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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?

Hi,
I have always fascinated the chemistry of phosphorescence.
Here there is a digitized copy of a spanish old book specialized on phosphorescent alkaline earth sulfides:
https://archive.org/details/revistadelareal01madgoog

In the PDF file watch the pages 145 to 188, 406 to 465 and 578 to 640

I made my sulfides with help of this book.
Basically I made a suspension in water or ethanol intimately mixing 100 gr. of alkaline earth carbonate, 30 g. of sulfur, 1 gr. of sodium carbonate, 0.5 g. of sodium chloride and 0.02 gr. or less of bismuth subnitrate. I heated this mixture to evaporate the solvent, and put 10 gr of mixture inside a covered crucible placed on a gas burner for 30 minutes.

The phosphorescence can be activated by sunlight or UV light. If the material is activated by exposure to sunlight five minutes, the phosphorescence lasts half an hour.
To get better results the calcination temperature must be about 1000 or 1200 ºC

In these other links you can find more information about making Inorganic phosphorescent materials:

Alkaline earth sulfide phosphorescent pigments:
http://www.google.com/patents/US2544507
Phosphorescent Calcium Sulphide:
http://calcium.atomistry.com/calcium_monosulphide.html
Luminescence (easy synthesis):
http://www.teralab.co.uk/Experiments/Luminescence/Luminescen...
Many pictures and recipes:
https://www.flickr.com/photos/28617364@N04/9457791470/in/pho...
Phosphorescent materials (Spanish):
http://www.cientificosaficionados.com/tecnicas/fosforos.htm
Luminescence properties of CaS:Bi phosphors with special reference to fluxes:
http://shodhganga.inflibnet.ac.in/handle/10603/39477

Mabus - 28-12-2015 at 12:16

Last year I added these cut steel pipes from some fridge scrubbers in some hydrochloric acid to get rid of the iron and get the copper. Little iron dissolved and instead I got these small light blue crystals, about 1 mm in size. They don't seem to be soluble in water and they're pretty well stuck on the iron pipe.

Methanol Distillation

Starcruiser - 29-12-2015 at 13:45

I haven`t identified locally no other source of methanol other than the commonly available windshield washer fluid. It should be about 40-45% conc. but I suspect that this is not really the case as long as the distillation started at a quite high temperature (~ 80 C).

The Volatile Chemist - 30-12-2015 at 16:44

Really nice photo of the luminol precursor, zts16. Quite a set-up starcruiser!

Rhodochrosite

Hegi - 31-12-2015 at 01:40

Maybe for those of you who love minerals

For more photos from my collection, visit my page CHEM.PIECEOFSCIENCE.COM

<img src="http://chem.pieceofscience.com/wp-content/uploads/2015/12/rodochrozit1.png" alt="rhodochrosite" height="600" width="900">

j_sum1 - 31-12-2015 at 03:05

Nice crystal Hegi. Kind of Christmassy with the red and green.

fluorescence - 31-12-2015 at 04:15

So I have no idea what this is probably some sort of Iron(III)Fluoride-Hydrate, due to its green color but could be anything else, too.

So I've been cleaning my crucibles that I brought home from the university over the holidays to clean them in the lab. That's how it started...

If you want to dissolve metals you boil them in Aqua Regia, if you want to get rid of Minerals, and that's what our expert for mineralogy at the university told me once, you boil it in a mixture of H₂SO₄ and HF. So I took the crucible added some conc. HF to it and left it for 2 days. It dissolved quite much of the compounds and probablya lot of my corundum crucible, too.

Then I changed it to the conc. HF + conc. Sulph Acid mixture and as that didn't help I took it a step further and even added some fluoride to it.
No it's hard to find anything on the solubility of HF in H₂SO₄ but it's something around 38% depending on the temperature. And as it was getting warmer I had a constant stream of dry, gaseous HF coming of for 48h ! So I checked it every few hours and it was always fuming off some HF. The mixture worked wonders and dissolved most of the stuff but I had placed everything in that stainless steel bowl and due to the cold temperatures it's quite humid in my lab so the bowl slowly filled with humid HF vapours ...

I have to say I used that bowl for like 4 years now. Boiled acids, Lyes and lot's of burning stuff in there. It has been heated to hot temperatures so often, left with conc. Acids for days and there was never ever any sign of rust, corrosion or any effect on it. It took the HF like 2 days to cause what you see on the pictures. After one day those green crystals had formed and back then those were really beautiful crystals not like the amorphous thing you see there.

