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theAngryLittleBunny
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Making KBrO3 using Ca(OCl)2 and bromide.
KBrO3 is a very interesting and useful salt, being a stronger oxidizer then even chlorates and providing a very convenient way to make bromine without
destillation. I found a quite easy way to make it using calcium hypochlorite and a bromide salt which might be interesting to anyone who doesn't wanna
deal with the black woodoo called electrochemistry.
I'll first explain the chemistry, so hypochlorite in solution reacts with bromide forming chloride and hypobromite (ClO- + Br- -> Cl- + BrO-). You
can easily see this, because when you dissolve a bromide salt in a hypochlorite solution (like bleach for instance) it will quickly turn orange, which
is hypobromite. Hypobromite is much less stabile then hypochlorite and quickly disproportionates into bromate and bromide at room temperature, which
is what we want here.
So I first dissolved calcium hypochlorite in a minimum amount of water and added over maybe 10 minutes a 1/3rd molar equivalent of sodium bromide
based on hypochlorite to it. The hypochlorite oxidizes the bromide to hypobromite which then turns into bromate and more bromide which gets oxidized
again until all the bromide is turned into bromate.
Ca(ClO)2 + 2 NaBr -> CaCl2 + 2 NaOBr
3 NaOBr -> NaBrO3 + 2 NaBr
Overall: 3 Ca(ClO)2 + 2 NaBr -> 3 CaCl2 + 2 NaBrO3
After all the bromide is added I heat it up in boiling hot water for maybe 30 minutes while slowly and carefully adding HCl until the solution is
slightly acidic (around pH 3 to 4), since commercial calcium hypochlorite contains a lot of calcium hydroxide/ carbonate. Obviously this produces a
bit of chlorine and should be done outside or with good ventilation. The solution has to be a bit acidic, otherwise the yield will be very poor.
After that I filter it, (the calcium hypochlorite contains impurities that won't dissolve) add a molar equivalent of a potassium salt to the bromide
and cool it in the fridge to crystallize out some nice white KBrO3. The yield was about 50% when the solution was made slightly acidic and about 15 to
20% when it was neutral/ slightly basic. The 50% yield was from my first try, so I'm sure it could be optimized. Probably 15% of the KBrO3 is lost in
solution since it still has a solubility of about 30g/L at 0°C. The theoretical yield is about 54g of KBrO3 for every 100g of 70% Ca(ClO)2, which
means in practice at least 27g for 100g. And since Ca(ClO)2 is rather cheap (8 to 10 euro per Kg) you could make 1Kg for KBrO3 for less then 40 euros.
The KBrO3 from that seems to be quite pure too, I used it to make bromine by dissolving it together with 5 molar equivalent of NaBr and acidifying it
with 6 molar equivalent H2SO4 and it worked perfectly (it's important not to use HCl, since the chloride will form a chlorodibromide (Br2Cl-) ion with
bromine keeping it dissolved.
5Br- + BrO3- + 6H+ -> 3Br2 + 3H2O
I hope this is helpful and if anyone is able to optimize this method to increase the yield I'd love to hear about it. Considering the KBrO3 that's
just lost in solution it should still be possible to get up to 40g KBrO3 per 100g Ca(ClO)2.
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Admagistr
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Quote: Originally posted by theAngryLittleBunny  | KBrO3 is a very interesting and useful salt, being a stronger oxidizer then even chlorates and providing a very convenient way to make bromine without
destillation. I found a quite easy way to make it using calcium hypochlorite and a bromide salt which might be interesting to anyone who doesn't wanna
deal with the black woodoo called electrochemistry.
I'll first explain the chemistry, so hypochlorite in solution reacts with bromide forming chloride and hypobromite (ClO- + Br- -> Cl- + BrO-). You
can easily see this, because when you dissolve a bromide salt in a hypochlorite solution (like bleach for instance) it will quickly turn orange, which
is hypobromite. Hypobromite is much less stabile then hypochlorite and quickly disproportionates into bromate and bromide at room temperature, which
is what we want here.
So I first dissolved calcium hypochlorite in a minimum amount of water and added over maybe 10 minutes a 1/3rd molar equivalent of sodium bromide
based on hypochlorite to it. The hypochlorite oxidizes the bromide to hypobromite which then turns into bromate and more bromide which gets oxidized
again until all the bromide is turned into bromate.
