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Chemist_Cup_Noodles
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[*] posted on 28-4-2016 at 07:54
Alum Crystals Thread


So I've just recently learned about what the group of compounds known as the alums are. I've seen the word alum thrown around a lot, most often preceded by potassium and/or aluminum, but never knew exactly what it meant. Thanks to @crystalgrower's neato website I now know exactly what they are! This group of compounds can certainly have a lot of beautiful combinations, so I'm very intrigued.

However, I do have one question about these. For an alum to form correctly, can it be comprised of any cation in a +1 and +3 oxidation state? If it can this opens a world of opportunities. If you can, could it be possible to maybe react a copper (I) sulfate and aluminum sulfate to get a copper (I) alum? I know that copper sulfate exists mostly with copper (II) but isn't copper (I) sulfate theoretically possible?

A few common +3 oxidation state cations I think would make some neat crystals are Co+3, Cr+3, Al+3, and Fe+3. Some +1 oxidation state cations I think might produce interesting results are NH4 +1, K+1, Na+1, and maybe even Cu+1. I know that the cuprous form of copper isn't super common, but it's not very hard to get it to the +1 state and I've seen it done several times on SciMad.
I have some ideas for some different combinations I want to try though. I'll list them: CoK, CrCu, FeNH4, CoNH4, CrNH4, or AlCu.

No idea if they'll work or not, but oh what the hell, why not.
If you have any experience with these, pictures, or perhaps produced rather interesting crystals, please share!

I will note though, I usually do a lot of reading and fact checking before making a post, reply, or topic but for this one I'm kind of throwing a healthy amount of conjecture at the wall.





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[*] posted on 28-4-2016 at 08:13


Cu(I) is only stable in the form of insoluble compounds- you aren't going to get an alum with it.



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[*] posted on 28-4-2016 at 08:43


Bear in mind also that the solubility of alums tends to decrease with increasing size of the M+ ion. For that reason many Na+/Li+ alums cannot be crystallised, because they are too soluble.

Other 'exotic' alums include Rb/Mn3+ and Cs/Ti3+.

Never heard of a Co3+ based alum. Co(+3) is a powerful oxidiser (oxidises chloride to chlorine), so may be not stable enough.




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[*] posted on 28-4-2016 at 08:56


Quote: Originally posted by DraconicAcid  
Cu(I) is only stable in the form of insoluble compounds- you aren't going to get an alum with it.

Damn, I was afraid it would be too unstable. If it weren't I'm sure it would have produced some beautiful crystals. Oh well, would a potassium cobalt (III) alum still be possible, or ammonium cobalt (III) alum? The +3 state of cobalt seems to be more stable than the +1 state of copper, but the +2 state of cobalt still seems much more optimal than the +3. I haven't seen where the +3 state can form crystals though so I'm losing hope for that. It seems like it forms a lot of different Werner complexes though, so that's interesting.




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[*] posted on 28-4-2016 at 09:24


http://pubs.acs.org/doi/abs/10.1021/ed079p958

Yes, you can make cobalt(III) alums. You just need a really big cation.

ETA: Why do you think Cu(I) would give beautiful crystals? Cu(I) is a d10 system and will be colourless in the absence of charge-transfer or Cu(II) impurities. That's why pure CuCl, CuBr, CuI, etc are all white powders.

[Edited on 28-4-2016 by DraconicAcid]




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[*] posted on 28-4-2016 at 09:27


Maybe that's a right time to attempt making some unconvetional alums ;).
I will let you know if I succesfully make some pretty alum :).




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[*] posted on 28-4-2016 at 10:41


Quote: Originally posted by DraconicAcid  
http://pubs.acs.org/doi/abs/10.1021/ed079p958

Yes, you can make cobalt(III) alums. You just need a really big cation.

ETA: Why do you think Cu(I) would give beautiful crystals? Cu(I) is a d10 system and will be colourless in the absence of charge-transfer or Cu(II) impurities. That's why pure CuCl, CuBr, CuI, etc are all white powders.

[Edited on 28-4-2016 by DraconicAcid]


Oh, I'm glad you've found an article on cobalt alums! Too bad you have to pay to see it though. I don't have an ACS membership yet so I can't even see the reduced price. But from the abstract it seems that a cobalt ammonium alum is possible, and I also found it on another neat site. This site also lists a that a potassium cobalt alum has been made before, by someone named "Marshall". So I know it showed a cesium cobalt alum in the article abstract, but now I don't know if a cation that big is really necessary if it's evidently worked with potassium (although it was only a sourceless blurb, so I'm a little skeptical actually).

But about Cu(I), I really didn't think or know that it was colorless. I just knew that Cu (II) makes a lot of great blues and greens, and that there was cuprous sulfite in Chevreul's salt that lent to the deep red color, but I suppose that is also due to the fact that cupric sulfite is present also. To be honest I've never worked with copper in it's first oxidation state. And like I said in my post, this is mostly me just throwing around conjecture, so thank you for correcting me (I hate how sarcastic that sounds in my head but it's sincere so oh well).

