Sciencemadness Discussion Board

Manganous-Ammonium alum and derivatives for use of toluene oxidation to benzaldehyde

Hilski - 18-10-2006 at 09:18

Inspired by Cycloknight's posts in the Toluene--->Benzaldehyde thread, I decided to have a go at making the manganous-ammonium alum descibed in US patent 808095. Basically the patent gives details on preparing manganese-ammonium sulfate from ammonium sulfate, manganese sulfate, and 60% sulfuric acid. To oxidize the compound to the desired manganous ammonium alum (MAA), the mixture then has current passed through it in an undivided cell, using lead electrodes. I believe, after looking at the numbers in the patent that the amount of current necessary to complete the reaction is ~0.12 amp hours per gram of oxidant.

Using scaled down ratios from the patent, I started with the following: (Note that the amounts listed are not 100% correct, going off of the patent. So anyone who has more patience than me can probably be much more precise in the measurements)

448g MnSO4 (2.97 mol)

204g (NH4)SO4 (1.79 mol) (The actual ratio in the patent is 2:1, but I am using about a .30 molar excess per CycloKnight)

690g H2O

1096g H2SO4 (~600ml 93% Rooto brand)

With stirring, I added the acid to about half the water, then dissolved all 204 grams of the ammonium sulfate in the water/acid mix. Then the MnSO4 was added, followed by the rest of the water. This mixture was stirred for a 24 full hours before all the MnSO4 dissolved. The yellowish precipitate did appear. However, the overall color of the mixture was a very pale pink color instead of the orangy-yellow color I was expecting. Im guessing it was from the MnSO4, but the drain cleaner has a bit of a brown tint to it, and that may have effected the color somewhat as well, but I'm not for sure. The (NH4)SO4 was 99% pure and the MnSO4 was hydroponic grade, and seemed to be free from any major contaminates. It was finely powdered similar to talc, with a whitish-pink color.

After the mixture had been stirring for over 24 hours, and everything was about as dissolved as it was going to get, I put the lead electrodes into the jar and hooked up the power supply. I am using a computer PSU which can supply 3, 5, or 12 VDC at 30 amps. (Only 10 amps for the 12v output) I am using the 5v output, which is actually about 4.8v across the cell. I am not using a power resistor in the circuit, since I am only using 5 volts. I am using an old ammeter from an automobile to measure the current (all my fancy electronic measuring equipment is at work at the moment). It is obviously not the most precise instrument in the world, but it's close enough for what I need it for. It is currently reading about 3.5 amps, and has been holding steady there for the whole time so far. I will try to remember to bring my electrical meters home with me tomorrow, so I can get a more accurate current measurement.

As soon as I switched on the power to the cell, a dark red color started pouring off of the anode, and within 5 minutes the entire mixture had turned a light red color. Within 15 miutes, the mixture was a suprisingly dark red color, and after about an hour, the mixture was a deep purpleish color. (See the photos below) Using the calculations I mentioned earlier, I figure it should take about 22.5 hours to complete the reaction. (652g oxidizer x 0.12 amp hours per gram = 78.24 amp hours. 78.24ah / 3.5 amps = 22.35 hours) I'll take some more pictures after the 22 hours is up, so we can see the difference in color.

Here is what the mixture looked like between 5 - 15 minutes after I switched on the power:





This is how it looked after an hour:





My caveman electrical test equipment; 4.88 volts across the cell



About 3.5 amps.....I think.



And last but not least........My ghetto fabulous 12v PC fan-harddrive magnet-2x4 lumber magnetic stirrer!



I'll be adding more to this in the near future.

Edit by Chemoleo: Changed title and merged threads!

[Edited on 24-1-2007 by chemoleo]

bio2 - 18-10-2006 at 11:34

Nice little setup! Just a wire shunt will give an accurate current
as voltage reading on a cheapo digital meter. That way your good meters won't end up coated with sulfuric fog which inevitably seems to happen.

What are the dimensions of your electrodes and how were they made?

Mr. Wizard - 18-10-2006 at 16:32

They sell a digital VOM at Harbor Freight for $3.99 that has a 10 Amp setting by changing where the leads plug in; if you have access to these cheap tool stores.

Using a wire shunt, as bio2suggested, is a great idea, but then you have to calibrate the shunt to know what the voltage reading represents. I have a couple of the old Wheatstone Bridges and can do that, but how would the regular amateur do it?

I'm also interested in the stirring device that looks like it was made from abiscuit fan. Any information on that would be nice.

Hilski - 18-10-2006 at 18:44

Quote:
Nice little setup! Just a wire shunt will give an accurate current
as voltage reading on a cheapo digital meter. That way your good meters won't end up coated with sulfuric fog which inevitably seems to happen.

What are the dimensions of your electrodes and how were they made?

Thanks. The electrodes are made out of 1/8 inch lead sheet from United Nuclear. The product name is RadMax (or something like that) and is sold as shielding material for radioactive element storage. It's pretty cheap at around 7 or 8 bucks for a 1 square foot piece. I cut the pieces out by cutting the surface with a utility knife, and then bending the sheet where it was cut to break it the rest of the way through. The cathode is 1.25 inches wide, and the anode, if I remember correctly, is 3.25 inches wide. The electrodes extend all the way to the bottom of the jar, which means about 6 inches of the electrode is actually in contact with the solution.

I didn't bother checking to see what the proper current densities were for this type of reaction, beacause I figured that it didn't matter much since this is basically a brute-force oxidation of a compound that has no possibility of being reduced at the cathode regardless. That little fact makes this electrolytic oxidation a piece of cake compared to a lot of others, that require membranes, temperature control, voltage/current control etc.

Quote:
They sell a digital VOM at Harbor Freight for $3.99 that has a 10 Amp setting by changing where the leads plug in; if you have access to these cheap tool stores.
Using a wire shunt, as bio2suggested, is a great idea, but then you have to calibrate the shunt to know what the voltage reading represents. I have a couple of the old Wheatstone Bridges and can do that, but how would the regular amateur do it?

I'm also interested in the stirring device that looks like it was made from abiscuit fan. Any information on that would be nice.



I thought about throwing together a shunt type ammeter, but I didn't have the proper resistors on hand to do it. I did try a 12v 50 watt RV lightbulb I had, and it worked. But it limited the circuit to about 2 amps, and I wanted to get at least 3-4 amps to cut down on reaction time. At 12 volts, without the lightbulb in the circuit, the cell behaved just like Mr. Ohm predicted it would, and drew the full 10 amps of current available from the 12v output. It would have heated the cell up way too much, and would have eventually tripped the power supply (or burned it out even). So, instead I decided to just use 5 volts, and wire in the automobile ammeter I had.

A shunt type ammeter can be calibrated by placing a trim pot in series in the shunt circuit, and placing another ammeter in the main circuit. (I know this defeats the purpose of using a shunted voltmeter in the first place, but it's really the only feasible way for most people to do it. Maybe one could borrow a real ammeter for a short time to use for this purpose) A load is applied to the circuit, and the trim pot is adjusted until the shunted volt meter matches the reading on the real ammeter. The value of the trim pot would be dependant upon what range of current one would be measuring.

Harbor Freight really is a great place to get cheap tools and equipment. I go there all the time. I work as an electrical tech for a large company, and have more crap to measure electricity with than I can shake a stick at. I guess it's kind of ironic that I use a dirt cheap multimeter from Home Depot when I'm at home. But I have an induction amp probe at work that would be pretty accurate for this purpose, and as soon as I measure the actual current again, I will post it here.

As for the stirrer, it is the simplest thing you will ever build that will do the most to help you out. I used a 12v muffin fan from an old computer, with the powerful little magnets from a computer hard drive mounted in the center of the fan. The magnets are strong enough to hold themselves to the fan without needing to use glue, or whatever to hold them on. The plate part is made out of a piece of thick Lexan with a hole drilled in the center of it big enough for the magnets to spin around in without hitting the sides. I only used Lexan because that's what I had on hand at the time, but this part can be made out of anything that won't interfere with a magnetic field. A lot of folks use aluminum for this, but some have said that aluminum bogs down the fan motor too much because of inductive loads. I can't say either way, because I haven't tried it with aluminum yet. The whole contraption is bolted to a couple of pieces of 2x4 lumber to hold it all together. It is powered by a 120v ac/dc adapter, but anything that provides 12vdc can be used as long as it has the proper current rating. The thing works beautifully, and will stir most anything as long as proper magnetic stirbars are used. Now I just have to figure out how to work a hot plate into this contraption, so I'll have something I can use for vacuum disillations, or other stuff that requires heating and stirring at the same time.

Here are some links that go into detail about building homemade magnetic stirrers:

http://www.home.earthlink.net/~jsbrewing/stirrer/stirrer.htm

http://www.instructables.com/id/E570N4YH3HERXTRU8J/?ALLSTEPS

http://brewiki.org/StirPlate

tumadre - 18-10-2006 at 19:40

a good rule of thumb is that a
10 gauge wire will have a resistance of one miliohm per foot.

and make sure not to use a stranded wire because it isn't 'exactly' 10,380 cir.mills

just measure the voltage with the meter set to the 200.0 milivolt range and you get a .1 to 200 amp range.

BTW the 10 amp shunt in the meter is .1 ohm for most cheap meters, so you will lose at least 1.6 volts at 10 amps due to the resistance of the leads+shunt.

for those chemists who commonly use molar excesses...
.0009972 ohms at 20C and will gradually reach .001114 at 50 C

Hilski - 19-10-2006 at 00:42

Quote:
a good rule of thumb is that a
10 gauge wire will have a resistance of one miliohm per foot.


There haven't been too many times when I've had to make my own ammeter, but the times that I did, I always had the appropriate sized resistor. Point being, I've never actually tried using a wire for a shunt, but I think I'll stick one in this circuit for curiosity's sake just to test the accuracy compared to a high end digital ammeter. I'm sure it's better than my old automobile ammeter anyway. :)

Also, I got to looking at the ratios in the patent again, and I think I need to add about another 110 grams H2SO4 to the mix to make it 60%. But seeing as how it's about 4:00am, my brain cannot be fully trusted. I think I'll do it anyway. I think too much acid is better than not enough in this case.

***Never mind. It's close enough to 60% already.

[Edited on 19-10-2006 by Hilski]

bio2 - 19-10-2006 at 14:00

...... I've never actually tried using a wire for a shunt, but I think I'll stick one in this circuit for curiosity's sake just to test the accuracy compared to a high end digital ammeter...........

Using a Fluke 87 DMM to calibrate copper wire shunts I have found that 5% accuraccy is atainable simply by measuring the corresponding wire lenght for 1, 10, or or 100 milliohms.

