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

Printable Version  
 Pages:  1  2  
Author: Subject: Dissolving germanium metal
ScienceSquirrel
Super Moderator
*******




Posts: 1863
Registered: 18-6-2008
Location: Brittany
Member Is Offline

Mood: Dogs are pets but cats are little furry humans with four feet and self determination! :(

[*] posted on 25-2-2011 at 05:11


Quote: Originally posted by blogfast25  
The combined solubility data from both sources are 1 g/100 ml at RT, 4.5 g/L at 25 C and 10.7 g/L at 100 C. Hardly a testimony to solubility, in my book...


I did not say it was soluble, quite the opposite, but it seems that refluxing will get it to dissolve a little.
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10334
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 25-2-2011 at 07:56


Assumimg these numbers are correct, that's quite remarkable IMHO...
View user's profile Visit user's homepage View All Posts By User
ScienceSquirrel
Super Moderator
*******




Posts: 1863
Registered: 18-6-2008
Location: Brittany
Member Is Offline

Mood: Dogs are pets but cats are little furry humans with four feet and self determination! :(

[*] posted on 25-2-2011 at 08:05


One of the uses for germanium dioxide is as a catalyst for making PET, it seems that making solutions is quite difficult so manufacturers supply ready made stable solutions.

http://www.polyester-technology.com/Publication/publication_...
View user's profile View All Posts By User
Mixell
National Hazard
****




Posts: 449
Registered: 27-12-2010
Member Is Offline

Mood: No Mood

[*] posted on 21-3-2011 at 12:28


I'm currently in the process of making germanium sulfate (Ge(SO4)2) by oxidation of germanium using copper ions in a solution. The process is quite slow but results can been seen (solid copper), can anyone tell me the characteristics of germanium sulfate meanwhile (couldn't find anything in Google...)?
I'll post some pictures latter, if I will manage to locate the cable from the camera.
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10334
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 21-3-2011 at 13:43


Quote: Originally posted by Mixell  
I'm currently in the process of making germanium sulfate (Ge(SO4)2) by oxidation of germanium using copper ions in a solution. The process is quite slow but results can been seen (solid copper), can anyone tell me the characteristics of germanium sulfate meanwhile (couldn't find anything in Google...)?
I'll post some pictures latter, if I will manage to locate the cable from the camera.


Where did you get this idea from?
View user's profile Visit user's homepage View All Posts By User
Mixell
National Hazard
****




Posts: 449
Registered: 27-12-2010
Member Is Offline

Mood: No Mood

[*] posted on 21-3-2011 at 13:48


By comparing the reduction potentials of germanium and copper. Why? Something isn't right?

EDIT: The potentials, from Wikipedia:
Ge4+ + 4 e− Ge(s) +0.12.
Cu2+ + 2 e− Cu(s) +0.34.

And its working, solid copper particles precipitate and the solution gradually becomes transparent.

[Edited on 21-3-2011 by Mixell]
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10334
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 21-3-2011 at 14:11


So the cell potential for:

2 Cu2+ (aq) + Ge(0) (s) === > 2 Cu(0) (s) + Ge4+ (aq)

… is about + 0.34 + (- 0.12) = + 0.22 V (thus ΔG < 0, reaction proceeds spontaneously), so plating out Cu with Ge metal is possible. Never occurred to me… But isolating your Ge(SO4)2 from solution to solid will be the problem, at least without hydrolysis. Ge is a half-metal, remember?



[Edited on 21-3-2011 by blogfast25]
View user's profile Visit user's homepage View All Posts By User
Mixell
National Hazard
****




Posts: 449
Registered: 27-12-2010
Member Is Offline

Mood: No Mood

[*] posted on 21-3-2011 at 14:54


Update: I got a nice clear solution of germanium sulfate, but decided to add more copper sulfate due to the fact that a lot of unreacted germanium was left in the vessel.
The reaction took approximately 6 hours at boiling point to react about 200 mg of germanium.
View user's profile View All Posts By User
Mixell
National Hazard
****




Posts: 449
Registered: 27-12-2010
Member Is Offline

Mood: No Mood

[*] posted on 21-3-2011 at 15:11


The hydrolysis is Ge(SO4)2 +2H2O--> GeO2 + 2H2SO4?
It would be helpful to find some information on germanium sulfate's decomposition/evaporation point. May be it is possible just to evaporate the water, or to dry with sodium hydroxide (if germanium sulfate is not too hygroscopic).
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10334
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 22-3-2011 at 06:16


I doubt very much if Ge(SO4)2 actually exists. In oxidation state +IV the element prefers to form anions like the germanate ion: GeO3(2-), not cations like Ge4+. In this it shows its metalloid character.

