Sciencemadness Discussion Board

Pretty Pictures (2)

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The Volatile Chemist - 30-8-2015 at 14:07

nice, nitro-genes. What method did you use to produce it?
Violet sin, how do you evaporate your solutions? I've had 250mL of Copper(II) acetate soltn. sitting for 2 weeks. It hasn't evaporated more than 50 - 100 mL. I don't really want super-nice crystals, just the product. But boiling hydrolyses the acetate to hydroxide and acetic acid. Any suggestions?

nitro-genes - 31-8-2015 at 12:10

Some pretty cool photo's in this thread, and thanks, it was a precursor to 4-diazo 2,6 dinitrophenol, described here: https://www.sciencemadness.org/whisper/viewthread.php?tid=63... :)

aga - 31-8-2015 at 12:37

photos. not photo's.

the ' signifies Possession in this context.

Do you want to be Possessed by the evil chemical spirit of a photo ?

[Edited on 31-8-2015 by aga]

nitro-genes - 31-8-2015 at 12:53

I AM possessed, you grammar nazi :P

aga - 31-8-2015 at 13:01

A : I mixed NO with a Lung
B : I mixed NO2 with a Lung
C : i mixed N2O with a Lung.

Which subject lived, which ones died ?

The one that got the TYPING correct or one of the others ?

Accuracy in typing your Own Language will give you more credence in what you write.

Getting it obviously wrong will tarnish what you type, even if what you're saying is amazing.

This becomes important when you want to make some serious $.

nitro-genes - 31-8-2015 at 13:25

All three would die or already have died, because mixing implies you either take out the lung, or have one in a jar, prior to you mixing them with NOx. "Exposed" is the word you are obvbiously looking for :)

And no, English (wow, look..a capital) is not my native language :)



[Edited on 31-8-2015 by nitro-genes]

aga - 31-8-2015 at 13:36

Wow !

Your English is incredibly good then, as many Natural English speakers commonly put an apostrophe (') in the wrong place.

What is your Mother Tounge ?

There are two ways to take a Correction :-

React aggressively and learn nothing

Accept the correction and learn something.

DraconicAcid - 31-8-2015 at 14:09

Quote: Originally posted by aga  
Wow !

Your English is incredibly good then, as many Natural English speakers commonly put an apostrophe (') in the wrong place.

What is your Mother Tounge ?

There are two ways to take a Correction :-

React aggressively and learn nothing

Accept the correction and learn something.


Mother tongue (not tounge) shouldn't be capitalized, and neither should correction.

aga - 31-8-2015 at 14:46

Very good catch for an American speaker.

Personally i like inappropriate Capitalisation, Lots.

Z is not really a substitute in Correct English.

Tire is an expression of weariness, and has absolutely nothing to do with Rubber.

[Edited on 31-8-2015 by aga]

anyway's'
im Tyre'd and need to got some's' sleeps's'es'

Back at you in 7~8 hours, like as not, methinks sirra'h.

[Edited on 31-8-2015 by aga]

Tsjerk - 31-8-2015 at 22:45

None of the three subjects died, as they all mixed one of the NxO/NOxs with a certain person called Lung.

Sulfur crystalization from DCM

Sulaiman - 1-9-2015 at 14:19

after reading the thread about sulfur in hot toluene
and since I just received 5l of 99.9% DCM, I thought I'd try sulfur in DCM ...

method: excess sulfur in 50 ml boiling DCM, filtered, cooled to room temperature.

Sulfur is not very soluble in DCM but enough to form a few small crystals:



DCM_Sulfur.jpg - 786kB

sorry about the poor quality photo'

P.S. I discovered that I do not like the smell of DCM.

[Edited on 1-9-2015 by Sulaiman]

j_sum1 - 1-9-2015 at 14:42

Yeah. You have a bit of camera shake there.
What I am finding interesting is the difference in the shape of crystals produced depending on the solvent used. I have not come across this phenomenon before. Does anyone have any insight into this?

(I am sensing a project looming.)

The Volatile Chemist - 4-9-2015 at 12:41

I recently realized that toluene is''''''''nt pronounced tooline...oops...
It's commonly noted about sulfur, but I don't recall any good explanation of why, or any other examples.

aga - 4-9-2015 at 12:47

You get nice transparent rhombic sulphur crystals in liquid disulphur dichloride exposed to air for a while.

Extracted from the liquid, they immediately form a dusting of powdery yellow sulphur on them.

