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Author: Subject: Everyday Chemistry (provisional)
artemov
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[*] posted on 15-11-2020 at 19:39


Quote: Originally posted by Sulaiman  

I don't like using conc. H2SO4 to seal joints


Why do you not like this? I use 60% H2SO4 when distilling HCl and low boiling alcohols. I hope I'm not doing things wrong ...
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[*] posted on 15-11-2020 at 22:36


nothing wrong, its just me, so that I do not forget an get it on me ... trivial really.



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[*] posted on 16-11-2020 at 14:15


Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.

20201114_162514_adj_30pct_small.jpg - 78kB
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[*] posted on 16-11-2020 at 21:40


I've been trying to extract ibuprofen from OTC pills. should be a simple base/acid extraction, but the filler made the first extraction completely insoluble and when I tried to recrystallize the acid precipitate from IPA, it has instead formed an oily bilayer. What a mess.



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[*] posted on 16-11-2020 at 21:59


Quote: Originally posted by Bezaleel  
Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.



Did you do that to reduce the solubility via the common ion effect?
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[*] posted on 17-11-2020 at 06:15


Quote: Originally posted by MidLifeChemist  
Quote: Originally posted by Bezaleel  
Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.



Did you do that to reduce the solubility via the common ion effect?
No, it's just a final purification step, to get rid of assumed co-crystallised chloride.
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[*] posted on 17-11-2020 at 08:07


Got it - ok I'm assuming the idea is that the iodide ions from the KI will replace the chloride ions. So why do you need to do this second step - to get the Ni(en)3I2 the first time, did you simply boil away or evaporate the solution, is that why you need to recrystallize? And how do you know this is the iodide and not the chloride? I couldn't find any references to Tris(ethylenediamine) nickel (II) iodide or to its solubility. It looks like an interesting compound that I may want to try to make one day. Thanks in advance for the info!

Quote: Originally posted by Bezaleel  
Quote: Originally posted by MidLifeChemist  
Quote: Originally posted by Bezaleel  
Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.



Did you do that to reduce the solubility via the common ion effect?
No, it's just a final purification step, to get rid of assumed co-crystallised chloride.
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[*] posted on 17-11-2020 at 10:03


Posted my answer here.
Quote: Originally posted by MidLifeChemist  
Got it - ok I'm assuming the idea is that the iodide ions from the KI will replace the chloride ions. So why do you need to do this second step - to get the Ni(en)3I2 the first time, did you simply boil away or evaporate the solution, is that why you need to recrystallize? And how do you know this is the iodide and not the chloride? I couldn't find any references to Tris(ethylenediamine) nickel (II) iodide or to its solubility. It looks like an interesting compound that I may want to try to make one day. Thanks in advance for the info!
Quote: Originally posted by Bezaleel  
Quote: Originally posted by MidLifeChemist  
Quote: Originally posted by Bezaleel  
Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.

Did you do that to reduce the solubility via the common ion effect?
No, it's just a final purification step, to get rid of assumed co-crystallised chloride.
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[*] posted on 18-11-2020 at 02:16


Quote: Originally posted by woelen  
Did some crystallization of a mixed ethylenediamine/ammine complex of nickel(II). Something like Ni(en)2(NH3)2(ClO4)2. Beautiful dark blue crystals. Of course I made a picture of the crystals.

In the past I have done experiments with Ni-(en) complexis, and the not fully coordinated Ni(en)2(H2O)2 complex is dark blue. I did not manage to crystallize that. I only could get the purple Ni(en)3(ClO4)2 complex. But now I think I have managed to make the ammonia analogue Ni(en)2(NH3)2(ClO4)2 from a dark blue solution of nickel perchlorate and ethylene diamine to which I added some ammonia.