Ca(ClO)2 + 2 NaBr -> CaCl2 + 2 NaOBr
3 NaOBr -> NaBrO3 + 2 NaBr
Overall: 3 Ca(ClO)2 + 2 NaBr -> 3 CaCl2 + 2 NaBrO3
After all the bromide is added I heat it up in boiling hot water for maybe 30 minutes while slowly and carefully adding HCl until the solution is
slightly acidic (around pH 3 to 4), since commercial calcium hypochlorite contains a lot of calcium hydroxide/ carbonate. Obviously this produces a
bit of chlorine and should be done outside or with good ventilation. The solution has to be a bit acidic, otherwise the yield will be very poor.
After that I filter it, (the calcium hypochlorite contains impurities that won't dissolve) add a molar equivalent of a potassium salt to the bromide
and cool it in the fridge to crystallize out some nice white KBrO3. The yield was about 50% when the solution was made slightly acidic and about 15 to
20% when it was neutral/ slightly basic. The 50% yield was from my first try, so I'm sure it could be optimized. Probably 15% of the KBrO3 is lost in
solution since it still has a solubility of about 30g/L at 0°C. The theoretical yield is about 54g of KBrO3 for every 100g of 70% Ca(ClO)2, which
means in practice at least 27g for 100g. And since Ca(ClO)2 is rather cheap (8 to 10 euro per Kg) you could make 1Kg for KBrO3 for less then 40 euros.
The KBrO3 from that seems to be quite pure too, I used it to make bromine by dissolving it together with 5 molar equivalent of NaBr and acidifying it
with 6 molar equivalent H2SO4 and it worked perfectly (it's important not to use HCl, since the chloride will form a chlorodibromide (Br2Cl-) ion with
bromine keeping it dissolved.
5Br- + BrO3- + 6H+ -> 3Br2 + 3H2O
I hope this is helpful and if anyone is able to optimize this method to increase the yield I'd love to hear about it. Considering the KBrO3 that's
just lost in solution it should still be possible to get up to 40g KBrO3 per 100g Ca(ClO)2. |
Hi, I find very attractive bromates Ba(BrO3)2 and Sr(BrO3)2 when crystallized from an aqueous solution there are small lightning bolts, miniature
electrical discharges. I'm going to try it... !
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theAngryLittleBunny
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| Quote: Originally posted by Admagistr |
Hi, I find very attractive bromates Ba(BrO3)2 and Sr(BrO3)2 when crystallized from an aqueous solution there are small lightning bolts, miniature
electrical discharges. I'm going to try it...!
|
Cool, but how do you wanna make Sr(BrO3)2? It has a solubility of about 300g per liter at room temperature, which is much higher then KBrO3.
One way would be to make AgBrO3 which has a rather low solubility in water (I think less then 1g/L) and then mix that with a solution of SrCl2 to get
AgCl as a precipitate leaving Sr(BrO3)2 in solution. Anyway, let me know how it goes.
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Admagistr
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| Quote: Originally posted by theAngryLittleBunny | | Quote: Originally posted by Admagistr |
Hi, I find very attractive bromates Ba(BrO3)2 and Sr(BrO3)2 when crystallized from an aqueous solution there are small lightning bolts, miniature
electrical discharges. I'm going to try it...!
|
Cool, but how do you wanna make Sr(BrO3)2? It has a solubility of about 300g per liter at room temperature, which is much higher then KBrO3.
One way would be to make AgBrO3 which has a rather low solubility in water (I think less then 1g/L) and then mix that with a solution of SrCl2 to get
AgCl as a precipitate leaving Sr(BrO3)2 in solution. Anyway, let me know how it goes. |
Here's a link I'm sure you'll be interested in, maybe you already know it?
https://illumina-chemie.de/viewtopic.php?f=18&t=4470
It doesn't address the production of Sr(BrO3)2, your idea is interesting but quite laborious and complex, perhaps inefficient to implement in
practice...I thought to neutralize HBrO3 by SrCO3, it would probably be easier to implement, but I would have to find out how HBrO3 is stable and
explosive/non-explosive...What do you think about it?
I bought my Sr(BrO3)2 from Chemcraft in Russia. But I bought it in a small amount because it's not cheap...I'll think about it some more...
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woelen
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HBrO3 is surprisingly stable, as long as you don't get too high a concentration. I have done experiments with it, up to 10% or so, using barium
bromate and sulfuric acid. Even boiling the solution does not destroy the acid.