I'm not even mad but just impressed how fast everything went. Not sure what it is but since there mostly fumes of HF present it should be some mixture of Iron - Sulphate, Fluoride, Hydrate whatever. Really beautiful. I dried it and it yielded some brown and light green powders that are insoluble in water. So looks like the Fluoride to me.

The Volatile Chemist - 31-12-2015 at 11:11

That's vague;y remnicient of my basement sink after my usage of it for a few years...

PHILOU Zrealone - 1-1-2016 at 13:01

Quote: Originally posted by stibium

image upload no compression

Quote: Originally posted by stibium

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

Thank you for those valuable infos...

crystal grower - 3-1-2016 at 13:15

upload a picture

image hosting over 10mb

My FeSO4.7H2O crystal druze .

its about 1,2 inch across.
image quality isnt wery good

ill try to make better photos next time.

Btw my 1st post

[Edited on 3-1-2016 by crystal grower]

[Edited on 4-1-2016 by crystal grower]

crystal grower - 3-1-2016 at 13:25

These are my first bismuth crystals:

host images

picture hosting

gif image hosting

bismuth is my no.1 metal

[Edited on 3-1-2016 by crystal grower]

aga - 3-1-2016 at 13:38

Very nice indeed.

Care to post your process ?

CharlieA - 3-1-2016 at 14:22

Very impressive Bi crystals!

crystal grower - 4-1-2016 at 00:28

Quote: Originally posted by aga

Very nice indeed.

Care to post your process ?

Thanks for compliment

The crystal growing process:

1. I created my Iron(II) sulfate by mixing 99.99% pure electrolytic iron with diluted sulfuric acid.(I didnt buy green vitriol bcouse I prefer making the compounds myself

).

2.There is a problem with making iron sulfate crystals becouse some brown silt creates in the solution after some period of time . it is probably some iron oxide so I come up with idea that you only need to add some extra sulfuric acid (it should react with the silt and create the FeSO4 again) - it looks it worked properly in my case .

3.I grew some seed crystals by cooling the heated saturated solution and hung the best of them onto nylon.

4.then I drowned it into solution and after several days the crystals have created

PS: sorry for my english feel free to correct me - I will learn something beside the chemistry at least

.

[Edited on 4-1-2016 by crystal grower]

[Edited on 4-1-2016 by crystal grower]

arkoma - 5-1-2016 at 17:53

Did an alkaline cleavage of piperine with KOH. My lil Kodak doiesn't have a macro mode darn it so the pics don't do justice to the BEAUTY of the potassium piperate.

[Edited on 1-6-2016 by arkoma]

JJay - 5-1-2016 at 23:32

It looks like you have several grams there. Does it taste like pepper?

arkoma - 6-1-2016 at 15:05

haven't tasted it, the piperine base sure did.

The Volatile Chemist - 6-1-2016 at 16:01

Nice picture arkoma. Perhaps you should be wary, though, of posting pictures containing your address, unless you don't mind.

arkoma - 6-1-2016 at 16:04

sulaiman was kind enough to U2U me about it, BUT, I really got squat to hide and it would take 3 whole MINUTES to edit the photo. Fort Meade ALREADY knows where we all live anyways. (Ft Meade=NSA HDQTRS)

The Volatile Chemist - 6-1-2016 at 16:58

Quote: Originally posted by arkoma

sulaiman was kind enough to U2U me about it, BUT, I really got squat to hide and it would take 3 whole MINUTES to edit the photo. Fort Meade ALREADY knows where we all live anyways. (Ft Meade=NSA HDQTRS)

Oh, yeah, they definitely do...

I already put lab pics in the lab folder, but I really think my lab is pretty

Anyways, Some chemical testing with ferro/ferricyanide and zinc cyanide.
Also the blue product is degraded Tetraamminecopper(II) acetate.

[Edited on 1-7-2016 by The Volatile Chemist]

NedsHead - 8-1-2016 at 21:14

Fresh batch of Ammonium Nitrate

greenlight - 9-1-2016 at 10:56

@Nedshead, nice looking AN.
Just curious, what method did you use to make it?

[Edited on 9-1-2016 by greenlight]

NedsHead - 9-1-2016 at 15:35

I used the acid-base method greenlight, I had a small amount of WFNA I needed to get rid of and I diluted it in some water and neutralised it with 25% ammonia solution, it's the best method I've tried so far, no filtering or recrystallising needed

crystal grower - 10-1-2016 at 03:16

My Vanadinite

print screen windows
I know that it isnt first in this topic but its just avesome.

