reaction will attack carbon 3 on the pentadione moity, removing an acetic acid residue and the other methyl ketone will be converted into a carboxylic acid salt adding two carbons to the chain. Example: 2,4-pentadione + 1 equivalent bromoethane + suitable base yields monoalkylated pentadione. Haloform reaction removes one of the methyl groups attached to the ketone and also removes an acetic acis residue. After acidifing, butyric acid can be produced. I have not tried it, but I plan to make a different acid due to butyric acid's awful smell. This is a possible way around the malonic ester synthesis of monocarboxylic acids.

[Edited on 12-3-2009 by benzylchloride1]

[Edited on 12-3-2009 by benzylchloride1]

Attachment: 2,4-pentadione synthesis of carboxylic acids.doc (17kB)
This file has been downloaded 1083 times

Sauron - 12-3-2009 at 20:44

Yes I understood all that from your previous posts.

But I see no advantage to that route as compared to the malonic ester synthesis of carboxylic acids itself, they are both active methylene compounds. But, whatever floats your boat. The haloform is easy on a small scale, but does not scale up well.

benzylchloride1 - 12-3-2009 at 20:55

Im not planning on scaling up the reaction, I would only want to make around 15 to 20 grams of product, I do not have much money to spend on chemicals and since the 2,4-pentadione is free in 60 mL quantities this synthesis is currently the best route for me. I would love to have some malonic ester, but that will have to wait. I used it once in organic chemistry lab for making barbituric acid; that was a fun synthesis, but they no longer run this lab at the university because the pre medical students are way to impatient and want out of the lab in an hour. They run labs that used to be professor conducted demonstations instead of fascinating labs like the malonic ester synthesis.

Magpie - 13-3-2009 at 10:07

Quote:

but they no longer run this lab at the university because the pre medical students are way to impatient and want out of the lab in an hour.

Who accredits the pre-med curriculum? The students? Perhaps the students would just like to have their medical degrees mailed to them. Then they wouldn't even have to spend an hour in the lab. :mad:

JohnWW - 13-3-2009 at 13:28

Quote:

Originally posted by Magpie

Quote:

they no longer run this lab at the university because the pre medical students are way to impatient and want out of the lab in an hour.

Who accredits the pre-med curriculum? The students? Perhaps the students would just like to have their medical degrees mailed to them. Then they wouldn't even have to spend an hour in the lab.

That does not surprise me - in my experience, medical students, including pre-meds, are arrogant and self-centered beyond belief, because of the inflated self-importance that becoming a medical doctor seems to give them. This may also be partly due to the fact that nearly all of them are highly privileged, being from very wealthy and indulgent families, of which many are an only child, with most of them being the children of medical doctors or dentists or of pharmacists who own their own retail shops.

Sauron - 14-3-2009 at 07:08

John, I know quite a number of MDs who did not come from money, were not only children, and if they had a doctor in the family it was an uncle or aunt rather than a parent.

The children of MDs that I went to school with that I can recall went into unrelated fields.

I know radiologists, opthalmic surgeons, internists, psychiatrists, othopedic surgeons, oncologists, pathologists, and the whole panoply and I see nothing to support your comments.

benzylchloride1 - 26-1-2010 at 22:07

I have a possible solution to the unavailability of diethyl malonate. If sodium metal can be obtained or high purity sodium ethoxide from the equillibrium between ethanol and sodium hydroxide via Nicodem's method. Vogel and Organic synthesis describe a synthesis of ethyl phenylmalonate from diethyloxalate and ethyl phenyacetate. The reaction produces ethyl phenyloxaloacetate which is then heated to about 150 C under vacuum causing it to eliminate carbon monoxide and results in the formation of ethyl phenylmalonate. Perhaps this can be applied to ethyl acetate, isolating the ethyl oxaloacetate which is usually sold as the sodium salt due to its instability, heating the ethyloxaloacetate to produce the diethyl malonate. I am planning on trying this using first sodium ethoxide produced from sodium metal to see if this route is feasible before trying the equillibriation method of making sodium ethoxide. Hopefully, more members will test the equillibriation procedure and work out an good high yielding method.

Another intriguing and exciting possibility, malic acid could be esterified with ethanol to produce diethyl malate, which could then be oxidized under controlled conditions in a non-aqueous solvent, acetone, to produce diethyl oxaloacetate. This could be isolated and decarboxylated to form diethyl malonate or may even decarboxylate during the oxidation, forming diethyl malonate. This is pure conjecture, but may be the answer to this problem.

[Edited on 27-1-2010 by benzylchloride1]

Theophrastus_2 - 31-1-2010 at 16:11

Dunno if anyone else posted something like this, as there has been some stuff on dialkylated diones, but a longer (but maybe...?) simpler route would be to perform a crossed aldol reaction between acetone and another readily available aldehyde (ie. benzaldehyde). In the case of my example, this would form benzalacetone. From here, reduction of the central carbonyl would produce something like 1,5 diphenyl penta 2,4 diene. Cleaving the double bonds with a common reagent like potassium permanganate would produce malonic acid and benzoic acid, which can be easily filtered off.

Pros: Aldol reaction is a very simple procedure
Potassium permanganate is quite available

Cons: The intrinsic problem is that the chosen procedure for the reduction of the central carbonyl would have to be selective, so as to keep the double bonds intact while reducing the central carbonyl. I have absolutely no idea of how this might be done. Thoughts? And perhaps some references for a means of reduction...

Attachment: Malonic Acid.sk2 (9kB)
This file has been downloaded 1122 times

Another means might be to cleave the benzalacetone directly to produce oxamalonic acid. Again, reduction is the problem, having to selectively reduce the central carbonyl, and keeping the carboxylic acid carbonyl, as is.

[Edited on 1-2-2010 by Theophrastus_2]

[Edited on 1-2-2010 by Theophrastus_2]

Pages: 1 2 3