Lion850 - 26-4-2020 at 02:43
The praseodymium sulphate turned out a beautiful green and green being my favourite color I wanted to make more Pr compounds (I already made the
iodide which was also green...but after 2 weeks in a clear vial it has gone light brown....). So I went for the chloride.
I calculated the weights of hydrochloric acid (33%) and praseodymium (iii, iv) oxide Pr6O11 using the following equation:
Pr6O11 + 18HCl = 6PrCl3 + 9H2O + O2.
The oxide reacted readily with the HCl at room temperature and a gas with no smell was bubbled off. I added the oxide in small amounts while the
solution was stirred, each time waiting for the black oxide suspension to disappear and a lovely clear green solution to appear before adding more.
All went well until the time around 12g was added. Then, chlorine gas started to bubble off! Seems the reaction in a more dilute medium changed to:
Pr6O11 + 22HCl = 6PrCl3 + 11H2O + 2Cl2.
Question: why did this not happen from the start but only in dilute solution?
This also means I would need a small bit extra acid to react with the 20g oxide. Sure enough after the final 2g of oxide was added some black
suspension remained. Another 2g of HCl dissolved all the oxide giving the green solution. However I wanted to have the oxide in excess to neutralise
as much acid as possible so I added another half gram of oxide. When the suspension of oxide remained after 10 minutes of stirring I filtered the
solution using a gravity filter and folder filter paper (lab quality paper). Expecting the clear green solution.
The solution struggled to get through the filter paper from the start, not sure if this was because some 30g of chloride was dissolved in 50ml
solution. Then it suddenly ran through the paper torn a hole! I thought I may have damaged the paper during folding and folded the next carefully. It
looked perfect. Again the solution dripped through very slowly....and again made a tear/hole! I can't figure out why this is, as there should not have
been any strong acid and it was near room temperature.
I added a few drops of HCl with stirring and waiting inbetween drops and it took some 5 drops to get the solution back to clear green. I did not try
to filter it again. It looked perfectly clean.
I boiled the solution down from 50ml to 30ml. There was no smell of acid in the fumes which was good. I then transferred it to a crucible and into the
sand bath - from what I read on Wikipedia and Atomistry PrCl3 was quite stable when dried out in air.
The solution bubbled in the sand bath, see photo. I stirred it with a stainless spatula until it reached a point where the boiling decreased and it
started to solidify on the spatula. See photo.
The crucible was removed and the contents agitated as it cooled. Eventually it was dry and was then crushed to some extend with a mortar. Nothing
stuck to the mortar, confirming how dry it became in cooling.
28.5 gram was recovered, not far off the expected 29 to 30. Although, the 29-30 would be a dehydrate and what I got was for sure hydrated to some
extend, but which hydrate I do not know.
The color is a light green, see photo of the vial. Not as pretty as the sulphate but not bad!
Next will be the nitrate!
Bedlasky - 26-4-2020 at 05:41
Very nice preparation! I love you reports on making uncommon compounds.
Reaction between HCl and Pr6O11 surprised me too. I would suspect releasing of chlorine gas until begining, because Pr(IV) have very strong oxidizing
properties and it should oxidized HCl to Cl2. But it can also oxidize H2O to O2. It's really strange that chlorine is released when concentration of
HCl is lower. It seems logical that chlorine should be released in higher rate at higher concentration. Maybe complex formation affects it? Oxidation
abilities of Ce(IV) salts depending on complex formation. In HClO4, which have low complexing abilities, have Ce(IV) high E° value, but in HCl, which
have strong complexing abilities, have much lower E° value. Maybe this is also case of Pr(IV), which maybe form some unstable chloride complex, which
preferably oxidized water than hydrochloric acid? And when chloride concentration decreases too much, complex formation stops and Pr(IV) can oxidize
hydrochloric acid. But this is only theory, not a fact.
Edit: I found this book. There is something mentioned about Pr redox chemistry.
[Edited on 26-4-2020 by Bedlasky]
Lion850 - 27-4-2020 at 23:33
Hi Bedlasky - thanks for the link. Wish I studied chemistry when I was younger! The other thing I am still thinking about is- what caused the filter
paper to fail? Something specific to the solution weakened it but what?
Lion850 - 28-4-2020 at 00:07
How can be subject be changed to read "Report on making praseodymium chloride & nitrate"? Short report on making Pr(NO3)3.xH2O:
70% nitric acid was reacted with Pr6O11 (the brown mixed valence oxide) in a beaker. For the reaction to proceed the solution must be heated. Once all
the oxide was dissolved into a clear green solution, 0.5g oxide was added to confirm all the acid was consumed; which was the case as some oxide
remained in suspension. This was then reacted by adding a few drops of acid to again get to the clear green solution. At this point the solution was
divided in 2 equal parts (by weight). By stoichiometry there should have been about 15g of anhydrous Pr(NO3)3 in each half. One part was poured into a
crucible and placed in a desiccator over NaOH, and under strong vacuum (low enough pressure to boil the solution at around 40C).
After 2 days in the desiccator, 17.5g of beautiful green crystals was recovered. The are still slightly sticky, thus not 100% dry, but I needed the
desiccator for something else. See pictures of the nitrate crystals.
I am still deciding what to make with the second portion of the nitrate solution....maybe a praseodymium fluoride, just to have a praseodymium salt
that is yellow 
