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Author: Subject: sulfuric acid from gypsum CaSO4
Anders Hoveland
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[*] posted on 16-6-2010 at 09:19
sulfuric acid from gypsum CaSO4


Also, came up with an industrial process for producing sulfuric acid from common and inexpensive gypsum (CaSO4). An old chemistry book I had said that scientists couldn't find a commercially viable process for turning this potential resource into the highly useful acid product. My process actually operates at a lower temperature than the SO2 oxidized by air process currently used. However, I figure H2SO4 is so cheap now, produced in China, that there probably wouldn't be much economic incentive to change processes now to a slightly better one.
The process involves distillation with boric oxide, then recovering the boric oxide (now borate) by mixing in NH4Cl. First solution is heated to separate out the NH3. Then the boric acid is heated to drive off water, leaving boric oxide again.
...then recovering NH4Cl by reacting waste NH4OH with NaCl and bubbling in CO2, precipitating out bicarbonate (less soluble in the excess basic NH4OH), and leaving NH4Cl to be used again. Na-Bicarbonate is a commercially valuable product, so this byproduct wouldn't go to waste. (this paragraph is the Solvay process which is no longer used).
In fact bisulfate is a byproduct of HCl industrial production (H2SO4 on NaCl), and ends up in dumps, though they try to pawn it off on the swimming pool industry. Heat bisulfate to get pyrosulfate, then heat boric oxide with pyrosulfate, and you can distill off SO3 at a surprisingly "not too hot" temperature, of around 300degC. leaving a borosulfate mess behind. that was my invention anyway.
Maybe if I had lived 120 years ago, they'd be producing sulfuric acid a different way now.
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[*] posted on 17-6-2010 at 03:23


Er, no!

The current process of oxidising sulphur dioxide with oxygen and then reacting it with water can produce very high purity acid or oleum very cheaply indeed.
A lot of energy can be recycled in the process by using the exhaust gases to heat the incoming gases and it is 100% atom efficent eg everything that goes in comes out as acid and nothing has to be recycled.
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[*] posted on 17-6-2010 at 15:05


Anders:

I must admit I've never run into this one, and I've done a lot of reading.
Have you actually tried this process?
Can you direct us to your reference? It sounds interesting.
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[*] posted on 17-6-2010 at 19:06


Quote: Originally posted by Anders Hoveland  
leaving a borosulfate mess behind.
If the boron content in the waste stream is significant, I can't believe that this method would be economical. Boron, which not hugely expensive, isn't hugely cheap either, certainly not like the existing feed stocks for industrial sulfuric acid production. If your boron cycle approaches a completely closed loop, it might not fail for that reason.

One of the reason that the contact process is so widespread is that is uses the sulfur-bearing feed stock as fuel, oxidizing it in the process and generating heat. All industrial sulfuric acid plant deal extensively with rejecting heat from its point of generation and moving it around, some to other parts of the plant and some to outside the plant as a form of cogeneration. The upshot is that a contact process plant is a net generator of heat. Your proposed process is a net consumer heat, and that's not looking good at all for economic viability.

The mineral sources of sulfur include not only raw sulfur, but also the exhaust of smelting sulfide ores, and it's not insignificant either, at about a quarter or third of world production. Here the acid plant is processing what would otherwise be a waste product into a separate income stream. It's hard to compete with a process where your main chemical feed stock has negative cost (an opportunity cost). Less common, but still common enough is the sweetening of natural gas by scrubbing out the sulfur. The point here being that non-oxidized sulfur is very plentiful, and there's no reason not to use the thermodynamic advantage of starting higher on the energy ladder.
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not_important
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[*] posted on 17-6-2010 at 21:20


Let us look at this proposal in a bit more detail, as a follow-on to WF's comments.

First we start with

CaSO4.2H2O => ( CaSO4.1/2H2O ) => CaSO4 + 2 H2O

2 H3BO3 => B2O3 + 3 H2O

Hmm .. bunch of energy input there.