Here is an interesting Ge mineral, Schaurteite, Ca3Ge(SO4)2(OH)6•4H2O

http://webmineral.com/data/Schaurteite.shtml

… but the structural formula is deceptive: it should probably be re-written as 2CaSO4.CaGe(OH)6.4H2O; a double salt of CaSO4 and calcium germanate.

The germanates can be re-written as: GeO3(2-) + 3 H2O = Ge(OH)6(2-), so CaGe(OH)6 can be rewritten as CaGeO3.3H2O and the total formula for Schaurteite as 2CaSO4.CaGeO2.7H2O!

So despite the sulphate groups, no actual Ge(SO4)2 in sight either!

On evaporating (more or less regardless of conditions) you’ll obtain plain old boring GeO2 again…



[Edited on 22-3-2011 by blogfast25]
View user's profile Visit user's homepage View All Posts By User
Mixell
National Hazard
****




Posts: 449
Registered: 27-12-2010
Member Is Offline

Mood: No Mood

[*] posted on 22-3-2011 at 09:06


Well, I do have a solution of some sort.
Possibly the germanium exists in the following equilibrium:
Ge(SO4)2 +2H2O <--> GeO2 +2H2SO4?
I'll check the solution for acidity tomorrow, and also try to evaporate some part of the solution and test the resulting product, then I'll try to dissolve the solid product again, and see what happens, if it will not hydrolyze immediately, I will check the solution for a presence of sulfate.
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10334
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 22-3-2011 at 10:05


Quote: Originally posted by Mixell  
Possibly the germanium exists in the following equilibrium:
Ge(SO4)2 +2H2O <--> GeO2 +2H2SO4?



Yes. In a nutshell.

[Edited on 22-3-2011 by blogfast25]
View user's profile Visit user's homepage View All Posts By User
Mixell
National Hazard
****




Posts: 449
Registered: 27-12-2010
Member Is Offline

Mood: No Mood

[*] posted on 22-3-2011 at 10:52


Sorry, but I lack the knowledge of the meaning of some phrases, can you explain to me what "in a nutshell" means?

And back to the topic, I isolated the solution, it has a very-very faint yellow color, almost completely clear.
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10334
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 22-3-2011 at 12:15


'in a nutshell' = 'in short', 'in summary', 'basically', 'essentially'...
View user's profile Visit user's homepage View All Posts By User
Mixell
National Hazard
****




Posts: 449
Registered: 27-12-2010
Member Is Offline

Mood: No Mood

[*] posted on 22-3-2011 at 12:38


Understood, thank you =)
View user's profile View All Posts By User
The WiZard is In
International Hazard
*****




Posts: 1617
Registered: 3-4-2010
Member Is Offline

Mood: No Mood

[*] posted on 22-3-2011 at 16:11


Quote: Originally posted by Mixell  
I'm currently in the process of making germanium sulfate (Ge(SO4)2) by oxidation of germanium using copper ions in a solution. The process is quite slow but results can been seen (solid copper), can anyone tell me the characteristics of germanium sulfate meanwhile (couldn't find anything in Google...)?

I'll post some pictures latter, if I will manage to locate the cable from the camera.

Me the Analogue guy — again.

The AG goes down the hallway and removes from a bookshelf :—

Greenwood & Earnshaw
Chemistry of the Elements
Pergamon Press 1984 [There is a latter ed.]

"An ustable sulfate Ge(SO4)2 is formed in a curious reaction
when GeCl4 is heated with SO3 in a sealed tube at 160o."

GeCl4 + 6SO3 ---> Ce(SO4)2 + 2S2O5Cl2.


