Sciencemadness Discussion Board - Methanol --> Formic Acid

Methanol --> Formic Acid

pHzero - 18-5-2009 at 09:45

I'm trying to make some formic acid, and I've got 25 litres of methanol, so I was thinking of oxidising it into HCOOH

The two methods I've heard of are using acidified Na2Cr2O7 or KMnO4 but I haven't really got a clue about them. Can anyone provide balanced symbol equations for these reactions so I know how much to use, what the byproducts are etc? Thanks

I read on wikipedia that when chromate's used as an oxidiser, the hexavalent chromium gets reduced to Cr(III), so I thought the formula might be something like this:
CH3OH+Na2Cr2O7-->HCOOH+Cr2O3+H2O+.... well then I had no idea where the 2Na+'es or the remaining 2 O's went.

vulture - 18-5-2009 at 09:53

If you're not capable of balancing equations, you're certainly not capable of carrying out large scale oxidation reactions using hexavalent chromium. Unless you like poisoning yourself or blowing yourself and your surroundings to hell in a gigantic fireball.

I suggest you study some more chemistry before attempting this.

Jor - 18-5-2009 at 10:04

I agree with vulture!

I would really not recommend to do an oxidation with hexavalent chromium on such a large scale! Cr 6+ is nasty.

I'm not sure, but if you use neutral or alkaline KMnO4, the oxidation stops at the carboxylic acid right? Or will the formic acid be further oxidised to CO2?

pHzero - 18-5-2009 at 10:07

It isn't, I don't think, a simple case of balancing the equation, is it? Cr(III) doesn't seem right
The best vaguely balanced equation i could make was this, but it still seems badly wrong
3H2HCOH + 2Na2Cr2O7 --> 2Cr2O3+NaOH+4H2O+3NaCOOH

Ok sorry for my stupidity, I've learnt most of what I know about chemistry from the Internet, so the majority of it's wrong - my chemistry teacher's particularly unhelpful and pedantic. Eg, when I asked him if you could distill sulfuric acid, he said "No, why would I want to do that?" then told me to go away. And when I asked him what happens to the dichromate ion when it oxidises an alcohol, he told me that it reduces it, then walked away. (Yes, I can tell it reduces it, but to what oxidation state?)
I hope to god that I dont get him teaching me AS level next year :s Aren't teachers usually pleased when someone takes an interest in their subject?

[Edited on 18-5-2009 by pHzero]

DJF90 - 18-5-2009 at 11:16

pHzero - Your teacher sounds like they're looking out for your best interest. They don't want to encourage some kid to distil sulfuric acid. You'll learn all about oxidation of alcohols in AS chemistry. You cant run before you can walk, and without being able to balance the equation then I think you're probably on thin ice attempting the practically by yourself. The best thing you can do is go an learn. It is important here to first ignore the spectator ions, as they are not participating in the reaction (as the name suggests...). Once you have balanced everything else then you just need to add in the appropriate quantities of each.

As I am not a cruel heartless bastard (merely looking out for your best interests, much like you reacher I assume...) I will point you in the right direction. Chemguide is an awesome site for helping with A-level chemistry:

http://www.chemguide.co.uk/

In particular you might want to look at this page...

http://www.chemguide.co.uk/organicprops/alcohols/oxidation.h...

I did notice that they did not show the half equation for the methanol/formaldehyde or methanol/formic acid couple. You can easily deduce these with a little knowledge of half-equations. I suggest you look at this link first:

http://www.chemguide.co.uk/inorganic/redox/equations.html#to...

panziandi - 19-5-2009 at 02:22

pHzero - Both formaldehyde and formic acid are quite distinct (and much unlike other aldehydes and carboxylic acids) in that they are powerful reducing agents. In fact formic acid will reduce sodium dichromate to chromium (III) and itself be oxidised to carbon dioxide. Oxidation of methanol using dichromate or permanganate will be a shoddy reaction. Not to mention the "scare" of the carcinogenic status of hexavalent chromium.

