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Author: Subject: microwave dehydration of sodium acetate
Polverone
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[*] posted on 14-6-2008 at 17:46
microwave dehydration of sodium acetate


This is a fairly trivial topic for a separate thread, but I didn't want to clog up the acetic anhydride thread with a digression about preparing anhydrous sodium acetate. Sauron suggested one reason that attempts to use sulfur halides with sodium acetate to produce acetic anhydride have failed may be inadequate drying of the sodium acetate. Some years ago I noted what seemed to be a convenient method for preparing anhydrous sodium acetate, and I wonder if others could confirm/refute my findings.

Recently I tried to reproduce what I recalled from memory, this time taking notes. I weighed out 10.0 grams of hydrated sodium acetate, believed to be approximately the trihydrate, prepared some years ago from distilled white vinegar and baking soda. The material is a light tan color, obviously containing traces of materials other than sodium acetate. The imperfect starting material is the greatest reason I doubt my results.

When this material is placed in a microwave oven set to high, no heating is apparent. If a few drops of distilled water are first added, the entire mass soon melts and boils vigorously under microwave irradiation until boiling abruptly ceases, leaving behind a "fluffed" mass of what appears to be anhydrous sodium acetate. The material is easily cut/powdered with a spoon. After the microwave treatment, my 10.0 g of starting material was reduced to 6.0 grams. To within the limits of my measurement capabilities, this corresponds with the water loss expected when the trihydrate is converted to the anhydrous salt.

As a crude check for traces of remaining water, I placed a few hundred milligrams of material in a test tube and heated it over a flame until it melted and just started to darken with carbon. I was unable to spot any condensation in the neck of the test tube.

It's still possible that this material contains traces of moisture that would render it unsuitable for especially sensitive reactions. It's also possible that my imperfect starting materials have concealed problems. I hope to repeat my experiments using lab-grade materials. At the very least, this seems like a nice shortcut for preparing near-anhydrous sodium acetate from hydrates as a prelude to treating it as described in (e.g.) Gattermann.

I believe I may have posted this information before in an off-hand way, either here or on the Hive, but I've been unable to find any other references to the microwave dehydration of sodium acetate. This technique may be applicable to other hydrated salts too. The novel thing seems to be that the solid hydrated form does not really seem to "couple" with the microwave radiation, so microwaving cannot effectively dehydrate it, but a small addition of liquid water permits local heating to occur which melts the whole mass into a form that strongly absorbs microwaves. When the water is gone and the material solidifies, the heating stops just as abruptly.

I've tried this same technique to prepare anhydrous magnesium sulfate, and it did not seem so elegant. Heating certainly occurred, but the material tended to violently pop as it dried, and the thermal stresses broke at least one pyrex dish.




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UnintentionalChaos
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[*] posted on 14-6-2008 at 17:52


This works for calcium chloride by the way. I also did it with magnesium sulfate but I let it run an extra minute or two after sizzling sounds ceased (which did little to the calcium chloride but make it somewhat hotter) and it got so incredibly hot in that time period that the bottom of the soda lime glass dish it was in melted and fused to the glass plate in the microwave. I would consider this a great way to generate immensely high temperatures (with good insulation of course) if decomposition wouldn't be an issue.

Sputtering and fusing is relatively unavoidable with these two salts. They'll need to be crushed afterwards which yields a fine powder or lumps depending on how you go about it.

[Edited on 6-14-08 by UnintentionalChaos]




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[*] posted on 14-6-2008 at 18:48


Hmm, I found it quite easy to make truly anhydrous fused sodium acetate from fusion of the trihydrate in a porcelain crucible.

The only issue is that it dehydrates in stages, not at once, and requires attention and care.

If the trihydrate is heated, it first melts cleanly, then starts to boil, and then quickly solidifies to a hard white mass by water loss.
Now, and this is important, this solid still contains a lot of water, and further dehydration occurs with a second melting.
The mass needs to be separated from the crucible walls and crushed a bit, and heat then has to be applied carefully to avoid overheating the liquid anhydrous sodium acetate that slowly forms in contact with the hot crucible walls and stays at the bottom of the crucible.
Everything has to melt completely again, and the resulting anhydrous melt is as mobile as water!
This is then poured out on a sheet of metal, allowed to solidify, and immediately coarsely powdered while hot and stored airtight.