[Edited on 10-1-2016 by crystal grower]

crystal grower - 10-1-2016 at 03:22

how to screen capture

Guess the name of this one

HINT: Its mineral ,nothing artificial.

mayko - 10-1-2016 at 10:21

Nickel, cobalt, and copper salts in various stages of crystallization and drying:

A moonscape left by the crust of a drying phosphate solution:

IMG_20150528_phosphateCorpse.jpg - 1.1MB

The reaction of zinc and sulfur, and the sponge-shaped zinc sulfide residue:

btw, I'm working on organizing my photo files, and I'm planning on putting together a portfolio on Wikimedia Commons soon!

[Edited on 10-1-2016 by mayko]

crystal grower - 11-1-2016 at 07:43

Diamond grinding wheel under magnification.

The Volatile Chemist - 12-1-2016 at 14:01

Nice photos mayko, what cobalt salt is that?

crystal grower - 21-1-2016 at 05:27

CopperII Sulphate (growth time: 2weeks)

[
Original post in chemicals for crystal growing.

[Edited on 21-1-2016 by crystal grower]

The Volatile Chemist - 21-1-2016 at 07:06

Quite a large crystal there. If you don't mind destroying it, dehydrating it in conc. Sulfuric acid might be interesting. Though of course that is a nice crystal!

crystal grower - 21-1-2016 at 07:59

Quite a large crystal there. If you don't mind destroying it, dehydrating it in conc. Sulfuric acid might be interesting. Though of course that is a nice crystal!

Oh that would be too big sacrifice destroying this one.
I will maybe try it with smaller piece.

MrHomeScientist - 21-1-2016 at 08:19

Holy crap, 2 weeks? That's huge! I spent several months growing one that only ended up about an inch across. How did you grow yours?

Shameless plug: I did the crystal dehydration experiment and posted a video to my channel: https://www.youtube.com/watch?v=J1zwFwmANw4
It probably didn't take nearly as long as I left it for, but after going white the shape never changed!

The Volatile Chemist - 21-1-2016 at 08:34

Quote: Originally posted by MrHomeScientist

Holy crap, 2 weeks? That's huge! I spent several months growing one that only ended up about an inch across. How did you grow yours?

Shameless plug: I did the crystal dehydration experiment and posted a video to my channel: https://www.youtube.com/watch?v=J1zwFwmANw4
It probably didn't take nearly as long as I left it for, but after going white the shape never changed!

Your video's the only reason I knew sulfuric acid could dehydrate CuSO₄. It's worth the watch. I'd be interested to see if it dehydrates other compounds non-destructively, though I suppose they would have to be sulfates.

crystal grower - 21-1-2016 at 08:46

Quote: Originally posted by MrHomeScientist

Nice video.
Procedure:
1.make seed crystal (I think its not neccessary to explain.)
2.when u have seed crystal attached on a string take it out of solution.
3.heat the solution and saturate it.
4. Put the crystal inside and let it slowly cool down.
5.Take the crystal out of solution and heat it again
6.saturate it ,put crystal inside and repeat the procedure.
I hope it will help you

arkoma - 21-1-2016 at 09:37

CrystalGrower--they are AWESOME.

crystal grower - 21-1-2016 at 11:34

Thnx

MrHomeScientist - 21-1-2016 at 13:20

Ah ha! So a more 'active' crystal growing process than usual, very nice. What sort of string do you use? I'd want something transparent with the same refractive index as CuSO4, so it'd be invisible.

Thanks for liking my video everybody

I need to get back to posting.

crystal grower - 21-1-2016 at 13:45

Quote: Originally posted by MrHomeScientist

I need to get back to posting.

Yeah I boost it a little

I use nylon string which is practically invisible.
But I must warn you, if you use this method, parasite crystals grow quite fast.(altough not always) and u must remove them carefully.

fluorescence - 22-1-2016 at 10:40

Had to prepare this at the University last week. Took a while to make it and only gave quite poor yield but it's really pure. The NMR for it is in the General Question's NMR post. Quite interesting stuff but not for me, our iron-based chemist commissioned it for his research.

If some are interested I can translate our preparation method if I have more time. The smell is really unpleasend !

crystal grower - 23-1-2016 at 10:07

lump of pure antimony.

gluon47 - 23-1-2016 at 14:20

detonation of MEKP.

kinda looks like a fire breathing dog if you examine closely.

Detonationology - 23-1-2016 at 15:47

Quote: Originally posted by gluon47

detonation of MEKP.

kinda looks like a fire breathing dog if you examine closely.