If there was a solvent that took the S2Sl2 off them before exposing to air, they'd be a Thing to behold - glittery sulphur ! (maybe).

aga - 4-9-2015 at 12:57

Quote: Originally posted by The Volatile Chemist  
I recently realized that toluene is''''''''nt pronounced tooline...oops..

Tool-een is probably a lubricant sold to car mechanics.

Tol-you-een (even Tol-oo-een) is more like it.

While i'm ranting, how did Soul-der-ing become Sod-er-ing ?

Soldering should never involve the same parts of the body as the word Sod implies.

Remember that the iron is Very HOT.

Edit:

Imagine the scene :-
You, having just spent a Lot of money on new copper pipes for your house, and have painstakingly placed them exactly where they need to be, and the plumber arrives to soLder them together for you.

Him : "Nice pipes. I'll go right ahead and Sodder them up for you."
You: "What ?!?! You're going to Bugger up my pipes ?!?!"

Doesn't work for me, at all.

[Edited on 4-9-2015 by aga]

DraconicAcid - 4-9-2015 at 12:58

I had a friend in grad studies who could not shake his habit of calling it "tol-you-leen".

aga - 4-9-2015 at 13:35

I had a friend once as well.

Lost long ago in a terrible drinking accident.

Praxichys - 10-9-2015 at 14:54

Refluxing S2Cl2

P1000225.JPG - 4.7MB

The Volatile Chemist - 11-9-2015 at 12:21

Ooh, pretty. Looks like really concentrated Potassium ferricyanide. :) Why are you refluxing it? Or is it combined with something else?

fluorescence - 18-9-2015 at 02:10

Literature said I should use the Acetate but that is really hard to make
if you want it very fresh, at least it's not working with Acetate for me and if you follow literature it takes ages to do that. So I thought I'd try it out with Citrate for a couple of reasons and it worked out quite well. Still not perfect but I need to heat in during the reaction and I did it now cold so seems to work.

So here are solutions I needed and since there wasn't much about them on the net I thought I'd share a picture. This is Cobalt-Citrate and on the right Chromium(III)Citrate. The cam didn't pick it up that well so I made a little square with the real color. That Chromium looks really cool.



Citrate.jpg - 68kB

The Volatile Chemist - 19-9-2015 at 14:01

Huh, very nice. The Cobalt Citrate is a similar color to Cobalt chloride?

fluorescence - 20-9-2015 at 06:39

Funnily it changed colour over night. I guess it slowely turned into
the Aqua Complex which is purple. Might have to add more citric acid, dunno. I'll have to prepare it fresh anyways so it's just waste now but it looks quite cool.

Here is the updated version:

From left to right:

- Nickel-Citrate

- Chromium(III)-Citrate ???

- Cobalt-Citrate



Citrate d-metal.jpg - 53kB

gdflp - 21-9-2015 at 05:14

Chromium(III) is notorious for taking obnoxiously long times to exchange ligands, so that's likely what happened. Perhaps some citrate ions coordinated to the chromium as well.

The Volatile Chemist - 22-9-2015 at 18:31

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 :)

Texium (zts16) - 22-9-2015 at 18:38

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.

Brain&Force - 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]

fluorescence - 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]

Squall181 - 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

fluorescence - 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

fluorescence - 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

Long TNP crystals

NeonPulse - 26-9-2015 at 17:28

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

fluorescence - 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.



whatsiit.jpg - 35kB

Eddygp - 30-9-2015 at 01:01

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

fluorescence - 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....



Im blue dabedidabedae.jpg - 34kB

NedsHead - 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]

Dmishin - 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

Detonationology - 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.

aga - 5-10-2015 at 11:29

That is a beautiful crystal.

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

The Volatile Chemist - 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.

Dmishin - 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:


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]

The Volatile Chemist - 11-10-2015 at 11:54

Very nice, good write-up.

violet sin - 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

NedsHead - 22-10-2015 at 23:40

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

DSC_1998.jpg - 1.2MB

greenlight - 24-10-2015 at 05:56

Magnalium polumna firecracker at sunset:

[Edited on 24-10-2015 by greenlight]

Screenshot_2015-10-24-20-15-00.png - 2.7MB

The Volatile Chemist - 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....

Dmishin - 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]

The Volatile Chemist - 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.

Aqua-regia - 26-10-2015 at 11:24

Potassium explosion on the top of water

potassium explos..jpg - 233kB

Dmishin - 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]

j_sum1 - 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?

Aqua-regia - 27-10-2015 at 09:51

Easy, cause was a handycam video very close I gained from this a foto with vlc media player

aga - 27-10-2015 at 10:26

Quote: Originally posted by NedsHead  
The Uranium glass bowl I found in a second hand shop for $2

How did you know it had Uranium in it ?