Then you can read the current directly throwing in the
adjustment factor if you want.

Or even better take the shunt out of one of those POS
$5.00 meters.

[Edited on 20-10-2006 by bio2]

chemoleo - 19-10-2006 at 16:28

Oh how annoying I can't access the patent. It requires some damn viewer, and after having spent 15 minutes on trying to get it to work, I give up. Bloody formats. Everytime I ahve to install lots of different softwares, making the system unstable ultimately and it still doesn't work in the end!

Anyway, Hilski, I imagine the formula of the manganese alum is NH4Mn(SO4)2? This is very interesting, I have never heard of this Alum to exist! I am interested in it in particular because it may form nice crystals. The solution, after the current treatment, is it clear or is it an emulsion? It certainly looks like it. Did you try to crystallise a small amount? What colour are the crystals? Are they stable?

[Edited on 20-10-2006 by chemoleo]

Hilski - 19-10-2006 at 21:15

Quote:
Originally posted by chemoleo
Oh how annoying I can't access the patent. It requires some damn viewer, and after having spent 15 minutes on trying to get it to work, I give up. Bloody formats. Everytime I ahve to install lots of different softwares, making the system unstable ultimately and it still doesn't work in the end!

Anyway, Hilski, I imagine the formula of the manganese alum is NH4Mn(SO4)2? This is very interesting, I have never heard of this Alum to exist! I am interested in it in particular because it may form nice crystals. The solution, after the current treatment, is it clear or is it an emulsion? It certainly looks like it. Did you try to crystallise a small amount? What colour are the crystals? Are they stable?


You can download the .pdf format here:

http://www.freepatentsonline.com/0808095.pdf

The formula given in the patent for the alum is (Mn2(SO4)3(NH4)2SO4)

After the electrolysis, the mixture is a dark reddish brown color, and the salts are actually in a suspension of sorts. CycloKnight was able to seperate some of the salts from the solution, and made a thick paste out of them. I'm not sure if it's possible to crystalize the alum, but the patent states that the precursor material, manganese-ammonium sulfate can be obtained as yellow anhydrous crystals. Its formula is MnSO.1/2(NH4)2SO4

You can probably learn much more about this compound from reading CycloKnight's posts than I can tell you. I've just began to experiment with it, and haven't tried to use it to oxidize anything yet. I took a couple of pictures of the mixture after the electrolysis, and will post them tomorrow hopefully.

Here's the thread with the pictures and posts made by CycloKnight:
https://sciencemadness.org/talk/viewthread.php?tid=2223&...

Quote:
Using a Fluke 87 DMM to calibrate copper wire shunts I have found that 5% accuraccy is atainable simply by measuring the corresponding wire lenght for 1, 10, or or 100 milliohms.


I would think 5% is more than sufficient for this purpose, and I will definately try use this simple method in the future. I measured the current today with a Fluke amp probe, and it read 4.1 amps, so the old automobile meter wasn't too far off the mark. I have a bunch of #10 wire, so I'll wire in a 1 milliohm shunt after I oxidize something and need to recharge the oxidant.

Rosco Bodine - 19-10-2006 at 23:28

Here attached is the pdf file for the patent .

This is a very interesting electrolytically regenerable oxidation reagent .

I am wondering if the oxidation product of ethanol
is acetaldehyde or acetic acid , or if either is possible
depending upon the temperature of the reaction .....
the patent doesn't say .

Attachment: US808095 Manganese Ammonium Sulfate Electrolytic Oxidation.pdf (247kB)
This file has been downloaded 2023 times


Hilski - 20-10-2006 at 04:52

Quote:
I am wondering if the oxidation product of ethanol
is acetaldehyde or acetic acid , or if either is possible

Supposedly, all one needs to do is adjust the acid concentration and/or the temperature of the oxidant solution to determine whether the aldehyde or acid is obtained as the final product. The patent does say that methanol will yield either formaldehyde or formic acid depending on the parameters I just mentioned. I don't see why the same won't hold true for ethanol ----> acetaldehyde/acetic acid. This will definately be something I will experiment with in the future, so if you or someone else doesn't try it first, I'll let you know how it goes.



What I like about this stuff is that it never gets used up. One could make an unlimited amount of whatever, and never have to worry about buying anything other than reactant. (Toluene, ethanol etc)

*Edited for spelling

[Edited on 20-10-2006 by Hilski]

Rosco Bodine - 20-10-2006 at 08:16

This electrolytically regenerated reagent scheme should
also be applicable to other polyvalent salts which may
work as well , or better in specific reactions for particular
compounds being oxidized .

For example the ferrous / ferric sulfates , also form an ammonium alum and the precursors would be cheap as for the manganese sulfate .

More expensive would be the nickel salts .

Some of these oxidants may be regenerable intermediates in conventional electrolysis which
is done continuously in the presence of the material
being oxidized , but using the oxidant in much smaller
amount in a continuous process , and in such a scheme
the more expensive polyvalent salts would become more practical and could have certain advantages in particular reactions .


Absent the ammonium alum formation which has usefulness in this specific oxidation scheme ....there may be alternate schemes where the polyvalent metal salt is useful alone .

And it also seems likely that the scheme could be reversed in order to arrange for reductions to be performed , probably requiring a different cell design
or different conditions for recycling the reagent .

The metal "ous" sulfate is a reducing agent with
regards to reactivity with certain other materials ,
while the metal " ic " sulfate is an oxidizer with respect to certain other materials . So an electrolytic cell
of appropriate design and electrode materials should
be able to accomplish many different reductions as
well as oxidations , by variations of this scheme .

Another interesting point of information is that the
" ous / ic " oxidation state of the polyvalent reagent can
be transitioned purely by chemical means and no electrolysis
is absolutely necessary . For example if one has a material
which needs to be reduced , and a second different material which needs to be oxidized , and both of these materials can be efficiently reduced or oxidized as desired by the
corresponding oxidation state of the polyvalent sulfate ....
then the reagent can be transitioned back and forth simply
by performing the reduction and oxidation reactions in the
correct sequence , using the reduction or oxidation of one
material to transition the reagent oxidation state , making it ready for use in the alternate reaction for the other material .

Ideally if one had two such precursors , one needing reduction and the other needing oxidation , the polyvalent reagent could cycled and recycled with most or all of the
electrical requirement eliminated . Where a scenario involving a huge scaleup is involved , this would increase the efficiency and economy greatly .

[Edited on 20-10-2006 by Rosco Bodine]

Hilski - 20-10-2006 at 10:20

Quote:
This electrolytically regenerated reagent scheme should
also be applicable to other polyvalent salts which may
work as well , or better in specific reactions for particular
compounds being oxidized .

For example the ferrous / ferric sulfates , also form an ammonium alum and the precursors would be cheap as for the manganese sulfate .

More expensive would be the nickel salts .

Some of these oxidants may be regenerable intermediates in conventional electrolysis which
is done continuously in the presence of the material
being oxidized , but using the oxidant in much smaller
amount in a continuous process , and in such a scheme
the more expensive polyvalent salts would become more practical and could have certain advantages in particular reactions .


There is definately a lot of room for experimentation here. Especially since there isn't much published research (that I have been able to find) on the production and use of these types of compounds for this purpose. The good thing about it is, that the processes are fairly simple and straightforward. And as you stated, with the exception of nickel or precious metal salts, the reagents are cheap and readily available for anyone who chooses to experiment with them.

Quote:
Absent the ammonium alum formation which has usefulness in this specific oxidation scheme ....there may be alternate schemes where the polyvalent metal salt is useful alone .


I'm sure you are right about this. But it also has me wondering whether an alum with the manganic oxidation state can be produced by using the same electrolytic process, only carried out in a divided cell as opposed to an undivided one? Would there even be any advantages to using a compound such as this, over a non-alum "ic" salt?

Rosco Bodine - 20-10-2006 at 11:07

I take it you mean the " ous " reduced salt as opposed to the " ic " oxidized salt ? Anyway the advantage of the non-alum form is greatly increased solubility over the alum form ,
much better contact between reactants in solution or emulsion , and a faster reaction rate ....a much more active system , which of course could be undesirable if a gentler reaction is required , in which case the alum would be preferable .

IIRC the alum form is used for commercial reasons
because the alum is less soluble , more easily purified
and crystallized , and in storage the alum of the " ous " form is less subject to atmospheric oxidation . But the
business part of the double salt is the polyvalent metallic
salt , so the alum may be considered in most cases just
a more convenient form of the polyvalent salt having
better handling and storage properties .

It may require a divided cell or simply different cathode material or a different pH for the reduction to occur preferentially to the oxidation . Electrochemistry is
extremely specific and even peculiar in how certain
combinations of conditions behave . It is a science
unto itself . I won't pretend to know the intricacies
of the theory on this even though I have done a few
electrolytic reactions , basically copying others work
or making slight variations .....so I very willingly defer to
others expertise or knowledge / suggestions concerning
suggested conditions and cell designs and electrode materials which would have the likely configuration needed for the reverse reaction , where the metal " ic "
salt would be favorably * reduced * to the " ous " salt .

[Edited on 20-10-2006 by Rosco Bodine]

Hilski - 20-10-2006 at 13:25

Quote:
I take it you mean the " ous " reduced salt as opposed to the " ic " oxidized salt ?


Doh. I need to proof read myself better. Yes that is what I was intending to write, instead of what I actually did write. Thanks for the correction.

And thanks for the lesson on the difference between alums and non alum polyvalent salts. This is all very interesting to me, and gives me something to experiment with that could actually turn out to be useful.

Hilski - 20-10-2006 at 20:08

I almost forgot. Here are a couple of pictures of the cell after about running at 4 amps for 22+ hours. The mixture has settled some, making it appear as though there are 2 phases. It was VERY thick at that point, and the magnetic stirrer was having a very hard time. I need to get some smaller stir bars for this, but ultimately I will rig up an overhead stirrer with either glass or nickle rods. But, for now I will be stirring by hand for the actual oxidation reaction to make sure there is complete mixing of the oxidant with the toluene (or ethanol, methanol, glycerine, or whatever)






Rosco Bodine - 20-10-2006 at 20:43

IIRC pure benzaldehyde is actually unstable and is subject to oxidation and polymerization unless it is inhibited with a trace of hydroquinone or something similar as an antioxidant , and kept sealed from exposure to air . A bottle of freshly made benzaldehyde can set up solid from a self-catalyzed polymerization probably via a peroxide of benzaldehyde which forms from exposure to the oxygen of the air acting as the catalyst . Keeping some vapor pressure of a volatile
solvent residue or deliberately added impurity like benzene or toluene will help the storage stability
compared to the absolutely pure benzaldehyde .