Detonation? You would have to have a high speed camera to capture that. Perhaps you meant deflagration?

violet sin - 23-1-2016 at 18:05

not 100% chem related, but it certainly opens doors

My dad had a nice gift for me. There was a storage container on his job site that the owner wanted cleaned out. 2x set of oxy/acet gauge, 2x welding head, 3x oxy/acet hose pairs.

I also got a bunch of copper. A whole copper downspout, 6' section of gutter, some random flashing and 2' of 1" pipe. They were cleaning up a 3 year project for a really wealthy man.

As you can see I have been leveling up my high temp ability

Disposable propane torches, refillable propane torch, disposable oxy/map, presto-o-lite acet/air torch with a few regulators n tips, 5x Oct/acet hoses, 5x gauge valve, 4x welding heads, 1x cutting head, 1 B-tank acet, larger oxy/acet tank pair w/ cart. Not pictured is a box of gauges, nitrogen/CO2 cylinder, and the lab bench top furnace w/ graphite crucible- Lindbergh solar basic.

Before registering here, I only had a disposable propane plumbers torch, and a newly acquired oxy/map pictured. Just goes to show what you can do over the course of a couple years, I'd you keep watching for deals. ~300$ total spent on the pictured items, $150 was the big tanks, with two welding heads, a cutting head, goggles, gauges, hose and the metal cart.

Done sourcing heat, next up shielding gasses.

fluorescence - 27-1-2016 at 10:44

Not really sure what it is but it's quite cool.

Very old literature (hard to find) talks about Berlin Green. Now since Berlin Blue is the German name for Prussian Blue it's the green version of it but not Prussian Green since Prussican Green unfortunately is a name for mixture of Dyes. So Berlin Green should form, according to the literature I use for my research on Cyanide chemistry by various methods, heating Prussian Blue in the dark, adding chlorine, nitric acid, ...

Now in the very old literature they talk about Fe^III[Fe^III(CN)₆] but according to a paper I found that isn't the case and there is no formula for this stuff. It's just an intermediate of possible oxidation states which is instable and will build prussian blue again. Still it looks quite cool.

On the left is prussian blue at 12°C and on the right at boiling temperatures. Both contain conc. Nitric Acid. I'm not sure if the same happens if no Acid is added but for the Acid thing it turns really bright green and if you leave it to cool for 10 min. it's blue again. I did it 3 times and it always changes to green while hot and then back to blue although I think the more often you do it the longer it will stay green. At least it looks quite different now than before.

Might try some other Oxidants tomorrow.

BlackDragon2712 - 27-1-2016 at 18:02

some nice acetaminophen crystals

Detonationology - 27-1-2016 at 18:15

Beautiful diamond shaped crystals! What solvent are you using?

BlackDragon2712 - 27-1-2016 at 18:45

95% ethanol, I hope to be posting tomorrow my attempts in making 3-nitro-4-acetamidophenol via nitrous acid

gluon47 - 28-1-2016 at 21:57

Quote: Originally posted by Detonationology

Detonation? You would have to have a high speed camera to capture that. Perhaps you meant deflagration?

Deflagration. Yeah that's what i meant. I can never remember that word

crystal grower - 31-1-2016 at 01:21

Sodium bisulfate (original post in chemicals for crystal growing).

Btw. I have become "hazard to self "

.

[Edited on 31-1-2016 by crystal grower]

Gooferking Science - 31-1-2016 at 08:37

Some pictures of the inner grid of my fusor project. This is not at full vacuum or voltage. When the turbomolecular pump gets up to full speed, the plasma extinguishes and it isn't near as pretty. So I took these pretty pictures at a slighly higher pressure.

crystal grower - 31-1-2016 at 09:28

Quote: Originally posted by Gooferking Science

Some pictures of the inner grid of my fusor project. This is not at full vacuum or voltage. When the turbomolecular pump gets up to full speed, the plasma extinguishes and it isn't near as pretty. So I took these pretty pictures at a slighly higher pressure.

WoW thats really impressive.

fluorescence - 2-2-2016 at 15:07

Quite hard to see but I think the blueish stuff should actually be Manganate(V). As you may know it can be prepared by simply adding Sulfite to a very alkaline solution of Permanganate at cold temperatures. Problem is it will appear green as is mixes with produced Mn(IV). And the Manganate(VI) is green, too. So quite hard to distinguish them. That's why I thought of something else. To make the Manganate(VI) you heat NaOH and Permanganate dry together. Now I added dry Sulfite to the mixture as well.