NedsHead - 27-10-2015 at 16:25

Quote: Originally posted by aga  
[/rquote]
How did you know it had Uranium in it ?


The shop had a heap of Uranium glassware, probably all came from the same collection, the bowl was an odd piece that didn't match any of the sets in the collection but looked to be the same glass

violet sin - 27-10-2015 at 16:56

in the past I have rewired a couple dollar store LED flashlights with UV LED's for identification purposes. They work great. Just unsolder the crappy white ones and replace with eBay versions. There are keychain version of UV lights that can be used as purchased, occasionally found randomly.

The Volatile Chemist - 27-10-2015 at 17:14

Quote: Originally posted by violet sin  
in the past I have rewired a couple dollar store LED flashlights with UV LED's for identification purposes. They work great. Just unsolder the crappy white ones and replace with eBay versions. There are keychain version of UV lights that can be used as purchased, occasionally found randomly.

I had one for a while from a kid kit, but at twelve, everything was better when hooked up to a nine volt battery...The LED turned a great violet color for a second, got very bright, then burnt out. Oh well, AS&S sells $8 UV lamps, which is crazy cheap. Gonna buy some when I get my next order.

violet sin - 27-10-2015 at 17:34

You can get a 50pc pack of UV LED's for less than that. I finally whittled my pile down to 5 or so. Time to buy more I guess. Took me quite a few years and many small projects.

arkoma - 27-10-2015 at 17:45

A Feller and his tripleneck
IMG_20151026_115521_795_zpsdzaz8v8r.jpg - 119kB
w00t
IMG_20151026_113732_299.jpg - 227kB
View out the "front" door
IMG_20151021_172452_600.jpg - 680kB

[Edited on 10-28-2015 by arkoma]

violet sin - 27-10-2015 at 18:38

Awe man, that.looks.like great fishing! Jealous a lill bit.

arkoma - 27-10-2015 at 19:05

33°18'14.66"N
93°43'50.55"W

Eddygp - 28-10-2015 at 06:51

He is back!! :D

Brain&Force - 28-10-2015 at 11:49

The absorption spectrum of holmium sulfamate, I used an LED and diffraction grating with the solution to get a spectrum and photographed it with my phone.

<blockquote class="imgur-embed-pub" lang="en" data-id="appmfzf"><a href="//imgur.com/appmfzf">Taken with a diffraction grating, camera, and LED light.</a></blockquote><script async src="//s.imgur.com/min/embed.js" charset="utf-8"></script>

JJay - 28-10-2015 at 12:49

Quote: Originally posted by arkoma  
A Feller and his tripleneck

w00t

View out the "front" door


[Edited on 10-28-2015 by arkoma]


How do you keep your thermometer out of the stirring vortex?

violet sin - 28-10-2015 at 12:56

What LED did you use to do that? Also, is the diffraction grating just a bit of CD cut to a convenient shape? It's pretty cool! I had wondered if there was a convenient file to burn to a disk to provide a solid background for the grating or if a blank CD is used. I remember the kits people bought that used a piece of CD/DVD for the grating, but of course there are readymade versions.

If you don't mind, could you post a pic of the setup you used? I find this interesting. Amongst other things I need to clean up, I have a 100ml or so of CoCl2 that would definitely be colored enough to show. I bet the solution I got from leaching Osage orange sawdust would be interesting as well. Bright yellow and it fluoresces green/yellow/orange depending on a few factors. Roots, bark or wood and concentration strength. Thanks for the share

The Volatile Chemist - 28-10-2015 at 15:00

Arkoma, nice place, great glassware :)
Quote: Originally posted by Dmishin  
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]

I was referring to the fume hood for catching gasses if I were to boil off the water. I hate having to do this if any of the reagents are acids.
If this procedure does work, I'll have to try it (with ammonium chloride and zinc sulfate. Not because I particularly desire ammonium chloride, but I do like the production of double salts, and this one sounds interesting.

nitro-genes - 31-10-2015 at 15:43

Added some hot ammonium dinitrosalicylate solution to an erlenmeyer containing some other waste solutions of nitrophenols, upon cooling down this was the result, covered the entire bottom. :)

Waste crystals.jpg - 129kB

[Edited on 31-10-2015 by nitro-genes]

Dmishin - 5-11-2015 at 13:45

Birefringence in a triglycine sulfate crystal.

biref-bigDSC01293.JPG - 338kB

[Edited on 5-11-2015 by Dmishin]

Tsjerk - 6-11-2015 at 05:45

Nice crystal Dmishin,

Would this crystal separate a non-polarized laser beam in two? Making the two beams polarized?

