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Brain&Force
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[*] posted on 22-9-2015 at 20:25


Here's samarium and holmium sulfamate in solution. As far as I know holmium sulfamate hasn't been characterized - Google results for "holmium sulfamate" turn up literally nothing.

<img src=http://i.imgur.com/gmSROtB.jpg width=800>
<img src=http://i.imgur.com/DGOTmIs.jpg width=800>

[Edited on 23.9.2015 by Brain&Force]




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[*] posted on 24-9-2015 at 04:41


Nice ! Couldn't find anything on it, either. But I don't have that much
literature on f-Metals. Maybe in a very old or very large book.


Due to the amount of Copper Sulphate I recentely bought I started
to do some copper chemistry. The combined rests were collected in
a 1L Erlenmeyer Flask and today it had that color. And the cam really
picked it up correctely for once :D The beautiful blue hue varies from the top to the very bottom. From a dark blue to a very bright blue and then a
greenish blue. Really nice color, guess due to the Hydroxdie an Ammonia.

The rest was an attempt to make Copper(I)Bromide from Cu(II), Sodium Sulfite and KBr. I got that slightley yellow ( not really good on cam) compound that now after a while turned green. So seemes to have worked but due to the contamination with Cu(II) I already put it to the waste.

Last there is an old experiment I wanted to do again, mixing a Cu(II) Salt with NaI which forms the instable Cu(II)Iodide that will form Cu(I) and elemental Iodine. I added a bit of Heptane and let the layers seperate. Since it wasn't good on cam I got a bit of the Heptane layer into a Pipette, so there is the purple layer with the Iodine dissolved on the top and a brown, now after washing it 3 times with Water white, cloudy solution.



Copper.jpg - 191kB

[Edited on 24-9-2015 by fluorescence]
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[*] posted on 24-9-2015 at 05:26


Titania Nanotubes (TNTs). Gotta love the acronyms people come up with for nano-structures.
These were anodically grown in an ethylene glycol electrolyte containing a small amount of ammonium flouride and water.

TNTs.jpg - 101kB TNTs2.jpg - 102kB
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[*] posted on 25-9-2015 at 06:39


Here is some Bismuth Chemistry:

From the left to right:

Bismuth(III)Iodide, a dark compound that formes when Iodide is added to a Bismuth(III)-Solution.

Tetraiodobismutate, an orange coordination compound that formes if you continue to add Iodide. At the bottom are the Bismuth(III)Nitrate Crystals I used as Bismuth Source, problem is that it's really insoluble. Luckily, traces of Bismuth are enough to from these colored compounds.

Bismuth-Chromate ? Dunno what that is. Added some Potassium Dichromate to a Bismuth(III) solution. Immedialtely yellow flakes appeared. Looks quite cool but I can't really find anything on that compound.

Colin MacKenzie says it's Bismuth(III)Chromate in his work "One thousand processes in manufactures and experiments in chemistry"
But that book is from 1825 and so I doubt the analytical background a bit.

The Pigment Compendium talks about a Basic Dichromate that is formed when the Dichromate is added. Since I used some Dichromate instead of Chromate I believe this compound is more acurate than the one before,
the formula here is given by Bi2O3 x 2 CrO3. I guess I had luck that I chose the Dichromate Bottle instead of the Chromate bottle (which was just closer at that time xD ) since only that compound is described in the book.





Bismuth Chemistry.jpg - 124kB
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[*] posted on 25-9-2015 at 07:15


And even more Bismuth Chemistry coming up, just the stuff I still had on my mind but I'll check literature for more.

