Synthesis of Sulfur Trioxide and Oleum: The persulfate method

This is a new synthetic pathway to sulfur trioxide, with completely OTC chemicals and without the need for fancy apparatus, like catalyst tubes and such.
The idea comes from the german wikipedia page about sulfur trioxide. The pictures and process description below are from me.

The only apparatus required is a ground- glass distillation setup, the same as you would use for the distillation of nitric acid.

Warning!
Sulfur trioxide is extremely corrosive. It instantly carbonizes any organic matter it touches, including skin. A drop placed on wood instantly makes a black spot, under strong fizzing and fuming. The same happens when a drop falls on the skin.

Sulfur trioxide also reacts with water with explosive violence.
When a drop of water falls into a flask containing SO3, the flask usually shatters because of the violent reaction and localized heating.

SO3 fumes in air very strongly.
The synthesis must be carried out outside or under a fume hood.

When sodium persulfate is heated, it loses oxygen to form sodium pyrosulfate.
2 Na2S2O8 ---> 2 Na2S2O7 + O2

The evolved oxygen contains a small amount of ozone, which is identifiable by its smell.

Sodium pyrosulfate further decomposes into sodium sulfate and sulfur trioxide, but the reaction only takes place when sulfuric acid is present.

Na2S2O7 ---(H2SO4)---> Na2SO4 + SO3

The sulfuric acid plays the role of a catalyst, it does not take part in the overall reaction, but without its presence, no SO3 is formed.

Experimental:


17,5g dry sodium persulfate were put into a 100ml round- bottom flask, and 2 ml concentrated sulfuric acid (of the highest obtainable strength, any water is very detrimental to the yield) were added.

The flask was attached to the distillation setup, and heated by means of a bunsen burner.
No water is run through the condenser, as the SO3 would crystallize in the condenser. If necessary, the receiver is immersed in ice water for cooling (not necessary with this small batch).
This was the setup:


On heating, the mixture starts to rapidly generate oxygen, with some ozone as byproduct.
When the oxygen evolution ceases, the mixture is heated stronger. It quickly melts into a clear liquid, and dense white fumes start to fill the setup.

Soon the first drops of liquid SO3 will fall into the receiver.
The thermometer registers a steam temperature between 45 and 50°C during the process. If it becomes significantly higher than that (e.g. over 60°C) heating should be reduced a bit, since some H2SO4 could start to come over.

Residues of grease in the apparature will be carbonized by the SO3, this is normal and nothing to worry about.
The SO3 will also usually be colored brown due to this.

The flask will reach a very high temperature during this step (around 300°C), this is necessary for the sodium pyrosulfate to decompose.

Care must be taken to avoid the crystallizing SO3 plugging the condenser outlet. This can be very dangerous. The solid SO3 can be molten by means of a hot air gun if it is found to crystallize at the condenser outlet.

When no more SO3 is coming over despite the reaction mixture still boiling, the reaction is over.

The SO3 will most likely have partially or completely solidified in the receiver.
Mine looked like this:



This shows the fuming of the SO3 (sorry, it is hard to capture it in a picture, and the smoke is constantly drawn away by the fume hood):



My yield was only 1,4g SO3 from the 17,5g persulfate, but this was mostly due to the very small batch size.
A yield of 100g SO3 from 200g persulfate is claimed from another person.
The ration of persulfate to H2SO4 can also be varied and its effects on yield studied.

The distillation flask can be emptied, recharged with fresh sodium persulfate and H2SO4 and distilled again to increase the amount of SO3 obtained.
The batches should not be larger than 40g persulfate at a time, because larger batches will be heated unevenly.

The combined portions of SO3 can be redistilled for better purity.
Pure SO3 boils at 44°C and solidifies at 16,8°C.

The still is left in the fume hood until no more fuming is observed (all the SO3 is turned into H2SO4 by atmospheric moisture), then it can be safely washed out with water.