:( My nice crystals of this complex have withered in the last few days. I had put them in a dry place to have perfectly dry crystals, but I think that they lost part of their ammonia. The crystals now look ugly with a pale blue/green powder, sticking to the crystals, and some crystals are halfway gone, simply disintegrated into powder. A solution of this material in water is turbid. Most likely, besides simple loss of ammonia, also some basic carbonate salt is formed, contaminated with ammonium perchlorate, due to absorption of CO2 from air.

So, the ammine-salts cannot be dried succesfully in air. Maybe next weekend I try again, but then with rinsing first with acetone, followed with a rinse of diethyl ether and then quickly transferring the crystals to a glass vial to keep them around without decomposition. Fortunately I still have the picture of the crystals. I'll make a short write-up with these pictures and share this.

[Edited on 18-11-20 by woelen]




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[*] posted on 18-11-2020 at 05:32


Yes, these ammine complexes really easily lose ammonia. If you make tetraamminecopper(II) sulfate, you must store him in ammonia atmosphere. I still have some of it, I made it few years ago. I put it in the glass jar, added few drops of concentrated ammonia and closed it. It's still perfectly dark blue.



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[*] posted on 18-11-2020 at 08:59


Nickel ammonia complexes are particularly prone to losing ammonia.



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[*] posted on 18-11-2020 at 11:06


Here follow pictures of the complex. One picture was made just after carefully pressing the crystals between filter paper and tissue to get most of the adhering absorbed in the paper. The other picture was made two days later. After that period the crystals were perfectly dry, but they do not look attractive anymore.



Ni_en2_nh3_2.jpg - 1.3MB



Ni_en2_nh3_2_dried.jpg - 1.6MB

These crystals have a size of appr. 0.5 mm. They are stable, but when heated in a flame, they deflagrate. Still, this experiment was quite interesting. It is possible to isolate a dark blue Ni(en)2-complexes. One more thing to try is to isolate a skyblue/cyan Ni(en) complex, just one (en) ligand attached. The color of this is very different from the bright green of aqueous nickel(II), but its lightness is comparable.



[Edited on 18-11-20 by woelen]




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[*] posted on 21-11-2020 at 09:46


Finally I am able to reduce tungstate in to W(V), but also in to W(IV) and W(III), which I thought was not possible in aqueous solution.

If you reduce tungstate by strong reducing agents in acidic solution (I tried zinc powder, various metal ions in low oxidation state, dithionate and ascorbic acid) you always end with insoluble tungsten blue.

But in the mixture of oxalic acid/hydrochloric acid you can reduce tungsten even in to very dark blue (maybe even darker than molybdenum blue) mixture of W(V) complexes.

If you dissolve tungstate in hot conc. hydrochloric acid and add some aluminium foil, firstly is tungstate reduce in to greenish blue W(V) (reduction must take place in hot conc. HCl, in cold it doesn't work, reaction is very vigorous and produce lot of HCl gas). If you add more aluminium foil, it is reduce in to green W(IV) and than in to reddish brown (and also very unstable) W(III). These conclusions aren't final, I try to find more sources about W(IV) and W(III), but according to my observations and books and articles that I read this is probably right conclusion, but still it need more research, because there are few differences between my observations and what books say.

All reduced tungsten species can be easily oxidized even by mild oxidizing agents like Cu2+ in to tungstate. Cu2+ is reduced in to Cu+ (when I added solutions upon oxidation in to KOH solutions, Cu2O precipitated out).

[Edited on 21-11-2020 by Bedlasky]




If you are interested in aqueous inorganic chemistry look at https://colourchem.wordpress.com/main-page/

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[*] posted on 22-11-2020 at 23:02


Nice!

My son and I were filming the various vanadium reductions for the online version of the standard vanadium lab, but the reduction with aluminum wire isn't reliable (and mucked up, giving us grey instead of violet, and the titration said it was V(III) instead of V(II)). We'll have to do that part over again- i hate the idea of using Hg amalgam, but I might have to.