Hilski - 21-10-2006 at 15:04

Quote:
Keeping some vapor pressure of a volatile
solvent residue or deliberately added impurity like benzene or toluene will help the storage stability
compared to the absolutely pure benzaldehyde .


I plan on doing 5 - 10 oxidations, using toluene or DMC for extraction after each. I should then have 300 - 400ml benzaldehyde in about twice as much solvent. After that, I will wash all the extracts and the it will be stored in the solvent used to extract it until I need to use it for something, at which time I will distill it to get the pure benzaldehyde.

Is there any reason that storing it this way would degrade the benzaldehyde? Or should I do one distillation to get the crude product and store it that way?

I had planned on doing it this way anyway, mostly because I don't have all the glassware I need for fractional distillation at this point.

[Edited on 21-10-2006 by Hilski]

Rosco Bodine - 21-10-2006 at 15:59

You can store the pure benzaldehyde in a tightly capped bottle indefinitely , just add per liter , about .1-.2 gram of hydroquinone in 10-20ml of benzene or toluene as a stabilizer and preservative if you are going to keep it for more than a week . An amber bottle and cool place are
a good idea too for long storage .

Hilski - 31-10-2006 at 09:32

Just an update: This stuff does indeed work for oxidizing toluene. Sorry, but I don't have any pictures to show the reaction at the moment. The mixture was stirred by hand, which is obviously less than ideal. But, I am currently working on an overhead stirrer which should make things a lot more 'automated'.

I'll post more pics etc. when I have them.

*****
And one more thing I forgot. ~30 inches of #16 stranded copper wire (which I am using) gives about .01 ohms of resistance. I tapped the + wire to the cell at two points 30 inches apart and tested the voltage drop, and the meter read .055 volts. I double checked this reading with an i1010 Fluke AC/DC amp clamp, and it was a nearly perfect match (after moving the decimal point in the DMM reading) at about 5.58 amps. I have tested this several more times with different voltages and currents and it is right on every time.



[Edited on 31-10-2006 by Hilski]

Rosco Bodine - 31-10-2006 at 10:55

An interesting potential for this cell relates to the possible
usefulness for production of pentaerythritol using denatured
ethanol and methanol as feedstocks .

If acetaldehyde and formaldehyde are formed under the same condition by oxidation of their respective alcohols
by the manganic oxidizer , then it may be trivial to make
the mixture of acetaldehyde and formaldehyde solution
needed for pentaerythritol . What I am thinking is that
a mixture of methanol and ethanol in a 4:1 molar ratio
could simply be stirred into the manganic reagent , and
then the liquid containing the acetaldehyde and formaldehyde products filtered , treated with sodium carbonate to precipitate any manganous impurity as carbonate and filtered again . To the mixture is then
added calcium hydroxide and it is stirred and gently heated
for two to three hours , treated with dilute sulfuric acid
to precipitate calcium sulfate from the calcium formate
byproduct , the mixture filtered again and evaporated
to obtain the pentaerythritol .

If this should work as planned , it is the sort of reaction
which could be performed using five gallon plastic buckets
as reaction vessels , and would be a technically simple and cheap source for pentaerythritol .

Just a few pictures

Hilski - 6-11-2006 at 07:18

Here is how the mixture looks after the reaction is complete:




If you look closely at picture #2, you can see that there is a pretty nasty emulsion. I'm not sure what this is from, but I'm hoping I will be able to get rid of it by filtering the top layer.
Cycloknight wasn't kidding when he said the whole house smells like benzaldehyde, because it damn sure does.

***Addon: After adding about 150ml toluene to the mix and stirring a bit with a glass rod, the emulsion cleared up for the most part and I was able to extract the benzaldehyde without too much problem. I know there is at least a little bit of benzoic acid left, because after I removed the benzaldehyde, I could smell it in the reaction vessel. I'm guessing that its from leftover toluene from the last reaction/extraction. I doesn't appear to have caused any problems at all, so I'm not too worried about it. I use 4 -5 amps when regenerating the oxidant, and it still regenerates in about the same amount of time as it did originally.

Here is the really ghetto mechanical stirrer I used to stir the mixture. It worked very well, and I will continue to use it for this purpose.



-Hilski

[Edited on 6-11-2006 by Hilski]

Rosco Bodine - 6-11-2006 at 12:20

That stirrer sure brings back some memories for me
because it is almost identical to the first one I ever
cobbled together . And if the reaction mixture is not
particulary corrosive , what I used for a stirring shaft
was the back half or so of an old fiberglass arrow ,
using the three plastic vanes for the agitator .
The hollow arrow shaft snugly fit the shaft on the
C-frame shaded pole motor and was secured with
a drop of glue . I didn't use the second bearing that
you have there for steadying the stirring shaft itself
away from the motor , but that is a good idea especially
for higher speeds and thicker mixtures which could
cause the stirrer shaft to flex and wobble without
the extra journal there to stabilize it .

Possibly another (easy) oxidizer for toluene to benzaldehyde?

Hilski - 23-1-2007 at 15:21

Here is an interesting patent that utilizes a manganic oxidizer made from oxidizing a mixture of manganese sulfate and either potassium or sodium sulfate electrolytically in an undivided cell in 8% H2SO4. Temps for the reaction are supposed to run no higher than 30C, and it's told to produce a high purity product with a minimum of tars and/or carboxylic acids. The example in the patent is oxidation of methoxytoluene to anisaldehyde, but according to the author, the oxidizer should work to produce other aromatic aldehydes and acids. I find the fact that only 8% H2SO4 is needed to be particulary appealing, along with the fact that no heating should be needed. (Large batches in plastic buckets should be feasible)
As a comparison, manganous ammonium alum needs 60% H2SO4 and 50C to react, although it does work well, and it is prepared in exactly the same way as the oxidizer in this patent.

http://www.freepatentsonline.com/3985809.html

Sorry if this has already been posted before.

Hilski

[Edited on 23-1-2007 by Hilski]

Rosco Bodine - 23-1-2007 at 18:07

Looks like another good method .

I would bet that this same manganic oxidation reagent
would also be an effective reduction reagent when it
is in the reduced "ous" manganous state .

For example it possibly would reduce a phenylnitropropylene
to a phenylisopropylamine :D ....if a reaction like that might possibly have similar usefulness :P

And it would possibly also reduce nitrobenzene to aniline ,
picric acid to picramic acid .....ect.

Of course the electrical regeneration to the reduced state
would require a porous cell around the anode and plain acid
anolyte .

Anyway , here is the patent you mention above attached

Attachment: US3985809 manganic sulfate oxidizer for production of aldehydes.pdf (341kB)
This file has been downloaded 2540 times


chemoleo - 23-1-2007 at 18:32

Please remind me, wasn't there a thread recently with excellent pictures on some manganese oxide derivative, which was used for oxidative purposes?
I'll happily delete this post once it's clarified.

Rosco Bodine - 23-1-2007 at 18:44

Quote:
Originally posted by chemoleo
Please remind me, wasn't there a thread recently with excellent pictures on some manganese oxide derivative, which was used for oxidative purposes?
I'll happily delete this post once it's clarified.



Yeah , the other manganic alum reducing reagent thread is

https://sciencemadness.org/talk/viewthread.php?tid=6882#pid7...

maybe even the threads should merge ?

chemoleo - 23-1-2007 at 18:57

Hilski, thanks first of all for your excellent contributions.
I merged this thread with the previous one on this matter.
Please send me a U2U if you aren't happy with the title.
I'd like to keep similar threads together because it will make searching for it so much easier, and it'll keep the forum more condensed.

PS Perhaps the info/pics on ammonium manganic alum should be written together for the pre-puplication forum?

Rosco Bodine - 25-1-2007 at 11:54

Just a bit of additional information which could be relevant and useful in strategies for original preparation of a manganic alum oxidation reagent , for adjustment / correction of the initial oxidation state , as well as an alternative chemical means for regeneration of
the spent reagent .

Potassium permanganate can be used to oxidize the
other manganous compounds present to the manganic
form . Of course by using additional manganate in the
form of permanganate , your quantity of regenerated
reagent will grow and accumulate . But this would provide a quick means of regeneration , in the case where time may be an issue and it is inconvenient to wait for perhaps days for an electrolytic regeneration to complete .

The proportions would have to be worked out , as applies
to the manganous alums regeneration , and the additional materials of ammonium or potassium donors needed to maintain the alum composition for the reagent ,
producible by such a chemical regeneration via permanganate .

Attached is a patent which shows how this proceeds
in the preparation / regeneration of a similar manganic
oxidation reagent . This may be of interest and adaptable to the alums as well .

Attachment: US837777 Usefulness of Permanganate for regenerating Manganous to Manganic reagent.pdf (236kB)
This file has been downloaded 1187 times


Hilski - 25-1-2007 at 20:22

Thanks for that post, Roscoe.
I had been thinking about ways one could go about regenerating the oxidizer in a more expedient manner. KMNO4 had came to mind, but I have no idea how to go about calculating proper proportions. I also wondered how adding another element (K) would effect the behavior of the oxidizer both during the reaction with toluene, and during electrolytic regeneration.
Perhaps a small specific amount of the permanganate could be used in addition to the electrolytic process to say, cut regeneration times in half. I also wonder if something like ammonium or potassium persulfate could be used to achieve the same results?
I guess seeing as how I already have about 2 liters of the alum prepared and fully functional, that I should be the one answering these questions instead of asking them. But what fun is that? :D

Rosco Bodine - 25-1-2007 at 21:11

I have wondered if even hydrogen peroxide , or
perhaps acetone peroxide , or MEKP could be used for regeneration of the manganous alum to the manganic alum . It sure would save hours of electrolysis to be able to stir in the stuff and boil off the inert byproduct .

Yeah I would think that the persulfates should also
work the same way with their respective potassium
or ammonium alums . You would have to keep an
inventory of your reaction mixture content with the
desired proportions of the respective ions , and
work out the amounts of what you add during regeneration , so that you keep the composition
of the reagent consistent with its proper stoichiometric proportions for the desired alum formation .

I haven't done the math , so I am not sure how the
numbers crunch on this idea ....it just looked like an
interesting possibility , particularly on the initial formulation of the potassium alum ....it looked like
that making the reagent in the manganic form ready
for use without electrolyzing for the first run , could
be done by forming some of the potassium sulfate and
manganic sulfate content from potassium permanganate
plus H2SO4 , eliminating or reducing the need for the
electrolysis . I think there is even some surplus oxidizing
capability there for oxidizing the bulk manganous sulfate
to the manganic state . So it appeared to be a way of shortcutting the workup for the initial alum ....if nothing else . What would be the economy of this as compared
with just using electrolysis alone on the manganous sulfate alum is another matter which could make it not worthwhile .