SO3 will quickly polymerize upon storage, evidenced by transformation into a white amorphous mass (the polymerization starts within a few minutes to hours after preparation, depending on purity), but it depolymerizes again at ca. 66°C. Therefore, simple distillation gives normal SO3 again.

The SO3 can be dissolved in conc. H2SO4 to give oleum of any desired strength, but be warned, this process is exothermic, and with the very low boiling point of SO3 it is very likely to start boiling. The liquid SO3 should be slowly dripped into stirred, ice- cooled H2SO4 to avoid this.
Polymerized SO3 also dissolves in H2SO4, but slower and with less heat evolution.



Sulfur trioxide and Oleum are exceptionally versatile reagents in the laboratory. Some of the most important uses shall be mentioned here:

With methanol or dimethyl ether, the powerful methylating agent dimethyl sulfate is formed. It is isolated by fractional distillation in vacuum.
Diethyl sulfate is prepared analogously, by distilling SO3 into dry diethyl ether, distilling away the ether and distilling the diethyl sulfate over Na2SO4 in vacuum.

By distilling SO3 into SCl2, the important chlorinating agent thionyl chloride is produced. This can be used for the production of acetyl chloride from acetic acid, and subsequently acetic anhydride.

By leading a stream of dry HCl gas through oleum until gas uptake stops, chlorosulfonic acid can be isolated by distillation.
Chlorosulfonic acid is used for chlorosulfonations, by reaction with benzene, benzenesulfochloride (C6H5-SO2Cl) is obtained.
Chlorosulfonic acid can also be used for an alternative preparation of dimethyl sulfate, look into "The war gasses" by Mario Sartori for a detailed synthesis.


Oleum is also employed as a condensing agent in organic chemistry, for example for the condensation of chloral with chlorobenzene to form DDT, the known powerful insecticide.



I hope that this synthesis is of use for you.
I am further developing the process and start to make larger amounts of SO3, as soon as I get more persulfate.
My batch was so small since I had only a small amount of Na2S2O8 left. I'll make pics of a larger batch when I get more.

Oleum, even if you can buy it, is horribly expensive from any of the known chemical suppliers (over 100$ per Liter).
Therefore, a good synthesis for SO3 is necessary.

[Edited on 14-3-2006 by garage chemist]


Edit by Dav: Topic name extended to aid differentiation between this method of synthesis, and Fleaker and NERV's catalytic V2O5 one.

[Edited on 3-8-2007 by The_Davster]

That's very nice work gc. Thanks for putting it into publication. I hope to make use of it someday as I've always wanted some thionyl chloride.

Aside: I enjoyed the blowup of your pictures. Especially the incongruity of the bottles labeled in German along side the ad for Chilly's bar & grill.

GC, this is a very nice result. I have a few ideas, which might be helpful for increasing the yield or using cheaper chems.

Could adding a small amount of H4P2O7 be beneficial for the yield? H4P2O7 can be prepared easily by heating H3PO4 (85%) until no water is coming off anymore and the liquid becomes somewhat syruppy.

I can imagine that adding 1 ml of this, 2 ml of 96% H2SO4 and 20 to 25 grams of Na2S2O7 can give a somewhat higher yield.

Another thing. Why don't you use the much cheaper and easier to obtain NaHSO4 (pH-minus for swimming pools, over here I pay just EUR 20 or so per 2.5 kilos). This then first needs to be preheated, until it has lost all its water. The then dry Na2S2O7 is mixed with the H2SO4 (and maybe H4P2O7) and then you perform the procedure as described above.

These are just my two ideas. Unfortunately I have no opportunity to do this at my home (no suitable room and outside to many little children's hands around ), but I'm very interested in your results.

I already got the persulfate. The electronics supplier delivered much faster than usual.

Now I have 2kg Na2S2O8, good for theoretically over 600g SO3.

I think it is worth investigating the use of only a very small amount of H2SO4 as catalyst, since H2SO4 always contains about 4% of water.
It's Oleum time!