Barium bromate can be made quite easily. It is not very soluble in the cold and if you have NaBrO3, then you can make it from a soluble barium salt
(e.g. BaCl2) and NaBrO3 and purify it by means of recrystallization. It can even be made from KBrO3 (which is less soluble than NaBrO3). If you want
to make a solution with HBrO3, containing only little amounts of metal ions, then you must recrystallize your barium bromate from boiling hot water.
Making pure HBrO3 solutions without barium ions or sulfate ions is quite laborious. I proceeded by first preparing an approximately 3M solution of
H2SO4. As I did not knbow precisely the concentration of my concentrated H2SO4 (could be anything between 95% and 98%) I decided to roughly make 3M
solution and then titrate against NaOH. I have 99+% pure NaOH and very quickly weighed this as it absorbs moisture from air quickly and made a
solution of known concentration with this. I think that I could achieve appr. 1% accuracy with this.
Weighing Ba(BrO3)2 is accurate, once you have a nice dry product. IIRC it exists as the monohydrate, when crystallized from water. I mixed solutions
of acid and a hot solution of barium bromate. I decided to take 1% extra of the acid, preferring a little sulfate as impurity over a little barium as
impurity. Getting it exactly matching is cumbersome and requires frequent probing. I did not take the effort to do that.
The biggest practical problem I ran into was that on mixing solutions of barium bromate and sulfuric acid, you get a lot of very fine white
precipitate, which is not easy to filter. Using paper is not advisable because the strongly acidic and oxidizing solution destroys the filter paper
quickly. I decided to boil the solutioin for a while to make the precipitate somewhat more compact and easier to settle and then allowed the
precipitate to settle. But settling took a long time and it did not really settle at the bottom. A tick white layer remained. I accepted the loss and
pipetted the liquid above the white precipitate.
After all this labor I finally had a solution of HBrO3, with a slight impurity of H2SO4.
I used my solution for oxidation experiments, but you could use that for making Sr(BrO3)2, by adding Sr(OH)2 or SrCO3 to the acid. Any cloudiness can
be allowed to settle (that will be SrSO4) and removed, the Sr(BrO3)2 will remain in solution, if it indeed is soluble at 300 grams per liter.
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Boffis
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@Angrylittlebunny; what an interesting idea! Well done. Did you find this reaction somewhere or invent it?
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Admagistr
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| Quote: Originally posted by woelen | HBrO3 is surprisingly stable, as long as you don't get too high a concentration. I have done experiments with it, up to 10% or so, using barium
bromate and sulfuric acid. Even boiling the solution does not destroy the acid.
Barium bromate can be made quite easily. It is not very soluble in the cold and if you have NaBrO3, then you can make it from a soluble barium salt
(e.g. BaCl2) and NaBrO3 and purify it by means of recrystallization. It can even be made from KBrO3 (which is less soluble than NaBrO3). If you want
to make a solution with HBrO3, containing only little amounts of metal ions, then you must recrystallize your barium bromate from boiling hot water.
Making pure HBrO3 solutions without barium ions or sulfate ions is quite laborious. I proceeded by first preparing an approximately 3M solution of
H2SO4. As I did not knbow precisely the concentration of my concentrated H2SO4 (could be anything between 95% and 98%) I decided to roughly make 3M
solution and then titrate against NaOH. I have 99+% pure NaOH and very quickly weighed this as it absorbs moisture from air quickly and made a
solution of known concentration with this. I think that I could achieve appr. 1% accuracy with this.
Weighing Ba(BrO3)2 is accurate, once you have a nice dry product. IIRC it exists as the monohydrate, when crystallized from water. I mixed solutions
of acid and a hot solution of barium bromate. I decided to take 1% extra of the acid, preferring a little sulfate as impurity over a little barium as
impurity. Getting it exactly matching is cumbersome and requires frequent probing. I did not take the effort to do that.
The biggest practical problem I ran into was that on mixing solutions of barium bromate and sulfuric acid, you get a lot of very fine white
precipitate, which is not easy to filter. Using paper is not advisable because the strongly acidic and oxidizing solution destroys the filter paper
quickly. I decided to boil the solutioin for a while to make the precipitate somewhat more compact and easier to settle and then allowed the
precipitate to settle. But settling took a long time and it did not really settle at the bottom. A tick white layer remained. I accepted the loss and
pipetted the liquid above the white precipitate.
