Sciencemadness Discussion Board - Solubility of elements without chemical reaction

Solubility of elements without chemical reaction

woelen - 10-2-2016 at 03:44

I am working on a periodic table display of the elements and for this I prepare samples of the elements in different forms. They, however, must remain present as elements, not as chemical compounds. For this purpose I also prepare some solutions of elements, which display the element's properties very nicely (e.g. a nearly saturated solution of Cl2 in CCl4, which very nicely demonstrates the color of Cl2 at high concentration and which remains stable indefinitely).

Now I wonder, for which elements I can make such display-solutions, which contain the element as free element and which are stable over long times. I identified the following:

+ Cl2 in CCl4
+ Br2 in CCl4
+ I2 in CCl4
+ I2 in CH3CH2OCH2CH3
- S in CS2
- Se in CS2 (??)
- P4 in CS2
+ C60 in CS2

The ones, listed with a "+" I actually made and ampouled.

This is all which I can imagine. Any other ideas? Metals, alloyed with other metals, do not count. I suspect this only works for non-metal elements. If you know of any other element which can be dissolved in a suitable solvent, then I would like to know.

I don't think I'll make P4 in CS2. This is just a colorless liquid.

DraconicAcid - 10-2-2016 at 08:24

Oxygen is very soluble in some freons. I wonder if ozone would similarly dissolve (and might be more blue than colourless).

I don't expect you're going to dissolve metals in anything but mercury, unless a suspension of colloidal silver or gold would count.

Magpie - 10-2-2016 at 08:42

I don't suppose Na dissolved in liquid ammonia would qualify?

Eddygp - 10-2-2016 at 08:47

Quote: Originally posted by Magpie

I don't suppose Na dissolved in liquid ammonia would qualify?

Na in liquid ammonia technically loses an electron so is not the "pure" element.

blogfast25 - 10-2-2016 at 08:52

Quote: Originally posted by DraconicAcid

Oxygen is very soluble in some freons. I wonder if ozone would similarly dissolve (and might be more blue than colourless).

And in those 'Fluorinert' things: https://en.wikipedia.org/wiki/Fluorinert

woelen - 10-2-2016 at 11:49

Quote: Originally posted by Magpie

I don't suppose Na dissolved in liquid ammonia would qualify?

No, this does not count for me, the color of these solutions is not due to the element dissolved, but due to the solvated electrons. Some time ago I made a deep blue solution like this with lithium metal in ethylenediamine.

--------------------------------------------------------------------------------------

I have done some research though and found that selenium and tellurium can be dissolved in a mix of ethanethiol and ethylene diamine:

http://pubs.rsc.org/en/Content/ArticleLanding/2014/SC/C4SC00...

On evaporation, the element appears again in crystalline form.
Ethylenediamine alone does not dissolve selenium, I tried that.

I, however, do not have ethanethiol, and I wonder whether I do want this compound in my lab (boiling point just 35 C while its odour threshold is appr. 0.5 ppb). I can buy it (appr. EUR 60 per 500 ml).

DraconicAcid - 10-2-2016 at 11:53

If you don't want ethanethiol in the lab, I suspect other thiols will have similar solvent properties.

careysub - 10-2-2016 at 12:05

C70 fullerene solutions are a different hue from C60. Showing that difference would be interesting.

careysub - 10-2-2016 at 12:19

I would say that you do not want ethanethiol in your lab. I was once affected by an ethanethiol spill evacuation in Long Beach, CA around 1999. ~100 grams evacuated a square mile area, and the place smelled awful for week after they let us come back.

Also the odor of ethanthiol from a wood pulp plant at Port St. Joe, FL can be detected on occasion at Ft. Walton Beach 100 miles away.

To significantly reduce the thiol odor problem you need a higher thiol, octanethiol perhaps (odor is described as "mild").

woelen - 10-2-2016 at 23:52

Quote: Originally posted by careysub

C70 fullerene solutions are a different hue from C60. Showing that difference would be interesting.

I know of the differences. Apparently, solutions of C70 in CS2 are bright red, while solutions of C60 in CS2 are purple with a bluish hue. I have 200 mg of C60, purchased that for EUR 13 or so. C70 is much more expensive though. I found a seller, who sells it at EUR 40 or so per 50 mg, which is too expensive to my taste.

deltaH - 11-2-2016 at 00:31

About the solutions of metals. I don't think having the metal in solution in an ionic form betrays much if you call this a 'solution periodic table' and display them in a standardised way. For example, metal oxides can be dissolved in anhydrous choline chloride/urea to show their colour and this would make a great table if you use the most stable metal oxide for each metallic element:

See how nice this works from this picture from p52 of the presentation:
https://wet.kuleuven.be/english/summerschools/ionicliquids/l...