Now I have the time to research SO3 production.

Yesterday, two experiments.

First one: usage of H2SO4 as catalyst.
23,8g sodium persulfate (0,1 mol) were added to a 100ml round- bottom flask. 1ml conc. H2SO4 was added.
Heated, a lot of oxygen was evolved, and a bit of smoke.
It melted into a liquid, and a very small amount of SO3 distilled over, just a few drops. I heated for 10 minutes on strongest flame, nothing more.
It was left to cool down a bit and 2ml H2So4 added, then reheated and a bit more SO3 distilled over.
Added another 2ml of H2So4 and heated for 20 minutes, the flask nearly glowed red. If it hadn't been made of Duran, it would have melted.

Total yield: 1,7g SO3 (theoretically, 8g should form).
PATHETIC!!!


Next experiment: using MgSO4 as catalyst.
MgSO4 hydrate (Epsom salts) were heated in a quartz dish for 10 minutes on hottest flame, until everything turned into a white powder (the last mol of crystal water of MgSO4 does not begin to split off until the temperature reaches 200°C, so intense heating is necessary).
3g anhydrous MgSO4 were added to 23,8g sodium persufate and heated in the distilation setup.
Again lots of oxygen, but this time a bit of the MgSO4 got blown out of the flask. This was due to too rapid heating.
The mix melted, and SO3 slowly distilled off.
It took maybe 30 minutes of highest heat until SO3 evolution stopped.

Yield this time: 2,8g of again theoretically 8g.
A bit better, but still bad.

Conclusion:

The MgSO4 method is definately the way to go, the yields with H2SO4 are too low to be of any use.

The next experiment will use 6g MgSO4 and again 23,8g sodium persulfate. If this gives a higher yield, another experiment with even more MgSO4 will be made.
When the optimal amount of MgSO4 has been found, it will be tested if additional drying of the sodium persulfate in an oven before synthesis does help (specifications for technical grade sodium persulfate specify up to 0,5% water content).



EDIT: if you choose to try this synthesis yourself, you have to get some advice first.
For example, the condenser must not be cooled by water, otherwise the SO3 will solidify in there.
Only the receiver must be cooled.
The SO3 has a very dangerous tendency to solidify in the outlet of the condenser, which can lead to pressure buildup. It has to be melted periodically with a hot air gun (without removing the receiver!).

The apparatus must be absolutely dry. If possible, dry it in an oven at over 100°C.
The vacuum connection of the receiver should be connected to a drying tube if possible (I didn't use a drying tube, it still worked, but some moisture can make it into your SO3 and decrease its percentage).

SO3 reacts explosively with water. A piece of solid SO3, thrown on water, hisses loudly and swims around like a piece of sodium, with lots of fuming.
Liquid SO3 makes a sharp crack each time a drop hits water.
Water added to SO3 produces powerful explosions and scatters the contents of the flask.




[Edited on 9-4-2006 by garage chemist]

SO3

After thinking about the procedure for a while I have produced a couple of suggestions:
1) Heat the persulphate on its own to evolve oxygen and remove any water content. This should produce the pyrosulphate:
2Na2S2O8 => 2Na2S2O7 + O2
2) Use some of the SO3 that you have already produced to make 100% H2SO4 as Sauron has suggested. Then use this 100% H2SO4 as the catalyst for this reaction. This should eliminate most of the water that would have otherwise been present.
3) Possibly use a mixture of anhydrous MgSO4 and 100% H2SO4 as the catalyst. I dont know if it will make any difference but surely its worth a try.

