Sciencemadness Discussion Board
Not logged in [Login ]
Go To Bottom

Printable Version  
 Pages:  1    3    5  ..  7
Author: Subject: Crystal Growing
Ioxoi
Harmless
*




Posts: 20
Registered: 24-9-2007
Location: Upstate NY
Member Is Offline

Mood: Pensieve

[*] posted on 18-2-2008 at 16:56
Favorite Crystal Growing Chemicals


Hi everyone,

I was just wondering what everyone's favorite crystal growing chemical was. Some people like the alums; others like monoammonium phosphate, copper sulfate or even sugars.

I'm posting this because I've got a new favorite: triethanolamine hydrochloride! I'm not a "pro crystal grower" (if that title even exists) by any means, but I think some of the really good ones out there should experiment with this chem and see what kinds of results they can get. I've pulled out some absolutely fantastic crystals out of some waste solution growing in one of my mom's plates that put crystals of other chemicals I've purposely tried to grow to shame.

As far as I've been able to tell, the most basic crystal shape (single crystal) is like the right angle prisms found in binoculars. I have one flawless crystal (it has a small internal striation, though) and when you look at it up close, it's unbelievable how perfectly smooth it's shape is. It actually looks like an acrylic prism from a precision optical device, except shrunk down to a couple mm on a side! I've also gotten some bigger clusters, but again they were growing in a dish unintentionally so I'm sure other guy out there could get much better results.

The main reason I like them better than others is
1. They're perfectly clear and
2. They're really stable. They don't dehydrate and turn opaque and they don't liquefy in humid air as far as I've been able to tell. I don't think they even form a hydrate at all.

You can make it in mass by buying a gallon of TEA from the ChemistryStore.com and then adding a big excess of muriatic acid. You can filter and dry the crystals (careful, heat makes solution outgas HCl) or just dilute the slurry down until it dissolves and use that as a crystal solution.

Anyone got any other favorites?




Chuck Norris does not uphold laws. He is the law.
View user's profile View All Posts By User
Nick F
National Hazard
****




Posts: 439
Registered: 7-9-2002
Member Is Offline

Mood: No Mood

[*] posted on 19-2-2008 at 05:10


I haven't tried it yet, but I'm going to try putting a sat. cobalt chloride solution into a vacuum desiccator. I hope that this will speed up crystal formation, while still making it slow enough for nice crystals to form. I just really love the colour of that compound! Then I'll probably seal it in something to keep it stable (maybe even cast it in some crystal-clear acrylic resin that I have to make a nice paperweight!).
Large potassium dichromate crystals are beautiful too, but apparenly they're not very easy to grow. And some lanthanide sulphates have nice colours.
View user's profile View All Posts By User
12AX7
Post Harlot
*****




Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline

Mood: informative

[*] posted on 19-2-2008 at 08:46


Incidentially, sodium bromate hydrate (monohydrate I believe) typically forms platelike crystals, reminiscient of BaCl2.2H2O and mica. Unfortunately they are air sensitive (turning white in this dry winter air).

It also forms thin crystals, probably the same habit but finer, when a boiling-hot solution is cooled with few nucleation sites. I had hot filtered a solution, half a day later discovering that it had congealed into a radiant suspension. Fortunately, such a suspension is easy to break up, because sodium bromate appears to have a relatively small solubility difference with temperature (indeed, a saturated solution at 15C is quite concentrated).

Tim




Seven Transistor Labs LLC http://seventransistorlabs.com/
Electronic Design, from Concept to Layout.
Need engineering assistance? Drop me a message!
View user's profile Visit user's homepage View All Posts By User This user has MSN Messenger
Magpie
lab constructor
*****




Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline

Mood: Chemistry: the subtle science.

[*] posted on 16-4-2008 at 12:23


Earlier in this thread YT2095 shows nice pictures of copper acetate he made. I presume this is cupric acetate.

I am attempting to make same, but I think my synthesis is going bad and I'm getting Cu(OH)2 instead. Here's what I did:

I reacted stoichiometric amounts of of Ca(OH)2 and white vinegar (5% HOAc) to form Ca(OAc)2 in lots of water. Then I added a stoichiometric amount of CuSO4*5H2O.

Ca(OAc)2 + CuSO4*5H2O --> Cu(OAc)2 + CaSO4*2H20 +3H2O

The gypsum was then removed by filtration, which is not hard when using a little diatomaceous earth added and a Buchner funnel.

Then I placed this navy blue translucent solution, which I assume to be aqueous Cu(OAc)2, on a hotplate-stirrer to evaporate off the water. But, it turned into a light blue slurry upon heating and vapors smell strongly of acetic acid. Hence my assumption that I'm forming a Cu(OH)2 slurry out of my dilute Cu(OAc)2. Right?