After all this labor I finally had a solution of HBrO3, with a slight impurity of H2SO4.
I used my solution for oxidation experiments, but you could use that for making Sr(BrO3)2, by adding Sr(OH)2 or SrCO3 to the acid. Any cloudiness can
be allowed to settle (that will be SrSO4) and removed, the Sr(BrO3)2 will remain in solution, if it indeed is soluble at 300 grams per liter.
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@WOELEN
Great, thank you for the valuable practical information! Have you tried crystalloluminescence of Ba(BrO3)2 and Sr(BrO3)2?!
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woelen
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I did not know at that time of the crystalloluminescence of these compounds. It sounds interesting and I still have Ba(BrO3)2, so I'll give it a try.
Is it just a matter of crystallizing Ba(BrO3)2 from water and watching little flashes in the liquid on/around the crystals?
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Admagistr
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| Quote: Originally posted by woelen | | I did not know at that time of the crystalloluminescence of these compounds. It sounds interesting and I still have Ba(BrO3)2, so I'll give it a try.
Is it just a matter of crystallizing Ba(BrO3)2 from water and watching little flashes in the liquid on/around the crystals? |
Yes, that's all that needs to be done! I did a similar experiment with a mixture of sodium sulphate and potassium sulphate, and the result was very
nice! It's like a thunderstorm in nature, where you never know ahead of time when and where the next lightning will come. Also, if it ever strikes
again, or if is over. The crystallization wasn't just flashes and tiny sparks, it was a sound effect too! You just have to have patience and time,
sometimes it can take an hour before it starts
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Admagistr
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| Quote: Originally posted by woelen | | I did not know at that time of the crystalloluminescence of these compounds. It sounds interesting and I still have Ba(BrO3)2, so I'll give it a try.
Is it just a matter of crystallizing Ba(BrO3)2 from water and watching little flashes in the liquid on/around the crystals? |
Here are references from a great German chemistry forum, at least You, Wilco could understand without a translator for the relative closeness of Your
native language to German.
https://illumina-chemie.de/viewtopic.php?f=18&t=4470
https://illumina-chemie.de/viewtopic.php?f=19&t=4331
https://illumina-chemie.de/viewtopic.php?f=18&t=3662
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Oxy
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| Quote: Originally posted by woelen | | As I did not knbow precisely the concentration of my concentrated H2SO4 (could be anything between 95% and 98%) I decided to roughly make 3M solution
and then titrate against NaOH. I have 99+% pure NaOH and very quickly weighed this as it absorbs moisture from air quickly and made a solution of
known concentration with this. I think that I could achieve appr. 1% accuracy with this. |
Besides water there is often a lot of carbonate in NaOH due to reactions with water and carbon dioxide from air. Which means that making a standard
solution just by weighting some sodium hydroxide and dissolving in water is actually bad idea. Carbonates have to be removed. Then the concentration
of NaOH should be titrated by acidimetry. And only then it may be used for quantitative analysis. Otherwise, you can't be really sure about the
concentration when it comes to analytical purposes.
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Admagistr
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| Quote: Originally posted by Oxy | | Quote: Originally posted by woelen | | As I did not knbow precisely the concentration of my concentrated H2SO4 (could be anything between 95% and 98%) I decided to roughly make 3M solution
and then titrate against NaOH. I have 99+% pure NaOH and very quickly weighed this as it absorbs moisture from air quickly and made a solution of
known concentration with this. I think that I could achieve appr. 1% accuracy with this. |
Besides water there is often a lot of carbonate in NaOH due to reactions with water and carbon dioxide from air. Which means that making a standard
solution just by weighting some sodium hydroxide and dissolving in water is actually bad idea. Carbonates have to be removed. Then the concentration
of NaOH should be titrated by acidimetry. And only then it may be used for quantitative analysis. Otherwise, you can't be really sure about the
concentration when it comes to analytical purposes. |
Well that's true, NaOH and KOH mostly contain water and carbonates, I read that their content can be as much as 10% of water and carbonates! Unless my
memory deceives me... If the such hydroxides were packaged in a vacuum, or in dry argon, they would be free of this problem...