BTW I found an analogous process is possible for chromium also , as is used here for manganese compounds ....but of course the chromium reagent
would be a way more expensive method .

I am still inclined to believe that the ferrous salt and
its alums would also work for this process , and also
other polyvalent metal salts .....although my interest
would be focused on the cheap ones like manganese
and iron :D

Organikum - 26-1-2007 at 01:16

Vanadium is told to work with H2O2 even in situ with no need for transferring something from one cell to another. The article should be in the toluene to b-dehyde thread.

/ORG

Hilski - 26-1-2007 at 12:54

I guess all in all, passing current through a non-divided cell for 24 hours isn't really all that bad. Unless one needs to regenerate the oxidizer like RIGHT NOW, then there just isn't a more economical way to do so. Here are a couple of pictures of the electrolytic cell regenerating the MAA recently. 4 amps have been passing through the cell for about 10-12 hours in the photos. I think it was still needing about 10-12 more to be fully regenerated. You can see the large quantity of yellow precipitate (MAS) still sitting at the bottom of the jar. This is because I have crammed so much oxidizer into this jar with not enough acid that the mag stirrer has a hard time keeping it all in suspension. A little poking around with a glass rod gets everything mixed up again though, so it really isn't a problem. If you look at the second picture, you can see a brown substance floating on top, in the center of the jar. This is mostly benzoic acid floating on top of the saturated acid solution. It is due to left over toluene and/or benzaldehyde from a previous reaction. It's hard to get every last little bit of organic material out of the mix, but it really doesn't affect anything. It just smells REALLY strong once too much builds up in the cell.





Quote:
Vanadium is told to work with H2O2 even in situ with no need for transferring something from one cell to another. The article should be in the toluene to b-dehyde thread.

/ORG

I recall seeing something about venadium sulfate (I think) being used for this purpose with H2O2. Not having to transfer anything back and forth is always a plus. That's one reason I like this MAA oxidizer so much. Once the oxidation is complete, and the aldehyde has been extracted, just drop in two electrodes and turn on the PSU. No membrane, or other special apparatus required.

[Edited on 26-1-2007 by Hilski]

Rosco Bodine - 26-1-2007 at 13:59

@Organikum

As I think I understand this process so far .....

That patent which I posted above , US837777 , is directly related .....but superior to :D
the " neograviton method " which you posted early on the Benzaldehyde from Toluene thread .

The earlier patent US780404 , by the same inventor Max Bazlen , is a patent which you listed among others in the other thread . The yield of benzaldehyde from toluene by oxidation with manganese (IV) sulfate , manganese persulfate , is reported as nearly quantitative , which generally means >95% , with no byproduct benzoic acid
from any overoxidation .

The subsequent patent US837777 simply advances the
same process using Potassium Permanganate for generating the dark brown colored
Manganese ( IV ) "per"sulfate , Mg(SO4)2 .

The (IV) valency variant is a yet still more active oxidizer than the also useful *green* colored Manganese ( III ) analogue , or Mangan " ic " Sulfate , Mg2(SO4)3 .

Reduction of the (IV) and/or the (III) compound during oxidation of toluene to benzaldehyde produces
the fully reduced Manganese ( II ) Mangan " ous " Sulfate , MgSO4 as the spent reagent , which is regenerated ,
by electrolytic oxidation , stepped back up to its (III) and
further to its original (IV) oxidation state for reuse .

CRC gives a range of colorations from pink through red for the ( II ) manganous sulfates depending upon the level of hydration , including the alum with ammonium sulfate , which BTW is a hexahydrate .

Oxidation to the ( III ) state is described as producing a
green color for the pure manganese(III) or mangan" ic "
sulfate .

I am thinking that the mixing of the green in with the incompletely regenerated red is what gives the darkening towards brown .....it could even pass through a tar black color first and then lighten slightly to more distinct brown for the final transition to the brown coloration of the manganese ( IV ) sulfate , manganese persulfate . IIRC the persulfate is only quasi-stable and only in a range of fairly concentrated H2SO4 .

Anyway I suppose the advantage realized for any
"hyper-regeneration" being applied chemically using permanganate , or perhaps peroxide if possible , or by electrolysis of the manganic sulfate to the even higher (IV) persulfate is that an additional amount of toluene may be converted to benzaldehyde by the same quantity of Manganese ion present , if its oxidation state is at maximum , which is the persulfate .

I haven't gone back and checked the details on the
regenerations by electrolysis which have been pictured
and described , to see if anyone allowed the regeneration
to keep running longer than required to produce the pink
or red material , even where the alum double salt is being
used ......to see if it will continue to oxidize further to the
solid brown color as would indicate the persulfate .

Edit: I just went back and reviewed this thread , and it looks like the ammonium alum manganese sulfate follows the US808095 patent description nicely , with regards to color
changes , and other details . Whether or not the yield or
the volume efficiency for this method could be improved by running the regeneration electrolysis longer , might be a worthwhile experiment . It might be necessary to have some
additional amount of manganese sulfate present which is
beyond that amount which will form the alum in combination with the ammonium sulfate present . The alum itself may be limiting with respect to formation of the persulfate .....and
of course the acid concentration would have to be within the allowable range also in order to permit its formation from
any added time of electrolysis . Color change and gas evolution from the anode would tell the tale on whether
this effort is worthwhile , along with of course comparison of the yield and how much toluene may be reacted with a given quantity of the regenerated oxidation reagent .

But the typical yield for the alum Mn(III) process is 80% , compared with >95% reported for the non-alum Mn(IV)
process , which should also have a better volume yield ,
although it would depend on the volume of acid ( I haven't checked this to be sure ) . So it would seem that using the
Bazlen patent method based on the Mn (IV) sulfate would be superior .

Quantitative yield is as good as it gets :D

[Edited on 26-1-2007 by Rosco Bodine]

Hilski - 26-1-2007 at 15:10

Quote:
to see if anyone allowed the regeneration
to keep running longer than required to produce the pink
or red material , even where the alum double salt is being
used ......to see if it will continue to oxidize further to the
solid brown color as would indicate the persulfate .

I have let it go longer than required, and a brownish solution does result, although I didn't know why the brown color was present. I haven't distilled the the aldehyde from the reaction with the brown(ish) oxidizer yet, so I don't know how it might have affected yields.

Rosco Bodine - 26-1-2007 at 15:14

The brown material is precisely what you want for maximum yield . And you can use more toluene
also .....I haven't done the math , but offhand
maybe 50% more toluene per reaction with the same amount of " hyper" regenerated brown reagent ....

I'll have to check the math on this to be certain .

Attached is the original Bazlen patent US780404
concerning near quantitative yield of benzaldehyde from toluene via oxidation by manganese (IV) sulfate ,
manganese persulfate , which may be regenerated
by electrolysis .


[Edited on 27-1-2007 by Rosco Bodine]

Attachment: US780404 Manganic(IV) Sulfate electrolysis regenerated.pdf (173kB)
This file has been downloaded 1407 times


Some experiments

Hilski - 31-1-2007 at 17:41

As I mentioned in another thread, I would like to set up a functional electrolytic cell that would produce Manganese(IV) sulfate aka (MnSO4)2 from MnSO4 in 55% H2SO4. I loosely followed the guidelines from This Patent to know what reagents/ratios etc to use. Also, one of my goals was to find out whether Tyvek (a commonly available material manufactured by Dupont) would make a sufficient membrane for a divided electrolytic cell. The following are some details and a few pictures of what I've come up with so far.

Chemicals and materials:

Anolyte:

MnSO4 - 50 grams disolved in 504 grams H2SO4 (55%)


Catholyte:

55% H2SO4

Cell construction:

Cathode compartment: A small HDPE container with holes drilled in the sides. Inside the HDPE container is placed a Tyvek bag which acts as a microporous membrane to allow for ion echange. The cathode is 0.5 inches wide X 6 inches long and is made of lead sheet. The compartment is filled with 55% H2SO4 and placed inside the glass beaker which will serve as the anode compartment.

Anode compartment: A 600ml Kimax beaker was used as the anode compartment. A lead anode, 1 inch wide X 6 inches long was placed in the compartment along with a solution of 50 grams MnSO4 dissolved in 504 grams 55%H2SO4. This compartment is where (hopefully) a solution of Manganese(IV) sulfate will be produced electrolytically.

The empty cell:



The cell after the pink MnSO4/H2SO4 solution has been added:



I hooked up the computer power supply I used for this project, and set the voltage to 5 volts. I chose this voltage because it's what I had used in the past to regenerate the manganous ammonium alum, and it has worked well. I switched on the power, and a dark red color, similar to that of the aforementioned alum, poured off of the anode.





I didn't use any kind of magnetic or mechanical stirring for this project, but if I decide to scale it up, I will definately add that to make the process a lot faster and more efficient. For this small cell I just used a glass rod to stir things up occasionally.

Here is the cell after about 10 minutes:



And again at about 1 hour:



After about an hour and a half, the solution in the cell had turned a very dark, opaque burgundy color and didn't appear to be making any further progress. I checked the current and saw that it was only pulling about 1.5 amps, and was generating almost no heat at all. So at that point I decided to switch from 5 volts to 12 to see if that made a difference. Immediately upon switching on the 12 volt power, a brown color was generated at the anode, along with more vigorous oxygen production. The hydrogen evolution at the cathode was enough that it started splashing the catholyte out of the compartment. I also noticed the the cell started to heat up a good bit (as expected) and the current draw shot up to over 6 amps.

Note the brown color starting to form around the anode



At that point, I decided to place a 50 watt DC lightbulb in series with the cell to help control the current. As of now, the cell is pulling about 3 amps, and gas evolution is not nearly as bad as it was before. There doesn't seem to be nearly as much heat produced as before either.



The cell has been running for a couple of hours now, and as soon as there has been significant change (or lack thereof) I will update the post with further results. At this point, I am optimistic that Tyvek does indeed make for a suitable membrane for divided electrolytic cells.

UPDATE:

At about the 3 hour mark I went to check on the cell and found what appears to be a successful experiment. Even with no stirring, the contents of the anode compartment were a deep brown, almost black color. The heat in the cell was a lot lower than I expected, less than 40C for sure. Now I just need to figure out exactly how long to this thing go to get 100% conversion to the persulfate.