At a minimum temperature of 500 C, more likely 600-700 C and possibly higher

B2O3 + CaSO4 => Ca(BO2)2 + SO3 => SO2 + 1/2O2

The temperature is needed to get a sufficiently low viscosity of the B2O3 and borate to give a reasonable reaction rate. Extremely fine powdering of the CaSO4, with the H3BO3 being added as a solution, might help but would need to be proven. Sulfur trioxide is less favored then SO2 and O2 at those temperatures, so unless the gases can be removed from the reaction zone and quenched very quickly, a contact process plant is still needed to regenerate SO3.

Note this also works with phosphates, producing P2O5 which is then reduced to give volatile elemental phosphorous, which escapes and drives the reaction.


Proposed is the following (inferred, no equations were given so I used the reference data I had at hand)

Ca(BO2)2 + 2 H2O + 2 NH4Cl => CaCl2 + 2 NH3 + 2 H3BO3

Note that there already is an oversupply of CaCl2, this is one factor in the decline of the Solvay process.


however Ca(B(OH)4)2.2H2O has a rather low solubility, especially at pH range of 8 to 13. Thus I would expect

Ca(BO2)2 + 6 H2O => Ca(B(OH)4)2.2H2O

to play an important role. NH3 can be driven off, helping force the CaCl2 + H3BO3 forming reaction. However the complexing action of CA(2+) and NH3/NH4(+) must be considered as well.

It would not surprise me that blowing superheat steam over finely ground Ca(BO2)2 would drive off H3BO3 and NH3, which could be condensed to a mixture of boric acid and ammonium borates, the exact proportions depending on conditions.

An alternative might be

Ca(BO2)2 + 2 H2O + CO2 => CaCO3 + 2 H3BO3

using excess superheated steam and CO2

An alternative might be treating the borate with MeOH, and distilling off the B(OMe)3-MeOH azeotrope. This would then be used to treat the gypsum to obtain the sulfate-borate mass

CaSO4.2H2O + 2 B(OMe)3 + 4 H2O => CaSO4 + 2 H3BO3 + 6 MeOH

with some boron remaining as the methyl ester in the MeOH exit stream, but that would be recirculated to break down further calcium borate and thus is not lost. Excess water would likely be needed as well, so the gypsum may remain as such while free water would hydrolyse the borate ester.


It should ne remembered that in the Solvay process, the post NaHCO3-precipitation stream contains some NaHCO3, NaCl, and considerable amounts of dissolved ammonia, as well NH4Cl.


Quote:
In fact bisulfate is a byproduct of HCl industrial production (H2SO4 on NaCl)


The majority of HCl production is byproduct of chlorination of organic compounds, production of NaOH, and incineration of chlorine-containing organic waste (often in-plant), further is produced by heating the metal salts resulting from metal cleaning ('pickling') to reclaim HCl for reuse and Fe2O3 for sale to iron smelters. In the U.S. 9/10 or more of the HCl production comes from chlorination of organics, although this may have declined in the last decade as government policies encouraged the emigration of U.S. heavy industry.

I believe that the main users of HCl from chlorides and H2SO4 are the food and pharmaceutical sectors, where freedom from trace organics is of importance.



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Anders Hoveland
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[*] posted on 18-6-2010 at 13:01


The main advantage of the technique is that it could operate at a lower temperature than the oxidation of SO2. The disadvantages are many, and include complexity, a need to recycle/recover the intermediate chemicals. Everything is recycled, however, so there are no byproducts. Theoretically, heat need not be wasted if it is carefully channelled through the the intermediate chemicals. Pyrosulfate and B2O3 is used, allowing for lower temperatures than either the decompostion of pyrosulfate or sulfate heated with B2O3 alone. It might be easier for the amateur, because oxidation of SO2 in a lab is much more difficult than in an industrial process.
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[*] posted on 18-6-2010 at 13:47


Quote: Originally posted by Anders Hoveland  
The main advantage of the technique is that it could operate at a lower temperature than the oxidation of SO2. ...