So in here is Sodium Sulfite, Potassium Permanganate and Potassium Hydroxide, all dry and then heated with a torch. Some yellow/brown Manganesedioxide has formed but on one side you see an intense blue color which is probably the Manganate(V). I have to admit that I have never seen it before. Might try different ratios till I get the mixture and temperature correctely.

j_sum1 - 2-2-2016 at 15:42

@fluorescence
Nice work. Pretty colours. If you refine the procedure, please post. as it is, this is a nice little demo.

fluorescence - 3-2-2016 at 06:55

Ground the components in a mortar, that is how it looked after some minutes. It started to turn dark green where the KOH powder took up water.

When I heated in it formed a liquid so I tried to cover the glass with it. It may look blue but its actually a mixture of blue-green. So might be because its transparent or simply because it still has too much Mn(VI) in there but looks really nice.

The Volatile Chemist - 3-2-2016 at 16:08

Indeed, certainly pretty colors!
If I remember, I have some nice bacterial pictures I took at 1000x, using a Grahm stain. Lots of fun.

wg48 - 5-2-2016 at 20:32

Pic of a glycerol complex of copper. Its more purple when viewed directly.

PHILOU Zrealone - 6-2-2016 at 03:44

Pic of a glycerol complex of copper. Its more purple when viewed directly.

Could you be more specific? What complex and weight ratio?

crystal grower - 6-2-2016 at 06:12

My little collection of minerals, synthetic minerals + some elements from my newly made collection.

Velzee - 6-2-2016 at 09:09

My little collection of minerals, synthetic minerals + some elements from my newly made collection.

Wow! I have to make me some K2Cr2O7 soon(when I get chromates). Do you have any radioactive minerals?

[Edited on 2/6/2016 by Velzee]

crystal grower - 6-2-2016 at 11:39

Quote: Originally posted by Velzee

My little collection of minerals, synthetic minerals + some elements from my newly made collection.

Wow! I have to make me some K2Cr2O7 soon(when I get chromates). Do you have any radioactive minerals?

[Edited on 2/6/2016 by Velzee]

No, I haven't . But I want to obtain at least some uraninite or torbenite in the future.
In the photo you can see my recently made k2cr2o7 crystals.
They aren't very big becouse I haven't enough p.chromate . I must buy at least 300g

.
(Sorry for the photo quality it was taken by my crappy mobile).

The Volatile Chemist - 9-2-2016 at 08:48

Pic of a glycerol complex of copper. Its more purple when viewed directly.

Any modes of precipitation/crystallization/collection known?

wg48 - 9-2-2016 at 09:50

Pic of a glycerol complex of copper. Its more purple when viewed directly.

Any modes of precipitation/crystallization/collection known?

I have found articles talking about powders but they only referenced the procedures for the production in an other article I did not have access too. But I have not given hope of finding the procedure.

Trienthanolamine complexes of copper in water apparently can be precipitated with ethanol. That may be applicable to the glycerol complex. Perhaps by adding more of the counter ion or lowering the ph would work.

I intend to try it with both complexes by adding the ethanol very slowly in an attempt to grow crystals. Very slow cooling from concentrated solutions may also work. One problem is copper sulphate is not as soluble as nitrates and I only have the sulphate.
A macroscopic crystals of either complex will be interesting.

The Volatile Chemist - 9-2-2016 at 09:57

Indeed. Perhaps acetone would work as a knock-out solvent for the complex?
It sounds like the complex has a similar color to the citrate-copper(2) complex I've made before.

wg48 - 9-2-2016 at 10:57

Pic of a glycerol complex of copper. Its more purple when viewed directly.

Could you be more specific? What complex and weight ratio?

Probably K2Cu(C3H6O3)2 ref: CHARACTERIZATION AND REACTIONS OF
COPPER (II)

wg48 - 9-2-2016 at 10:59

Pic of a glycerol complex of copper. Its more purple when viewed directly.

Could you be more specific? What complex and weight ratio?