I would also like to congratulate you on writing up this experiment, I too think that SO3 production for the home chemist is important as it is a very useful chemical. I myself have some sodium persulphate that I used to etch a PCB for my physics coursework project, and would like to use it to produce some SO3 but unfortunately I have yet to get any labware, and doubt if I will have the money to do so anytime soon

Thats amazing news! Any idea when the write-up will be complete? I can't wait to read this

Thanks gc, I now have something to look forward to reading next week. It's a shame I can't read german

I have been off and on working on building a small electric furnace for the last several weeks so that I can produce SO₃ among other things. The furnace was made from kiln bricks (2.5 inches thick), angle iron (2" X 2"), an old clothes dryer heating element (ca. 15 gauge nichrome wire) and an infinite switch from an old kitchen range for temperature control. The infinite switch was rated for 240V, but seems to work fine at 120V. It is the current rating, in this case, which is of greater importance I think. I have a decent digital thermocouple meter but the best probe I have has insulation rated to only 500C. I have ordered a probe that is rated for up to 1250C, but I don't have it yet. The 4000-6000W at 240V dryer heating element was reduced in length slightly (cut) and was run at 120V instead of 240V. The current to the element was measured as about 9.5A when the infinite switch was in the on state. This means the furnace element is putting out about 1140W (P=IV). With the infinite switch turned to the maximum setting and the cover on the furnace an internal air temperature of 500 Celsius could be reached in about 2 minutes. I didn't actually time it, but the temperature rises very quickly. Until I have the higher temperature thermocouple probe I won't know exactly how high the temperature can go. I used the following video as a guide to build my furnace, but I also read David J. Gingery's book, "LI'L BERTHA, A Compact Electric Resistance Shop Furnace", which had in it a lot of useful information.

http://www.youtube.com/watch?v=en4yhzLuD9A

Below are pictures of a trial run. I did two trial runs and led the SO₃ into a Florence flask which contained what was left in the bottom of a bottle of drain cleaner sulfuric (91% H₂SO₄). I obviously still need to do some work on the collection part of the apparatus, as some SO₃ was escaping and water was undoubtedly getting in and being absorbed, but the furnace works wonderfully and produces SO₃ from NaHSO₄ with ease. I have read enough to know that I need a moisture protection tube with anhydrous calcium chloride or some other suitable desiccant in it.

The quartz tube that a glass blower made for me is 30mm id. I had the side arm made rather long so as to avoid troublesome connections and also to act as an air condenser. About 130-140g of sodium bisulfate (pool pH down) was put into the tube for each run. It was just a trial run as I said before, but after the first run the 91% sulfuric acid in the flask was ~94% and after the second ~97% as measured by titration. I didn't measure the volume at the start because I was mostly focused on the furnace for the initial runs, but the final volume of 97% sulfuric acid was 166mL. This is likely not the best use of SO3, since the acid could be brought up to at least 95% reasonable efficiently by distilling off water. I also have a much bigger quartz tube on the way. I am trying for a 50mm tube, with a little greater length also, which would allow me to make about 3 times as much SO₃ per run.

Air temperature in the furnace was kept between 300 and 400C for about an hour where it seemed mostly water came over. Once the air temperature in the kiln got up over 500C it was not long until much more dense white fumes appeared and the drops coming out of the side arm made hissing sounds that could be heard from 50 feet away when they hit the lawn below. The drops instantly carbonized the lawn they fell on. This is when I moved the collection flask into place and began collecting SO₃.




[Edited on 20-8-2014 by Hennig Brand]

Thanks Garage Chemist, I will definitely use your suggestions on my next batch.

Edit:

Ok, I ran another batch. I used the suggestion above and put some glass wool around the end of the side arm/air condenser so that it fit reasonably snuggly in the neck of the Florence flask. I left it a little looser than I probably should have, mostly because I was a little nervous of sealing the system off too tightly. I also used ice around the Florence flask as suggested and didn't add the concentrated sulfuric acid until after all the SO₃ had been collected.