You say this, yet my readings and some experience suggest 500 C as a minimum temperature, and likely higher temperatures needed. If you have evidence to the contrary, perhaps consider presenting it. As it stands, this is hotter than the conventional V2O5 catalysed oxidation method, especially if uses Cs2SO4 in place of K2SO4.

At higher temperatures
CaSO4 + C => CaCO3 + SO2
is another route.

Electrodialysis of Na2SO4 or (NH4)2SO4 using ion selective permeable membranes would be a low temperature method; unfortunately the membranes are not easy for an amateur to obtain.


[Edited on 18-6-2010 by not_important]
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[*] posted on 18-6-2010 at 18:19


Quote: Originally posted by Anders Hoveland  
Everything is recycled, however, so there are no byproducts. Theoretically, heat need not be wasted if it is carefully channelled through the the intermediate chemicals. [...] It might be easier for the amateur, because oxidation of SO2 in a lab is much more difficult than in an industrial process.
There's a huge difference between "is recycled", "can be recycled", and "can be recycled economically". If you're going to claim recycling as an advantage, you need to do more than simply state some formulas and shout "recycles!!" You must (as not_important illustrated for you) consider side reactions, material properties, etc., at minimum. And the recycling must be energetically efficient, which leads me to my next point...

When you say "heat need not be wasted if it is carefully channelled", I see red flags in abundance about someone who does not appreciate how quickly cascaded heat engines, with their maximum Carnot efficiency determined by temperature differences, lose out in maximum efficiency compared to a single heat engine with same total temperature difference. If you're going to argue that this is thermally efficient, I want to see the calculations.

Finally, you originally argued that this was a manufacturing process, and now you've switched to saying that it could be a viable amateur process. OK, show me. Do the reaction, even just the step that generates SO3, and report your yield of H2SO4. It's not for me to show that your process works, it's for you. Only you.
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[*] posted on 18-6-2010 at 19:32


And you failed to respond to my questioning the 2nd step, the processing of the calcium borate. Further evidence that step is bogus that

there are a number of studies on the properties of ammonium borates
http://dx.doi.org/10.1016/j.tca.2004.04.027

they exist in nature
http://www.galleries.com/minerals/carbonat/larderel/larderel...

are commercial products
http://www.borax.com/pdfs/dist/Profile_Ammonium_Pentaborate....

and boric acid is used to capture ammonia in the Kjeldahl method
http://www.coleparmer.com/techinfo/techinfo.asp?htmlfile=Kje...

There's little purpose served by a thumbnail sketch of a process when so little information is given that it is not clear what the reactions are intended to be, where there is no discussion of details of the steps to illustrate that they are real, and where it appears that several of those steps (as well as can be guessed) do not work.

Present the concept with enough information that it can be evaluated, even say something like "I'm not sure about this one, any input" or"how might this be improved", is something concrete that can be discussed in a useful way.
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[*] posted on 19-6-2010 at 17:41


3Ca(BO3)2 + 6NH4Cl heated--> 3CaCl2 + 2H3BO3 + 6NH3gas
I have a beginning chemistry book that says B2O3 will displace SO3 from BaSO4 at 400C. Pyrosulfate decomposes at 460C. Combining the two reactions together seems to reduce the temperature required, though I am uncertain exactly by how much. 2Na2S2O7 + B2O3 --> Na4S2B2O11 + 2SO3(gas). It is possible that someday all the economical reserves of sulfur will be exhausted, or possibly a country that has an embargo against it, without access to cheap sulfur, but with abundant source of gypsum. WWI is responsible for making the Germans develop NH3 production, when supplies of Chile salt peter were cut off.
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[*] posted on 20-6-2010 at 04:15


Ah, heating it until the NH4Cl fluxes the mixture. Several hundred C, the H3BO3 will dehydrate to an appricable extendt, the water released would carry away remaining H3BO3, which would combine with NH3 in the condensate. Some HCl would likely escape too. Corrosive as well. And more energy input, the first and second step both take a fair amount.