Probably K2Cu(C3H6O3)2 ref: CHARACTERIZATION AND REACTIONS OF
COPPER (II)

last part of above post

wg48 - 9-2-2016 at 11:01

COPPER (II) Probably K2Cu(C3H6O3)2 ref: CHARACTERIZATION AND REACTIONS OF
COPPER (II) GLYCEROL COMPLEX
by
HAZIMAH ABU HASSAN
July 1 998

for synthsis see ref: Preparation of Ultrafine Copper Powders with Controllable Size via Polyol Process with Sodium Hydroxide Addition
Pattanawong Chokratanasombat and Ekasit Nisaratanaporn
ENGINEERING JOURNAL Volume 16 Issue 4, ISSN 0125-8281 (http://www.engj.org/)

quote:In a typical synthesis, copper (II) nitrate trihydrate (Cu(NO3)2∙3H2O, Carlo Erba) was dissolved in the solution of sodium hydroxide (NaOH, Mallinckrodt) and glycerol (C3H8O3, Carlo Erba) with varying the molar ratio of NaOH:Cu(NO3)2∙3H2O in the range of 0:1 to 5:1 and
Cu(NO3)2∙3H2O:glycerol at 0.02:1

I used 3:1 K not Na

[Edited on 9-2-2016 by wg48]

[Edited on 9-2-2016 by wg48]

wg48 - 9-2-2016 at 11:47

Indeed. Perhaps acetone would work as a knock-out solvent for the complex?
It sounds like the complex has a similar color to the citrate-copper(2) complex I've made before.

I think acetone has a better chance of working. The Nickel complex is a green colour not very intense but Nickel chloride appears to be much more soluble in glycerol (by volume) than copper sulphate so perhaps more chance of crystals. Perhaps copper chloride would be too.

The Volatile Chemist - 9-2-2016 at 14:00

Indeed. Perhaps acetone would work as a knock-out solvent for the complex?
It sounds like the complex has a similar color to the citrate-copper(2) complex I've made before.

Definitely try it! Looking forward to pictures.

fluorescence - 11-2-2016 at 08:17

An old photo I found a bit off topic but still something quite unique. I think only Pok will be able to understand that but those two books are "Chemische Experimente die Gelingen" from Roempp. A very important book for private chemistry which helped some generations with their experiments and might be the origin of some experiments you see here on the forum.

The interesting fact about these two books is the date when they were published. The one on the left is my personal copy from 1941 printed during the second world war.

The one on the right is quite hard to get. I think it was even Pok who told me about it's rarity but I might be mistaking here. I borrowed that from a library, not sure from which since I gave this as an order for my local library to find it somewhere in Germany. Actually I wanted to find the very first version but the sent the wrong book. To my great luck they sent me this one. It's from 1947 so shortly after the end of said world war and printed in a completely different time. The first page says

That is was published under the supervision of the military government (US) and was limited to 5000 copies only.

So this book was censored by the military government after world war II and it was limited.

If you look at newer versions of that book in the 1960s for example there will be chapters about radiation and nuclear energy since this topic came up. The last updated versions like 30 years ago were quite censored because of dangerous experiments. The one from 1947 was inbetween war and nuclear energy so there is no additional chapter. It's just shorter than the other copies.

So mine, the one on the left is one of the last really uncensored versions. I'd like to have one from 1939, would be interested to know whether it differs from the 1941 version. But mine has the additional war-chapters like

-"Charcoal as livesaver"
-"A model of a gas mask"
-"Why you can't fertilize with sodium chloride"
-"The secret of the knock-out-gun"
...

The picture shows both of them before I had to return the 1947 edition so close to each other like the have never been before and will probably never ever be, telling two completely different stories it's just incredible to think about this.

[Edited on 11-2-2016 by fluorescence]

Attachment: phpqefW0n (45kB)
This file has been downloaded 848 times

[Edited on 11-2-2016 by fluorescence]

Follow up of the copper glycerol complex

wg48 - 13-2-2016 at 02:49

Indeed. Perhaps acetone would work as a knock-out solvent for the complex?
It sounds like the complex has a similar color to the citrate-copper(2) complex I've made before.

Definitely try it! Looking forward to pictures.

No acetone does not drop the complex but its only 5% miscible with glycerol. But I checked the mixture today and the acetone has gone blue and there is a slight white precipitate in the glycerol layer.

So I put a drop of blue acetone solution on a slide in the hope it would evaporate and leave a blue compound.

The acetone evaporated leaving what I assume was mostly glycerol with a slight bluish colour. So I tried to evaporate the glycerol above a flame. Before all the glycerol had evaporated what looked like a copper precipitate had formed. See the pic below. Sorry its not very good pic. The lighter region of precipitate is still wet with glycerol.

So the complex can be extracted in to acetone. Perhaps vacuum distillation could remove the acetone and glycerol without decomposing the complex.

Incidentally the original purpose of the complex was to producer a nano copper powder for a glass artist.