Included are a few pictures that I thought were interesting. The charge of sodium bisulfate containing feed used was 130g, which came in a 4kg pail as pH down for swimming pools. After SO₃ generation had gotten quite slow, and increasing the furnace temperature no longer seemed to help, the flask was removed and covered tightly with a polyethylene plastic bag. It only took me a few minutes to go and come back with 100mL of sulfuric acid, but in that time the SO₃ vapors had turned the polyethylene cover completely black. The flask holding the SO₃ was left in the ice bath, while I went for sulfuric acid, so the vapor pressure would have been quite low. When the 100mL of sulfuric acid was poured into the flask holding the SO₃, the SO₃ dissolved very quickly (approximately a few seconds) and the flask got hot to the point of being almost uncomfortable to hold.

Yield

About one gram of acid was measured out for each titration using scales accurate to hundredths of a gram. A sodium hydroxide solution was first standardized against a known sample of HCl. Phenolphthalein was used. Only one titration was performed for the sulfuric acid before SO₃ addition and one for after. Performing more than one titration for each and taking an average could have increased accuracy, but I was very careful. As stated before the scales used were only accurate to hundredths of a gram and the amount of acid used for each titration was about a gram. Ideally much more accurate scales should be used, but a larger quantity of acid could be used per titration as well.

Volumes of acid were measured with a 250mL graduated cylinder. The sulfuric acid density table from Perry's was used.

Pre SO₃ Addition: Titrated as 94.4% H₂SO₄
0.944(100mL)(1.832g/mL) = 172.94g H₂SO₄
(1-0.944)(100mL)(1.832g/mL) = 10.26g H₂O

Post SO₃ Addition: Titrated as 97.2% H₂SO₄
0.972(114mL)(1.8364g/mL) = 203.49g H₂SO₄
(1-0.972)(114mL)(1.8364g/mL) = 5.86g H₂O

H₂SO₄ added = 203.49g - 172.94g = 30.55g
SO₃ added = 30.55g (80.066g/mol / 98.079g/mol) = 24.94g
H₂O added = 30.55g - 24.94g - (10.26 - 5.86) = 1.21g

Theoretical Yield

According to the MSDS for the pH minus product used its composition is 91.5-94.7% sodium bisulfate and 4.8-8% sodium sulfate.

For 130g of Feed (assuming 91.5% NaHSO₄)
0.915 * 130g / 120.06g/mol (80.066g/mol)(1 SO₃ / 2 NaHSO₄) = 39.67g SO₃

For 130g of Feed (assuming 94.7% NaHSO₄)
0.947 * 130g / 120.06g/mol (80.066g/mol)(1 SO₃ / 2 NaHSO₄) = 41.05g SO₃

Percentage Yield of SO₃

Assuming 91.5% NaHSO₄ Feed = 24.94g / 39.67g * 100% = 62.9%

Assuming 94.7% NaHSO₄ Feed = 24.94g / 41.05g * 100% = 60.8%

There is a purity versus quantity trade-off here. If collection had been started earlier more SO₃ would have been collected, but the product would have contained a larger proportion of water. It looks as though the glass wool did a great job of keeping the water out of the collection flask. I still noticed a bit of SO₃ escaping from the flask, through the glass wool, but the wool could have been packed much more tightly and a lot more ice could have been used around the receiver as well (I was low on ice).




Just to clarify, the orange/red color of the sulfuric acid, in the last picture, is due to the dye or other contaminants in the sulfuric acid drain cleaner used. Pure, highly concentrated, sulfuric acid is a clear viscous liquid.


[Edited on 16-8-2014 by Hennig Brand]

I have a few more pictures which I thought might be appreciated by someone trying to follow the experiment.

Once the furnace cooled down the quartz tube was removed and the sodium sulfate residue/plug was easily removed by turning the tube upside down and gently tapping it on a soft surface like wood. Sometimes the sodium sulfate plug sticks a bit, but pouring a little water into the tube and letting it soak down past and around the plug frees the plug and allows it to be simply dumped out by inverting the tube.




Here are a couple pictures I found on my phone from when I was building the furnace. The first picture shows the frame just after the welding was completed. The second picture was taken during the middle of heating element groove filing and grinding. A metal file and a rotary cutting bit and cordless drill were used to make the element grooves in the kiln brick. The third picture shows the infinite switch that is used. It is the old fashioned type dial controller for the burner on an electric cook stove. This infinite switch was for one of the large burners on an old stove.