OK, there is a difference between calcium and barium, looking up the decomposition temperature of MgCO3 and expecting BaCO3 to act identically will bring you grief.

400 C is a satisfactory temperature for the V2O5 contact process. Commercial units see a temperature rise to as high as 600 C due to the energy release, but that temperature reduces conversion. The following stage operate in the 400 to 500 C range, interstage cooling being used to drop the gas temperatures.

The temperature can be lowered a bit further with proper catalyst tweaking, most industrial plants don't bother. Even so, practical contact process temperatures run the same to 100 C above your claimed but untested temperature. And the contact process avoid complicated multistep fiddling using reactions known to have material losses.

Now I've smelled sulfur oxides in the flue gases of a pottery kiln when due to not having Gerstley borate on hand a potter used plaster of Paris and boric acid. But this was a firing in the range of cone 06 to 04, in the region CaSO4 decomposes on its own and well above your 400 C.

And the thing is, industry already uses waste CaSO4 as a SO2 source. I gave on reaction earlier, try searching for Müller-Kühne and see what that brings up. So they already can and do use CaSO4 under certain circumstances.

As for running out of sulfur, as already pointed out that smelting sulfide ores is a significant portion of SO2 production, while desulfuring natural gas is another - there are parts of Canada where there are huge piles of sulfur from that source. As your process takes energy input, if the natural and coal supplies run out your process may have serious trouble being implemented, and until then those fuels will provide sulfur.

A country under embargo might have different evaluations of cost. But there are more suppliers of sulfur than borates, and there are those existing processes that use CaSO4 without the complexity of yours, and with working installations vs "I think it will work, but I haven't tried it"


Quote:
2Na2S2O7 + B2O3 --> Na4S2B2O11 + 2SO3(gas)

Interesting compound there, have a reference on it? No?
I propose the reaction would be

Na2S2O7 + B2O3 => 2 NaB + S2O3 + O7

At this time I've presented as much information as you regarding a reaction of pyrosulfate and B2O3.

Time for you to break out the scales, notebook, burner, and so on. We expect a report on this reaction before any further WAGs get posted.




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[*] posted on 20-6-2010 at 07:14


Quote: Originally posted by Anders Hoveland  

I have a beginning chemistry book that says B2O3 will displace SO3 from BaSO4 at 400C. [...] It is possible that someday all the economical reserves of sulfur will be exhausted, or possibly a country that has an embargo against it, without access to cheap sulfur, but with abundant source of gypsum. WWI is responsible for making the Germans develop NH3 production, when supplies of Chile salt peter were cut off.
I want to elaborate on the non-analogy between CaSO4 and BaSO4. The sulfate is ionically bonded in both cases. The ionization potentials (they're on Wikipedia; go look them up) are the difference between 4s and 6s electrons. Your energy partition function from thermal activity needs to get adequate population above that ionization energy for the reaction to succeed irreversibly. This alone should have been a warning sign to you that this reaction did not occur under the same conditions.

not_important didn't mention it explicitly, but if you've got coal, you've also got a lot of sulfides. It's the major source of sulfur impurity in coal. Already mentioned that if you've got oil and/or gas, you've also got sulfur. So if you have an industrial energy source, you've got sulfur.

To review: First you touted ordinary manufacturing. Then you touted amateur access. Now you have touted countries under embargo. I'm sure that there are populations of ever-greater unlikelihood you could propose, but please don't. Go do some thermodynamics instead.