So here is on the left a very famous compound, it's a Bismuth-Thiourea-Complex usually formed during the Bismuth-Slide. I dunno if that experiment is still common in the lab but it's very good to find Bismuth in contaminated mixtures. What you prepare is a filter paper and you fold that so you get like a little canal where a liquid can travel throught. Then you add your substance in the middle and the rest is stacked on top creating a little pile, you add Sodium Fluoride ( to get away Aluminium and Iron ) then a Chloride ( to get away silver and mercury ) and Sodiumpotassium Tartrate (Antimony and Tin ) and finally Thiourea. Then you add a bit of Nitric Acid and let it slide though the canal into the mixture. What's happening is that all other Elements that could form a yellow compound with Thiourea will form a stable salt or coordination compound like the Iron Flouride and thus create a yellow color ( if you ever did the experiment where you add Fluoride to an Iron-Thiocyante Complex you know how stable the colorless Iron-Fluoride Complex is). And so only Bismuth will hopefully show a yellow color. I did that with just adding a bit of Thiourea and Nitric Acid to a bit of Bismuth Nitrate solution.


On the right is the Dimethylglyoxime-Bismuth Compound formed, when those are added in Ethanol/Water and afterwards a bit of Ammonia is added. It will take some time but after a while and if the conc. of Ammonia is high enough you'll get a yellow solution.



Bismuth 2.jpg - 55kB
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[*] posted on 26-9-2015 at 17:28
Long TNP crystals


take a look at the size of ths picric acid crystal garden i grew. these crystals are huge. the longest one is 4in plus.
This synthesis has become one of my favourites to perform recently.

IMG_1292.JPG - 1MB




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[*] posted on 28-9-2015 at 03:56


What a beautiful purple color...if I only knew what it was.
That happens when you take the wrong metals...should have
taken Manganese but took Chromium accidentally when that formed.

Dunno. Should be some kind of K3[(CN)5NO]. At least it would have been it if I used manganese. There is that Cr(I)-Complex but that one is supposed to be green, not purple and my Manganese one should be purple but is green....I have no clue what's going here.

But the Educt was highly saturated with Cyanide usually not attacked by Bases and yet the insoluble green to yellow Cyanide Salt changed in a second to a soluble purple solution...too fast to be the Aqua-Ligand and if there is Cyanide in it I really doubt it anyway...but what is it xD.



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[*] posted on 30-9-2015 at 01:01


Organic chemists be like "picturez of white substancez much".



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[*] posted on 1-10-2015 at 05:43


Had to do some Iron Chemistry for a Presentation. One of it was to make
Prussian Blue. That's all the wastes combined....



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[*] posted on 3-10-2015 at 23:41


Today I found a Uranium glass bowl in a second hand shop for $2 as well as some chemistry books, if I can find my UV light I will post a pic of it fluorescing

DSC_1865.jpg - 2.5MB

[Edited on 4-10-2015 by NedsHead]
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[*] posted on 5-10-2015 at 11:14


Ferric alum. I made it from ferrous sulfate, baking soda, bleach, sulfuric acid and ammonia solution.
First precipitated FeCO3 by soda, oxidized it with bleach, washed sediment of Fe(III) oxides/hydroxides, then dissolved it in 43% acid, added (NH4)2SO4 and re-crystallized.

Alas, I was unable to get clear crystals like I saw somewhere on this forum. Never expected such color from Fe(III) salt.

square-DSC01330.JPG - 307kBfront-singleDSC01337.JPG - 459kB
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[*] posted on 5-10-2015 at 11:21


Quote: Originally posted by Dmishin  
Ferric alum. I made it from ferrous sulfate, baking soda, bleach, sulfuric acid and ammonia solution.
First precipitated FeCO3 by soda, oxidized it with bleach, washed sediment of Fe(III) oxides/hydroxides, then dissolved it in 43% acid, added (NH4)2SO4 and re-crystallized. Alas, I was unable to get clear crystals like I saw somewhere on this forum. Never expected such color from Fe(III) salt.

This is absolutely gorgeous! How long did that thing grow for? Looks like the size of a golf ball.




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[*] posted on 5-10-2015 at 11:29


That is a beautiful crystal.

Please post a step-by-step procedure so we can all make one.