The kiln bricks were expensive for me, mostly because I live in a lower populated area of the country and anything remotely specialized often needs to be ordered. The box of 12 kiln bricks was about $40 and because they came from halfway across the country shipping was about $35. The steel was purchased from a scrap yard, as new steel, for about $10. I removed about 8 infinite switches from old stoves and about 6 dryer elements from old clothes dryers, which I was allowed to take free of charge from the same scrap yard. The element will need to be replaced from time to time, especially if operating the furnace at very high temperatures (>1000C for instance), but they can easily be found free (old clothes dryers) and with this design replacement is a fairly simple procedure.

Attached below is a wiring diagram for an infinite switch.


Attachment: 120 & 240V Infinite Switch Wiring Schematic.pdf (75kB)
This file has been downloaded 1231 times


[Edited on 1-8-2014 by Hennig Brand]

Several days ago I was running the furnace and producing SO3 when it got dark. I took a few pictures as it was getting dark, and after dark, which I think are very beautiful. Attached are a few of the better ones.





[Edited on 21-12-2014 by Hennig Brand]

A Simple Char Test for Determining The Approximate Time Collection Should Start

According to "Small Scale Synthesis of Laboratory Reagents", from the Sulfur Trioxide and Oleum chapter summary:

"
• SO₃/oleum is prepared by the pyrolysis of NaHSO₄ to a maximum temperature of 820°C.
• A 500°C cut gives 86% oleum at 90% yield. A 580°C cut gives 94% oleum at 65% yield.
• 100% volatilization of H₂SO₄ from NaHSO₄ shows that the residue is entirely Na₂SO₄.
• The reaction can be conducted in an unprotected quartz tube. Na₂SO₄ does not attack quartz at the reaction temperature.
"

From page 182 in the same chapter:

"Figure 20.2 shows data from the present experiment for the percentage of H₂O and SO₃ volatilized in a slow distillation as a function of temperature. From this, it is seen that the separation of the two volatile components is reasonably efficient, so that at 500°C 90% H₂O and only 10% SO₃ has volatilized, corresponding to an oleum strength of 86% free SO₃ at 90% yield in the upper cut. A cut at 580°C gives 94% oleum at 65% yield."

The graph described above was scanned and used to draw tangents to the water and SO₃ curves at 450°C, 475°C and 500°C in order to estimate the rate of water and SO₃ production at those temperatures. From those rates the concentration of the product produced at those temperatures was estimated. The graphs are attached. The following are the associated calculations that go with the graphs.


Basis: 1 mole SO₃ & 1 mole H₂O Produced in Total

Description of method used:
Instantaneous rates of change in terms of percent volatilized with respect to temperature were estimated by graphically determining the slope of the lines tangent to the curves at each particular temperature. Those values were converted to moles per degree Celsius by assuming a basis of 1 mole of water and 1 mole of sulfur trioxide produced. The flow rates found, for H₂O and SO₃, represent the compositions of the product leaving the generator at the three temperatures examined. The x and y axis of the graph were found to be on different scales, so they were measured and a factor calculated to relate the two.
(F (factor) = 3.744 %/°C)


At 450°C

H₂O Instantaneous Rate of Change: (3.39%/°C) / 100 * 1 mole = 0.0339 moles/°C
SO₃ Instantaneous Rate of Change: (0.890%/°C) / 100 * 1 mole = 0.0089 moles/°C

SO₃ + H₂O ----> H₂SO₄

0.0339 - 0.0089 = 0.025 moles of water left over & 0.0089 moles H₂SO₄ (anhydrous)

Total Mass flow = 0.025moles/°C * (18g/mol) + 0.0089moles/°C * (98.1g/mol) = 1.323g/°C