Lastly, the Haber process for ammonia went live on an industrial scale in 1913, a few years before World War I. It was originally making fertilizer, shifting into nitrate production only later.
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[*] posted on 22-6-2010 at 08:35


Quote: Originally posted by not_important  

Proposed is the following (inferred, no equations were given so I used the reference data I had at hand)

Ca(BO2)2 + 2 H2O + 2 NH4Cl => CaCl2 + 2 NH3 + 2 H3BO3
[...]
Ca(BO2)2 + 2 H2O + CO2 => CaCO3 + 2 H3BO3
[...]
CaSO4.2H2O + 2 B(OMe)3 + 4 H2O => CaSO4 + 2 H3BO3 + 6 MeOH
I've quoted the borate recycling reactions here to focus on that part of things. It seems clear enough that this cycle isn't viable for acid from CaSO4. But it did get me thinking that an internal borate recovery cycle might be good for some other purpose. The example here is calcium-focused, but Ca doesn't seem particularly unique here. Nothing occurs to me offhand, but I thought that others here on the board might have brainstorming suggestions.
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[*] posted on 27-6-2010 at 21:24


OK, preliminary report on attempting the 1st step, under somewhat primitive conditions.

An old box of Plaster of Paris sufficed as a source of CaSO4. Lacking was B(OH)3, so I had to make some from muriatic acid and borax. I used a bit of old pipe with end cap I found in the alley as a reaction vessel.

A mixture of roughly equal parts by volume of the plaster and boric acid were mixed and placed in the tube. Using a bent-up coat hanger as a clamp, the tube was held over a gas flame. Occasionally the tube's open top was tested with damp pH and turmeric papers, and sniffed as well.

At first steam was driven off, then boric acid was seen crystallising near the top of the tube; the test papers giving a weak acid and strong borate response. There was a faint order, hard to place, but unlike either common oxide of sulfur.

Continued heating produced some additional B(OH)3 crystals, which partially melted and partially volatilised with the steam being released as heating was increased.

Nothing further of interest happened for some minutes, until the lower portion of the tube reached dull red heat, at which time a faint odor of SO2 could be detected and the pH paper showed a moderately acidic response of a pH of around 3 or 4.

I would say that this first experiment indicates that the reaction between CaSO4 and B2O3 does not happen at temperatures below 500 C, however the experiment had some shortcomings. I need to repeat this with some changes.

First I need to clean the tube and really cook it well to drive off anything volatile; some of the odors may have been from oil or such on its surface, heating to a bright orange should take of that.

The iron of the tube could have reacted with the acidic reagents and/or products, preventing the reaction or capturing the products. After the high temperature bake, I'll coat the inside with a CaSO4 slurry, dry, and heat to ~300 C to reduce its contact with the sample.

I need to dehydrate the CaSO4 before mixing with the boric, I suspect that the water released by the aged PoP, which may have picked enough water to be near CaSO4.2H2O, may have caused excessive loss of B(OH)3 through its being carried off with the steam.




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[*] posted on 29-6-2010 at 20:04


Reaction of B2O3 with sodium sulfate begins at 500 C, and at 1200 C, 59% of the sulfate has decomposed. H2O-vapor quickens the action: 2 B2O3 + Na2SO4 = Na2B4O7 + SO3. M.G. Levi, Garavini (Gazz. 41 I [1911] 756). It is like the decomposition of NaCl with B2O3 yielding borate and Cl2.

I have doubts about BaSO4 working at a lower temperature, its decomposition temperature alone is higher than Na2SO4, but lower than CaSO4. In either case, what can happen is like above, at around 500 C or so ("decomposition begins at"), you get some reaction, but to get significant decomposition within a shorter rate, you may end up needing to heat around blow torch temperatures, increasingly forming SO2.

Also, remember the SO3 equilibrium. Some metals and cations will decompose SO3 (if they form it) to SO2 and O2 upon heating past certain temperatures. As mentioned in the NaHSO4 to SO3 thread by garage chemist, Na+ is not one of those.

The vessel to do this in is another problem. With SO3 in mind, hot SO3 will attack iron and steel, etc. The heated B2O3 could react with glass and quartz.

Decomposition of CaSO4 to harvest SO2 has already been used industrially. The processes made use of the mentioned carbon reduction. With C and CaSO4, at 1200 C, decomposition begins 2.3%, at 1300 C, 4.1%, then at 1375 C, nearly complete decomposition occurs.