I suspect this would make something very special when done with all the metals :cool:

woelen - 11-2-2016 at 01:06

This indeed looks very nice, but to my opinion it is quite arbitrary how this looks like. If I use aqueous solutions of ionic compounds of metals, then the set of colors becomes quite different. Then there is the issue of different oxidation states. There can be coordinating agents in the solution, which again change the color.

If I do such a thing, then I will do this for the transition metals in all their common oxidation states, dissolved in water, assuring that no ligands other than water are available in the solutions, but I am still pondering on whether I will do that or not. The number of samples will be large, and still there is the feeling of randomness (e.g. for chromium, should I only display sky blue Cr(2+) and grey/violet Cr(3+) solutions, or should I add yellow CrO4(2-) and orange Cr2O7(2-) as well?). Also, the space, available in my display, is limited. Only appr. 7 cm per element and a height of appr. 10 cm. Having many element boxes crammed full with vials is not attractive.

[Edited on 11-2-16 by woelen]

j_sum1 - 11-2-2016 at 01:19

Interesting deltaH. But it is definitely a different kind of table. To me, this has more appeal than compound collecting or acid collecting that have been recently suggested. I have long been fascinated by transition metals and their assorted colours. A display showing all stable oxidation states would be really cool. I am reminded of strontiumred's work. Not complete but a really good start with some nice preparations to follow. (Not dissimilar to some sections of your pages either, woelen. I am thinking of where you show different oxidation states of n and V and others.)
I intend to play with some Ti at some stage soon. Beginning with dissolving in HCl and then producing the red peroxide complex.

[edit]
I love SR's gorgeous photo of Mn(V) that he uses as a slide at the top of his pages. That's on the to do list too.
Same photo in the thread below.
http://www.sciencemadness.org/talk/viewthread.php?tid=14644&...

[Edited on 11-2-2016 by j_sum1]

MeshPL - 12-2-2016 at 10:15

I'm surprised nobody mentioned amalgams! Plus there are some other alloys liquid in room temperature. Most alloys are not compounds. Technicaly brass, bronze, cast iron (carbon solution

, who needs fullerenes?), all the jewlery and some on, are solutions.

DraconicAcid - 12-2-2016 at 10:26

Quote: Originally posted by MeshPL

I'm surprised nobody mentioned amalgams! Plus there are some other alloys liquid in room temperature. Most alloys are not compounds. Technicaly brass, bronze, cast iron (carbon solution

, who needs fullerenes?), all the jewlery and some on, are solutions.

That's because alloys were ruled out in the original post.

Texium - 12-2-2016 at 11:47

Regarding the choline chloride/urea eutectics... unfortunately the picture deltaH used upthread was actually depicting solutions of metal oxides in choline chloride/malonic acid. The power point presentation that it is from mistakenly labeled it as choline chloride/urea. I have made the choline chloride/urea eutectic before, and it does not dissolve a significant amount of any metal oxides. Claims about it in that presentation are unfortunately false.

See this other presentation for a comparison of the solubilities of several metal oxides in choline chloride/urea versus choline chloride/malonic acid: http://www.erc.arizona.edu/seminar/Current-2011/DineshThanu_...

[Edited on 2-12-2016 by zts16]

MeshPL - 13-2-2016 at 04:42

Ok, sorry, alloys were ruled out , but only as metal-metal mixtures.

Quote:

Metals, alloyed with other metals, do not count.

But still, you can alloy many non-metals with metals, although they will kinda form compounds. Not sure if in all cases though. Like do boron, silicon or arsenic containing alloys contain compounds of respective element?

Also, do gaseous non-metals liquified under pressured and mixed together count? Like I'm pretty much sure liquid nitrogen, oxygen, helium, neonium, argonium, kryptonium and xenonium will form liquid mixtures under certain conditions.

careysub - 13-2-2016 at 09:40

Quote: Originally posted by woelen

Quote: Originally posted by careysub

C70 fullerene solutions are a different hue from C60. Showing that difference would be interesting.