(0.873g/°C) / (1.323g/°C) * 100% = 66wt% H₂SO₄ Is Produced at 450°C


At 475°C

H₂O Instantaneous Rate of Change: (2.85%/°C) / 100 * 1 mole = 0.0285 moles/°C
SO₃ Instantaneous Rate of Change: (2.42%/°C) / 100 * 1 mole = 0.0242 moles/°C

SO₃ + H₂O -----> H₂SO₄

0.0285 - 0.0242 = 0.0043 moles/°C H₂O & 0.0242 moles/°C H₂SO₄ (anhydrous)

Total Mass flow = 0.0043moles/°C * (18g/mol) + 0.0242moles/°C * (98.1g/mol) = 2.451g/°C

(2.374g/°C) / (2.451g/°C) * 100% = 97wt% H₂SO₄ Is Produced at 475°C


At 500°C

H₂O Instantaneous Rate of Change: (1.73%/°C) / 100 * 1 mole = 0.0173 moles/°C
SO₃ Instantaneous Rate of Change: (3.11%/°C) / 100 * 1 mole = 0.0311 moles/°C

SO₃ + H₂O ----> H₂SO₄

0.0311 - 0.0173 = 0.0138 moles of pure SO₃ & 0.0173 moles of H₂SO₄ (anhydrous)

Mass Flow of SO₃ = 0.0138 moles SO₃ * (80.1g/mol) = 1.105g
Mass Flow of H₂SO₄ = 0.0173 moles H₂SO₄ * (98.1g/mol) = 1.697g

Product Concentration at 500°C = 1.105g / (1.105g + 1.697g) * 100% = 39% Oleum (well over 100% H₂SO₄)






What exactly is considered a good trade-off will of course depend on one's goals. My usual objective is to bring ca. 95% sulfuric acid up to 98-100% sulfuric acid; so the goal is to raise the concentration while at the same time add as much weight of sulfuric acid as possible. From the data obtained from the graphs, it is clear that starting collection when the acid produced by the generator is between 95-100% is a good trade-off in terms of purity and yield, if increasing the concentration of concentrated sulfuric acid is the goal. When the generator is producing 97wt% H₂SO₄, approximately 87% of the water has been removed and 93% of the SO₃ remains to be collected. The concentration increases very rapidly at this point, in relation to temperature, so even if collection is started at 95% H₂SO₄ it will be much higher in short order. Using a char test is a simple way to determine when the acid concentration being produced is over 95%. It is quite easy to determine the difference between 95% sulfuric acid and even 98% sulfuric acid by this method with a little experience. There is a huge difference in the degree and rapidity of charing between 95% and 98% H₂SO₄ on wood or paper.


The Char Test Procedure

The following simple test has been used the last couple of times in order to determine the appropriate time to start collection:

Small pieces of scrap cedar have been used to determine when it was time to start collecting from the SO₃ generator. A drop of 95% acid can be placed on the test wood and watched to determine the rate and degree of charring. Once the furnace air temperature has risen to about 500°C the output from the generator is periodically checked with the cedar block to see if the rate of charring is faster than the 95% reference. When the output is determined to be over 95% sulfuric acid collection is started.

The yields obtained have been very noticeably better since following this new procedure. I will have to get back with a yield, when the generator is run again, since the quick notes taken and measurement techniques used are insufficient to use to report an accurate yield. Also, without knowing the exact temperature that collection was halted, in both cases, it is difficult to know exactly how much difference that may have made to the yields. In any case, the ability to fairly accurately determine an appropriate time to start collection, without any special equipment, will improve yields and reproducibility when producing SO₃ by this method.


[Edited on 18-1-2015 by Hennig Brand]

Quote: Originally posted by Oscilllator

I notice that everybody who has tried the persulfate decomposition has used glass apparatus of one form or another. Is there any reason that you couldn't use a stainless steel container for this? It should easily be able to withstand the temperatures involved, and the decomposition from the metal walls of the container would be fairly minimal, I would have thought.