Now my question for YT2095: How please did you make your nice dark green crystals? :o

And, if anyone else can offer a method for making the Cu(OAc)2 please do. I searched but could not find one.
View user's profile View All Posts By User
UnintentionalChaos
International Hazard
*****




Posts: 1454
Registered: 9-12-2006
Location: Mars
Member Is Offline

Mood: Nucleophilic

[*] posted on 16-4-2008 at 14:21


Why even bother with the messy calcium step? Use dilute NaOH and make Cu(OH)2 (carbonate, bicarbonate (some carbonato complex will exist solvated no matter what so don't expect the liquid to go colorless) or even very slow careful addition of ammonia, stopping when the ppt. begins redissolving as dark blue tetraamminecopper (II) ions) from your CuSO4, filter out and wash the ppt. then dissolve in HOAc. Boil down, with an excess of acetic acid until the solution is intensely dark blue and chill down as much as possible without freezing and keep it there for a few hours (It doesn't mind hanging out supersaturated while crystals slowly form), collect crystals, then boil down again (I found a little more than halving the remaining solution worked out well.

Your slurry is likely calcium sulfate which has a nasty habit of easily forming supersaturated solutions and slowly crystallizing out (happened when I did an analagous procedure using CuSO4 and CaCl2). The boiling probably agitated it enough to force at least some of it to crystallize out.




Department of Redundancy Department - Now with paperwork!

'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
View user's profile View All Posts By User
garage chemist
chemical wizard
*****




Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline

Mood: No Mood

[*] posted on 16-4-2008 at 14:45


Or even better, use CuO for reaction with the acetic acid- it's way easier to filter than Cu(OH)2 and has a stochiometric composition and no water of hydration.
Just heat the Cu(OH)2 suspension from CuSO4 and NaOH until completely black and the CuO begins to settle.

I thought you had some pure acetic acid by now, Magpie? Why are you using vinegar?




www.versuchschemie.de
Das aktivste deutsche Chemieforum!
View user's profile View All Posts By User
Magpie
lab constructor
*****




Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline

Mood: Chemistry: the subtle science.

[*] posted on 16-4-2008 at 14:57


GC, that method looks very straight forward.

Quote:

I thought you had some pure acetic acid by now, Magpie? Why are you using vinegar?


Because I am a skinflint. :D

And thanks Un. Chaos for the insight on the CaSO4 formation. I guess that batch is toast. :(

Maybe not toast. I turned off heat, let it cool, then turned off stirrer. White solids (CaSO4) quickly settled out leaving a dark blue translucent supernate with no apparent Cu(OH)2 ppt gel. pH is slightly acid. So maybe some Cu(OAc)2 is recoverable here. I will filter again and go from there.

[Edited on 16-4-2008 by Magpie]

[Edited on 16-4-2008 by Magpie]

[Edited on 16-4-2008 by Magpie]
View user's profile View All Posts By User
chemoleo
Biochemicus Energeticus
*****




Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline

Mood: crystalline

[*] posted on 16-4-2008 at 16:28


Quote:
I'm posting this because I've got a new favorite: triethanolamine hydrochloride! I'm not a "pro crystal grower" (if that title even exists) by any means, but I think some of the really good ones out there should experiment with this chem and see what kinds of results they can get. I've pulled out some absolutely fantastic crystals out of some waste solution growing in one of my mom's plates that put crystals of other chemicals I've purposely tried to grow to shame.


Just now I wanted to write, triethylamine HCl doesn't crystallise well (as I've found that it forms needles which are very hygroscopic) and dang, there it is, you are writing on triETHANOLamine HCl :(
Where did you get the triethanolamine from?

On the matter of CuAc2 - I've made liters of the stuff using this very method, with CaSO4 - CaCO3 was dissolved in 25% HAc until it would dissolve no more, and to this the calculated amount of CuSO4 was added. This was filtered, rinsed more with cold H2O until the CaSO4 on the filter went pale green blue. The supersaturated solution started crystallising finest CuAc2 in no time, as glittering dark green crystals, which were so small that it was essentially a powder that could be filtered off and dried. It keeps well for several years. The CuAc2 at the time was used to produce lead acetate with Pb metal.
Anyway, I'd perhaps omit the boiling step, and work right away with highly concentrated solution. THat should allow you to filter off a significant amount of beautiful pure CuAc2.




Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
View user's profile View All Posts By User
DJF90
International Hazard
*****




Posts: 2266
Registered: 15-12-2007
Location: At the bench
Member Is Offline

Mood: No Mood

[*] posted on 16-4-2008 at 17:37


Potassium triethanedioxatoferrate (III) crystals are bright green, and looked spectacular. I can remember making some at school when we were learning about transition metals, shame I didnt take any pictures :( Anyway heres the equation if anyone wants it:

FeCl3 + 3K2(C2O4) => K3[Fe(C2O4)3] + 3KCl

Basically we made weighed out stoichiometric amounts of the Iron (III) chloride and the Potassium oxalate. The FeCl3 was dissolved in half of the amount of water as the Potassium oxalate, and the two resulting solutions were mixed in a small beaker and covered. Unfortunately I have a feeling that the compound may be light sensitive, as the textbook has made the note "leave a sample exposed to sunlight or powerful light".
View user's profile View All Posts By User
UnintentionalChaos
International Hazard
*****




Posts: 1454
Registered: 9-12-2006
Location: Mars
Member Is Offline

Mood: Nucleophilic

[*] posted on 16-4-2008 at 21:25


The above compound would be more easily understandable if named as potassium trisoxalatoferrate (III). Here's something that popped up on google about photosensetivity... http://adsabs.harvard.edu/abs/1975JSSCh..12...92S



Department of Redundancy Department - Now with paperwork!

'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
View user's profile View All Posts By User
DJF90
International Hazard
*****




Posts: 2266
Registered: 15-12-2007
Location: At the bench
Member Is Offline

Mood: No Mood

[*] posted on 16-4-2008 at 23:21


Sorry theres a typo... It's potassium triethanedioatoferrate (III), not triethanedioXatoferrate (III). It is equally understandable, its just that instead of naming it using the traditional name (oxalate), its named using the IUPAC accepted name (ethanedioate). Sorry for any confusion caused by my typo.
View user's profile View All Posts By User
Magpie
lab constructor
*****




Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline

Mood: Chemistry: the subtle science.

[*] posted on 17-4-2008 at 09:52


re: Cu(OAc)2 from vinegar, Ca(OH)2, and CuSO4*5H2O

I filtered off the second crop of CaSO4 then placed the filtrate in 2 shallow dishes appropriated from the kitchen, w/permission. This morning there were hundreds of very small dark crystals forming evenly all over the bottom of the dishes.

This procedure does have its drawbacks in dealing with the CaSO4 ppt. The best removal of this most likely would be done when the suspension is near boiling due to CaSO4*2H2O retro solubility with temperature.

[Edited on 17-4-2008 by Magpie]
View user's profile View All Posts By User
YT2095
International Hazard
*****




Posts: 1091
Registered: 31-5-2003
Location: Just left of Europe and down a bit.
Member Is Offline

Mood: within Nominal Parameters

[*] posted on 18-4-2008 at 01:10


Quote:
Originally posted by Magpie

Now my question for YT2095: How please did you make your nice dark green crystals? :o


I made a vinegar soln with GAA and tap water (roughly a 10% soln).
I then simply dissolved copper carbonate into it until all the Fizing stopped and some carbonate was left unreacted.
this was then filtered into a beaker and left to evaporate on a warm power transformer.
it took about a week to be be completely crystalised in the way you see in the pictures I posted.




\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
View user's profile View All Posts By User
Magpie
lab constructor
*****




Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline

Mood: Chemistry: the subtle science.

[*] posted on 18-4-2008 at 10:29


Thanks YT2095 for the method on your large crystals.

Mine have been growing for about 2 days now. I have two shallow dishes sitting in a window (eastern exposure) to warm in the morning sun as shown.

cupric acetate.jpg - 63kB
View user's profile View All Posts By User
Magpie
lab constructor
*****




Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline

Mood: Chemistry: the subtle science.

[*] posted on 19-4-2008 at 21:15


My cupric acetate is nearly finshed crystalizing. Upon examination I find quite a bit of contaminating crystals that formed right at the end. They are clear needles and I assume that they are calcium acetate. There is also a gummy material mixed in with the crystals. I suppose this is the sugars/misc organics that come in with the vinegar as mentioned in another thread. So I can't really recommend this method. Next time I will use the method recommended by garage chemist, especially with regard to using GAA instead of vinegar.
View user's profile View All Posts By User
16MillionEyes
Hazard to Others
***




Posts: 153
Registered: 11-3-2007
Location: 16 Million Eyes, US
Member Is Offline

Mood: No Mood

[*] posted on 20-4-2008 at 21:05


I tried making Cu(AcO)2 the stupid way ( NaAcO + CuSO*5H2O) and although I did get very few dark colored crystals they seem to be more a matter of odd luck than an actual process. What I do get in relatively copious amount though, is some light turquoise solid that is just slightly less soluble than the CuSO4 or the NaSO4 so that it forms a thin layer on the bottom of the beaker. I don't know what it is although I suspected it to be Cu(AcO)2 before seeing actual pictures (some sites described it as what I have so I felt confident). Hopefully someone can tell me.