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[*] posted on 5-10-2015 at 12:06


Quote: Originally posted by zts16  
Quote: Originally posted by The Volatile Chemist  
Interesting. I was slightly suprised to find barium had an insoluble citrate. This fact salvaged an attempt to make barium hydroxide (which didn't precipitate well enough to collect). Like the bad chemist I am, I just threw some citric acid into solution and fitered the precipitate. So I at least got one compound for my barium collecton. Which is starting to get some size :)
You have a barium collection? Must be very white! :P
Try making barium manganate if you haven't yet though. I've tried it a few times with partial success. It's supposed to be dark blue, but mine keeps turning brown about a day after I wash and dry it.

Yeah, I'm pretty white too. :P I have Barium Potassium ferrocyanide crystals, which are straw colored. Yes, mostly white, though, but I have made only about 7 compounds, it's hardly a collection. By the way, will nitrate oxidise sulfite readily? I made some 'barium sulfite' from barium nitrate, but I realized later the sulfite could've just oxidized...
I don't have any manganate salts yet, haven't purchased any. I recently acquired a few more chemicals from the hardware and grocery stores nearby (NaOH, for the first time, finally! (Though I couldn't get Ba(OH)2 to precipitate :/), also CaCl2 as pickling agent, anhydrous, and Citric acid.) I'm trying to exhaust my lab of things I can make before getting any more reagents, so I probably won't try that till next summer. But It's a good Idea, thanks. Hopefully I can get it to work.

Fluorescence, I'm not well educated on the 'nitroferrocyanide' (?) complex ion. Is it formed when a nitrate salt is introduced to a solution of ferrocyanide or ferricyanide? I made some barium potassium ferrocyanide from barium nitrate and potassium ferrocyanide, but now I wonder if the crystals precipitated contained this ion instead. The literature called for barium chloride, which I did not have.




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[*] posted on 7-10-2015 at 10:44


Quote: Originally posted by Detonationology  

This is absolutely gorgeous! How long did that thing grow for? Looks like the size of a golf ball.

Thanks! Not that long: only 3 weeks, and I took measures to slow down growth (covered the beaker). This compound can grow really fast thanks to it great solubility: 124g/100ml.

Quote: Originally posted by aga  

Please post a step-by-step procedure so we can all make one.

Nothing special, actually. I explained it briefly above. But, as you wish.

My goal was to use only compounds, available outside of specialized chemicals store. So, my preparation is surely not the easiest way to get this salt.

Source compounds:
  • Ferrous sulfate FeSO4*7H2O, fertilizer grade
  • Baking soda NaHCO3, food grade
  • Bleach NaClO solution, 5-15% according to labeling.
  • Sulfuric acid 43%, electrolyte
  • Ammonia solution 10%, from a drug store (the most limited reagent for me).


Prepare (NH4)2SO4:
Accurately and quickly pour calculated amount of H2SO4 into ammonia solution while stirring constantly. With 10% ammonia and 43% acid, solution becomes very hot (80 - 90 C), but not boiling. Beware boiling, burns, cracking glass and NH3 fumes while reaction is not yet complete.

Prepare Fe2(SO4)3
1. Dissolve FeSO4 in the minimal amount of hot water, add soda in small portions. A lot of CO2 is produced, beware excessive foaming. The sediment is ferrous carbonates and hydroxides.
2. Without washing, oxidize the sediment with bleach. Add bleach in big portions (it is quite dilute). Sediment becomes brown, gas (CO2) is produced (FeCO3 oxidizes to Fe(III) oxides/hydroxides and CO2 is left). The reaction is slow (hours), heat it mildly on water bath. After finishing reaction, let it stay, decant the excess solution and add new portion of bleach. Stop, when decanted liquid starts smelling chlorine. The reaction is:

2Fe(II) + NaClO -> 2Fe(III) + NaCl

For me, it required 400% of calculated amount of bleach. Either my bleach was too dilute, or there were side reactions, consuming hypochlorite (ClO3 formation?).

3. Wash the sediment, remove the excess water, then add sulfuric acid (stoichiometric amount + small excess). It dissolves very badly: took 2 days. Heating on hot water bath improves it. Filter the product.
You'll get dark brown solution of Fe(III) sulfate.

Prepare the double salt
Mix ammonium and ferric sulfates in equimolar amounts, let the solution evaporate (avoid strong heating, it causes hydrolysis). Ferric alum forms almost colorless crystals. Collect them and utilize the last small portion of solution that contains excess chemicals and impurities.