On another note, I searched up the CsTi alum, and apparently it has paramagnetic properties which is pretty cool. I saw a thread elsewhere on SM where a guy tried but failed at making a titanium alum, so I didn't think it would work. And could manganese alums be something else to look into? I didn't initially think they would work because I did remember how strong of an oxidizer Mn3+ was, and it's rather uncommon to see it in that state. But I was reading the page in Encyclopedia Britannica on it, and it had manganese listed so maybe it does work.




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[*] posted on 28-4-2016 at 11:52


Yes, you can make an alum with Mn(III).

https://books.google.ca/books?id=xnsAAAAAMAAJ&pg=PA621&a...

I have a prep for manganese(III) sulphate at work- I'll give you the reference later.

ETA: Vogel's Quantitative Inorganic Analysis sez...

Preparation and Standardization of manganic sulphate solution.

To prepare a c. 0.07 N solution...

To 50 mL of a solution of A. R. manganous sulphate (15.1 g in 1 L of 6N-sulphuric acid) add 3 mL of conc. sulphuric acid with water-cooling; then add 12 mL of 0.5N potassium permanganate solution, 2 mL at a time, at intervals of about three minutes. After 8 mL and again after 12 mL of the permanganate solution have been added, cautiously introduce a further 2 mL of conc. sulphuric acid. Store the solution for 4 hours before use.

Note that a 0.5N solution of potassium permanganate corresponds to a 0.1 mol/L solution, or about 16 g per litre.

[Edited on 28-4-2016 by DraconicAcid]




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[*] posted on 28-4-2016 at 11:53


Quote: Originally posted by Chemist_Cup_Noodles  
I saw a thread elsewhere on SM where a guy tried but failed at making a titanium alum, so I didn't think it would work. And could manganese alums be something else to look into? I didn't initially think they would work because I did remember how strong of an oxidizer Mn3+ was, and it's rather uncommon to see it in that state. But I was reading the page in Encyclopedia Britannica on it, and it had manganese listed so maybe it does work.


That 'guy' was me. It would have worked had I tried the Cs/Ti alumn. The K/Ti didn't work and the NH4/Ti was a bit inconclusive. But forget about making nice, big crystals from these: that would be much harder to do.

Despite Mn(+3) being a powerful oxidiser, reports on Rb/Cs Mn3+ alums can be found.




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[*] posted on 28-4-2016 at 15:15


I think I'm going to attempt to make ammonium iron (III) sulfate first. It doesn't seem too awful to make, all I need to do is first make iron (II) sulfate with iron + excess sulfuric acid, then add nitric acid that I just made to that to oxide it to iron (III) sulfate Fe3(SO4)2. Making ammonium sulfate should be straightforward enough. Then I'll just stoichiometrically combine the two, and crystallize for the alum. Shouldn't be too bad.

But does anyone have any ideas on how to test for iron (II) vs iron (III)?




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[*] posted on 28-4-2016 at 16:19


Quote: Originally posted by Chemist_Cup_Noodles  
I think I'm going to attempt to make ammonium iron (III) sulfate first. It doesn't seem too awful to make, all I need to do is first make iron (II) sulfate with iron + excess sulfuric acid, then add nitric acid that I just made to that to oxide it to iron (III) sulfate Fe3(SO4)2. Making ammonium sulfate should be straightforward enough. Then I'll just stoichiometrically combine the two, and crystallize for the alum. Shouldn't be too bad.

But does anyone have any ideas on how to test for iron (II) vs iron (III)?


The easiest way to make sure your iron is iron(III) is to precipitate it as the hydroxide and leave the suspension exposed to air for a while. Then you can dissolve it in sulphuric acid.




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[*] posted on 28-4-2016 at 17:03


Quote: Originally posted by DraconicAcid  

The easiest way to make sure your iron is iron(III) is to precipitate it as the hydroxide and leave the suspension exposed to air for a while. Then you can dissolve it in sulphuric acid.


That's a lot slower than you might think. It also requires constant stirring of the slurry.

There's a preparatory thread on FAA:

http://www.sciencemadness.org/talk/viewthread.php?tid=5650#p...




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[*] posted on 30-4-2016 at 12:08


Quote: Originally posted by Chemist_Cup_Noodles  

If you have any experience with these, pictures, or perhaps produced rather interesting crystals, please share!


Here's a crystal of 'ordinary' KAl(SO4)2.12H2O that started as a seed perhaps a year ago... in the meantime it has grown a few more mm :D

The other alum I've grown crystals from is chrome alum, and this can be mixed with aluminium alum in all proportions, just like a dye.

Those more exotic alums are very interesting, especially cobalt! I have some cobalt(II)chloride which I'll have to convert to the sulphate first.


alum_in_hand2.jpg - 228kB
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[*] posted on 30-4-2016 at 12:24


@cucurbit
Your alums are absolutely gorgeous!
Looking forward to cobalt alum ;).