There always is some degree of darkening during the second melting phase- it's unavoidable. The melt becomes brown to black, but it must not start to give off visible amounts of gases, this would mean overheating.
Impurities in the starting materials are responsible for excessive darkening- with 5% "distilled vinegar", the food product, it will be very bad. Don't use this!
Use technical or better reagent grade 60 or higher percent acetic acid for preparing the sodium acetate. Those stronger grades of acid are usually prepared by e.g. oxidation of acetaldehyde instead of fermentation, which means they are purer.




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[*] posted on 14-6-2008 at 20:18


A microwave oven heats dielectrics such as water with heating power delivered being roughly

power = P0 x sample volume x H2O density in sample

when H2O thinkness is of order wavelength or less, and provided power << max power of oven.

It is a very bad way to eliminate the final few VOLUME% H2O from a sample because the power dissipated per volume drops drastically at those levels. For NaAc with a molar mass of 82 the volume% H2O is almost 5 times less than molar%. For example at 25% molar water in the NaAc the MW power input will be reduced by a factor of 20 with respect to that at full power - hardly enough to boil off that 25% in almost all cases.

The method garage chemist suggests is much better I think.

[Edited on 15-6-2008 by len1]
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[*] posted on 14-6-2008 at 20:31


But that assumes the material being heated is a good dielectric, which it may not be.

Silicon carbide, magnetite or carbon composite crucibles are a natural option here. You can easily melt salt, an excellent dielectric, in such a crucible. In fact, you can easily melt silver, an excellent reflector at any temperature, given the same amount of insulation of course.

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[*] posted on 14-6-2008 at 20:38


That is similar to how I made 'hopefully anhydrous' sodium acetate. (I never did math to check)
Baking soda and vinegar, followed by repeated microwaving to boil all the water and eventually to a solid powder. I never observed the transition from viscous liquid to solid. But it worked just fine to distill out glacial acetic from.
Great to hear it was anhydrous when I used it.



[Edited on 14-6-2008 by The_Davster]




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[*] posted on 14-6-2008 at 21:14


Molecules in a solid dielectric cant 'move' as they can in a liquid hence solid dielectrics - in general - are much less efficient in terms of heating compared to liquids - compare H2O and ice. NaOAc and NaCl are ionic solids in the latter category and absorb far less energy from the em field - in fact the microwave walls probably absorb as much. You can use a suitable (lossy-dielectric) container as with any heating form or even a focusing device, but you dont have good control. You lose all the 'conceptual' advantage of the microwave method where you heat the specimen just enough to eliminate the water but no more. So why bother. Just thermally soak it at a fixed temperature in an oven as per gc.

The difference between Ac2O and AcOH is 15% H2O by weight. NaOAc adequately dry for AcOH synthesis can still be far from dry for Ac2O purposes.

[Edited on 15-6-2008 by len1]
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[*] posted on 15-6-2008 at 03:36


Wouldn't the easiest way be to go via the NaOH/EtOH saponification of ethyl acetate? The only issue would be whether the NaOAc is sufficiently thirsty to dehydrate the alcohol (and the NaOEt produced in equilibria therewith), although that would mean that anhydrous sodium acetate is a fucking handy dehydrating agent for aliphatic alcohols and I can find nothing to suggest that to be true. If it cannot dehydrate the alcohol, then the azeotropic removal of water, ethanol and ethyl acetate will give some awfully close to dry sodium acetate in a lot less time (and in a whole lot purer form) than from vinegar. Possibly a good bit dearer (depends how much 10L of household vinegar is, doesn't it?).

EDIT

A similar process is discussed here

http://jas.fass.org/cgi/reprint/33/6/1310.pdf

[Edited on 15-6-2008 by LSD25]




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[*] posted on 15-6-2008 at 05:56


Unfortunately, for most people, ethyl acetate is rather expensive. Also, the rising cost of ethanol, energy and chemicals "derived" from oil ($134USD a bbl) is raising the cost. The price of EtOAc almost doubled a few years ago.

Right now, bulk EtOAc purchased by the 55gal drum is about $7 a gallon (not including the drum).

This sucks. I wonder sometimes how chemistry will survive without cheap petroleum derived chemicals.




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[*] posted on 15-6-2008 at 13:05


BTW. Sodium acetate bacomes anhydrous slightly above 60°C and melting this salt is optional.
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[*] posted on 15-6-2008 at 17:23


I'm thinking that Polverone may be onto something useful, for several reasons.

The first is that the microwave method looks as if it may be much less likely to overheat the acetate, as heating drops off when liquid water is gone.

The second is that the product is obtained in an easy to powder form, described as a fluffy mass.