Onyxmet.com has it a bit cheaper: EUR 30 for 100 mg.

deltaH - 15-2-2016 at 05:32

Quote: Originally posted by zts16

Regarding the choline chloride/urea eutectics... unfortunately the picture deltaH used upthread was actually depicting solutions of metal oxides in choline chloride/malonic acid. The power point presentation that it is from mistakenly labeled it as choline chloride/urea. I have made the choline chloride/urea eutectic before, and it does not dissolve a significant amount of any metal oxides. Claims about it in that presentation are unfortunately false.

See this other presentation for a comparison of the solubilities of several metal oxides in choline chloride/urea versus choline chloride/malonic acid: http://www.erc.arizona.edu/seminar/Current-2011/DineshThanu_...

[Edited on 2-12-2016 by zts16]

Excellent moderation zts16! I didn't understand how metal oxides dissolved in ChCl/urea, but an acid makes perfect sense! So these are actually nearly anhydrous solutions of metal malonates. Well that's not so interesting, metal oxides dissolve much better in conc. HCl and it's OTC

Texium - 15-2-2016 at 07:39

Yeah, it's very disappointing.

Although, choline chloride is a pain to work with anyway. It's so hygroscopic that even the reagent grade stuff from Fischer that my school purchased was quite wet before the jar was even opened! I'm kinda glad I've moved onto other projects.

woelen - 15-2-2016 at 08:07

If you want an interesting set of metal ions in solution, then I would not use HCl, but preferrably HClO4, or dilute HNO3. These do not coordinate to the metal ions and then you get solutions which really represent the color of aquated ions, and not the color of ions, coordinated to some anion.

E.g. an iron(III) solution in dilute HClO4 or dilute HNO3 is nearly colorless (actually very faint violet), while in hydrochloric acid such a solution is bright yellow with a brownish hue (due to formation of the FeCl4(-) complex). Copper(II) tends to form green solutions when chloride is present at moderate concentrations, while it is pure sky-blue in dilute HClO3 or dilute HClO4.

Such a display would indeed be nice. The first row of transition metals has the following simple aqueous ions, of which Ti(3+), Cr(2+) and FeO4(2-) are borderline stable and I doubt whether solutions of these can be kept around in ampoules for years.

Sc(3+)
Ti(3+) / TiO(2+)
V(2+) / V(3+) / VO(2+) / VO2(+)
Cr(2+) / Cr(3+) / Cr2O7(2-) / CrO4(2-)
Mn(2+) / Mn(3+) / MnO4(2-) / MnO4(-)
Fe(2+) / Fe(3+) / FeO4(2-)
Co(2+)
Ni(2+)
Cu(2+)
Zn(2+)

You could also argue about whether you want stuff like Cr2O7(2-) or MnO4(-) in the display.

blogfast25 - 15-2-2016 at 10:01

Quote: Originally posted by woelen

E.g. an iron(III) solution in dilute HClO4 or dilute HNO3 is nearly colorless (actually very faint violet), while in hydrochloric acid such a solution is bright yellow with a brownish hue (due to formation of the FeCl4(-) complex).

Interesting. I always believed the 'yellow with brownish' hue was due to [Fe(OH)(H₂O)₅]<sup>2+</sup> but I'm probably conflating.

Thinking about it, dilute solutions of ferric alum don't appear yellow, only slightly brownish, depending on concentration, of course. And they're a bit thermochromic: darken on heating, presumably the equilibrium in pushed to the hydroxy complex at higher T.

woelen - 15-2-2016 at 23:45

If you dissolve an iron salt like Fe(ClO4)3 or Fe(NO3)3, then you get brown solutions. This brown color indeed is due to [Fe(OH)(H2O)5](2+) and possibly further deprotonated ions. When chloride ion is added to the game, then the color shifts more to yellow due to ligand replacement by chlorine ions. At low pH, you get FeCl4(-).

Pure [Fe(H2O)6](3+) is nearly colorless. Try dissolving some Fe(NO3)3·9H2O or Fe(NH4)(SO4)2·12H2O in dilute HNO3 or dilute HClO4. You get a nearly colorless solution, not brown, nor yellow. Next, add a little HCl or NaCl and the liquid at once turns bright yellow.

Bezaleel - 16-2-2016 at 08:18

Sulphur is reported to be soluble in toluene, and can be recrystallised from it. Doesn't add a new element to the list, though.

chornedsnorkack - 16-2-2016 at 09:28

Quote: Originally posted by woelen

Such a display would indeed be nice. The first row of transition metals has the following simple aqueous ions, of which Ti(3+), Cr(2+) and FeO4(2-) are borderline stable and I doubt whether solutions of these can be kept around in ampoules for years.