[Edited on 13-2-2016 by wg48]

3DTOPO - 15-2-2016 at 23:51

These crystals grew in front of my eyes the first time I applied power to a molybdenum disilicide (MoSi2) heating element. Some kind of oxide?

alexleyenda - 16-2-2016 at 07:23

Results of 20 hours of inorg chemistry in my lab classes.

12696173_674706312672111_118206457_n.jpg - 12kB

12746568_674706282672114_935311776_n.jpg - 10kB

PbCl6(pyr)2, PbI2, recristalised PbI2, (Ni(NH3)6)Cl2, Ni(NH3)4(NO2)2, (tren)Ni(NO3)2, Bn3SnCl, SnCl4(DMSO)2

[Edited on 16-2-2016 by alexleyenda]

CharlieA - 16-2-2016 at 16:18

You are a busy beaver! Good work.

Noflers - 16-2-2016 at 23:24

These are my first bismuth crystals:
http://postimg.org/image/jcr0ngeev/full/
host images

http://postimg.org/image/tlkyw458n/full/
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bismuth is my no.1 metal

[Edited on 3-1-2016 by crystal grower]

Very nice, Bismuth is definitely my #1 as well. Here's a few of mine. I figure at least one of these phone photos will look good on the big screen.

PHILOU Zrealone - 17-2-2016 at 08:10

Results of 20 hours of inorg chemistry in my lab classes.

PbCl6(pyr)2, PbI2, recristalised PbI2, (Ni(NH3)6)Cl2, Ni(NH3)4(NO2)2, (tren)Ni(NO3)2, Bn3SnCl, SnCl4(DMSO)2

[Edited on 16-2-2016 by alexleyenda]

How did you do this one Ni(NH3)4(NO2)2?
Isn't the Ni hexacoordinator? --> Ni(NH3)6(NO2)2?
Like in your Ni(NH3)6)Cl2 and (tren)Ni(NO3)2?
With N2H4, Ni(NO3)2 fixes 3 N2H4 --> Ni(H2N-NH2)3(NO3)2

alexleyenda - 17-2-2016 at 08:47

Haha the Ni(NH3)4(NO2)2 was quite a long process. First, NiSO4 hexahydrate is reacted with KHCO3 to make the carbonate, then acetic acid is added to form Ni(OAc)2(H2O)4 (both reactions bubble CO2). Then you mix NaNO2 and NH4CH3COO in large excess to make a good NO2 donor (NH4NO2) and add it to the nickel complex with conc. NH3 (aq) to give Ni(NH3)4(NO2)2. After mixing 10 min at room temp, you let it precipitate for around 3 hours, which gives the red solid that precipitates out of the blue solution. You must be careful not to add too much water, or the complex will not precipitate.
Metal centers do not always hexacoordinate. I guess the complex is tetrahedric or square planar, which is common when metal centers are d8 (8 electrons in the d orbitals) (search about it, it is very common). For exemple d8 paladium is almost always square planar.

DraconicAcid - 17-2-2016 at 11:38

Haha the Ni(NH3)4(NO2)2 was quite a long process. First, NiSO4 hexahydrate is reacted with KHCO3 to make the carbonate, then acetic acid is added to form Ni(OAc)2(H2O)4 (both reactions bubble CO2). Then you mix NaNO2 and NH4CH3COO in large excess to make a good NO2 donor (NH4NO2) and add it to the nickel complex with conc. NH3 (aq) to give Ni(NH3)4(NO2)2. After mixing 10 min at room temp, you let it precipitate for around 3 hours, which gives the red solid that precipitates out of the blue solution. You must be careful not to add too much water, or the complex will not precipitate.
Metal centers do not always hexacoordinate. I guess the complex is tetrahedric or square planar, which is common when metal centers are d8 (8 electrons in the d orbitals) (search about it, it is very common). For exemple d8 paladium is almost always square planar.

I expect that it is octahedral, with the two nitrites coordinated along with the ammonias. Nitrite is a much better ligand than nitrate (I have a batch of K4[Ni(NO2)6] on my shelf).

alexleyenda - 17-2-2016 at 13:37

Quote: Originally posted by DraconicAcid

I expect that it is octahedral, with the two nitrites coordinated along with the ammonias. Nitrite is a much better ligand than nitrate (I have a batch of K4[Ni(NO2)6] on my shelf).