Those two points may be resulting in a high quality product. In Inorganic Synthesis V2 is a procedure for producing porous boron oxide. The normal method is to fuse boric acid, at the tepid temperature of 600 to 1000 C; then chilling it to a hard, difficult to powder glass. The alternative method in I.S.v2 is to pull a vacuum on the boric acid, then slowly raise the temperature to 200 C; this results in a lightly sintered porous product that rather energetically rehydrates. On a large scale a multi-stage fluid bed dryer can be used.

The open structure of the dehydrating acid/oxide in the alternative gives much more easy loss of moisture, allowing far lower temperatures to be used. I would not be surprised if the microwave dehydration of sodium acetate follows in a similar way, as water is driven off an open sponge of solid anhydrous sodium acetate acetate forms, to which the remaining liquid water plus acetate clings. This open structure allows water vapour to readily escape from the drying mass (applying a vacuum late in the operation might speed it up and insure full dehydration).
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[*] posted on 16-6-2008 at 00:53


Not quite on topic, but I'm fucked if I know which one to put it in - if this is the wrong place, please shift it

<a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA314&ci=1 8,843,982,822&source=bookclip"><img src="http://books.google.com.au/books?id=TFA6AAAAMAAJ&pg=PA314&img=1&zoom=3&hl=en&sig=u3n4IjWRCiOBnZ81PCa8UKJlLcM&ci=18,843,98 2,822&edge=1" border="0" alt="Text not available"/></a><br/><a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA314&ci=1 8,843,982,822&source=bookclip">The Elements of Experimental Chemistry By William Henry</a>

<a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA315&ci=1 8,18,975,1604&source=bookclip"><img src="http://books.google.com.au/books?id=TFA6AAAAMAAJ&pg=PA315&img=1&zoom=3&hl=en&sig=xdirZiyFcefe5-pwSW_0mVW4YF0&ci=18,18,975 ,1604&edge=1" border="0" alt="Text not available"/></a><br/><a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA315&ci=1 8,18,975,1604&source=bookclip">The Elements of Experimental Chemistry By William Henry</a>

<a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA316&ci=2 2,4,978,1639&source=bookclip"><img src="http://books.google.com.au/books?id=TFA6AAAAMAAJ&pg=PA316&img=1&zoom=3&hl=en&sig=_tHl6hysQXRjgs2cBWwKc1-Miek&ci=22,4,978, 1639&edge=1" border="0" alt="Text not available"/></a><br/><a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA316&ci=2 2,4,978,1639&source=bookclip">The Elements of Experimental Chemistry By William Henry</a>

<a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA317&ci=0 ,22,1000,1643&source=bookclip"><img src="http://books.google.com.au/books?id=TFA6AAAAMAAJ&pg=PA317&img=1&zoom=3&hl=en&sig=CyPchkq8-RYqYJrumzfnxH9Zzws&ci=0,22,1000 ,1643&edge=1" border="0" alt="Text not available"/></a><br/><a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA317&ci=0 ,22,1000,1643&source=bookclip">The Elements of Experimental Chemistry By William Henry</a>

<a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA318&ci=6 8,36,932,1629&source=bookclip"><img src="http://books.google.com.au/books?id=TFA6AAAAMAAJ&pg=PA318&img=1&zoom=3&hl=en&sig=qFV1eWv4jKOXm76DT07lk-Vo6rk&ci=68,36,932 ,1629&edge=1" border="0" alt="Text not available"/></a><br/><a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA318&ci=6 8,36,932,1629&source=bookclip">The Elements of Experimental Chemistry By William Henry</a>

Now, pardon me if I am wrong, but an inflammable acetic acid from distillation of an acetate? Hmmm, whatever could that be?

<a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA319&ci=1 1,15,989,1650&source=bookclip"><img src="http://books.google.com.au/books?id=TFA6AAAAMAAJ&pg=PA319&img=1&zoom=3&hl=en&sig=bLuN-8e3WFaQPjusmpNKLNkw2b8&ci=11,15,989 ,1650&edge=1" border="0" alt="Text not available"/></a><br/><a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA319&ci=1 1,15,989,1650&source=bookclip">The Elements of Experimental Chemistry By William Henry</a>

<a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA320&ci=1 8,18,982,1647&source=bookclip"><img src="http://books.google.com.au/books?id=TFA6AAAAMAAJ&pg=PA320&img=1&zoom=3&hl=en&sig=UtcI04DweJnYL12y_mV6Egir1Y0&ci=18,18,982 ,1647&edge=1" border="0" alt="Text not available"/></a><br/><a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA320&ci=1 8,18,982,1647&source=bookclip">The Elements of Experimental Chemistry By William Henry</a>