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theAngryLittleBunny
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| Quote: Originally posted by Admagistr | | Quote: Originally posted by theAngryLittleBunny | | Quote: Originally posted by Admagistr |
Hi, I find very attractive bromates Ba(BrO3)2 and Sr(BrO3)2 when crystallized from an aqueous solution there are small lightning bolts, miniature
electrical discharges. I'm going to try it...!
|
Cool, but how do you wanna make Sr(BrO3)2? It has a solubility of about 300g per liter at room temperature, which is much higher then KBrO3.
One way would be to make AgBrO3 which has a rather low solubility in water (I think less then 1g/L) and then mix that with a solution of SrCl2 to get
AgCl as a precipitate leaving Sr(BrO3)2 in solution. Anyway, let me know how it goes. |
Here's a link I'm sure you'll be interested in, maybe you already know it?
https://illumina-chemie.de/viewtopic.php?f=18&t=4470
It doesn't address the production of Sr(BrO3)2, your idea is interesting but quite laborious and complex, perhaps inefficient to implement in
practice...I thought to neutralize HBrO3 by SrCO3, it would probably be easier to implement, but I would have to find out how HBrO3 is stable and
explosive/non-explosive...What do you think about it?
I bought my Sr(BrO3)2 from Chemcraft in Russia. But I bought it in a small amount because it's not cheap...I'll think about it some more...
|
HBrO3 is only stabile in a solution of I think below 20%, the silver bromate methode seems easy because you recover the silver as just silver chloride
which can be easily turned into silver metal. 100g of silver costs I think 40 dollars and doing one run with this would in theory yield 159g of
Sr(BrO3)2, and you will get most likely a close to theoretical yield. And thanks, I'll take a look.
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theAngryLittleBunny
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| Quote: Originally posted by Boffis | | @Angrylittlebunny; what an interesting idea! Well done. Did you find this reaction somewhere or invent it? |
No I just tried it to see what happens, the reagents are all cheap anyway. I knew that Br- gets oxidized to BrO- by ClO- so I was pretty sure it would
work.
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Admagistr
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| Quote: Originally posted by theAngryLittleBunny | | Quote: Originally posted by Admagistr | | Quote: Originally posted by theAngryLittleBunny | | Quote: Originally posted by Admagistr |
Hi, I find very attractive bromates Ba(BrO3)2 and Sr(BrO3)2 when crystallized from an aqueous solution there are small lightning bolts, miniature
electrical discharges. I'm going to try it...!
|
Cool, but how do you wanna make Sr(BrO3)2? It has a solubility of about 300g per liter at room temperature, which is much higher then KBrO3.
One way would be to make AgBrO3 which has a rather low solubility in water (I think less then 1g/L) and then mix that with a solution of SrCl2 to get
AgCl as a precipitate leaving Sr(BrO3)2 in solution. Anyway, let me know how it goes. |
Here's a link I'm sure you'll be interested in, maybe you already know it?
https://illumina-chemie.de/viewtopic.php?f=18&t=4470
It doesn't address the production of Sr(BrO3)2, your idea is interesting but quite laborious and complex, perhaps inefficient to implement in
practice...I thought to neutralize HBrO3 by SrCO3, it would probably be easier to implement, but I would have to find out how HBrO3 is stable and
explosive/non-explosive...What do you think about it?
I bought my Sr(BrO3)2 from Chemcraft in Russia. But I bought it in a small amount because it's not cheap...I'll think about it some more...
|
HBrO3 is only stabile in a solution of I think below 20%, the silver bromate methode seems easy because you recover the silver as just silver chloride
which can be easily turned into silver metal. 100g of silver costs I think 40 dollars and doing one run with this would in theory yield 159g of
Sr(BrO3)2, and you will get most likely a close to theoretical yield. And thanks, I'll take a look. |
Thank you, I'll try it!
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woelen
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| Quote: Originally posted by Admagistr |
Well that's true, NaOH and KOH mostly contain water and carbonates, I read that their content can be as much as 10% of water and carbonates! Unless my
memory deceives me... If the such hydroxides were packaged in a vacuum, or in dry argon, they would be free of this problem... |
Well-packaged NaOH, which has not been exposed to air too much, actually can be fairly pure. In my titration, I certainly did not have significant
quantities of Na2CO3 in my NaOH. My NaOH does not produce any bubbles, when added to dilute acids. Commercial general lab grade NaOH (e.g. from
laboratoriumdiscounter.nl) is sold as 99% NaOH and I think that is a reasonable claim.