Then grow the crystals in the regular way. My preparation is lengthy, but it avoids uncommon reagents, hazardous gases, need to boil down large amounts of solutions and separate significant impurities.

[Edited on 7-10-2015 by Dmishin]
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[*] posted on 11-10-2015 at 11:54


Very nice, good write-up.



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[*] posted on 13-10-2015 at 00:17


got a couple nice potassium trisoxalatoferrate(III) crystals. all it took was some patients and my previous solution yielded much better samples. I noticed a few clear crystals in the mix, so reagent use wasn't dead on. but that doesn't bug me in light of the results :) . the color is much more intense in person, but my phone was balancing it poorly against the background of a white door. nice emerald green

K,Fe oxalate 1.jpg - 775kB K,Fe oxalate 2.jpg - 822kB k,Fe oxalate 3.jpg - 790kB
and a zirconia disk I will use for a standoff while torching precious metals
zirconia disk.jpg - 899kB
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[*] posted on 22-10-2015 at 23:40


The Uranium glass bowl I found in a second hand shop for $2 under UV light

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[*] posted on 24-10-2015 at 05:56


Magnalium polumna firecracker at sunset:

[Edited on 24-10-2015 by greenlight]

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[*] posted on 24-10-2015 at 12:44


NedsHeads, nice! The potassium trisoxalatoferrate(III) crystals are huge! We made some nice ones in AP chem, but the teacher made us use a mortar on them....



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[*] posted on 25-10-2015 at 01:11


Monoclinic crystals of ammonium zinc sulfate hexahydrate. Belongs to the family of Tutton's salts. Shiny as brilliants, my photos can't show it fully.

(NH4)2Zn(SO4)2·6H2O

There is one trick to get very clear crystals: after preparing saturated solution, add baking soda (NaHCO3) to it, 0.5g of soda per 100g solution, and stir. The reaction will start, producing some gas (CO2) and sediment (mostly ZnCO3). Let it stay for a day, then filter the solution and use it for growing. This procedure drastically improves quality of crystals.
Authors of this procedure assumed that it precipitates impurities (Al and Fe3+), though I have some doubts. Maybe, crystal growth is affected by lower pH, or by ions of Na.

DSC01385.JPG - 596kB DSC01439.JPG - 530kB

[Edited on 25-10-2015 by Dmishin]
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[*] posted on 25-10-2015 at 14:48


Those are crazy big... and I have a lot of zinc sulfate... Too bad I don't have a fume hood to quickly make some ammonium sulfate.



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[*] posted on 26-10-2015 at 11:24


Potassium explosion on the top of water

potassium explos..jpg - 233kB
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[*] posted on 26-10-2015 at 12:16


Quote: Originally posted by The Volatile Chemist  
Too bad I don't have a fume hood to quickly make some ammonium sulfate.


If you are going to try 2NH3+H2SO4, then it is not really needed, reaction finishes instantly, if acid is sufficient and reagents are mixed well. However, if your reagents are concentrated (at least, more concentrated than 10% NH3 and 43% H2SO4, I used), mix could start boiling.

Ammonium sulfate can also be bought as fertilizer (haven't seen though), and of course, chemicals shops have it, in good purity.

If you have a lot of ZnSO4, and don't mind wasting half of Zn, using fertilizer ammonium nitrate (or chloride) is an option. Double Zn-NH4 sulfate is much less soluble than its constituents. Mix concentrated, hot solutions of ZnSO4 and NH4NO3, and double salt will precipitate:

2ZnSO4 + 2NH4NO3 + 6H2O -> (NH4)2Zn(SO4)2*6H2O ↓ + Zn(NO3)2

(haven't tried exactly this reaction, but it must work)

[Edited on 26-10-2015 by Dmishin]
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[*] posted on 26-10-2015 at 18:00


Quote: Originally posted by Aqua-regia  
Potassium explosion on the top of water

That is pretty cool. How did you capture that shot?
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