Here are some links to high res copies of the last 3 images. The small ones are a little hard to see.

http://www.sciencemadness.org/scipics//Hilski/persulfate_cel...
http://www.sciencemadness.org/scipics//Hilski/persulfate_cel...
http://www.sciencemadness.org/scipics//Hilski/persulfate_cel...

Questions, comments or suggestions about any of this are welcome. Ideas for scaling up would be great as well.

-Hilski









[Edited on 1-2-2007 by Hilski]

Rosco Bodine - 31-1-2007 at 22:54

That flouropolymer braid fly line would probably be good for lacing the tyvek snugly around the top of a wide mouth bottle ....but there is another idea that may be better .

Get one of those clear polyethylene food storage cylinders that has the snap on poly lid and test it to
see if you put a sheet of tyvek across the top and
depress some slack in the middle of it , if the lid
will still snap down into place , sandwiching the
tyvek in the seal groove in the lid and securing it
in place . The tyvek will probably wrinkle a bit ,
but if the sealing groove in the lid will seat over the
tyvek , and secure it in place .....then you can use
the lid as a retaining ring to secure the tyvek liner
bag to the rim of the container . You can take a
razor knife or boxcutter and cut a big hole in the
lid , but leave the outside perimeter intact for use
as a retaining ring . Some dots of hotmelt glue might work for securing the tyvek to the inside of the storage cylinder . Leave slack in the tyvek where it loops
over the sealing edge , so that when you seat the ring it doesn't pull the liner taught and pop it loose . If you
make the lead cathode cylindrical , maybe 1/4 to 1/3
the diameter of the inside of the catholyte compartment would be about right , bubbles will have room to escape
without the foam climbing the inside of the tyvek . A plastic trivet like they use in the bottom of those lettuce keepers would make a good bottom spacer , or just use
another smaller diameter plastic lid on which to rest the bottom of the cathode , bore holes in it too if you like .

Heavy solid copper wire could probably be used to make
a tripod suspension bridle for the cathode and catholyte
compartment , hanging it from a suspension harness as if it were an ornamental hanging basket planter . I would
form a ring of the solid wire all the way around the plastic
rim of the container underneath that stiffening ridge and
make attachment loops to that wire ring , rather than
just going through small holes made in the plastic .....
because that plastic softens and weakens when it gets hot , and the attachment points might stretch and tear loose .

You could do a similar thing with some polyethylene
braided cord if you can find some small enough diameter
to work with .

Another strategy is heat welding .....tyvek could stick very
well to the rim of a polyethylene container which was gently
melted at the edge , just to the point of inflammation then quickly blown out like a candle , the tyvek pressed quickly
and firmly against the melt as it solidifies . Or just the hot
tip of a soldering iron might be used to " spot weld " the
tyvek at points like a dotted line , to " stitch " it to the
rim of the container . Poking holes through with a hot
nail and actually stitching it with a sailmakers needle
and the fluropolymer line is another option for lashing it down
to the rim of the cathode compartment . Slippery stuff to
handle , but there's a few ways of managing it .

[Edited on 1-2-2007 by Rosco Bodine]

Hilski - 1-2-2007 at 05:58

Those are all good ideas, and the type of thing I would need to address if scaling up. But one thing I would like to try is eliminating the plastic cathode container altogether. I want to make sort of a sleeve to put the cathode in that is tall enough to cover most the cathode and that has enough volume to hold enough catholyte to keep the cathode mostly covered all the time. Then, after filling the pouch with the 55% H2SO4 solution, I would use a zip tie to secure the top of the pouch around the top of the cathode to attach the tyvek in place directly to the cathode.
If I do end up using a plastic container, another thing I had though about was to use one of those tall, flat HDPE bottles that hand lotions and shower gel etc come in. The bottles are tall enough to be able line with tyvek and fit a large cathode in. But they are flat enough not to take up too much space when stood upright inside a larger glass or plastic container, in an arrangement similar to what I used in my small scale experiment.

Rosco Bodine - 1-2-2007 at 10:31

I definitely believe you should have some sort of rigid
compartment to keep everything positioned coaxially ,
so that hydrogen gas evolution from the cathode can
be managed smoothly , without " gurgling " and splashing .

Those red plastic press on cover " coffee cans " that
13 ounces of coffee is sold , are HDPE , 4" diameter
and about 5" tall might make a good compartment .

Something smaller in diameter and taller form cylinder
would be like a shampoo bottle . But the more narrow
and tall the container , the more of a geyser of liquid
will be carried up by the hydrogen bubbles , so your choice is going to set a limit on the current you can
manage .

Hilski - 1-2-2007 at 11:14

Quote:
Those red plastic press on cover " coffee cans " that
13 ounces of coffee is sold , are HDPE , 4" diameter
and about 5" tall might make a good compartment .

My thoughts exactly. I just salvaged one of those from the pantry. You are right about the bubbling and foaming being worse with a narrower container, but I want to be able to accomplish effective stirring within the anode compartment, which the cathode compartment will be sitting in. So unless I use a pretty large container for the anode compartment (like an 8 or 12 liter plastic bucket) then I need to keep the cathode compartment as small as I can. Using a plastic bucket for this would be great, except for the fact that I have no way of heating the contents of said bucket to 50C for the actual toluene oxidation. I could transfer the persulfate solution to the 2 liter glass jar like I use for the alum, and heat it the same way, but it would be nice not to have to transfer the solution around if I don't need to.

Thanks for the suggestions.

Hilski

Rosco Bodine - 1-2-2007 at 11:52

It would probably be better to put the tyvek around the *outside* of the perforated plastic form , as opposed to
having it inside like a liner bag . I just think that using
the rigid compartment as an inner form around which the
tyvek is stretched .....would simplify things greatly all around .....instead of " all inside " :D

This would definitely be a better arrangement with the smaller diameter cathode compartments .

Since this is the cathode also , a carbon rod or a piece of
copper or stainless rod or pipe might work fine , as no worries about oxidation and the hydrogen blanket on the surface should provide a lot of " anti-erosion " by " cathodic protection " .

[Edited on 1-2-2007 by Rosco Bodine]

Eclectic - 1-2-2007 at 16:03

Maybe one of these would work:



http://www.maycocolors.com/shapes/shapeDetails.cfm?shape=101...

Low fired porous ceramic bisque.

moxsnixs - 9-2-2007 at 03:58

In cycloknights original post he reacted MnO2 with HCL then with H2SO4 to get MnSO4. Anyone have a more detailed procedure?

jimmyboy - 9-2-2007 at 16:57

So overall - what is the best method - opinions? toluene/persulfate or reflux/oxidation of cinnamaldehyde - this is abit more involved but not too hard - some manganese - electrodes - power supply - it looks like you really need that stirbar - actually the supplies are easier to get - you have to look a little for cinnamon oil - i am just looking at it from a yield standpoint

Hilski - 9-2-2007 at 19:41

Unless huge quantities are needed, oxidation of toluene by manganese compounds would be the simplest, cheapest and most OTC in my opinion. I have never tried the cinnamon oil route, due to the fact that the stuff is actually more expensive than just buying benzaldehyde. A one gallon container will easily hold enough oxidizer/H2SO4 to produce 200g or more of benzaldehyde or benzoic acid, formaldehyde, formic acid, acetaldehyde (maybe) etc.

Quote:
In cycloknights original post he reacted MnO2 with HCL then with H2SO4 to get MnSO4. Anyone have a more detailed procedure?

Although making the MnSO4 from MnO2 isn't hard, it is much better to just order some fertilizer grade MnSO4 from the net. (NH4)SO4 can be found cheaply and in large quantities in any decent garden center or nursery supply type store.

Here is a thread that goes into detail on some different methods for making MnSO4

https://sciencemadness.org/talk/viewthread.php?tid=6777#pid7...

[Edited on 10-2-2007 by Hilski]

moxsnixs - 10-2-2007 at 04:07

OK thanks I added the Na2CO3 and got the MnCO3 but the final step with H2SO4 wasnt really hit on. Do you add it to the MnCO3 or filter the MnCO3 and add it to that?
I would buy it but not in the USA and to buy it here requires a 3hr drive one way, get a hotel room then try to find where its sold.So making it would save alot of BS .Thanks

not_important - 10-2-2007 at 04:24

Quote:
Originally posted by moxsnixs
OK thanks I added the Na2CO3 and got the MnCO3 but the final step with H2SO4 wasnt really hit on. Do you add it to the MnCO3 or filter the MnCO3 and add it to that?
I would buy it but not in the USA and to buy it here requires a 3hr drive one way, get a hotel room then try to find where its sold.So making it would save alot of BS .Thanks


Wash the MnCO3by to get rid of all the sodium salts, cover with water, and slowly add H2SO4 with stirring. Might be best to pre-dilution the sulfuric acid if its concentrated. You want enough water to dissolve the MnSO4 formed.

If you don't protect the wet MnCO3 from air it will be oxidised and turn dark. Not a big deal if you're make the Mn(3+)-alum from it.

If your MnO2 wasn't pure, or if your HCl has iron in it, it might be best to set aside about a tenth of the MnCO3. Add just enough acid to the remaining 90% to get it into solution, you'll not want an excess. Bring it to a gentle boil to drive off CO2, cool to it's just warm, add the remaining MnCO3 with stirring, continue to stir and bring to a boil for a minute or two. Then let it cool, filter off the precipitate which is MnCO3, iron carbonates/hydroxides, and hydrated oxides of manganese in 3+ and 4+ oxidation state.

moxsnixs - 10-2-2007 at 20:29

And then evap the MnSO4 solution to obtain the powder correct? because I did and got a faitly pink\white powder

Hilski - 10-2-2007 at 21:29

Quote:
And then evap the MnSO4 solution to obtain the powder correct? because I did and got a faitly pink\white powder

Correct.
Just FYI MnSO4 can be found on ebay. I think I remember seeing it sold for use on palm trees or something similar.

moxsnixs - 10-2-2007 at 21:56

yeah I know but Im in another country and customs dont like chems. And although it was interesting to do,if I was to need it again I'll drive the 3hrs one way and buy it.The trouble is the beautiful women there........can't leave till I get one :P

moxsnixs - 11-2-2007 at 20:46

Could ammonium chloride be used instead of the sulphate ?

