Sciencemadness Discussion Board

crystallising nickel nitrate

Antwain - 23-11-2007 at 12:11

I have had some nickel nitrate solution laying around for quite some time now and have finally got around to trying to crystallise it. It never crystallised on its own, which is not surprising since it is hygroscopic. The only data I have found gives a mp. of 56*C and a bp. of 137*C, I assume this is for the hexahydrate, but the Merck does not specify. This makes it difficult for me to reduce its volume and obtain crystals, since (assuming mp and bp are for hydrate) I could easily drive off water of crystalisation and presumably form a solution with more than one Ni(NO3)2 for every 6 H2O, again making an assumption that the boiling point of water will be elevated by a large proportion of dissolved solid.

nickel nitrate is soluble in alcohol as well, so I can't precipitate it that way. I also *think* it may decompose if I try to make it anhydrous by heating. Can anyone suggest a method of obtaining a solid of known composition?

Xenoid - 23-11-2007 at 12:29

....... and on evaporation the solution deposits green monoclinic crystals of the composition Ni(NO3)2.6H2O. These are very soluble in water (94g/100g water at 20oC.). On heating [the crystals], some of the water of crystallisation is driven off, but the nitrate decomposes before it becomes anhydrous .......

It's very soluble, you may just need to concentrate your solution a little more!

Edit: That's 94g/100g H2O of the theoretical anhydrous salt Ni(NO3)2 not the hexahydrate!

Regards, Xenoid

[Edited on 23-11-2007 by Xenoid]

Nickel nitrate

chloric1 - 23-11-2007 at 14:30

You know twice in my life I observed nickel nitrate crystals. First time in high school second time when I bought some. I remmber that even upon their delivery they where quite wet and they continued to be that way. On a warm and humid day you can put a crystal on some paper and end up with a green wet spot on the paper. IMHO it is not worth trying to isolate solid Nickel nitrate. But if you are so inclined, use moderate heat(not boiling!) and slowly evaporate your solution into a syrup then place in an air tight vessle that would make a good desicator and support the syrup in a crystallizing dish over concentrated sulfuric acid or potassium hydroxide flakes. If you cannot get these then calcium chloride might help you. You can regenerate calcium chloride I think too.;)

12AX7 - 23-11-2007 at 15:12

Vacuum also comes to mind.

If it doesn't have to be nitrate, you could add maybe ammonium sulfate and crystallize nickel sulfate. Hmm, I'm not sure if ammonium sulfate is less soluble than nickel sulfate.


Eclectic - 23-11-2007 at 15:20

I had a concentrated syrup of nickel chloride that took WEEKS to crystallize in in a sealed plastic container. Eventually, I got very large emerald green crystals. The nitrate may behave the same way.

not_important - 23-11-2007 at 17:25

Originally posted by 12AX7
Vacuum also comes to mind.

Reduced pressure is a great help with inorganics that 'melt' (self-dissolve) at low temperatures. One of the best investments for the amateur chemist is a recirculating aspirator pump, either bought or built. All plastic construction doesn't care about acid vapours, you can toss some marble chips in the water tank to keep the pH under control.

There were a number of predecessors to the rotovap, and they were often used for inorganics. Basically a pair of flasks were connected by a wide glass tube (through stoppers), then set in a pair of waterbaths. One bath was heated, the other had cold water slowly flowing through it to a drain. The flasks would be rigged in some fashion to rotate, sometimes a secondary tube would be used to evacuate them and then be closed off before starting rotation. It's the sort of apparatus you can cobble together for not much money.

If it doesn't have to be nitrate, you could add maybe ammonium sulfate and crystallize nickel sulfate. Hmm, I'm not sure if ammonium sulfate is less soluble than nickel sulfate.

Actually I think the mixed sulfate would crystallise out, its solubility is fairly low - something like 10 to 30 g/100 cc. Like the corresponding ferrous and cobaltous sulfates, nickel ammonium sulfate crystallises reasonably well and isn't horribly hygroscopic.

If you really want the nitrate you might try concentrating a small amount of the solution to near saturation and then slowing add some concentrated nitric acid, the common ion effect might cause nickel nitrate to drop out. If it does then perhaps repeating on the main part of the solution would be in order.

Or you could evaporate to a syrup or even mush, give it time to crystallise in mass, and then recrystallise from alcohol (Me, Et, or i-Pr depending on the solubility in each).