A much better route for you to use would be the glycerol-oxalic acid method whereby glycerol and oxalic acid are heated together forming glycerol oxalate followed by in situ decarboxylation and if I recall correctly transesterification with additional oxalic acid leads to distillation of formic acid. Other by-products are possible and can be favoured and I'm sure this will be a much better method for you. Both reagents are easily obtained in the UK.

I wouldn't distil sulphuric acid, you can of course boiling point is ca 300*C and is 98%, have you seen how vicious conc sulphuric is at room temperature? Now imagine its vapour and liquid at 300*C ... this is exactly what you teacher is saying! You have to use all glass set up for the distillation, and what is the point? It is cheap and readily available anyway! Only reason one would distil it is to produce high purity sulphuric acid free of metal ion traces for critical processes.

EDIT: why would you want to distil sulphuric acid if you have 2.5L of 98% AnalaR grade ... mate keep those reagents despite the purity they are very handy reagents for organic chemistry, washing, salting out, extraction etc!

[Edited on 19-5-2009 by panziandi]

DJF90 - 19-5-2009 at 03:03

I don't quite think he's capable of organic chemistry just yet...

pHzero - 19-5-2009 at 08:38

Quote: Originally posted by panziandi

A much better route for you to use would be the glycerol-oxalic acid method whereby glycerol and oxalic acid are heated together forming glycerol oxalate followed by in situ decarboxylation and if I recall correctly transesterification with additional oxalic acid leads to distillation of formic acid. Other by-products are possible and can be favoured and I'm sure this will be a much better method for you. Both reagents are easily obtained in the UK

I read about the oxalic acid+glycerol method, but i read that an hour of heating and condendsing yields about 10ml of HCOOH

Quote:

I wouldn't distil sulphuric acid, you can of course boiling point is ca 300*C and is 98%, have you seen how vicious conc sulphuric is at room temperature? Now imagine its vapour and liquid at 300*C ... this is exactly what you teacher is saying! You have to use all glass set up for the distillation, and what is the point? It is cheap and readily available anyway! Only reason one would distil it is to produce high purity sulphuric acid free of metal ion traces for critical processes.

EDIT: why would you want to distil sulphuric acid if you have 2.5L of 98% AnalaR grade ... mate keep those reagents despite the purity they are very handy reagents for organic chemistry, washing, salting out, extraction etc!

[Edited on 19-5-2009 by panziandi]

It's cheap and readily available if you have a company credit card, but unfortunately I don't

The H2SO4 which I have got is from my mum's work and once its gone, its gone. Anyway, I have got an all-glass distillation kit - I've got several quickfit flasks, a still head and 2 condensers

I can't imagine distilling H2SO4 would be too much different from distilling something like water, apart from the higher temperatures, as long as you don't boil off the H2SO4 faster than the condenser can condense it.

And DJF90, maybe you're right, perhaps I'll stick with inorganic for the time being xD
Since Na+ is a spectator ion in the reaction, I expect I'll end up with NaCOOH, which I'd then have to acidify with H2SO4 and then I'd add too much and dehydrate it to CO+H2O.... that'd be a right kerfuffle xD (Those who've never watched Little Britain, ignore that last bit)
That chemguide website looks quite interesting/useful though, thanks for pointing me there

DJF90 - 19-5-2009 at 11:35

Seriously how much H2SO4 do you think you'll use??! I expect 2.5L to last me a LONG time (possibly life...also make sure you store your chemicals SAFELY). You should not use conc H2SO4 to acidify HCOONa (notice where the H is... writing NaCOOH is misleading as it appears you've lost the carbon's hydrogen). However a "dilute" solution (i expect something like 50%?) should yield you no problems (or use another acid)

The oxalic acid-glycerol synthesis will not yield you 10ml of formic acid if you do it on a larger scale. I cant remember how much reactants the preparation I read used but it will vary. If you do decide to go this route then watch out for the acrolein that may sometimes be produced (in small quantities).

panziandi - 19-5-2009 at 13:47

Yes I fully agree with DJF90, you will likely never get to the end of that 2.5L unless you use it to unblock your drains or if you seriously scale up your reactions! I think I have literally gotten through less that 250mL in the past two years!