I know that NaOH is not a good standard for titration, but in my experiments I did not need utmost accuracy, but an acceptable result. And I think
that using my 99% NaOH would be more accurate than using H2SO4 of somewhat unknown concentration as a standard. But I agree that for good accuracy,
other standards must be used.
NaOH also usually does not contain much water, as opposed to KOH. With KOH you can have as much as 15% by weight of water (pellets or flakes). The
little granules of NaOH contain less water.
[Edited on 9-12-21 by woelen]
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Admagistr
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| Quote: Originally posted by woelen | | Quote: Originally posted by Admagistr |
Well that's true, NaOH and KOH mostly contain water and carbonates, I read that their content can be as much as 10% of water and carbonates! Unless my
memory deceives me... If the such hydroxides were packaged in a vacuum, or in dry argon, they would be free of this problem... |
Well-packaged NaOH, which has not been exposed to air too much, actually can be fairly pure. In my titration, I certainly did not have significant
quantities of Na2CO3 in my NaOH. My NaOH does not produce any bubbles, when added to dilute acids. Commercial general lab grade NaOH (e.g. from
laboratoriumdiscounter.nl) is sold as 99% NaOH and I think that is a reasonable claim.
I know that NaOH is not a good standard for titration, but in my experiments I did not need utmost accuracy, but an acceptable result. And I think
that using my 99% NaOH would be more accurate than using H2SO4 of somewhat unknown concentration as a standard. But I agree that for good accuracy,
other standards must be used.
NaOH also usually does not contain much water, as opposed to KOH. With KOH you can have as much as 15% by weight of water (pellets or flakes). The
little granules of NaOH contain less water.
[Edited on 9-12-21 by woelen] |
In my country before 1990, chemicals were packed in non-hermetically sealed containers as well as alkaline hydroxides. When I filled the empty bottle
from them with water and made very violent movements with it on my hand, always a few drops of water would squirt out... All the chemicals in these
bottles were "baked" and I had to get them out using a hammer and a big pair of scissors, or a screwdriver. They were made by Lachema in Brno. Depends
on to how hermetically sealed or leaky bottles the chemical manufacturer packed...Under socialism, there was only one state-designated manufacturer
for a particular product, like those bottles, and when the production was poorly handled, it went like this...
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Fantasma4500
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interesting, interesting.
TCCA, trichlorocyanuric acid and NaOH forms NaClO, concentrated
now, dump KBr into that and you have some neat yields
i suppose they still sell out TCCA?
this is especially interesting as KBrO3 acts like KClO3 in many ways, but may give very sensitive compositions, with red phosphorus is almost ignites
on contact
what about copper bromate? copper chromate is insoluble in water
and lets not forget about lead bromate, whats neat about lead salts is how heavy they are, super easy decantation action. on second thought, it also
drops out chlorides, and with copper- hydroxides can be turned into copper oxide with a bit of heating
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theAngryLittleBunny
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| Quote: Originally posted by Antiswat | interesting, interesting.
TCCA, trichlorocyanuric acid and NaOH forms NaClO, concentrated
now, dump KBr into that and you have some neat yields
i suppose they still sell out TCCA?
this is especially interesting as KBrO3 acts like KClO3 in many ways, but may give very sensitive compositions, with red phosphorus is almost ignites
on contact
what about copper bromate? copper chromate is insoluble in water
and lets not forget about lead bromate, whats neat about lead salts is how heavy they are, super easy decantation action. on second thought, it also
drops out chlorides, and with copper- hydroxides can be turned into copper oxide with a bit of heating |
I made another post about an explosion I had with TCCA. Mixing TCCA with bases is dangerous since it can form NCl3. If you add TCCA to an NaOH
solution it will start fizzing because of the NCl3 that is decomposing as it is formed. There may be no explosion hazard in solution, but I wouldn't
risk it. You could use the TCCA instead to make chlorine gas and lead that into a cold NaOH solution to make NaOCl which can be used like the
Ca(OCl)2.
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Boffis
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I was so intrigued by the OP's process I decided to have a go at producing sodium bromate by simply evaporating down the calcium chloride /sodium
bromate solution that results from the original reaction mixture. Since sodium bromate is the least soluble of the four possible combinations of ions,
has a steep solubility curve and both sodium and calcium chlorides are very soluble and difficult to crystallise it is possible that sodium bromate
will crystallise out in preference to other phases.