It is true that it might be, I did not take time to think about it and made the too rapid asumption that because NO2* is negatively charged it was only the counter ion, but it can be coordinated too. The nickel is d8 tho, so it "likes" to have a coordination of 4. I'll have to sit and think about it for longer :p

[Edited on 18-2-2016 by alexleyenda]

DraconicAcid - 17-2-2016 at 14:40

It is true that it might be, I did not take time to think about it and made the too rapid asumption that because NO3 is negatively charged it was only the counter ion, but it can be coordinated too. The nickel is d8 tho, so it "likes" to have a coordination of 4. I'll have to sit and think about it for longer

NitrIte, not nitrAte. Nitrate is a lousy ligand; nitrite is a fairly good one.

Nickel(II) is d8, but it will only be square planar with ligands that are strong field splitters (cyanide, or ones like DMG). Apart from low-spin complexes, it has no more tendency to be 4-coordinate than cobalt(II) or iron(III). With halides, it will form NiCl₄<sup>2-</sup> or NiBr₄<sup>2-</sup>, but that's more to do with the size of the halides than the electron configuration.

DraconicAcid - 17-2-2016 at 17:01

Copper(II) lactate trihydrate

[Edited on 18-2-2016 by DraconicAcid]

alexleyenda - 17-2-2016 at 23:12

Quote: Originally posted by DraconicAcid

Well you are right, it makes sense. This inorganic theory is quite fresh for me, I have to really take my time to really think about it and not forget anything to get to the good conclusion :p And yeah, i'm aware of the difference between nitrate and nitrite as ligands (I just wrote NO3 by mistake), my error was to forget that it had to be low spin to favorise the coordination of 4. Now the question is on which atom is it coordinated even tho (i'm not sure but if my memory is good) most of the time it is on the oxygen. I'll have to look at my IR spectrums, maybe tomorrow ^^ Nice cristals by the way.

[Edited on 18-2-2016 by alexleyenda]

[Edited on 18-2-2016 by alexleyenda]

PHILOU Zrealone - 18-2-2016 at 07:44

Thank you for the process.

Seems like Cu also form such a complex: Cu(NH3)4(NO2)2
For your Ni complex I thought the NO2 was nitrite anion, but it takes part into the hexacoordinated complex and this explains why two positions are not available anymore for NH3 ligands.

Those complex may be energetic since NO2 is an oxydiser and NH3 a fuel...
-->Could you test a few mg in contact with a flame?

By experience with Ni(N2H4)3(NO3)2, Cu(N2H4)3(NO3)2, Co(N2H4)3(NO3)2, Ni(NH3)6(NO3)2, Cu(NH3)4(NO3)2, Co(NH3)6(NO3)2, Ni(en)3(NO3)2, Cu(en)2(NO3)2 and Co(en)3(NO3)2...
The addition of some CH3OH, CH3-CH2OH may help isolation and crystalization from watery solutions followed by washing with ether ((Eth)2O).

PHILOU Zrealone - 18-2-2016 at 07:49

Pic of a glycerol complex of copper. Its more purple when viewed directly.

Could you be more specific? What complex and weight ratio?

Probably K2Cu(C3H6O3)2 ref: CHARACTERIZATION AND REACTIONS OF
COPPER (II)
last part of above post

COPPER (II) Probably K2Cu(C3H6O3)2 ref: CHARACTERIZATION AND REACTIONS OF
COPPER (II) GLYCEROL COMPLEX
by
HAZIMAH ABU HASSAN
July 1 998

for synthsis see ref: Preparation of Ultrafine Copper Powders with Controllable Size via Polyol Process with Sodium Hydroxide Addition
Pattanawong Chokratanasombat and Ekasit Nisaratanaporn
ENGINEERING JOURNAL Volume 16 Issue 4, ISSN 0125-8281 (http://www.engj.org/)

quote:In a typical synthesis, copper (II) nitrate trihydrate (Cu(NO3)2∙3H2O, Carlo Erba) was dissolved in the solution of sodium hydroxide (NaOH, Mallinckrodt) and glycerol (C3H8O3, Carlo Erba) with varying the molar ratio of NaOH:Cu(NO3)2∙3H2O in the range of 0:1 to 5:1 and
Cu(NO3)2∙3H2O:glycerol at 0.02:1

I used 3:1 K not Na

Thank you!
Interesting.
I knew that Pb(OH)2 can form a glycerolate with mild energetic properties...
Pb(OH)2 + HOCH2-CHOH-CH2OH -heat/reflux-> HOCH2-CH(-O-*)-CH2-O-Pb-* + 2 H2O
So apparently Cu(OH)2 also form a glycerolate :-)

alexleyenda - 18-2-2016 at 09:50