Electroplating of Nickel, waterless / organic solvent / possibly ionic liquids.

I've been experimenting for around two years with electroplating of metals in orgainc liquids. I've had partial successf in saturated carboxylic acids and alcohols. My chemistry background is weak; just undergrad courses have been taken.

Recently, I decided to try an easier metal, Nickel; which in water can be plated with just table salt and distilled vinegar. Even severe overvoltage (12+ volts) doesn't stop it from plating; although there is continual hydrogen bubble production during plating which I don't want for my final process. (Hence the reason I'm trying to move to polarized organic liquids where the hydrogen is weakly ionizable). My goal was to verify that using organic liquids without water didn't in some way affect even nickel, and interfere with plating.

My first two tests failed, which is surprising to me.

I've tried electroplating in pure glacial acetic acid 2cc's. Since this is not conductive by itself, I added a drop of formic acid to cause acid ions to form. I also (later) tried a drop of menthol crystals, dissoved by heat, since this is able to plate sodium non-electrolytically, via Magnesium metastasis; I thought Menthol might be less reactive than formic acid.

This solution ought not be sensitive to moisture because Nickel plates even in water, fine.

But, I ended up with precipitation and no deposit. Even putting it under inert gas, or under kerosine to block of oxygen access, made no difference.

On my second test runs, I switched to a weaker acid, propionic as the main carrier, and then added a drop of acetic acid -- which is the normal acid that works in water. I'm counting on the very low ionizability of propionic acid to act as a substitute for water. There is only a tiny amount of acetic acid in the solution. Ions go into solution, very well. It forms the same characteristic (?ligand?) color of nickel in water, Green. The conductivity of pure propionic acid (by itself) is sub 1 miliamp, before introduction of acetate; even with 100V applied at less than 4mm separation.

The propionic solution is not acidic/ionized enough to clean tarnished nickel off if dipped in it, cold.

So if it is ethching, I surmise the rate is less than a micron per day. eg: I think this is not be able to clean off deposits from the (-) cathode. Since even the (+) anode remains tarnished for over 24 hours. Even after adding a small amount of impurity, Acetate, or salt, the solution becomes conductive. Again, the solution does not etch of the tarnish from the (+) anode, but only dissolves where the metal is clean. The tarnish is formed by just placing it in formic acid, and applying a high voltage burst before drying it and placing in the test tube. Its' formic acid leavings, and it *WILL* etch off in water/acetic acid.

After 24 hours of running it in glacial acids (2 cc's is very small sample), the color is far darker green than I have ever obtained in water, it's well distributed, but it's not plating.

After centerfuging, the solution is clear, and conductivity remains the same. (Picture not shown). So, the colloidal haze seen in the following picture is not the main mechanism of conduction, but dissolved nickel is the main mechanism.

The following picture is after 24 hours, and is a microscope zoom. The test tube is 12mm on the inside. The graphite electrodes (-), are a high current and low current pair, in order to check if current density is the issue; Neither of them shows nickel plating even after running for an additional 24 hours.

I have heard of Kolbe reactions, and realize that it might be causing fatty acids or carboxylic acids to decompose at the (+) supply anode. But, I wouldn't expect that to stop a reaction from occurring at the (-) supply.

I would love to hear any ideas of why this isn't working. I've had luck with zinc, iron, and minor luck with aluminum and titanium (0.001 inch thick, max).
But I'm stumped as to what it is that is affecting the cathode in these solutions that limits the thickness of metal that can be plated.

I'm using pulsed D.C. 60Hz, resistor limited, with 15mA forward currents and 1mA reverse currents for cleaning; and then I stop reverse currents to check if it might be etching the nickel off. Doesn't seem to matter.

I can add boric acid, benzoic acid, lithium carbonate, or potassium hydroxide (and a dozen other impurities), to increase conductivity to any desired level. But none of these make a difference either. (Note: I don't think table salt will dissolve in glacial carboxylic acids; so I didn't try it; But, I didn't think the particular salt to increase conductivity in the analogous water processes mattered ... does it? )

What types of chemical side reactions would you expect to happen when attempting to plate in glacial acetic acid, or glacial propionic acid?

Thanks. (Stumped.)
(Btw: My experiment count is up to 487 over the last two years, I AM determined to figure this out. )




[Edited on 19-11-2023 by semiconductive]

No one has commented yet, so I'll give it a go....

Your persistence is impressive (400 attempts?!) but your narrative is a bit confusing. What is your anode material? Are you simply placing a Ni anode and a graphite cathode in an acetic acid plating bath, and applying 12V? It's not quite clear what your bath chemistry is, and what your anode is...

Acetic acid will tarnish nickel plated objects. That's why you don't use nickel-plated forks and spoons - they tarnish.