I would stick with H2SO4 and pyrosulfate (or persulfate).
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[*] posted on 1-7-2010 at 19:18


Is the decomp temp for BaSO4 lower than CaSO4? Certainly for MCO3 and M(OH)2 the decomposition temperature increases as Mg<Ca<Sr<Ba. The numbers I have for BaSO4 indicate decomposition ~1600 C at 1 bar and 1300-1400 C in vacuum.

Yes, ferrous alloys are crappy containers, but if it had worked as claimed then some SOx should have escaped and been noticeable; I detected none.

The point of my trying this was it was easy, nothing complicated or hard to obtain. The originator should have tried at least the first step before starting this thread.



Ah - watson.fawkes :



[Edited on 2-7-2010 by not_important]

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Anders Hoveland
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[*] posted on 2-7-2010 at 00:15


I tried the reaction with Na2S2O7 and B2O3, and has indeed able to get a small amount of SO3, but not enough to get anything substantial. A few little droplets of H2SO4 condensing was all. I used a stainless steel metal pipe that fed into a glass container for collection. I used a propane burner, but it was obvious that the heat was having trouble reaching the inside of the pipe. A spot on the pipe got to a very dull red heat. B2O3 has a very high boiling point, I was unable to get it to boil, even with a burner directly on it and it glowing orange. On its own, I was unable to heat the pyrosulfate hot enough to decompose in a metal container.

Perhaps, trying BaCl2 and Na2S2O7 fused together, might make SO2Cl2, since BaSO4 is so stable. ?

BaCl2 + 2Na2S2O7 -->BaSO4 + 2Na2SO4 + SO2Cl2

I no longer have Ba(NO3)2, so I cannot do this reaction.
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[*] posted on 8-7-2010 at 22:44


Quote: Originally posted by not_important  
Is the decomp temp for BaSO4 lower than CaSO4? Certainly for MCO3 and M(OH)2 the decomposition temperature increases as Mg<Ca<Sr<Ba. The numbers I have for BaSO4 indicate decomposition ~1600 C at 1 bar and 1300-1400 C in vacuum.


For BaSO4, most references in Gmelin state in an electric oven in dry air, at 1500 C, begins the decomposition. Some others noted decomposition of BaSO4 at 1100 C, 1200 C, or 1300 C.

I have seen the decomposition temperature of CaSO4 given as 2000 C. This value was from Chemie Selbst Erlebt by E. Grosse, which mentions J. Müller as 'recognizing the decomposition temperature of CaSO4 could be decreased from 2000 to 1200 C through coke'. I'm not sure of possible differing crystal forms. Ullmann's seems to be saying a rotary kiln can be used to decompose CaSO4 between 900-1400 C.

Both compared to sodium sulfate: Na2SO4 in a dry air stream decomposition begins between 1200 and 1220 C, and goes quickly between 1330 and 1350 C (Cobb, J. Soc. Chem. Ind. 29 [1910] 399). At white glowing heat, it volatilizes (Boussignault, Ann. Chim. Phys. [4] 12 [1867] 427).

Quote: Originally posted by Formatik  
The vessel to do this in is another problem. With SO3 in mind, hot SO3 will attack iron and steel, etc. The heated B2O3 could react with glass and quartz.


Never mind that! B2O3 attacks glass pretty bad. I was doing another experiment, heating sulfur flowers with B2O3 in a distillation flask under the Bunsen burner. The hope was that the B2O3 would preferentially oxidize the sulfur, and form SO3: B2O3 + S = SO3 + 2 B. The mixture did darken as expected, but the sulfur mostly just volatilized. After heating for something like 30 minutes to an hour, no SO3 condensed over, and rest in piece, distillation flask.:( Pulverizing B2O3 is also a pain in the hands, it's not much different than pulverizing glass after you've made it from H3BO3.
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