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[*] posted on 30-4-2016 at 13:25


Thanks Crystal Grower! Same specimen as in your nice thread in 'miscellaneous'.
I now have 4 big ones still growing, it's like having a pet or a plant that requires a bit of maintenance but nothing too demanding ;)

The cobalt variety will take some time, as usual with crystals... but I'm very curious about the color. Chrome alum is fascinating: in cold daylight it's violet, but in candle light it's blood red, like a ruby! Who knows what cobalt is going to do. But they will never be as big as my current Big Four!
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[*] posted on 30-4-2016 at 14:53


I wish I'd bookmarked some of the double-salts I'd ran into recently - I assume you don't mind talking about double salts in general because they, being the broader type of solid which includes alums, sometimes have the same ease of crystal-making. I could've swore I saw a magnesium-something double salt with chloride or sulfate which made big crystals. It was also a kind of mineral...

Ah, Magnesium Potassium Sulfate (cite). Also, magnesium potassium chloride and zinc ammonium chloride, both with 1:1 of the two cations. I haven't tried any of these, though. And ferrous and Ferric ammonium sulfates are cool, I own small samples of both. Very pretty.




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[*] posted on 10-5-2016 at 14:15


Quote: Originally posted by The Volatile Chemist  
I wish I'd bookmarked some of the double-salts I'd ran into recently - I assume you don't mind talking about double salts in general because they, being the broader type of solid which includes alums, sometimes have the same ease of crystal-making. I could've swore I saw a magnesium-something double salt with chloride or sulfate which made big crystals. It was also a kind of mineral...

Ah, Magnesium Potassium Sulfate (cite). Also, magnesium potassium chloride and zinc ammonium chloride, both with 1:1 of the two cations. I haven't tried any of these, though. And ferrous and Ferric ammonium sulfates are cool, I own small samples of both. Very pretty.


Hm, another thread on this board where someone wondered if a copper and ferrous sulfate double salt got me searching around. While I knew that is was not possible, my searching around lead me to a wikipedia page on Tutton's salts which seem fairly similar to the alums. The page also has a quite wonderful table of a variety of the salts with plenty of sources. Many of the exotic salts that seem a little out of reach of the average amateur chemist though. The magnesium potassium sulfate you've cited seems to actually occur as a natural mineral too, neato.




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[*] posted on 17-5-2016 at 13:13


Thanks for that link about Tutton's salts, very useful!

I knew about copper ammonium sulphate because of an old book, my 'bible' when I was a kid, 70's... "Scheikunde thuis" by H.L. Heys. My mom threw it away after an explosion in the garage, but the chapter about crystal growing is etched into my brain!

Just made a bit of cobalt sulfate solution (from CoCl2.6H2O and excess Na2CO3, dissolve the carbonate in dilute sulfuric acid)... the beautiful red filtrate is now evaporating.
Next step: combine with potassium sulfate to get the double salt.

Just discovered a video about growing crystals of cobalt(II) ammonium sulphate - never mind the delicious heavy accent ;)
https://youtu.be/nsL_EYIEoG4
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[*] posted on 17-5-2016 at 15:25


Quite the interesting double salt, seemingly vaguely analogous to the Chromium(III) Potassium Sulfate double salt with regards to its properties and coloration.



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[*] posted on 28-5-2016 at 09:34


Hello everyone,
I've got problem with my potassium aluminium alum growing I need help with. My solution is always getting after few days some white floating weird stuff at surface. I think it's mold. I tried warming the solution and filtering it afterwards and with no change. I also tried UV-B light and also nothing.

My idea is to add little bit of sulfuric acid, but I'm little bit scared of interrupting the crystalization.
I would really appreciate help.
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[*] posted on 28-5-2016 at 11:27


A little H2SO4 is very unlikely to disrupt crystallisation but likely to disrupt bacterial growth.



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[*] posted on 28-5-2016 at 12:13


Sulfuric acid present in the solution often cause crystallization of lower ,hydrate than one which would crystallize from normal water solution.
But I don't know exactly how it would affect the crystallization of KAl(SO4)2.




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[*] posted on 28-5-2016 at 14:10


I don't think it's mold or bacteria, they should have nothing to feed on in the solution. More likely it's aluminium hydroxide hydrate gel, due to hydrolysis.
I had some of these flakes in my solution but it didn't seem to harm the crystals much. Sudden temperature rises are more disastrous :o

To be sure: a few drops of sulfuric acid should help acidify the solution a bit - not too much. And regular filtering, once or twice a week!
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[*] posted on 29-5-2016 at 00:15


Okey, thanks for ideas, I'm gonna try little bit of sulfuric acid.



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[*] posted on 3-6-2016 at 12:30


Quote: Originally posted by Cucurbit  
I don't think it's mold or bacteria, they should have nothing to feed on in the solution. More likely it's aluminium hydroxide hydrate gel, due to hydrolysis.
I had some of these flakes in my solution but it didn't seem to harm the crystals much. Sudden temperature rises are more disastrous :o

To be sure: a few drops of sulfuric acid should help acidify the solution a bit - not too much. And regular filtering, once or twice a week!

Mold can grow in a lot of stuff, and can get sufficient resources from dust, depending upon its location. If you don't add too much H2SO4, then the lesser hydrate, I wouldn't think, would crystallize out. You could always do a prep run.




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