Overall, nickel plating is complicated. You chose the 'hardest one' to play around with. Maintaining a proper nickel plating solution is complicated. Cleanliness of the bath is always a headache. So is the pH. Getting proper adherence of the nickel is also a pain in the butt - many things can interfere with a solid, smooth nickel surface. Nickel plating starts with a "strike" phase at higher voltage, and then backs off for the main phase. You are messing around with literally the hardest, most complicated plating process. Zinc plating is much, much easier.

You stated that your goal is to minimize hydrogen gas at the cathode? You do that through voltage control and pH control of the bath.

Just my initial impression.... but don't give up!

[Edited on 21-11-2023 by Johanson]

Today's experiment, #490

I've got some plating activity on the higher current electrode.
Not sure if it's metal or not.

I decided to try reducing the amount of acid available, and replace it with an alcohol.

14 drops acetic acid
7 drops sulfer-oxide bearing molecule (not acid / no hydrogen ).
14 drops iso-position alcohol

The polar sulfer-oxide molecule takes the place of salt; this generally tends to reduce bubbling at the cathodes, and biases acid attacks more toward the anode.

This is the solution after 15minutes of plating at 10mA average current.
It's on about 7mm length of 0.9mm graphite.
I touched the graphite accidentally before putting it in, and I can see where my fingerprint was.



Only a very tiny bit of nickel shows up on the low current electrode.
I'm getting some floating solids, which is typical with acetic acid. I think the glacial acid has probably picked up some moisture from the air from the number of times Ive opened the flask.

I'll centerfuge the tube ... and do some calculations to figure out what molecule percentage of reactions have taken place after letting this run for 12Hours.

Unusually nice.
This morning, all solds were dissolved. Centerfuging, not needed.

From the dimensions of the test tube and what I estimate the height to be, a drop of fluid is closer to 16 [uL] or maybe even 20 [uL].

Acetic acid, glacial, is 17.5 Molar.
So, 17.5 [Mol/L] * 16 [uL] * 96320 [ Coulombs/Mole ] ~= 377 coulombs.
Acetic acid is only singly ionized.

Last night, the start current was at 10 [mA]. It's run 9 hours, and the current has dropped to 3.3 [mA].

Conservatively, then, 3.3 [mCoulomb/second] * 60 [seconds/minute] * [60 minutes/hour] * [9 hours] ~= 107 [ Coloumbs ].

If it was running at 10 [mA] most of the night, then ~= 325 [Coulombs]
I think this means between 1/3 and 8/10's of the acetic acid has been cycled.




I don't see anything on the web, doing Google searches, for "Basic Nickel Acetate"; although the floating stuff I hypothesize probably is a half hydroxide, half acetate, nickel material floating around in solution (and settling.)

But, I did find this paper:
https://www.researchgate.net/publication/338612394_Synthesis...

Which does suggest that it's possible to half-aceate - half hydroxide a nickel molecule. I doubt what I've got floating around in there is a quadra-hydrate, because there simply isn't enough water to form that. All chemicals are 99.9+% pure.

I'll let the experiment run for another three days, and then consider how to repeat the experiment to learn something more using the Scientific method. Hypothesis, welcome.... (anyone?)


[Edited on 23-11-2023 by semiconductive]

Okay, your claim, "my chemistry background is weak" is obviously a bit exaggerated. Nice video BTW. I should have had you nickel plate my wife's chrome bmw wheels, instead of having a local vendor here in L.A. do it - they wouldn't have flaked off so bad! I'm afraid your experiments above are beyond my level of expertise. Someone with a lot of experience running Ni plating baths is going to have to help you sort through this

Quote: Originally posted by Rainwater

This suggestion is expensive and dangerous, really leaning towards stupid. High pressure can drastically change the solubility of ionic compounds. One of the factorys I work with cleans aircraft parts in a water bath, under 30k psi. (15 tons) Turns out grease will dissolve in water under this high pressure. The water comes out discolored but a single phase and stays that way without treatment. Perhaps the same principle could be applied to your problem.

It's something to think about. But I'm not sure how I could do it.

I have noticed that pressure, even small amounts, can change the solubility of things.

Note:
The last experiment dropped below 1[mA] so I pulled out the anode to check it, and sure enough it was coated with a green polymer. Most likely the alcohol + acetic acid.

So, that version of the experiment is already finished/no good. There is metal or/metal hydroxide on the cathode, and it is adhered. But the coating is so thin, that it's not representative of the average chemistry of the bath; it's basically a dud experiment.

On the other hand; I've found that oils and polars often do mix when as strong as I'm using. I have gotten oils with 20 [mA] currents due to dissolved ions, easily.