<a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA321&ci=2 9,58,971,1567&source=bookclip"><img src="http://books.google.com.au/books?id=TFA6AAAAMAAJ&pg=PA321&img=1&zoom=3&hl=en&sig=DIouXOsJ7kGk4YcW3yrUikam-jY&ci=29,58,971 ,1567&edge=1" border="0" alt="Text not available"/></a><br/><a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA321&ci=2 9,58,971,1567&source=bookclip">The Elements of Experimental Chemistry By William Henry</a>

<a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA322&ci=0 ,0,1000,1661&source=bookclip"><img src="http://books.google.com.au/books?id=TFA6AAAAMAAJ&pg=PA322&img=1&zoom=3&hl=en&sig=I7A3R6NaUTjo22VhjM56jk3T6A8&ci=0,0,1000, 1661&edge=1" border="0" alt="Text not available"/></a><br/><a href="http://books.google.com.au/books?id=TFA6AAAAMAAJ&lpg=PA316&ots=3EOFWOBxod&dq=distillation+acetic+acid+charcoal&pg=PA322&ci=0 ,0,1000,1661&source=bookclip">The Elements of Experimental Chemistry By William Henry</a>

This is about the only way I could think of to describe the find - I even downloaded the entire thing to find that the specific page that was most interesting (and unless read in context, less than compelling) did not come out.

[Edited on 16-6-2008 by LSD25]




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[*] posted on 16-6-2008 at 06:35


Note that this is early 19th century chemistry, at that time the formula for water was HO atomic weights were frequently half or double their current value, and quantitative organic analysis was close to using a dartboard. This can lead to difficulty interpreting some of the reactions and results.

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[*] posted on 16-6-2008 at 08:50


I always use the microwave to dry Epson salts - magnesium sulfate. When the powder is microwaved continuously to dryness, the result is a rock hard fused mass that is very difficult to use. My work around is to watch the powder carefully as it heats up. When it reachs the point where it just starts to fuse, I stop the heating and pound the mass to a powder with a hammer between a folded over sheet of heavy cloth or plastic. Then reheat, repowder and repeat. I end up with a much lighter, fine, white opaque powder which I have assumed is anhydrous. I believe that when the sulphate has reached a point where it no longer heats up, it must be free of water molecules. I have not tested this assumption because the dried powder has served my needs adequately.
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[*] posted on 16-6-2008 at 14:39


I like the idea of this very much!
Did anyone try things such as
H2SO4 (preferably in a modified MW with a gas outlet)
H3PO4


and the many salts and oxides that would work just so much better if it was dehydrated properly?
How about using this as a method of dehydrating plaster of paris? Would be interested to hear of ideas, and known successes.
Of course heat can be used, but this seems a lot easier, and perhaps gentler in some ways.




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[*] posted on 16-6-2008 at 20:34


This seems an appropriate thread to add a few comments re microwave heating versus conventional methods. I have recently been performing some tests on this subject. I have been using two old microwave ovens, one a cheap one rescued from a dumpster (made 2005), and the other an ancient but still good combination microwave/ forced convection model dating from 1983, only recently retired from long and honorable cooking service. The older has a defrost cycle. The dumpster model is a crude cheapie with painted steel (China), the other a large refined model (Japan) with stainless steel interior and a ceramic tray. Both made for GE. The quoted power consumed is 950 w for the cheapie and 1.6 KW for the other. (Note: These are probably KVA rating, and certainly not power delivered.)

Power delivered to Load

In order to test this ordinary tap water was (a) brought to boiling temp (assuming a slight superheating to 104C); (b) The quantity evaporated in a timed interval (~ 10 - 15 mins) measured by weighing. The first measurement is fraught with inaccuracies. The time is short and the heat absorbed by the vessel containing the water is difficult to estimate. Method (b) is more reliable. The latent heat of water greatly exceeds the heat required to bring the water to boiling, which is estimated (and added). The results of several trials with different positions of the container (the older model did not have a rotating plate but the microwaves are somewhat randomized by the metal fan used in force air convection in the traditional heating mode. The newer one had one but the drive did not function. The result was, averaged, that about 775 watts of heat was delivered to the load from a rated input of 1.6 KVA (assumed, not measured). A test on the smaller over gave 515w delivered to the load (net efficiency is thus 48-53%). The turntable was non-functional in the cheapie.

The small oven had a ‘power control’ of 0.1 – 1.0 max power, achieved by duty cycle of about 13 secs on followed by about 100 off for the minimum and pro rata for higher levels.