By the time you have gotten to the end of the bottle I bet you will be in uni and spending so much time in a lab you wont have time to do anything at home, or will have sources to replenish your stock. Either way you will be at a point where you would cringe at yourself for having thought about distilling sulphuric acid (especially from commercial OTC sources... YUCK!)

Anyway, a quick google revealed a Rhodium synthesis for formic acid using the glycerol-oxalic acid technique:

Quote:

Required: Glycerol 70ml; oxalic acid 40g. Since glycerol is a very hygroscopic substance, it is necessary first to ensure that the sample used is anhydrous. For this purpose, place about 70ml in a porcelain evaporating-basin, and heat it carefully over a gauze (preferably in a fume-cupboard), stirring it steadily with a thermometer until the temperature is 175-180°C then maintain this temperature for a further 5 minutes. Allow the glycerol to cool, but while it is still warm (i.e., before it becomes viscous) pour 50ml (63g) into a 250 ml distilling flask containing 40g of powdered crystalline oxalic acid. Fit a thermometer in the flask so that the bulb is completely immersed in the glycerol mixture, and then fit a water-condenser to the flask. Heat the mixture carefully over a gauze so that the temperature rises to 110-120°C, and then adjust the heating so that the temperature remains within these limits. A vigorous effervescence of carbon dioxide occurs, and the aqueous formic acid begins slowly to distill over. When the effervescence tends to subside, remove the Bunsen flame and allow the temperature to fall to 70-80°C: then add a further 40g. of powdered oxalic acid, and continue the heating as before. Ultimately 25-30ml of distillate is obtained, the total period of heating being about 1 hour.

It would be nicer synthetically and cheaper than the dichromate method. Both glycerol and oxalic acid are cheap and available in bulk. Glycerol especially so since it is a by-product from bio-diesel.

Alternatively consider a Cannizzaro reaction on formaldehyde although if you can't get formic acid you may struggle with formaldehyde.

Ever considered distilling formic acid from kettle descaler brand Kilrock? Or even from ants, was the method of choice for many years! Much nicer than distilling sulphuric acid

497 - 19-5-2009 at 14:49

If I understand correctly, you just need concentrated H2SO4 right? So why is there talk of distilling it? Panziandi is right, you really don't want to deal with that.

Since it is so much less volatile than water you can just boil the water out of it in an open borosilicate flask.. Its not hard, just make damn sure you do it outside with the breeze away from you.

Unless lead-acid battery electrolyte is for some reason unavailable in the UK, just save up some change and go buy some from a car parts store. It's cheap, not suspicious and probably as pure as you'll ever need. If they won't sell it to a minor, get a your parents or a friend to buy it..

I was boiling down sulfuric when I was 15.. As long as you're smart about it, it's not that bad. There are many worse things a teenager could be doing with his time.

Good luck.

Jor - 19-5-2009 at 15:52

Boiling down of sulfuric acid, I regard that as very dangerous, especially in the end, when you approach the 90+%. Have you ever realised what 300C conc. H2SO4 will do your skin? Even thinking of the fact that there is a small chance of breakage

Use a good beaker.
The fumes are carcinogenic.

I agree, sulfuric acid is a key chemical, but it lasts a long time. I use hydrochloric acid nitric more often to be honest, they are more useful in inorganic synthesis. I use H2SO4 a lot to acidify in test-tube chemistry, but hardly in synthesis.
But 250mL in 2 years! I must say that use it more. I think I have used about 300-400mL in 14 months. I have probably lifetime supply. Once we had a great threat in end 2008 that acids were about to be regulated and unavailable for amateurs (I'm not sure if the threat still presists), I quickly bought more. I know have a total 4L H2SO4, 2,5L HNO3 and about 7,5L HCl. However, i moved most of it to another place, to a family member who lives outside of the. I dont like to have so much acids around.

panziandi - 19-5-2009 at 16:37

Well boiling down dilute sulphuric acid isn't THAT dangerous! To be fair you would only be dealing with sulphuric acid at ca 100*C or so, which certainly will be quite nasty should it spill, but would be much much nicer than distilling suphuric acid at 300*C (and I'm sure someone will suggest vacuum distillation at some point).

