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IrC
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Quote: Originally posted by Bezaleel  | | That's highly interesting, IrC. I wonder whether you tried to detect any iron in your nitrates? If so, by what methods? |
None, I'm not a chemist. Electronics has always been my real profession. Mad Science is my hobby for over half a century and I have studied chemistry
just enough to aid in whatever it was I was trying to create at the time. My conclusion stems from reading in patents about problems caused by Iron
combined with the fact my glow powders started working when I quit using a hacksaw and went to carbide cutting wheels on fiber disks. Not such a
stretch to come to my conclusion. As to microwave kilning it should work but I built a quartz tube furnace to run H2 only through the material. Yet
another idea from reading patents. Enhanced by observations of the finished material once I stopped letting O2 be in the atmosphere inside the kiln.
Glow lasts longer when reduction takes place, from far less time to not more than a few seconds if any O2 is present when making the powders. I have
not seen Fleaker post in a long time but if he was still around he could tell you they work very well, being the only member on SCM I ever sent some
of my work to study.
I must not be paying enough attention to threads lately. Out of curiosity I searched and Fleaker is still around, having posted in March. Good to know
some older members are still around. There are so many I have not seen post in a long time.
Brain&Force "He used microwave synthesis, not a kiln. The reagents were purified by recrystallization, but the aluminum nitrate nonahydrate
appeared to have some FeCl3 in it."
Yeah the Iron will ruin the effect. I went and looked at some of the NurdRage video. Going to have to try that out of curiosity. I wonder how hard it
will be to create a non conductive sealed vessel filled with H2. Need to study that one because I believe if he had tried that there would have been a
big difference, assuming he gets all the elements detrimental to glow persistence out of the reaction.
[Edited on 4-30-2014 by IrC]
"Science is the belief in the ignorance of the experts" Richard Feynman
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elementcollector1
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Speaking of nitrates, just distilled some nitric acid. Going to use most of it on the mix of RE metals I have from mischmetal (in hopes of separating
the cerium), and we'll see how that goes.
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blogfast25
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EC1:
Remind me where you got the mischmetal from? Why waste nitric when strong HCl will dissolve it too?
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elementcollector1
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Whoops! I meant to say "oxides". The mischmetal was from Coleman magnesium firestarter blocks.
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blogfast25
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Still: why nitric?
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elementcollector1
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Was it not discussed earlier that nitric acid would not dissolve cerium compounds, but would lanthanum and others?
To be fair, I probably should've used sulfuric acid, then precipitated the hydroxides, *then* used nitric, but we'll see how it goes. So far, nothing
much seems to be happening at RT.
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Brain&Force
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Nitric acid (concentrated) will oxidize cerium, as well as a portion of the terbium, out of solution. Cerium nitrates are generally insoluble
according to Wikipedia.
At the end of the day, simulating atoms doesn't beat working with the real things...
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blogfast25
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| Quote: Originally posted by elementcollector1 | Was it not discussed earlier that nitric acid would not dissolve cerium compounds, but would lanthanum and others?
To be fair, I probably should've used sulfuric acid, then precipitated the hydroxides, *then* used nitric, but we'll see how it goes. So far, nothing
much seems to be happening at RT. |
That would have been my strategy but this might work too. What concentration is your nitric?
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elementcollector1
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I didn't make any measurements, but rough estimate of 50 mL water to maybe 20 mL HNO3 (white fuming).
EDIT: After an hour or so refluxing, the solution looks like this:
Presumably that's cerium messing with the color - if so, darn!
[Edited on 5-4-2014 by elementcollector1]
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Brain&Force
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Try adding some hydrogen peroxide to a small amount of the solution. It should reduce the cerium(IV) present and evolve oxygen.
At the end of the day, simulating atoms doesn't beat working with the real things...
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elementcollector1
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Thanks, B&F - turned it water-clear, with a tiny amount of precipitate.
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Töilet Plünger
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That, elementcollector1, was one of my wild guesses. 
I read that the cerium(III) tetrakis(dibenzoylmethide)triethylammonium complex is blood red. I wonder if a cerium complex caused the color.
Am I assuming correctly that oxygen was evolved on addition of peroxide? And what color is the precipitate?
[edit] This is my (Brain&Force's) other account. Sorry if I confused you.
[Edited on 2014-5-5 by Töilet Plünger]
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elementcollector1
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There was very minor gas production - barely anything, to be honest. Given that and that the precipitate was so small in mass, this leads me to
believe that the solution was, in fact, mostly lanthanum, and only a very small amount of cerium contributed to the dark color - rather like the
permanganate ion. The precipitate appears to be some mix between tan and white, my eyes fail me.
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Brain&Force
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Here's where I got the idea from:
https://www.youtube.com/watch?v=IFmAhhiam9g
Samarium and praseodymium in acid are quite interesting. I'd like to try it myself.
At the end of the day, simulating atoms doesn't beat working with the real things...
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blogfast25
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| Quote: Originally posted by Töilet Plünger | I read that the cerium(III) tetrakis(dibenzoylmethide)triethylammonium complex is blood red. I wonder if a cerium complex caused the color.
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Everything points to quite a bit of Ce. He dissolved the mishmetal in nitric, so that would give Ce (IV) which is red to yellow depending on
concentration.
Adding peroxide in acid conditions reduces the Ce (IV) to Ce (III), hence it goes back to colourless.
No complexes needed.
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elementcollector1
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Quite a bit? But the precipitate was so small... Perhaps there is more hiding in this colorless solution?