12AX7 - 11-2-2007 at 22:06

Chloride is oxidized by Mn(III) or (IV).

moxsnixs - 12-2-2007 at 17:18

For the washing of the extracted benz. is it DH2O only? thanks

Hilski - 13-2-2007 at 09:27

Quote:
For the washing of the extracted benz. is it DH2O only? thanks

I have used both plain DH2O as well as a saturated NaHCO3 solution to wash the benzaldehyde/toluene mixture. This is done to remove as much acid and other byproducts as possible to reduce the amount of red tar produced during distillation. I haven't noticed much difference in the results between water or NaHCO3. Once I get the up-scaled MnSO4(IV) cell going, I think I may try using steam distillation to get the product out instead of doing extractions.

CycloKnight - 13-2-2007 at 15:31

Nice work Hilski, lots of potential with this manganous-alum route. Great pics BTW !
This thread sure brings back memories :D


[Edited on 13-2-2007 by CycloKnight]

Eclectic - 13-2-2007 at 21:00

Ok, I got a case of these, and they seem to be perfect for electrochemistry experiments. When filled with water, the outside gets wet and "sweats" on the lower half after about 5 minutes. The clay is an off white low iron type.




http://www.maycocolors.com/shapes/shapeDetails.cfm?shape=101...

Low fired porous ceramic bisque.

[Edited on 14-2-2007 by Eclectic]

Rosco Bodine - 13-2-2007 at 23:54

Hehehe ....looks like the mother of all porous cups :D

It should work very well too , *if* it holds up okay in the
pretty strong H2SO4 which is used for the electrolyte in this process .

Buying a case of them , hmmmm ...you must be confident .

Let's just hope it isn't a "case" of only .......
"hope springs eternal" and/or
advance christmas shopping for that perfect gift
for all your wino friends :D

I just had an idea that maybe a 100 ml mixing cylinder is placed inside , it would make a perfect water cooled cathode form for wrapping sheet lead around .....fitted with a two whole stopper with one long tube nearly to the bottom , and one short tube .....circulate cooling water through the cylinder , and you could really put some Amps through the thing . It would keep the electrode nicely centered too .

[Edited on 14-2-2007 by Rosco Bodine]

Eclectic - 14-2-2007 at 05:38

$42 for 6 including S&H. Google for the best price. They are made out of clay fired at cone 4, so they are just as chemically resistant as any other ceramic.
I was thinking they would also make a good substrate for lead dioxide. If you could get a good heavy plating, then the connection could be made through an electrolyte and electrode on the inside.

[Edited on 14-2-2007 by Eclectic]

Rosco Bodine - 14-2-2007 at 06:34

Outstanding . You could build a serious electrolytic cell
of coaxial electrode arrangement , having that component as the porous cup . This is great , it looks like the tyvek works for a divided cell , and this will probably work fine too

Hmmm. I'm thinking maybe as this progresses we should have a new specific topic category for Electrochemistry , for all related reactions and equipment , literature , ect.
I know it is included as a part of " Technochemistry " now
but that is sort of an ambiguous name ....
since *all* chemistry is actually pretty " technical " :P

What do you fellows think about this ?

[Edited on 14-2-2007 by Rosco Bodine]

Hilski - 14-2-2007 at 09:10

Quote:
Hmmm. I'm thinking maybe as this progresses we should have a new specific topic category for Electrochemistry , for all related reactions and equipment , literature , ect.

I'll second that. Electrochemistry is the branch I have always been the most interested in, so I would love to have it broken out into it's own forum.

And those wine coolers that Eclectic has look like they would work well for large cells. I have some large Pyrex graduated cylinders that would work for the electrode cooling jacket inside those containers. If you try to plate PbO2 onto those, I would definitely like to have some details on how that project goes.

Quote:
Nice work Hilski, lots of potential with this manganous-alum route. Great pics BTW !
This thread sure brings back memories

Thanks! I have to give you the credit though. I got the idea and most of the information I needed to do this (the alum) from your posts and pictures. I am anxious to try out the Mn persulfate on a similar scale, so I can make a yield comparison with the alum. I am working on a 3 liter divided cell which will have overhead stirring, so hopefully, I will have some aldehyde production results to post within the next week or two.

roamingnome - 21-2-2007 at 22:18

as i was saying
i have been swayed by the anode side of the force...young experimenter

but to clarify, in this divided cell only MnSO4 is used, with acid of course

the ammonium sulphate is not needed now to protect the charge.
the jedi arts are proliferating

Hilski - 22-2-2007 at 20:07

Quote:
but to clarify, in this divided cell only MnSO4 is used, with acid of course

the ammonium sulphate is not needed now to protect the charge.

Yep. That is correct.

Rosco Bodine - 22-2-2007 at 22:25

There is a smaller part that may work also
with a smaller scale cell .

This is 3.25"D X 4"H

http://www.maycocolors.com/shapes/shapeDetails.cfm?shape=203...

Also see this page for all the Mayco bisque .

http://www.ceramicssf.com/maycobisqueware.htm

About halfway down the page on the right is a " message bottle " and it looks like it might make a better substrate
for PbO2 and using an electrolyte inside for the connection .
I am not quite sure if in the case for this being used as the anode , that there may not be some gassing on the surface of the electrode in the interior electrolyte due to the slight
voltage drop ....or what electrolyte would be useful . It seems possible that some series voltage differences might
be problematic for making the connection internally via an
electrolyte . That whole idea is very interesting , but not guaranteed to work because of the unknowns ...I think :D

[Edited on 23-2-2007 by Rosco Bodine]

Hilski - 7-3-2007 at 17:53

Since the site got hacked, and was going to be down for a few days, I figured I should try something constructive in the meantime. So I decided to make some more Mn(SO4)2 aka. Manganese persulfate, or Mn(IV), and oxidize something with it. So I preceded to use a terracotta cylinder I picked up at a local craft type store and make a divided cell.
I dissolved 450 grams MnSO4 in about 2.5 liters of 55% H2SO4, and placed it in the anode compartment of the cell. I decided to use the inside of the terracotta vessel as the anode compartment, because it was just more convenient. I electrolyzed the mixture at 12 volts, 3 amps for about 24 hours, with current control being provided by a 50 watt DC light bulb in series with the cell. There was no heat created at all, so next time I will add another 50 watt light bulb in parallel with the other one to up the current to around 6 amps. I stirred the mixture vigorously 4 or 5 times during the 24 hours, and that seemed to be enough.
After the electrolysis was complete the mixture was kind if thick, and a very deep brown color. Most of the Mn(IV) was dissolved in the acid, but there was some black sludge at the bottom of the cell that needed to be stirred around in the acid so it could be poured out into the reaction vessel.
I used a 1 gallon glass jar for the reactor, which worked pretty well. To the reactor I added the 450 grams Mn(IV) in about 2.5 liters 55% H2SO4, along with about 210 ml toluene. I decided to try the reaction at room temperature, so a homemade overhead stirrer was set up and the mixture was stirred vigorously overnight (probably about 12 hours).
When I went back to check on things, I was somewhat surprised to see the the reaction had proceeded nicely and was totally complete. The deep brown solution had turned to a light brown sand color, and the the smell of benzaldehyde inside the vessel was almost overwhelming . The scent of toluene was not detectable at all, so I am guessing that nearly 100% conversion was achieved. This is not a good thing though, since some benzoic acid was produced, and was floating in a layer between the acid and the benzaldehyde as can be seen in the picture. This problem can be remedied by simply using a larger volume of toluene, so that there is an excess compared to the oxidizer.
Extraction was a little bit tedious because of the benzoic acid, but not too bad. I added another 100ml of toluene to mixture and decanted most of the product, and used a syringe to remove the remainder that I couldn't pour off. I haven't distilled anything yet, so I can't give any info on yield. But I am guessing that it will be pretty good, probably between 150 and 250 ml. I will update this post once I do have the final results.

I forgot to take a picture of the Mn(IV) before I started the reaction, but here it is after it has been reacting for a few hours. You can see some of the lighter colored brown oxidizer at the bottom that has already been reduced by the toluene. Remember that the mixture was totally black before I started.




Here is the reactor after the oxidation is complete. You can see the organic layer, which was a bright yellow color after it was extracted.



So far I am pleased with the results of this experiment, and if yields turn out to be pretty good, I would definitely consider this a good, relatively simple method for producing aldehydes.





[Edited on 8-3-2007 by Hilski]

AEM’s in general

roamingnome - 8-3-2007 at 11:01

What are the short comings of the tyvex membranes? I rather liked the concept…
Thinking about using Teflon tape PTFE maybye too.

Reading about the preparation of these membranes, it seems that net positive charges are impregnated in the polymer matrix as they make it….

Is the SiO2 or some component of the terra cotta pots giving it a net positive charge to repel cations?

Also what is the cathode reaction?

From: http://ocw.mit.edu/NR/rdonlyres/Materials-Science-and-Engine...

PbO2 + 4 H+ +2e- + SO42- ---> PbSO4 + H2O


Alittle benzoic acid doesn’t seem to cancel the show :) good stuff....

Hilski - 8-3-2007 at 16:11

Quote:
What are the short comings of the tyvex membranes? I rather liked the concept…
Thinking about using Teflon tape PTFE maybye too.

The electrolysis takes place fine using tyvek as a membrane, but it was rather akward to work with, when trying to cover a large vessel with it as I was. I also noticed that over time, the electrolyte (the anolyte in my case) inside the tyvek covered vessel would slowly seep through into the catholyte. This wasn't a huge deal since the catholyte was the same composition as the anolyte, and could simply be added back into the anode compartment. I chose to use terracotta because I found some 1.5 liter cylinders that did not have a hole in the bottom, at a local craft store for $3.00 and they worked very well for these purposes.

Quote:
Reading about the preparation of these membranes, it seems that net positive charges are impregnated in the polymer matrix as they make it….

Is the SiO2 or some component of the terracotta pots giving it a net positive charge to repel cations?

I have no idea. Roscoe might have to answer this one for ya.

Quote:
Also what is the cathode reaction?

As far as I know, the reduction reaction is the same as with a lead acid battery; I have seen this written as:
PbO2 + HSO4- + 3 H+ + 2 H+ --> PbSO4 + 2 H2O
The reaction you posted is probably correct as well.

The benzoic acid really isn't a big deal. Most of it can be strained out with a small SS strainer with no problem. It doesn't dissolve in the saturated H2SO4/Mn solution, and floats on top after the benzaldehyde has been removed. Organic material that remains in the cell supposedly lowers current efficiency during regeneration, but I could care less about that.



[Edited on 9-3-2007 by Hilski]

jimmyboy - 11-3-2007 at 12:56

so you placed a terra cotta container inside another larger one - and used the ceramic as the membrane (generating the alum inside the pot instead of outside) - the larger container catching most of the bubbling? this worked very well from what it looks like.. :)

Hilski - 11-3-2007 at 13:40

Quote:
so you placed a terra cotta container inside another larger one - and used the ceramic as the membrane (generating the alum inside the pot instead of outside) - the larger container catching most of the bubbling? this worked very well from what it looks like..