View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10334
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 23-3-2011 at 05:47


Interesting WiZ. Interesting also how the germanium transmutates to cerium! :D
View user's profile Visit user's homepage View All Posts By User
The WiZard is In
International Hazard
*****




Posts: 1617
Registered: 3-4-2010
Member Is Offline

Mood: No Mood

[*] posted on 23-3-2011 at 06:47


Quote: Originally posted by blogfast25  
Interesting WiZ. Interesting also how the germanium transmutates to cerium! :D


Through what Einstein (1947) called spukhafte Fernwirkung
(Spooky action at a distance).
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10334
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 23-3-2011 at 12:49


Quote: Originally posted by The WiZard is In  
Quote: Originally posted by blogfast25  
Interesting WiZ. Interesting also how the germanium transmutates to cerium! :D


Through what Einstein (1947) called spukhafte Fernwirkung
(Spooky action at a distance).


While maintaining the oxidation state too! Spookier than a direct hit on an MI6 safe house!
View user's profile Visit user's homepage View All Posts By User
Mixell
National Hazard
****




Posts: 449
Registered: 27-12-2010
Member Is Offline

Mood: No Mood

[*] posted on 25-3-2011 at 13:33


Well, I'll post that in this topic, silicon and germanium have similar properties.
Anyway, I got my hands on about 50g of ultra-pure silicon (the kind used to make computer chips), any suggestion what to do with it? Its a pretty inert element, but I think it still can be used for something interesting.
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10334
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 25-3-2011 at 13:46


For a start, dissolve it in strong alkali: it forms silicates.
View user's profile Visit user's homepage View All Posts By User
Mixell
National Hazard
****




Posts: 449
Registered: 27-12-2010
Member Is Offline

Mood: No Mood

[*] posted on 24-4-2011 at 07:59


I dissolved some germanium in nitric acid, and I got a very faint green solution (with some germanium left at the bottom).
I added some hydrogen peroxide which immediately turned the solution from very faint green to brown-yellow (urine colored).
At the moment the germanium at the bottom is giving off a good amount of bubbles, but I think its just the hydrogen peroxide decomposing (the decomposition rate gets bigger with time).
The interesting thing, is that my germanium sulfate solution is in the same color as this one, maybe its the color of the Ge4+ cation, so the hydrogen peroxide must have oxidized the former germanium cation (presumingly Ge2+) to Ge4+ ?
Or it formed some peroxo complex with a similar color to the Ge4+ cation?
View user's profile View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10334
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 24-4-2011 at 10:34


Dissolving Ge in nitric acid should, as far as I know, lead only to insoluble white GeO2 being formed (like its family member Sn, but Ge is even less inclined to make water soluble compounds).

The green, later brown yellow colour? Check the quality of your nitric acid: these colours are likely due to a contamination.
View user's profile Visit user's homepage View All Posts By User
sternman318
Hazard to Others
***




Posts: 121
Registered: 21-4-2011
Member Is Offline

Mood: No Mood

[*] posted on 24-4-2011 at 11:01


Quote: Originally posted by blogfast25  
Quote: Originally posted by Mixell  
Possibly the germanium exists in the following equilibrium:
Ge(SO4)2 +2H2O <--> GeO2 +2H2SO4?



Yes. In a nutshell.

[Edited on 22-3-2011 by blogfast25]


I may not follow or understand exactly, but if its in some equilibrium, couldn't you just add sulfuric acid to shift the reaction towards germanium sulfate?
View user's profile View All Posts By User
Mixell
National Hazard
****




Posts: 449
Registered: 27-12-2010
Member Is Offline

Mood: No Mood

[*] posted on 24-4-2011 at 11:28


The hydrogen peroxide and the nitric acid (as everything else) are laboratory grade (CP). And the germanium is clean too. So I don't know what, if any impurities could of caused that.I noticed some companies are selling germanium nitrate, so it must exist...
And about adding sulfuric acid to shift the equilibrium, germanium sulfate only exist in solution, so I'll just get a solution of sulfuric acid and germanium sulfate.
View user's profile View All Posts By User
 Pages:  1  2  

  Go To Top