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blogfast25
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I don't understand why the precipitate is significant to you.
Correct me if I'm wrong: you dissolved the mishmetal in nitric acid and obtained a reddish solution. I believe the colour to be due to Ce (IV), having
seen it several times. Other Ln(III) nitrates are also present but are colourless. Then you added hydrogen peroxide, we know that reduces Ce(IV) to
Ce(III) (done it) and the colour disappeared, in accordance with that reduction.
I'm not sure where or why the precipitate came into play? Ce (IV) and Ce (III) nitrates are highly soluble.
Remember my thread on CeO2 and soluble Ce salts?
http://www.sciencemadness.org/talk/viewthread.php?tid=24638#...
It showed unequivocally that Ce(NO3)4 is well soluble in water.
[Edited on 7-5-2014 by blogfast25]
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Zyklon-A
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Sorry to disrupt this discussion, slightly.
I hope to dissolve the flints in sulfuric acid soon.
Lanthanum sulfate is nearly insoluble, while the sulfates of Iron, Cerium, Neodymium, Praseodymium, and Magnesium are soluble. So I will separate the
Lanthanum from the precipitated sulfate.
Then I'll add acetic acid to precipitate the nearly insoluble cerium acetate. (The acetate salts of the other cations are soluble).
Perhaps phosphoric acid to precipitate insoluble magnesium phosphate. Does anybody have the solubility data for Neodymium and Praseodymium phosphate?
I might neutralize the acid with NaOH until their hydroxides precipitate, and try to separate them later.
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elementcollector1
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| Quote: Originally posted by blogfast25 |
I don't understand why the precipitate is significant to you.
Correct me if I'm wrong: you dissolved the mishmetal in nitric acid and obtained a reddish solution. I believe the colour to be due to Ce (IV), having
seen it several times. Other Ln(III) nitrates are also present but are colourless. Then you added hydrogen peroxide, we know that reduces Ce(IV) to
Ce(III) (done it) and the colour disappeared, in accordance with that reduction.
I'm not sure where or why the precipitate came into play? Ce (IV) and Ce (III) nitrates are highly soluble.
Remember my thread on CeO2 and soluble Ce salts?
http://www.sciencemadness.org/talk/viewthread.php?tid=24638#...
It showed unequivocally that Ce(NO3)4 is well soluble in water.
[Edited on 7-5-2014 by blogfast25] |
True, but I was rather hoping that Ce(III) nitrate wasn't, due to the quote on the Wikipedia article.
| Quote: | | The oxides are dissolved in nitric acid that excludes one of the main components, cerium, whose oxide is insoluble in HNO3. |
I had thought that meant unreactive, but this is clearly not the case.
My main concern is separating lanthanum: I have more than enough cerium oxide to make cerium compounds without mischmetal.
Zyklonb: In my experience, this doesn't count for much; cerium sulfate was always near-insoluble (only a very little bit dissolved before it got
saturated). At 0 C, the difference between the two (~22 g/100 mL and 3 g/100 mL) is large enough that this could be possible, but I'm not sure this
will work.
[Edited on 5-7-2014 by elementcollector1]
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Zyklon-A
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Cerium sulfate has an interesting solubility curve, at ~0°C 20 grams dissolves in 100 mL of water. As the temp rises, solubility decreases.http://en.wikipedia.org/wiki/Solubility_table#S
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blogfast25
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| Quote: Originally posted by Zyklonb | Lanthanum sulfate is nearly insoluble, while the sulfates of Iron, Cerium, Neodymium, Praseodymium, and Magnesium are soluble. So I will separate the
Lanthanum from the precipitated sulfate.
Then I'll add acetic acid to precipitate the nearly insoluble cerium acetate. (The acetate salts of the other cations are soluble).
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I'd plan very carefully if I were you: the Ln(III) sulphates are ALL poorly soluble in hot, and not greatly soluble in cold. La is one of the poorer
soluble ones but a great separation your idea will not make.
And dissolving mixed lanthanides in hot sulphuric acid is a recipe for getting all kinds of unexpected precipitates, due to the inverse
temperature/solubility relationship for Ln(III) sulphates.
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blogfast25
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EC1:
”The oxides are dissolved in nitric acid that excludes one of the main components, cerium, whose oxide is insoluble in HNO3”
Don’t conflate: this statement says nothing about the solubility of cerium nitrate. Wiki solubility table for Ce(NO3)3: 234 (!) g / 100 g water. See
also my photo of Ce(NO3)4 solution.
It’s also non-sensical: whether an oxide dissolves in an acid or not depends largely on the thermal history of the oxide.
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aga
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| Quote: | | depends largely on the thermal history of the oxide |
Oxides have Memory ?!??!?!?
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Brain&Force
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It's called calcination; some oxides will become inert when strongly heated.
At the end of the day, simulating atoms doesn't beat working with the real things...
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Brain&Force
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From several research papers I have read, the best idea for tracking terbium in mixtures is to add in a planar ligand as a complexing agent, which
increases the fluorescence of the ion. Dipicolinic acid appears to be suitable for the job. I'll be trying this as soon as I can get my hands on a lab
and some DPA.
I don't know if this works with other rare earth ions, however: I know that the hexakis(antipyrine)europium(III) ion is only weakly fluorescent. As
for DPA, I have no idea.
If someone is processing fluorescent phosphors, this may be part of an extraction method, as well as taking advantage of the redox chemistry of both
europium and terbium ions.
At the end of the day, simulating atoms doesn't beat working with the real things...
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