Correct. I used a big glass jar for the outer container, and set the clay vessel inside of it in a coaxial type arrangement. The glass jar could just have well have been a plastic bucket, or anything else that is acid resistant. Pretty much all the foaming took place in the outer container like you said, and presented no problem at all. I'll post a picture of the contraption as soon as I get a chance. I still need to get off my arse and distill the stuff so I can see what the actual yield was. But I'm crossing my fingers that it will be 70% to 80% or better. :cool:


*Edit* I just noticed that you were referring to this oxidizer as 'alum'. This stuff is just MnSO4 oxidized to Mn(SO4)2. No alum this time.

[Edited on 11-3-2007 by Hilski]

Coldfinger - 13-3-2007 at 12:40

would benzyl alcohol be able to be oxidised to the aldehyde also? or would it over-oxidise to benzoic acid? maybe less time in the reactor would do the trick? 6 hours as opposed 12 maybe?

roamingnome - 13-3-2007 at 14:30

In the PAIRED ELECTRO_OXIDATION
J. of applied Electrochemistry 1987...

jiin-jang etal.
proposed

C6H5CH3 + 4M(n+1)+ + H2O ------- C6H5CHO + 4M(n+) + 4H+

basically a 4 time molar amount of Mn is needed is what im gathering there.


they were also using 6M acid which would be very conveinent since i think this is very close to H2SO4 with a specific gravity of 1.256 which refreshes regular batterys.

Are you useing 55% acid to disolve more salt per given volume? Or other papers use differnet concentrations....

Hilski - 13-3-2007 at 23:31

Quote:
would benzyl alcohol be able to be oxidised to the aldehyde also? or would it over-oxidise to benzoic acid? maybe less time in the reactor would do the trick? 6 hours as opposed 12 maybe?

There is a procedure in US patent 780404 for oxidizing benzyl alcohol with Mn(IV) sulfate. It describes adding the oxidizer drop by drop to an emulsion of the alcohol with H2SO4 at 50C. One might be able to skip the tedious dropwise additions, and just put everything in the reactor at once if things are done at room temp or cooler. But I have no way of knowing for sure. Patent 0808095 mentions the oxidation of methanol with the alum to obtain methylal, formaldehyde, or formic acid but I don't know of a written procedure for oxidizing benzyl alcohol with the alum.

Quote:
C6H5CH3 + 4M(n+1)+ + H2O ------- C6H5CHO + 4M(n+) + 4H+

basically a 4 time molar amount of Mn is needed is what im gathering there.


they were also using 6M acid which would be very conveinent since i think this is very close to H2SO4 with a specific gravity of 1.256 which refreshes regular batterys.

Are you useing 55% acid to disolve more salt per given volume? Or other papers use differnet concentrations....

Judging by the reaction from the journal article you referenced, it would appear as though a 4:1 molar ration is needed. But none of the patents I have read that pertain to the oxidation of toluene with Mn salts use that ratio.
For example, the Paired Electro Oxidation article you mentioned, uses a pretty dilute anolyte at only 0.1m Mn2+. This equates to only 5.5g manganese per liter. (I assume the author meant 0.1M elemental Mn. If he meant MnSO4, then the ratios are even worse than I thought.) With the experiment described in the paper, the authors used 250ml of the 0.1m Mn3+ solution to oxidize 50ml (43g) toluene. The molar ratio in that experiment was 1:18.8 Mn:toluene. How that could have possibly worked is beyond me. Especially when in the same article the author says that a 4:1 ratio is needed. The Max Bazlen patent for production of Mn(SO4)2 and oxidation of toluene, recomends 600g Mn(IV):300g toluene, which equates to about a 1.4:1 molar ratio of Mn:toluene. The maganese-ammonium alum patent states that a ratio of nearly 12:1 w/w of oxidizer to toluene should be used, which is a ratio of roughly 5:1 Mn:toluene. The only thing that is common to all the patents and articles is that it is recommended to carry out the reaction at ~50C.
The 6M acid solution in the same paper you mentioned is about 46.5% by weight. The specific gravity would be ~1.36. The concentration I have used is about 8.1M or 55% by weight. The reason I used that particular concentration is because that is what is called for in both Bazlen patents, as well as in the alum patent. A 6M concentration may work fine, but I remember Cycloknight having issues with too much benzoic acid being produced if the acid was too weak, so I haven't tried anything weaker than 55%.

Coldfinger - 3-4-2007 at 23:32

Quote:
Originally posted by Hilski I haven't distilled anything yet, so I can't give any info on yield. But I am guessing that it will be pretty good, probably between 150 and 250 ml. I will update this post once I do have the final results.


is there any results regarding yields from your latest experiment?

raiden - 27-5-2009 at 19:53

Sorry for reviving such an old thread - but I'm going to give this a swing.

At the moment I am waiting for the MnSO4 to dissolve, am using the amounts in the original post. My mag stirrer seems to be having trouble getting any agitation at all, do you guys think it will dissolve by itself if left to its own devices? Going to be doing my run with a pool chlorinator MMO anode at around 12 amps.

manimal - 28-5-2009 at 12:12

No. Use overhead stirring (nickel or HDPE stirrer).

Sedit - 13-6-2009 at 15:36

I have recently been toying around with MnSO4 in the divided cell and getting some success with it. I was truely delighted when the material in the cell started to turn pink the proceeded to a deep purple and later a dark brown.

Simple alcohols seem to be converted into there aldahyde with ease which is nice because it will be great to have a steady source of formaldahyde and whatever other use I find for it.

A question that I have it how will this work with double bonds. Will materials such as styrene, or allyl benzene be converted to there aldahyde/carbonyl... or something else altogether? The reaction that I am really confused with is how substances such as trans-β-butylene will behave. Will it be converted to an aldahyde or will it place the carbonyl across the double bond?

It appears to be a gental yet effective oxidiser and im just looking at the different uses this may have as far as oxidising various compounds.

Also I wanted to know what it the storage of this oxidiser? If one wanted to extract the solid from the cell and store it how well will it hold up?

[Edited on 13-6-2009 by Sedit]

Impurities in Pottery-Grade MnO2

DyD - 31-7-2009 at 08:35

I'm planning to attempt to construct an MAA electrochemical oxidation cell to oxidize toluene and methanol. In my attempts at sourcing manganese compounds, I've constructed the following chart:
--------------------
Source Material PricePerPound Shipping Purity PercentMetal
CC MnCO3 5.80 15usd/15lb 100% 45-47%
CC MnO2 1.50-.8 15usd/15lb 96.8-81% 51-55%
1-5% Fe2O3
1-5% SiO2
1-7% Al2O3
0.2-2% BaO
EBay MnO2 3.99 5.99usd/1lb 99% 63%
--------------------
My question is this: As one can see, the cheapest source of manganese per mol is the manganese dioxide from "Columbus Clay" at 1.50usd for one pound, and eight dollars for ten pounds. However, the manganese dioxide is, at worst, 81% pure, with Fe2O3 and Al2O3 making up the most relevant impurities. Would either of these compounds or their respective sulfates negatively impact the performance of a MAA cell?
Thanks!

not_important - 1-8-2009 at 05:05

I wouldn't expect Al2O3 to dissolve in moderately strong H2SO4, certainly not SiO2, so those two should cause problems. Iron wouldn't seem to, but I have no information one way or another.

You can remove most of the iron by adding an excess of the crude MnO2 to the H2SO4 used to dissolve it, and boil it for awhile. Allow to cool, decant/filter from remaining solids, and split into two portions of 9/10 and 1/10 of the whole. Treat the 1/10 portion with NaOH to precipitate the manganese as Mn(OH)2, filter and wash that well, and then stir the Mn(OH)2 into the larger portion and boil for a few minutes. Cool and filter, most of the iron will be in the precipitate.

Alternatively, MnSO4 can be purified by evaporating to dryness, then heating to red heat. It melts at 700 C and decomposes around 840 C, while almost all other d block sulfates decompose below 700 C. Cool, dissolve in water, filter to remove the contaminating metals as oxides or basic sulfates. Only the alkali metals, Mg, and low levels of Ca will remain as contaminates in the MnSO4; it's not likely that those would be in the starting MnO2, and they except for Mg could be washed out of that before treatment with H2SO4.



felis - 22-8-2009 at 21:03

Is it possible to use tin electrodes instead of lead in this cell? Should be easier to handle since tin salts are not poisonous; in addition, pretty pure tin can be easily obtained OTC.

behemoth - 23-8-2009 at 08:08

This thread is a goldmine! In some posts there was mentioned that the mixture oxidizes Alkohols as well. Does anybody have some results for the oxidation of aliphatic alcohols to aldehydes?

[Edited on 23-8-2009 by behemoth]

My first go

Formula409 - 3-10-2009 at 06:39

Ok so I've finally gotten around to making my first cell. I decided to just make a cell the same size as Cycloknight's in order to get a feel for the reaction conditions, and then do a bucket-size scale up once I am confident. The cell runs at around 3.5A as measured by my digital multimeter. Naturally I won't be conducting any oxidations in the plastic container. The contents of the cell will be transferred into a glass jar which will allow it to be heated with toluene. Here are some pictures:

The computer power supply I am using. Capable of 40A at 5V:

The overall setup

The purple colour from the anode

The MMO electrodes from a pool chlorinator. They are rated at 25A over 25 years at an extremely low chloride concentration, would be very exciting to try out a chlorate cell with them! I do not know their composition, but the grey one is 1m long and flexible and the black one is 50cm long and rigid. In the cell the black electrode is the anode and the grey one is the cathode.

The lid of the container consists of a rigid outer section and a removable centre.



I turned off the electrolysis as I noticed that after a couple of hours no more purple was being produced. It was also noted that the colour of the liquid had not shifted to the expected dark purple after a period of time. I put this down to not allowing the MnSO4 time to dissolve, so the cell will be stirred overnight.

What follows is the overhead stirrer I made from a beater I obtained at a charity shop which sells second hand goods (called Lifeline in Australia), a pen and a coke bottle. I used a hacksaw to liberate the metal part of the part of the beater that "beats" and used hot glue to attach this to the casing of a cheap pen which was almost a perfect fit. A coke bottle was then cut up with a Stanley knife and a section of plastic hot glued to the pen lid to serve as an agitator. I originally planned to use the part of the pen which allows it to be hung from a pocket for this purpose but it was too streamlined. The only downside to this stirrer is that it obviously cannot be used in the more plastic-unfriendly solvents, so it will not be able to stir the mixture during the actual oxidation. But manual stirring should be sufficient for this.