[Edited on 21-4-2008 by __________]
View user's profile View All Posts By User
12AX7
Post Harlot
*****




Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline

Mood: informative

[*] posted on 21-4-2008 at 00:46


Could it be a I,II compound salt perhaps?

Crystallizing ternary mixtures like that can be a bitch. Instead of fractional crystallization, you're probably better off precipitating the important stuff (e.g., NaOH giving Cu(OH)2 and NaOAc and Na2SO4 in solution) and starting over. Outside chance you could seperate one or two phases from the dried, ground product with another solvent.

Tim




Seven Transistor Labs LLC http://seventransistorlabs.com/
Electronic Design, from Concept to Layout.
Need engineering assistance? Drop me a message!
View user's profile Visit user's homepage View All Posts By User This user has MSN Messenger
16MillionEyes
Hazard to Others
***




Posts: 153
Registered: 11-3-2007
Location: 16 Million Eyes, US
Member Is Offline

Mood: No Mood

[*] posted on 21-4-2008 at 05:59


I took a picture of it, for a while I did consider it being a Cu(I) salt but I have no way of confirming it. One thing I did notice though, the solution smells strongly of acetic acid and I can't see why it should.


Any ideas?
View user's profile View All Posts By User
YT2095
International Hazard
*****




Posts: 1091
Registered: 31-5-2003
Location: Just left of Europe and down a bit.
Member Is Offline

Mood: within Nominal Parameters

[*] posted on 21-4-2008 at 08:13


that Almost looks like the copper carbonate I start with :P



\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
View user's profile View All Posts By User
woelen
Super Administrator
*********




Posts: 7734
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 21-4-2008 at 08:33


This green stuff is a basic acetate, or a mix of basic acetate and basic sulfate.

NaOAc partly hydrolyses, because HOAc is a weak acid:

Dissolving: NaOAc --> Na(+) + OAc(-)
Hydrolysis: OAc(-) + H2O <---> HOAc + OH(-)

The hydroxide ions, formed in this hydrolysis reaction, precipitate with the copper(II) ions. No true Cu(OH)2 is formed, but a mixed Cu(OH)(OAc) salt, or a mix with sulfate ions as well. The remaining solution most likely is quite acidic and smells of vinegar.




The art of wondering makes life worth living...
Want to wonder? Look at https://woelen.homescience.net
View user's profile Visit user's homepage View All Posts By User
16MillionEyes
Hazard to Others
***




Posts: 153
Registered: 11-3-2007
Location: 16 Million Eyes, US
Member Is Offline

Mood: No Mood

[*] posted on 22-4-2008 at 06:23


I see, thanks. I did imagine the acetate was deprotonating the water and thus forming the hydroxide, the problem was I couldn't picture where the hydroxide would go so I didn't know exactly how to explain the vinegar smell. Now that you mention the possibility of a basic acetate it makes sense. :D
Anyhow, is this stuff good for anything useful/interesting?
View user's profile View All Posts By User
chemoleo
Biochemicus Energeticus
*****




Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline

Mood: crystalline

[*] posted on 11-5-2008 at 08:11


For the record, here's a picture of triethylamine hydrochloride...not a good target for growing nice crystals :(

[Edited on 13-5-2008 by chemoleo]

triethylamine-HCl.jpg - 72kB




Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
View user's profile View All Posts By User
12AX7
Post Harlot
*****




Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline

Mood: informative

[*] posted on 11-5-2008 at 09:31


Nice picture of what looks like awful crystals! Pointy and thin, like urea. Man, I'm recrystallizing urea, my fingers keep getting poked when I stir around in the solution!

Tim




Seven Transistor Labs LLC http://seventransistorlabs.com/
Electronic Design, from Concept to Layout.
Need engineering assistance? Drop me a message!
View user's profile Visit user's homepage View All Posts By User This user has MSN Messenger
The_Davster
A pnictogen
*******




Posts: 2861
Registered: 18-11-2003
Member Is Offline

Mood: .

[*] posted on 11-5-2008 at 10:42


Also has remarkable resemblance to the crystals of phenol which grow under the cap by sublimation.



View user's profile View All Posts By User
12AX7
Post Harlot
*****




Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline

Mood: informative

[*] posted on 11-5-2008 at 19:23


Ah, indeed! Only, those are fucking impossible to scrape off the sides! ARGH! :P

Tim




Seven Transistor Labs LLC http://seventransistorlabs.com/
Electronic Design, from Concept to Layout.
Need engineering assistance? Drop me a message!
View user's profile Visit user's homepage View All Posts By User This user has MSN Messenger
 Pages:  1    3    5  ..  7

  Go To Top