Jor, the last thing on my mind, whilst distilling water from sulphuric aicd or distilling sulphuric acid itself, would be carcinogenicity of sulphuric acid fumes! Oh my god you have cheered my evening up 100-fold, reading about ATP binding domains of ABC transporters was killing me! That is hilarious, are you joking? There is a point when the carcinogenicity of chemicals is taken to the extreme... much like expenses

BUT - this thread was ABOUT methanol to formic acid using Cr6+ oxidant. I have suggested some other additional preparations which I feel personally are likely to be nicer, especially for a novice.

EDIT: I am at university so don't do much chemistry but sulphuric acid is usually a catalyst and I don't tend to require much of it.

[Edited on 20-5-2009 by panziandi]

497 - 19-5-2009 at 17:56

Quote:

Boiling down of sulfuric acid, I regard that as very dangerous, especially in the end, when you approach the 90+%. Have you ever realised what 300C conc. H2SO4 will do your skin? Even thinking of the fact that there is a small chance of breakage

Jesus, man, do you expect people to parade around with beakers of 300*C H2SO4 in their hands or what?? While the things you said are true, Hot conc. H2SO4 is nasty stuff I won't debate that, the risks are so easy to eliminate by simply not getting near it while its hot...

OK no more derailing for me..

PS. I'm curious what the hell you need multiple liters of formic acid for?

Back to the topic,

Sedit - 19-5-2009 at 18:58

Oxidation using DLF electrodes are discussed in comparison of the use of Ozone or the Fentons reagent for the oxidation of alcohols to carboxylic acids.

For me Ozone looks like it would be very nice as it is reported to produce 80% yeilds from the alcohol IIRC.

I did manage to produce some sodium formate and sodium acetate using Fe(II)sulfate and H2O2 but feel this method has some drawbacks when it comes to cost and yeilds.

Attachment: Electro-diamond[1].pdf (883kB)
This file has been downloaded 5475 times

panziandi - 20-5-2009 at 07:58

Sedit nice post, thanks for sharing. I am not sure how much formic acid pHzero is after. 25L of methanol, did you want it all converted to formic acid? What sort of chemistry are you planning, perhaps there are nicer alternatives to formic acid etc? Methanol itself is a very useful reagent and solvent. I certainly wouldn't want to convert the whole lot to formic acid, seriously repetitive hector-scale

pHzero, perhaps consider vapour phase dehydrogenation of some of the methanol over a metal catalyst? Producing formaldehyde which you could collect in water. A useful chemical if you can't buy it.

cnidocyte - 2-2-2011 at 05:43

Yesterday I added some K2MnO4 to 15mL of pure methanol and noticed that no reaction occured at room temperature but I placed the beaker in a bowl of hot water and fairly soon the purple solution turned clear. I don't know what formic acid or formaldehyde smell like but whatever I made it smelled kind of fruity. I'm guessing I ended up with mixture of methanol, formic acid, formaldehyde and MnO2. There were big black chunks floating around in the beaker, I'm guessing that was MnO2.

DJF90 - 2-2-2011 at 07:02

If it smelt fruity you probably got methyl formate; - formic acid is very irritating to the eyes, nose and throat, while formaldehyde is not much nicer. With a bp of 32C, I doubt any ester would hang around long enough for you to observe a separate layer, especially in a hot water bath.

1281371269 - 2-2-2011 at 07:21

Fruity is the classic aldehyde smell. Carboxylic acids are supposed to smell more acrid.

Panziandi: are you sure that methanoic acid is oxidised all the way to CO2? What makes it so much stronger a reducing agent than longer-chain carboxylic acids?

[Edited on 2-2-2011 by Mossydie]