Sciencemadness Discussion Board - Production and storage of europium(II) compounds

Production and storage of europium(II) compounds

Brain&Force - 23-3-2014 at 13:34

Yup, back at the f-block again...

I don't currently have any europium at the moment, but I'm interested in studying the fluorescence of europium(II) derivatives after my work on terbium wraps up. I'm wondering about the best method of storing these because they have a reduction potential of -0.35V for the reaction Eu³⁺ + e^- → Eu²⁺, which implies that soluble salts of said cation will be unstable with respect to oxidation in acidic solution.

As shown in the lanthanide corrosion test, europium corrodes very quickly in air to form a yellow oxide of europium(II), which gradually turns white as it oxidizes to europium(III) oxide.

<img src="http://www.elementsales.com/re_exp/re020904.jpg" width="600">
Day 9

<img src="http://www.elementsales.com/re_exp/re030304.jpg" width="600">
Day 32

<img src="http://www.elementsales.com/re_exp/re20050908.jpg" width="600">
Day 586 (the oxide was removed from display and later replaced in order to compare it to the other corrosion products)

I've seen an image of this yellow oxide fluorescing blue under UV light, I'll try digging it up.

Anyway, what would be the best way to produce europium(II) salts without having them oxidize in air, or become oxidized to europium(III) in production?

Quick note: I think I found out why europium is extremely reactive. Europium appears to be naturally divalent at STP, and the potential for Eu²⁺ + 2e^- → Eu_(s) is pretty high - -2.80V because it has an extremely favorable f⁷ configuration. For a direct oxidation to the trivalent state the potential is the lowest among lanthanides, a measly -1.99V (most lanthanides have a potential around -2.2V). But what is meant when it is said that the metal (not its compounds, just the metal itself) is in a divalent state? Does it bond to two other europium atoms? My question arises from this.

[Edited on 23.3.2014 by Brain&Force]

Mailinmypocket - 23-3-2014 at 13:38

Don't have time to comment more right now but, beautiful experiment!!!

[Edited on 23-3-2014 by Mailinmypocket]

Brain&Force - 23-3-2014 at 13:45

Sorry if I confused you but this isn't my work - it was done by Dave Hamric at Metallium.

smaerd - 23-3-2014 at 13:50

Inert gas such as argon could help for storage (denser then air). Perhaps glass vaccutainers? Probably need to sparge/purge all solutions before doing anything with the material. I'm sure people who work with things like this in academia use purged glove-boxes for storage and synthesis. Not sure about the actual protocal, those are just some of my off the top of the head ideas.

Really cool pictures though.

Brain&Force - 23-3-2014 at 20:20

smaerd, I'm about to go to a research university so I should have access to a lot of equipment, but I'd like something more amateur friendly in case my plans don't work out. Vacuum storage in a dessicator seems like a good idea. I'm also curious as to what the solubility of europium(II) salts is in methanol or ethanol. I assume water solutions degrade easily.

I'll try to find out about the stability of EuSO₄ is in air and under water. This compound appears to be similar to its barium analog and is highly insoluble in water. If it can be somehow converted to the hydroxide as trivalent lanthanides can, it may be a good starting point for europium(II) compounds.

blogfast25 - 24-3-2014 at 13:58

Quote: Originally posted by Brain&Force

I'll try to find out about the stability of EuSO₄ is in air and under water.

Have you looked up the reduction potential Eu(III)/Eu(II)? If you can't find it I'll look it up in m CRC.

Töilet Plünger - 24-3-2014 at 14:29

It's already up there, the potential is -0.35V.

Brain&Force - 25-3-2014 at 08:25

Here's the photo I was talking about :

It's more greenish than blue, but still fluorescent. I've also read Eu₃O₄ exists, I wonder what color its fluorescence is.

MrHomeScientist - 25-3-2014 at 08:43

Neat! Is that also from Metallium? I have two europium samples from them, one heavily oxidized, so I'll definitely have to try shining a black light on them.

chornedsnorkack - 25-3-2014 at 13:25

What is the pKa of Eu3+?
The reduction potential of Eu3+/Eu2+, at -0,35 V, is close to the potential of Cr3+/Cr2+, at -0,41 V. Producing Cr2+ in solutions and solid salts is a standard exercise. How about europium?

Brain&Force - 25-3-2014 at 14:13

MrHomeScientist: It appears to be. This sample is not mine, though. If you can, post a picture of your europium sample fluorescing!

chornedsnorkack: I don't know. If this helps, I have a simple Pourbaix diagram:

And a more detailed one for the trivalent state:

NurdRage's video shows the yellow divalent form as the hydroxide when europium is added to water.

chornedsnorkack - 26-3-2014 at 02:08

Quote: Originally posted by Brain&Force

As shown in the lanthanide corrosion test, europium corrodes very quickly in air to form a yellow oxide of europium(II), which gradually turns white as it oxidizes to europium(III) oxide.