{Newer and more expensive models may use a switched power supply, with consequent improvement in efficiency, less weight (the give-away) and far more rapid PWM for power control. The current culinary version in use, a Panasonic, is like this.}

Some Results and comments

1) The MW proves very good at drying desiccants. Those tried were silica gel, montmorillonite clay, Na2SO4, CaCl2. The last was completely liquid, density 1.33 saturated solution. The final product from the CaCl2 was a pure white low density solid rather like expanded polystyrene. I cannot guarantee all water is removed but the samples seemed to be essentially anhydrous.. To be sure, use of the oven on convection heat would have been a good idea after the evaporation was complete. Covering the sample with a watch glass or plate stops spattering which becomes a problem toward dryness.

Have not tried CaSO4.

(2)Concentrating H2SO4. A sample of about IN acid produced by electrolysis was evaporated to about 4N with no problems. The sample was too small to continue, but this could be continued until the acid exhibits noticeable vapor pressure, I feel. Volatile acids are a no-no for obvious reasons of corrosion. This is rather limits to sulphuric and phosphoric acids, and perhaps some organic ones.

(4) Certain inorganic salts (notably iodides and possibly bromides, sulphides, selenides or tellurides), some metal oxides (especially in higher oxidation states, such as MnO2, Fe3O4, etc) and also hydroxides are either electronic or ionic conductors and semiconductors. Attempts to dry these can result in melting and/or incandescence. I tried to dry a sample of impure wet ZnO and got glowing hot spots forming at a white heat. Likewise metal powders are unsuitable. Bulk metal, such as Al or Cu foil, causes no problems unless sharp edges are present, because they are excellent conductors. But lossy materials such as Fe with a large skin depth might cause problems with excessive heating.

(5) The oven begins to operate on near no load as the sample finally dries. This is not a good thing as the magnetron may suffer excessive heating and consequent damage. It is interesting that no attempt is made in these ovens to copper plate the steels used for the walls, which are quite lossy. This may be to ensure a minimum subsidiary load when running empty. SS is particularly lossy at 2.458 GHZ.

(6) It is well known that organic reactions, especially in polar solvents with polar reactants, can often be carried out with much greater rapidity using MW heating. This is thought to be due to the oscillatory motion imparted to the dipoles of the reactants by the EM field, the same being responsible for the heating effect in polar liquids. Certain inorganic reactions, especially oxidations and some reductions, are notoriously slow. It will be interesting to try this. I intend to, and will keep you posted if results warrant.

(7) Neither convection oven nor MW heating should be attempted in organic polar solvents of low flash point, for obvious reasons.

(8) Never ever open a MW oven interior unless you know what you are doing and used to high voltage (c. 5KV). If you do, unplug it, else we’ll hold a requiem mass for you.

(9) Before I forget, you cannot measure the temperature in a MW oven with a mercury or any other thermometer I can think of. A mercury thread is too lossy and would either vaporize and expode or even ionize. Electronic means such as semiconductor junctions, thermocouples etc., will experience heating or destruction due to the high E field.
Regards,

Der Alte

[Edited on 16-6-2008 by DerAlte]
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[*] posted on 16-6-2008 at 20:50


Of course CaCl2.6H20, Na2SO4.10H20, epsom salts etc can be dried substantially with respect to the initial state.

CaSO4.1/2H2O with a low water content and high temperature of final dehydration is much more like drying NaOAC for use in Ac2Othough - and just like the latter I dont think it will work.

[Edited on 17-6-2008 by len1]
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aliced25
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[*] posted on 17-5-2013 at 21:55


Anyone tried reactivating 3A/4A sieves in a microwave? There are patents on the dehydration of phosphoric acids to meta- and polyphosphoric acids using MW, as well as meta- and polyphosphoric acid salts.

[Edited on 18-5-2013 by aliced25]




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[*] posted on 18-5-2013 at 00:35


Quote: Originally posted by aliced25  
Anyone tried reactivating 3A/4A sieves in a microwave? There are patents on the dehydration of phosphoric acids to meta- and polyphosphoric acids using MW, as well as meta- and polyphosphoric acid salts.

[Edited on 18-5-2013 by aliced25]


Thats considered standard procedure by most people I know. Staged predrying in the convection oven and storage, Final activation in the microwave directly before use, 3 to 5 minutes on max suffices nicely.

@ DerAlte:
There are household microwave ovens which come with a thermocouple to measure temperature in the heated medium, I own an older MIELE from 1988 incorperating this feature, see the attached picture.

The measuring by thermocouple is rather simple, see the attached patent.


/ORG



[Edited on 18-5-2013 by Organikum]

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