The beater itself has three different speeds, but the slowest is usually more than enough. The stand is also very handy as it provides a solid base and eliminates shaking. I cut a hole in the lid which allows the shaft to pass through but stops any stray acid coming out. Overall pretty successful for $15.00!

The part that does all the work, quite sophisticated!



It can be seen that the stirrer achieves adequate agitation. All of the off-white precipitate is in suspension when it is running. Unlike magnetic stirring:

Stirrer

Magnetic




One odd thing I did notice was that upon addition of a couple of drops of 6% hydrogen peroxide solution, the electrolyte turns an odd yellow colour, which is replaced by the normal "strawberry milkshake" colour after several minutes of electrolysis. I found this by seeing if any purple alum would be made upon addition of oxidiser, as has been previously posted.



Would someone like to throw a guess at what this material is? I've tried mixing the peroxide with both acidified ammonium and manganese sulfate solutions by themselves and have not been able to replicate the precipitate. Obviously this leaves the conclusion that it is an oxidation product of the manganese ammonium sulfate, but wouldn't the oxidation product be the purple compound we are seeking?

Another thing I noticed was that when the 1m long electrode was used as the anode, the reaction would not proceed at all and only 0.04A would be recorded on the multimeter. I thought this was strange as that electrode would presumably be used as the anode in the chlorinator cell it was designed for as it is the longest electrode and therefore has more surface area. It doesn't affect the experiment, though as the electrodes have shown no wear in the past when they were used in similar electrolytes.

If anybody wants some more specific pictures, don't hesitate to ask!

Formula409


Formula409 - 6-10-2009 at 01:23

Hmm, I appear to be having an issue. My cell has been running for around 150 amp hours but it is still the dark purple colour. It's meant to be going for 120 in order for the oxidant to be fully generated! What the hell?

Could this be because I'm using an MMO anode? I thought any anode would work!

Formula409.

[Edited on 6-10-2009 by Formula409]

Hilski - 21-12-2009 at 15:53

Quote: Originally posted by DyD  
I'm planning to attempt to construct an MAA electrochemical oxidation cell to oxidize toluene and methanol. In my attempts at sourcing manganese compounds, I've constructed the following chart:
--------------------
Source Material PricePerPound Shipping Purity PercentMetal
CC MnCO3 5.80 15usd/15lb 100% 45-47%
CC MnO2 1.50-.8 15usd/15lb 96.8-81% 51-55%
1-5% Fe2O3
1-5% SiO2
1-7% Al2O3
0.2-2% BaO
EBay MnO2 3.99 5.99usd/1lb 99% 63%
--------------------
My question is this: As one can see, the cheapest source of manganese per mol is the manganese dioxide from "Columbus Clay" at 1.50usd for one pound, and eight dollars for ten pounds. However, the manganese dioxide is, at worst, 81% pure, with Fe2O3 and Al2O3 making up the most relevant impurities. Would either of these compounds or their respective sulfates negatively impact the performance of a MAA cell?
Thanks!

Hydroponic grade MnSO4 is the best thing to use in the cell. Its readily available online.

sabbath06 - 6-9-2010 at 06:35

I've been trying to duplicate this cell and my oxidizer is now a dark purple like in the pictures, but when I add the toluene, it doesn't turn back to the original color. I've left it to react for 12 hours with maximum magnetic stirring and it was the same color as when I first added the toluene, although I can smell some benzaldehyde, mostly I just smell toluene.

I was thinking maybe it has something to do with a rust inhibitor that may be in my hardware store toluene, should I have distilled it first?

EDIT
Just needed to heat it a little more

[Edited on 6-9-2010 by sabbath06]

Sedit - 6-9-2010 at 09:22

All the information posted on how to use this oxidiser is pretty much wrong. Everyone go back and read the patent more closely and see how to use this oxidiser. It is not a simple case of just dumping the oxidiser into toluene and waiting. The oxidiser is suppose to be added dropwise into warm toluene over a period of time.

I found this out after many dissapointing yeilds and decided to reread the patent as closely as possible to see what was amiss. I have not rerun this because I had since disposed of my oxidiser and did not have the time or money to rerun the reaction at the time. Toluene is also harder for me to get so I never reran it but it would be great to hear of the results from someone running it the way it should be done.

sabbath06 - 6-9-2010 at 09:28

Very interesting. Although I smelled almonds I'm pretty sure I got a crappy yield as well. Will try what you suggested and report back.

Sedit - 6-9-2010 at 09:34

Thats the same results I was getting. A drop left to evaporate would smell like BnO strongly but I was never to isolate a decent amount at all. I feel this is the reason for it in that it must be done as stated in the patent to get the yeilds given.

Hexagon - 7-9-2010 at 03:57

So, back to the manganese sulfate + ammonium sulfate cell?

Sedit - 7-9-2010 at 20:16

No the so called Mn(IV) patent is the one that spoke of dropwise addition. I can not comment due to no first handexperiance on the MAA.

Hexagon - 10-9-2010 at 16:57

I've got no direct experience with neither of them, but the original manganese alum procedure seems more simple, higher yielding (based on the substrate) and easier to up-schale...

roamingnome - 10-9-2010 at 17:34

@ sabbath06
"I've been trying to duplicate this cell and my oxidizer is now a dark purple like in the pictures, but when I add the toluene, it doesn't turn back to the original color. I've left it to react for 12 hours with maximum magnetic stirring and it was the same color as when I first added the toluene, although I can smell some benzaldehyde, mostly I just smell toluene. "


Are you reacting a 50 degrees Celsius?

I used 2 aquarium type test tube heaters to achieve this temp

Also the ratio of Bz made to toluene left over will be low
What is the best work up?

[Edited on 11-9-2010 by roamingnome]

Sedit - 10-9-2010 at 19:48

BnO not BzO, I was corrected myself sometime ago and I will gladly share the same.

Hex if it does work as patented as per dropwise addition Mn(IV) oxidiser would be better then MAA. Not to mention cheeper startup cost.

roamingnome - 12-9-2010 at 13:45

I had a chance to play around with "royal purple" again this weekend

I tried to drip the purple into the toluene but ultimately the kinetics of the reaction are to slow to for this to be of use Sedit

For kicks I bubbled ozone and later hydrogen into a smaller mix of Tolu and Mn.
The ozone began to form white solids or benzoic acid.

The point is, this reaction needs an activator probably for the toluene because when ozone forms white solids instantaneously that's a reaction

So i have 5 gallons of purple stuff .. any ideas out there





sabbath06 - 15-9-2010 at 06:48

Yeah, by heating it more I got all of the purple oxidizer to turn back to it's original color, toluene layer was tinted yellow but does not smell like benzaldehyde at all. This run was in a 400ml beaker. Will try scaling up to about 1 L and watch the temps more closely

behemoth - 15-9-2010 at 10:19

Quote: Originally posted by sabbath06  
Yeah, by heating it more I got all of the purple oxidizer to turn back to it's original color, toluene layer was tinted yellow but does not smell like benzaldehyde at all. This run was in a 400ml beaker. Will try scaling up to about 1 L and watch the temps more closely


Are you refering to the manganese alum procedure or to the manganese (IV) in the divided cell setup?

[Edited on 15-9-2010 by behemoth]

Sedit - 15-9-2010 at 16:45

Quote: Originally posted by sabbath06  
Yeah, by heating it more I got all of the purple oxidizer to turn back to it's original color, toluene layer was tinted yellow but does not smell like benzaldehyde at all. This run was in a 400ml beaker. Will try scaling up to about 1 L and watch the temps more closely


Im pretty keen on the idea of dripping the oxidizer at a rate and Toluene temperature that pretty much converts it as its dripped in and stirred always allowing an excess of Toluene.

atomicfire - 22-4-2011 at 04:16

I've decided to test out this cell but I have a question.

I dissolved the ammonium sulfate in the 60% sulfuric like stated. I added the amount of manganese sulfate and let it stir for 24 hours. I still have a 'pink milkshake' of a solution with no visible yellow precipitate. Is this normal?

I don't have the power supply for the electro part yet, so I can't see if this will turn red with electryolsis.

One question I had is that I can obtain manganese sulfate in either anhydrous or monhydrate form. I haven't seen anyone debate this but I wonder if that extra water does something. I used anhydrous.

Also, whats the stoichiometry of this reaction? If you add toluene and turn that into benzaldehyde you are adding an O atom, where does that come from?

Thanks

A little enlightenment.

DetachablePenis - 4-12-2014 at 09:49

Hello,

I constructed and ran such a cell some years ago (about 6 liters total volume of all reagents) and after much the same results as I read now here regarding low yields, it finally occurred to me that what was occurring was that the toluene was simply evaporating out of the top of the open cell before it was able to be reacted.

Upon sealing the cell and using a conventional condenser to manage pressure build up I was able to achieve the claimed yields.

If you look at the illustration in the old patent where they used several hundred KG of reagents, you can see that the 'tardis' or I guess even phallic shaped steel (I assume) reactor they used was a closed unit with a condensing pipe running from the top back down into the main body of the reactor.

Hope this helps everyone with their frustrations :)

P.S. Computer power supplies connected in parallel are a cheap source of lots of amps if you can scrounge old discarded PC's which is very easily done as you may be aware.

I had about 10 or so CPU power supplies (about 3 plugged into each household circuit) in parallel at up to 35 amps a piece... I did not need any heating in the cell :) I did need thick, thick copper braid !

This regenerates fast.

Lead is easily obtained on the cheap from your metal recyclers, melted by a propane flame and poured onto ceramic tiles to create beautiful sheets of almost pure lead that you can cut up with tin snips or whatever for your anode etc.

Overhead stirring using a medium speed/high torque motor is a must and a bitch to seal. Failing going pro and purchasing a proper lab quality nipple setup, you can use the brass/s steel combo bearing out of the bottom of a food blender or similar appliance and use silicone or epoxy to jig it up.

Such a setup requires constant replacement of silicone or ptfe based lubricant or the sulfuric will eat the SS. You couple/join the SS center originally from the blender to a glass stirring shaft and paddle with a plastic "shaft coupler" from a hobby/model shop (looks like a hollow cylinder standing on its short end with two horizontal plastic hex head screws on the side that grasp the two vertical shafts) as anything else will work loose in my experience.

I am sure others will think of better ways, this was just what occurred to and worked for me.

Chemosynthesis - 4-12-2014 at 20:46

Excellent insight. Thank you very much for sharing.