Firmware21 - 6-11-2015 at 06:33

The joy of playing in a lab that makes tons of fluorescent compounds. These are NDIs :D

*Obviously under U.V light.



[Edited on 6-11-2015 by Firmware21]

Dem TLC.JPG - 1.3MB

The Volatile Chemist - 10-11-2015 at 15:22

Both of the last two photos are sweet. Chromatography of fluorescent compounds... :)
I'm going to try to make some thin plates for chromatography this weekend.

greenlight - 17-11-2015 at 02:57

Potassium chlorate crystals found in my cell electrolyte after being in the fridge for two days:

[Edited on 17-11-2015 by greenlight]

20151117_165843.jpg - 3MB

NedsHead - 17-11-2015 at 05:07

Nice Potassium chlorate crystals greenlight, I haven't seen chlorate crystals like that before, the structure looks similar to Bismuth

greenlight - 17-11-2015 at 06:35

They definitely are Potassium chlorate. After electrolysis of Potassium chloride the cell contents were left at room temp for 2 days and then put in the fridge and left for a further three days to precipitate the rest.
The whole top layer was a blanket of these joined crystals.

The Volatile Chemist - 17-11-2015 at 19:47

Nice. Sounds like a successful prep. Try decomposing some in a flame, if you haven't already :)

Zephyr - 17-11-2015 at 21:54

Copper phosphate finally dried out:


IMG_2498.JPG - 202kB

TheIdeanator - 18-11-2015 at 00:04

Quote: Originally posted by greenlight  
Potassium chlorate crystals found in my cell electrolyte after being in the fridge for two days:

[Edited on 17-11-2015 by greenlight]


They make hopper crystals? Cooooool.

The Volatile Chemist - 19-11-2015 at 10:28

Quote: Originally posted by Pinkhippo11  
Copper phosphate finally dried out:

How'd you make it? Or do you own phosphoric acid...?

Texium (zts16) - 19-11-2015 at 15:25

Quote: Originally posted by Pinkhippo11  
Copper phosphate finally dried out:
Are you sure that's copper phosphate? I've never seen it make crystals like that... or at all really. I've made it a couple times and it's always precipitated as a fine, blue powder.

Zephyr - 19-11-2015 at 16:34

I made it by reacting copper carbonate with phosphoric acid and drying it slowly in a desiccator over sodium hydroxide.

Texium (zts16) - 19-11-2015 at 17:54

Perhaps it is a soluble acid salt then, like CuHPO4
Analyzing the ratio of phosphate ions to copper ions in it would clarify it.

The Volatile Chemist - 20-11-2015 at 13:42

Nonetheless, very pretty. And if you can stoichiometrically analyze it (titration with NaOH?) then you've got a fine compound to vial. As long as it's all the same salt, of course...

arkoma - 20-11-2015 at 20:29

18514d56-4c5d-44aa-9ffa-63c2721f1da4_zpsdewfdojl.jpg - 366kB

first reflux in my new space. pecan shells in acetone.

The Volatile Chemist - 21-11-2015 at 06:15

Cool 'hotplate' ! Wish mine had two plates, though I've seen ones with more too.
By the way, pretty pictures (1) ended at 40 pages. This thread has 40 now. Is it time for another, or did new software fix the problems with having a 40pg. thread?

plastics - 24-11-2015 at 06:23

Just trying out my 'new' reconditioned Edwards 2 stage vacuum pump





IMG_0907.JPG - 203kB

[Edited on 24-11-2015 by plastics]

The Volatile Chemist - 28-11-2015 at 08:58

Wow, very nice. Is that just the readout, or the entire device?
Did you purchase it new, or did you 'acquire' it used?

CrystalCage - 1-12-2015 at 18:40

The first pic is the flat alum crystal
The second one is a perfectly octahedral alum crystal :P
And the third one is sodium bicarbonate (baking soda crystal)

The crystals are homemade. :))

738372_559768857395932_1320046660_o.jpg - 102kB1606207_643642112341939_2064944133_o.jpg - 143kB10682377_789568847749264_1265765932151861459_o.jpg - 296kB

The Volatile Chemist - 2-12-2015 at 12:32

Nice. What's the second crystal inside of?