Can you confirm that the yellow and white corrosion products were oxides, seeing that there is no sign of either air moisture or carbon dioxide being excluded?

How would the same experiment look like if, besides Eu, the alkaline earth metals (the longlived ones) and magnesium were exposed to air?

Brain&Force - 4-4-2014 at 14:14

It is highly likely the carbonates formed, based on my experience with terbium compounds. I'm sure the divalent state is even more basic than the trivalent state.

Beryllium barely corrodes at all, magnesium corrodes only slowly, calcium corrodes pretty rapidly (I only know this because we have calcium at my school covered in the white oxide/hydroxide). Not sure about strontium and barium, but they would corrode as rapidly as europium (barium and europium(II) are quite comparable in general).

chornedsnorkack - 5-4-2014 at 04:02

Quote: Originally posted by Brain&Force

It is highly likely the carbonates formed, based on my experience with terbium compounds. I'm sure the divalent state is even more basic than the trivalent state.

Beryllium barely corrodes at all, magnesium corrodes only slowly, calcium corrodes pretty rapidly (I only know this because we have calcium at my school covered in the white oxide/hydroxide).

How do you know whether it is calcium oxide, hydroxide or carbonate?

Quote: Originally posted by Brain&Force

Not sure about strontium and barium, but they would corrode as rapidly as europium (barium and europium(II) are quite comparable in general).

Well, between calcium, strontium, barium and europium, that is a matter of measurement!
The properties and reactions of calcium oxide, hydroxide and carbonate should be pretty familiar. Like the rapid and violent exothermic reaction of quicklime with water, the slow reaction of quicklime with air humidity and air carbon dioxide, slow reaction of slaked lime with air carbon dioxide...
By contrast to the violent slaking of quicklime, hydration of magnesium monoxide is pretty sluggish.
Europium monoxide should be easily produced by heating europium with a small amount of air.
With europium monoxide, you have two options of reaction:
1) EuO+H2O=Eu(OH)2
2) 2EuO+4H2O=2Eu(OH)3+H2
Which of these tends to happen on contact of europium monoxide and water? Is the reaction sluggish as with MgO or violent as with CaO?

Artemus Gordon - 6-4-2014 at 13:41

B&F, for preservation of samples look into wine preservation kits. They use small cans of Nitrogen or Argon, or a combination, to displace the air in wine bottles. Also, the corrosion test you linked to suggested mineral oil.

[Edited on 6-4-2014 by Artemus Gordon]

MrHomeScientist - 8-4-2014 at 06:07

I briefly looked at my Europium under fluorescent light this weekend and did not notice any fluorescence. I still need to try it in a very dark room and see if that makes any difference.

Brain&Force - 8-4-2014 at 09:15

I don't think it'll fluoresce unless it's pretty corroded. Also, depending on air exposure, some of the fluorescence may be red (or even magenta if there's a mix of oxides). I would recommend you remove the sample from the oil, but that's probably going to make the corrosion worse. Is the oxide on the sample yellow or white?

When europium is added to water the divalent hydroxide appears to form, as evidenced by its yellow hue (europium(III) compounds are white).

Artemus, I'm going to university soon, so I think storage won't be as big an issue as I thought - I'll have access to tons of equipment. I'm currently looking into research opportunities, and currently drafting ideas for a research proposal.

[Edited on 8.4.2014 by Brain&Force]

chornedsnorkack - 9-4-2014 at 04:01

Quote: Originally posted by Brain&Force

When europium is added to water the divalent hydroxide appears to form, as evidenced by its yellow hue (europium(III) compounds are white).

The suspension is overall yellow. I did not quite get whether the yellow colour is in solid, solution or both.
Looking at the comparison of Eu(II) with alkaline earths, limewater is saturated at about 1,5 g/l at 25 Celsius, strontium hydroxide dissolves to 18 g/l at 20 Celsius, and barium hydroxide to 39 g/l at 20 Celsius. What is the saturated (clear yellow!) solubility of europium (II) hydroxide in water?

aga - 9-4-2014 at 15:35

Wonderful photos.

Brain&Force - 11-8-2014 at 09:38

Quote: Originally posted by MrHomeScientist

I briefly looked at my Europium under fluorescent light this weekend and did not notice any fluorescence. I still need to try it in a very dark room and see if that makes any difference.

Did you use a shortwave or longwave UV light? I think only shortwave light can excite lanthanides.

MrHomeScientist - 11-8-2014 at 11:11

Whichever one comes from a standard fluorescent tube-style black light. I suppose its also possible the glass of the vial blocked most of the UV, although the picture you posted apears to be in one.