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[*] posted on 23-11-2020 at 06:34


For making V(II) is sufficient zinc powder, you don't need AlHg amalgam. I did it thousand times by this way. Just mix vanadate, HCl or H2SO4, zinc powder and mix it well. I did it just cca two weeks ago when I needed some reducing metal ions for molybdate and tungstate reduction.



If you are interested in aqueous inorganic chemistry look at https://colourchem.wordpress.com/main-page/

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[*] posted on 23-11-2020 at 07:20


Has anybody heard of urea being converted to ammonium formate? I'm trying to find a renewable source.
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[*] posted on 29-11-2020 at 11:31


I mentioned in another thread that I was trying to expand the number of possible metals in my college's qualitative analysis lab. I was amazed to find that copper hydroxide did indeed redissolve in excess sodium hydroxide- I've used hydroxide to precipitate copper dozens of times, and have never seen that happen before.



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[*] posted on 10-12-2020 at 09:45


Can anyone shed some light on what "FG" means in the linked paper? Also, no developments on my previous reply so any ideas would be appreciated.

https://sci-hub.st/https://doi.org/10.1021/ol8004326
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[*] posted on 10-12-2020 at 11:06


Quote: Originally posted by njl  
Can anyone shed some light on what "FG" means in the linked paper? Also, no developments on my previous reply so any ideas would be appreciated.

https://sci-hub.st/https://doi.org/10.1021/ol8004326


Functional group?
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[*] posted on 10-12-2020 at 11:54


I tried searching the paper for FG, F.G., F G etc and nothing came up... can you quote the sentence in question?



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[*] posted on 10-12-2020 at 17:10


Had to have been drunk. Thought I posted here about the propyl and formate esters I've been fooling with last few daze.

*edit* n-propyl alcohol and salicylic acid cooking as I type

*edit* propyl formate ain't real nose friendly

[Edited on 12-11-2020 by arkoma]

[Edited on 12-11-2020 by arkoma]




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[*] posted on 10-12-2020 at 17:41


njl,

FG does indeed mean "functional group." The abstract graphic is not all the clear so I can see your confusion.

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[*] posted on 10-12-2020 at 17:59
Purification of 925 sterling silver


18.58 g of 925 sterling silver (SS) was added to a solution of 40 mL HNO3 70%, 20 mL H2O2 and 60 mL dH2O.
The solution was left for 48 hours, at which point undissolved SS was still present and so mild heat was applied until all silver was in solution.
A slight excess (11 g) of NaCl dissolved in 50 mL of dH2O was added to the silver nitrate solution and a thick white precipitate of AgCl immediately formed.
The AgCl was then vacuum filtered in a Buchner funnel and rinsed with dH2O until the AgCl was perfectly white.
The AgCl was then added to 100 mL of dH2O and 10 g of NaOH was added slowly while stirring with a glass rod. After 5 minutes of stirring a few white flecks still persisted so another 5 g of NaOH was added under stirring until a consistent dark precipitate of Ag2O was observed in the reaction vessel.
The precipitate was rinsed once with dH2O then a further 400 mL of dH2O was added prior to adding 150 g of dextrose slowly and in portions while stirring.
No reaction was observed so the solution was heated to boiling while stirring.
After 5 minutes the precipitate started to become lighter in colour and after a further 5 minutes all traces of dark Ag2O precipitate had been converted to silver powder.
The silver powder was rinsed 10 times with dH2O until neutral with litmus.
The powder was then dried and weighed.
Total mass of silver powder was 17.82 g which corresponds to a yield of 101%. Presumably the SS was a little more than 92.5% silver.
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[*] posted on 11-12-2020 at 13:17


propyl salicylate smells pretty good



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[*] posted on 11-12-2020 at 14:21


Arkoma: How propyl salicylate smells like? I plan to make propyl benzoate or salicylate, I am not sure which one I'll pick.



If you are interested in aqueous inorganic chemistry look at https://colourchem.wordpress.com/main-page/

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