Sc(3+)
Ti(3+) / TiO(2+)
V(2+) / V(3+) / VO(2+) / VO2(+)
Cr(2+) / Cr(3+) / Cr2O7(2-) / CrO4(2-)
Mn(2+) / Mn(3+) / MnO4(2-) / MnO4(-)
Fe(2+) / Fe(3+) / FeO4(2-)
Co(2+)
Ni(2+)
Cu(2+)
Zn(2+)

I´m also unsure about V²⁺ and Mn³⁺.

While nitrate is a poorly complexing anion, it also tends to decay to NO₂, which is yellow. Dilute perchlorate is more stable, but Cl₂ and ClO₂ are also yellow.

How complexing is sulphate? E. g, can iron(III) be observed to turn from brown although still clear in a partially hydrolyzed solution to clear as sulphuric acid excess is added?

woelen - 16-2-2016 at 12:12

Mn(3+) is stable at very low pH. I made this myself and could keep it around in a test tube for several days. After that I used it for further experiments.
V(2+) might be unstable. I can imagine that it reduces water and that hydrogen is formed.

For many anions, sulfate is suitable as well. Sulfate does not coordinate to copper(II), cobalt(II), nickel(II). It does coordinate to Cr(3+) though. With iron(III) I never tried, but I expect that in dilute sulphuric acid the iron(III) ion also looks nearly colorless.

Dilute nitrate (less than 1 M) certainly does not tend to form NO2 when no suitable reductor is present. Dilute solutions of nitrates of nearly all metal ions, even the somewhat reducing ones, are stable. Even better indeed are perchlorates. These are very stable in aqueous solution, even more so than sulfate. Only at high temperatures or in near anhydrous situations the perchlorates and perchloric acid become strongly oxidizing.

chornedsnorkack - 18-2-2016 at 14:00

Quote: Originally posted by woelen

Mn(3+) is stable at very low pH. I made this myself and could keep it around in a test tube for several days. After that I used it for further experiments.
V(2+) might be unstable. I can imagine that it reduces water and that hydrogen is formed.

For many anions, sulfate is suitable as well. Sulfate does not coordinate to copper(II), cobalt(II), nickel(II). It does coordinate to Cr(3+) though. With iron(III) I never tried, but I expect that in dilute sulphuric acid the iron(III) ion also looks nearly colorless.

Dilute nitrate (less than 1 M) certainly does not tend to form NO2 when no suitable reductor is present. Dilute solutions of nitrates of nearly all metal ions, even the somewhat reducing ones, are stable.

And NO₂ dismutes in dilute solutions, where nitric acid is reduced to NO instead (which is colourless).
How do dilute nitrate solutions of strongly reducing metal ions behave - Sn²⁺, V²⁺, Cr²⁺, Eu²⁺, U³⁺?

woelen - 18-2-2016 at 23:36

Sn(2+) hydrolyses, formation of a precipitate.
V(2+) and Cr(2+) probably cannot be kept for a long time in water, regardless of the presence of nitrate.
Eu(2+) definitely cannot be kept in water for more than a few minutes, regardless of nitrate presence.
U(3+) I don't know. Does this exist in aqueous solutions? I only know of the yellow uranyl and the green U(4+).

chornedsnorkack - 19-2-2016 at 13:56

Quote: Originally posted by woelen

U(3+) I don't know. Does this exist in aqueous solutions? I only know of the yellow uranyl and the green U(4+).

This:
https://en.wikipedia.org/wiki/Jones_reductor
with link to
Mendham, J; Denney, R.C; Barnes, J.D.; Thomas, M. (2000). Vogel's Textbook of Quantitative Chemical Analysis (6th ed.). Pearson Education Ltd. pp. 446–448. ISBN 0-582-22628-7.
lists cations produced by Jones reductor as:
titanium(III), vanadium(II), chromium(II), molybdenum(III), niobium(III) and uranium(III).

I have not heard much of Nb(III) or Mo(III) elsewhere.
What do Ta and W do in Jones reductor?

ScienceHideout - 22-2-2016 at 19:18

Quote: Originally posted by MeshPL

Technicaly brass, bronze, cast iron (carbon solution

, who needs fullerenes?), all the jewlery and some on, are solutions.

There is a professor at my university who does a lot with fullerenes! I am pretty sure they dissolve in toluene...