CrystalCage - 3-12-2015 at 05:21

It is potassium alum crystal. :)

LargeV - 4-12-2015 at 13:51

Quote: Originally posted by CrystalCage  
It is potassium alum crystal. :)

He asked what you were storing the crystal in the second picture in

CrystalCage - 6-12-2015 at 00:12

Oh, my apologies. The red one holding the crystal is the plastic base from graduated cylinder (my graduated cylinder broke :'( ) contained in a glass jar.

https://static.fishersci.com/images/FS106312~wl.jpg


[Edited on 6-12-2015 by CrystalCage]

stibium - 10-12-2015 at 07:45

Phosphorescent alkaline earth sulfides activated with bismuth.
From left to right:
BaS / SrS / CaS

phosphorescent sulfides.jpg - 1.1MB

Dmishin - 10-12-2015 at 13:09

Quote: Originally posted by CrystalCage  

The second one is a perfectly octahedral alum crystal :P

Nice one! But I should warn you that putting do-not-eat bags with crystal defeats the purpose of putting it into sealed container.

For sulfates, most common problem is not absorption of moisture from air, but dehydration.
Alum is very stable, but in extremely dry environment it could start loosing water and turn into white powder. For storing crystals, I would recommend do the opposite: put into the sealed jar a piece of tissue, wet with saturated solution of the compound. It would create just right humidity level. Used this method to preserve uncoated crystal of ferric alum - works great, crystal has not changed for months.
Quote: Originally posted by CrystalCage  

And the third one is sodium bicarbonate (baking soda crystal)

I have very strong suspicion that it sesquicarbonate (trona), not bicarbonate. Did you grew it from hot solution? Interesting result, by the way.

[Edited on 10-12-2015 by Dmishin]

CrystalCage - 11-12-2015 at 08:23

The alum crystal is kept for 2 years in that conditions. But anyway, I will do your method in preserving my copper acetate crystal, which is still growing. :)

I grew that star-shaped crystal in a hot solution... But that crystal is very fragile. It crushed when I picked it up.

arkoma - 19-12-2015 at 11:00

I love element 29. reduced Cu, CuSO4, and CuAc across top, and crystallizing CuCl2 on the glass. My lil Kodak doesn't focus well close up *sigh* as the chloride is bee-you-ti-ful

100_1369.JPG - 2.3MB

Brain&Force - 20-12-2015 at 13:56

Element 28 is calling to me right about now...

here's some nickel sulfamate I'm probably going to use for electroplating. I just need to find a nickel electrode.

<img src="http://i.imgur.com/JHxrO6B.jpg" width=800>

Sulaiman - 20-12-2015 at 14:04

For small scale nickel electrodes, http://www.ebay.co.uk/itm/30PCS-Pure-99-96-Low-Resistance-Ni...

cyanureeves - 20-12-2015 at 14:49

Quote: Originally posted by Sulaiman  
For small scale nickel electrodes, http://www.ebay.co.uk/itm/30PCS-Pure-99-96-Low-Resistance-Ni...
been waiting for 1 1/2 months for delivery of those nickel strips.have to wait until jan, 01 2016 before i ask for money back.

PHILOU Zrealone - 21-12-2015 at 04:58

Quote: Originally posted by stibium  
Phosphorescent alkaline earth sulfides activated with bismuth.
From left to right:
BaS / SrS / CaS

@Stibium,
Magnificent!
How did you do/make those? Via United State Patent n°US2544507A or else?
How is that phosphorescence activated, sunlight, UV light/laser and how long exposure?
How long does the phosphorescence lasts?

Texium (zts16) - 22-12-2015 at 21:47

This evening I made some 3-nitrophthalic acid for future production of luminol. I noticed that one of the soluble byproducts of the nitration exhibits yellow-green fluorescence. I got a good picture of it:

Fluorescence.jpg - 559kB
My tablet's camera makes the fluorescence appear more yellow than it does in person, it's actually quite green. By any chance, does anyone know what the cause of it is?

PHILOU Zrealone - 23-12-2015 at 14:09

Quote: Originally posted by zts16  
This evening I made some 3-nitrophthalic acid for future production of luminol. I noticed that one of the soluble byproducts of the nitration exhibits yellow-green fluorescence. I got a good picture of it:


My tablet's camera makes the fluorescence appear more yellow than it does in person, it's actually quite green. By any chance, does anyone know what the cause of it is?

Could it be m-nitrophtalic acid or m-nitrophtalate anion?

Texium (zts16) - 23-12-2015 at 17:20

There's definitely a substantial amount of 4-nitrophthalic acid in that solution, but I don't know if that's what is causing the fluorescence. I haven't been able to find much out there about its properties. It might also be a trace amount of a dinitrophthalic acid, as I did unfortunately allow the temperature to rise too high for about half a minute during the nitration.
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