To test this I ran an experiment using the OP basic method with 20g of sodium bromide being added to a solution of 67g of calcium hypochlorite
(theoretically only about 58g are necessary but my bleaching powder is very old so my be less than the 75% claimed) in 300ml of water. Once most of
the solids had dissolved the solution was warmed and the pH was adjusted by adding 28% hydrochloric acid dropwise to about 5-6. To start with each
drop of acid caused an intense orange colour to form but this faded rapidly. After the addition a significant amount of acid (about 50-70ml not
measure accurately on this occasion) the colour change was accompanied by the liberation of a little chlorine and the pH was approaching 5. The
addition of acid was stopped at this point and the mixture simmered for 10 minute to complete the conversion of hypobomite to bromate and the cooled a
little. While still hot circa 40-50 C the mixture was vacuum filtered and the filtrate left to cool overnight. A slight film of crystals floating on
the surface of the 350-400ml of liquor had formed but nothing else.
The cold mixture was filtered again with the addition of a little Keiselguhr (celite) and then slowly evaporated down to about 150-170ml and allowed
to cool. A white crystalline ppt formed, this was removed by filtration but not isolated further. The evaporation was continued and further white ppt
formed then quite suddenly at about 70-80ml the slurry turned orange and began to effervesce, it smelled of a halogen but the gas proved to be mainly
oxygen. A little cold water was added to quench the reaction and it was allowed to cool to room temperature. The slightly orange crystalline ppt was
filtered off and sucked dry but not washed.
The last filtrate was mixed with 25ml of saturated KCl brine and left to stand, a small amount of glassy crystalline material began to crystallise
out.
The orangish filter cake (including the initial cake) was dissolved in 50ml of hot water treated with celite and filtered hot. I am still waiting to
see if anything crystallises from the filtrate before I try adding KCl to it.
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Fantasma4500
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you may try to knock it out by adding EtOH
http://chemister.ru/Database/properties-en.php?dbid=1&id...
as you see here CaCl2 is rather soluble in EtOH
http://chemister.ru/Database/properties-en.php?dbid=1&id...
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Boffis
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Well; once I had got the orange coloured cake into solution, in 50ml of hot water, nothing crystallised on cooling. The solution was mixed with an
equal volume of saturated KCl solution but still nothing crystallised out, even after 24hours at 5 C. When a small parts was mixed with a little
dilute sulphuric acid a little bromine separated so some bromate was present but not enough to ppt KBrO3. This should not be taken as an indication
that the technique doesn't work since I am pretty sure that I grossly overheated the mixture during evaporation. The viscous calcium chloride filtrate
was also treated with 25ml of saturated potassium chloride solution as mentioned above and about 1.1g of crude K bromate slowly crystallised out.
I am going to try this again using the OP method using a direct ppt of KBrO3 with saturated Potassium chloride solution and also with barium chloride
solution to get the sparingly soluble barium bromate.
[Edited on 16-12-2021 by Boffis]
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Fantasma4500
International Hazard
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thats the kind of yield you should write into a note on achievements
how doable is it to make calcium hypochlorite yourself? they dont sell it in europe
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Fery
National Hazard
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Location: Czechoslovakia
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Antiswat - what, EU restricted to sell Ca(OCl)2 too? I'm not watching bad news at all.
In my country I'm still able to buy large packs, like 10 kg 70% Ca(OCl)2 for 50 EUR
https://fichema.cz/chlornan-vpenat-caclo2-cas/1410-chlor-sok...
How much stable is Ca(OCl)2? Maybe the right time to buy it while it is still available. NaClO is still sold here too - in big quantities, for
swimming pool owners (at least every second house with big enough garden in my town has a swimming pool). With which will it be substituted if EU
denies hypochlorites? Then using persulfates, perborates, percarbonates, ozonization devices, UV lamp devices?
Decades ago I made KIO3 from KI and KMnO4. I wonder if this route works for KBrO3 too or not (I mean alkaline medium, not acidic where Br2 is produced
instead)?
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woelen
Super Administrator
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Ca(ClO)2 is not restricted in the EU, as far as I know. In NL you also still can buy this without any problem.
It, however, is replaced by TCCA and Na-DCCA more and more, for a good reason. The latter are much more stable and safer to handle (although they also
can be quite risky).
I myself had 2 kg of Ca(ClO)2, but it decomposed and became a nasty sticky mess in a year or so. It also caused corrosion of a lot of nearby items and
I decided not to buy it anymore, unless I have an immediate use for it.
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