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Author: Subject: ParaFormaldehyde preparation from formaline
Vikascoder
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[*] posted on 29-3-2012 at 09:28
ParaFormaldehyde preparation from formaline


Please spread your precious gems on me by guiding me how to make paraFormaldehyde from formaline i was just able to found on wiki that it is formed due to polymerization of formaline. when formaline is kept in cold white product precipitates and that is paraFormaldehyde
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barley81
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[*] posted on 29-3-2012 at 11:41


I left a bottle of formalin in my lab during winter. It polymerized. The white residue on the sides of the bottle was very hard. I'm not sure if there's any formaldehyde left in solution.
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mnick12
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[*] posted on 29-3-2012 at 19:40


The simplest way to prepare paraformaldehyde from a formalin solution is through vacuum distillation.

I have made paraformaldehyde three or four different times using this method, and it is quite simple. Measure out the desired volume of formalin (I usually use 100ml), set up for vacuum distillation and distill at least half the original volume off. Allow things to cool so that they can be handled, and pour the remaing liquid into a beaker. As it cools the liquid will solidify into a thick gell (usually overnight), and after drying over MgSO4 one obtains a pure white crumbly product.

This method produces a pure product which can be used in a variety of interesting reactions, such as the formylation of phenols.
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[*] posted on 30-3-2012 at 09:34


Does boiling , evaporating or freezing will work
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mnick12
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[*] posted on 30-3-2012 at 14:58


Well, boiling a formalin solution under atmospheric pressure will result in massive amounts of formaldehyde gas and little to no product. Remember formaldehyde is a gas at stp, and formalin is a solution of formaldehyde. So unless you alter the conditions in which you wish to boil the formalin solution, you will only succeed in gassing yourself. Evaporation might work I personally have never tried it, however leaving a dish of formalin to evaporate wouldn't be very prudent considering its toxic and carcinogenic properties. As for freezing, cooling solutions of conc formalin does indeed produce paraformaldehyde however this is pretty much useless as the yields are tiny. By far vacuum distillation is the easiest and safest way of concentrating formalin solutions to produce paraformalehyde.
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Vikascoder
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[*] posted on 30-3-2012 at 21:46


What could be the possible yield of paraformaldehyde by evaporating 500 ml of 37% formaldehyde solution
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[*] posted on 31-3-2012 at 08:58


Quote: Originally posted by barley81  
I left a bottle of formalin in my lab during winter. It polymerized. The white residue on the sides of the bottle was very hard. I'm not sure if there's any formaldehyde left in solution.


I also have a very old bottle of formalin which has been stored incorectly, in an unheated space during all seasons. There's a large amount of white solid at the bottom, but most of the formaldehyde is still in solution.
For most uses, you can simply shake up the precipitate and use the suspension as if it were 37% formalin. The paraformaldehyde reacts in the same way as the aqueous solution.

A simple method to determine formalin strength is to add an excess of aqueous ammonia to a weighed sample. Leave the solution to evaporate and weigh the residue of hexamine, and from this calculate how much formaldehyde was in the weighed sample.
Paraformaldehyde also quantitatively reacts with ammonia.




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RonPaul2012
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[*] posted on 31-3-2012 at 18:53


Quote: Originally posted by Vikascoder  
What could be the possible yield of paraformaldehyde by evaporating 500 ml of 37% formaldehyde solution
Are you trying to synthesis paraformaldehyde or do you just want some as a reagent ?

If it's the latter I know a US supplier who sells it for $12 a pound (I can't vouch for purity though) but I'm not sure if he does international or not.

If you're interested I could give you the link.
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barley81
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[*] posted on 31-3-2012 at 19:27


Quote: Originally posted by garage chemist  
Quote: Originally posted by barley81  
I left a bottle of formalin in my lab during winter. It polymerized. The white residue on the sides of the bottle was very hard. I'm not sure if there's any formaldehyde left in solution.


I also have a very old bottle of formalin which has been stored incorectly, in an unheated space during all seasons. There's a large amount of white solid at the bottom, but most of the formaldehyde is still in solution.
For most uses, you can simply shake up the precipitate and use the suspension as if it were 37% formalin. The paraformaldehyde reacts in the same way as the aqueous solution.

A simple method to determine formalin strength is to add an excess of aqueous ammonia to a weighed sample. Leave the solution to evaporate and weigh the residue of hexamine, and from this calculate how much formaldehyde was in the weighed sample.
Paraformaldehyde also quantitatively reacts with ammonia.


Thanks! That's good to know. I was planning to make a little hexamine from it.
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[*] posted on 31-3-2012 at 19:30


I am going to synthesize only a little bit . I only just want 50 grams . And i have lots of formaline so i am thinking to make from it without distillation so if you have any method tell me
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[*] posted on 5-11-2017 at 22:32


Sorry to revive a 5 year old thread:

Since formaldehyde is a gas, and as mnick12 states heating formalin at atmospheric pressure will release the gas, my first assumption would be that vac distilling would only cause more formaldehyde to come out of solution. Is this not the case? And if not, why?
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[*] posted on 19-11-2017 at 07:37


Don't act so shocked! :)

As I was reading this thread I also wondered the same thing. Phenomena like this are not so uncommon in chemistry. Let's check it out.

Formaldhyde is indeed a gas in it's pure form at stp, and water isn't, but why? Formaldehyde is certainly the heavier molecule, so other factors must be affecting volatility. My first suspects are hydrogen bonding and van der wals forces. Formaldehyde has a single hydrogen bond acceptor, but no donors, because of the charge state of the carbon. So it doesn't hydrogen bond to itself (like water does, with two acceptors and one donor). This is why pure formaldehyde is a gas.

But we are talking about a solution, and in a solution with water, formaldehyde is the heavier molecule, and is flanked on all sides by hydrogen bond donors. Another factor that could contribute is that formaldehyde might be polymerizing in solution at lower temps in a vacuum, whereas heating at atm might be causing the polymer to decompose back into the more volatile monomer.

But you shouldn't be surprised, you can see quirky relationships like this all the time. For example, HCl and water. At concentrations below 20%, water is the more volatile component. Also, more closely related to the topic, formic acid: formic acid has a boiling point almost identical to water (100.8°C). It forms an azeotrope with about 23% water which boils at 107°. But under vacuum, it is fairly easy to separate excess formic acid from the azeotrope, even with though the two components have a boiling point within 1° of each other. Strange stuff, no? Now, formic acid is an anomaly in that it forms a dimer in the gas phase, and doesn't obey the ideal gas laws, and this is probably the reason for it's "unexpected" behavior. But still, this just goes to show that chemists are in the business of expecting the unexpected.

I hope this helps clear things up, chemistry can make one's brain hurt...
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