I don't want to repeat failed stuff, so I'll try something new (random).
To demonstrate, I'm going to redo experiment #490 again.
The only variable I'm going to change is that I'll replace the alcohol with ether.
I think this will drop the ionic character of the liquid enough that oil will mix with it.

7 drops sulfur polar molecule
14 drops glacial acetic acid
24 drops ether

I would have (last year) expected this mixture to be totally polar and not mix with oil. In the last experiment, the kerosine, floated on top.
I'll now put silicone oil on top of this mixture; which is 'almost' as oily as kerosine.




There is no oil above the liquid; the silicone went into solution and mixed with vinegar, ether, and polar sulfur-oxide bearing molecules.

Notice also:
You can clearly see the Nickel Anode bubbling; there are *no* bubbles forming on the graphite cathodes (-).

This is the kind of solution I was talking about that doesn't release hydrogen at the cathode. I can get solutions similar to this (but different chemicals which I won't name) to plate Iron, Titanium, and aluminum, but only in thin layers.

I would think nickel or iron or zinc to be easier.

None the less: The solution always stops working after a few days.

In this new attempt:
The current, starts at only 1 [mA]. It's not very good.

eg: The resistivity is 10x higher than the previous solution using an iso-alcohol.
Since the voltages are 'high' 12-120V; that means there are 10x less ions present in this solution.

What advantage do you think pressure might bring to this type of experiment?
Is pressure likely to increase the number of ions? or only increase the probability that salt ions dissolve into solution?

Because if it's the latter (making oil and water/ions mix) I am clearly already doing it by just using glacial acetic or propanoic acids.

But if pressure can increase the number of ions present in solution, that would be interesting.

Just to follow up:
How thick is the 'bomb' that mixes water and oil at the company you work at?
Does 15Tons require wall thickness on the order an inch or so of steel?

I could imagine building something to put in a hydraulic press to manually squeeze small amounts of fluid to 15 Tons. But I'm not sure how I'd be able to seal it.

I definitely wouldn't be able to put electrodes inside it while it was being pressed.


[Edited on 24-11-2023 by semiconductive]

I'm still following this b/c plating/electrochem fascinates me for some reason. HOWEVER, I'm still baffled by the direction you are going. Your STATED GOAL in paragraph 2 of your original post was to develop a process that eliminated H2 gas as the cathode. That's what you stated. Why aren't you just using an electroless nickel bath? It's none of my business, but it seems like you're spending a lot of time messing around with things that were tried and discarded 150 years ago. Look at the old electroplating books from turn-of-the-century for example; they knew an awful lot about nickel plating, and they dismiss acetate bath-chemistry as totally unhelpful. But I'm sure you know more about that than me after 2 years of playing around with it

If I had the time to mess around with nickel processes, I would be focusing on REMOVING nickel from various surfaces and RECOVERING it for reuse, not plating it on surfaces. The area of battery recycling is going to be HUGE in the near future. Recovering metals from secondary batteries is almost impossible from an economic-payback standpoint, and a breakthrough would pay off handsomely I would imagine. Anyway, keep going!

Bubbling has accelerated, current remains at 1[mA]. Experiment #490, the "ether" variation.
Noticeable at the bottom is a white precipitate. There is also a whitish film on the nickel.

No plating is evident, just traces on the right (high current) cathode.



Since the color is not green, I don't think this is the same coordination complex as the last experiment had. There's only tiny amounts of green on the nickel anode itself, which is consistent with traces of water from the air haven gotten into the glacial acetic acid.

Whether the white stuff is just anhydrous nickel acetate, or possibly nickel acetate adduct with either ether or a sulfuric oxide bearing molecule, is unknown. But it's interesting.

Runing the experiment to exhaustion will give clues; I expect between 5 and 10 days according to the calculation I ran before. That's acceptable. I will report back, in later posts, when I get all the data.


Edit: Current rose to 1.5 [mA] this afternoon; 16 hours in. After gently tapping the anode to dislodge the white film which fell to the bottom as flakes; current rose to 2.0 [mA]. So, closer to 5 days. This is a slightly unusual change; most current increases happen within the first hour of starting an experiment. It's atypical for current to rise *after* precipitation has begun.

Edit: Mixed up a second test tube just to check solubility without Ether.
7 drops sulfer based polar molecule
14 drops glacial acetic acid
21 drops silicone (Dimethicone) oil, cosmetic grade, 1Cst.

Initially there was no mixing of silicone at all. The acetic acid + sulfer polars sank to the bottom. Upon heating tube (gently) with a 25Watt soldering iron, a substantial amount of silicone seemed to be absorbed by the polar mixture; maybe 1/3rd of it. 1/2 to 2/3rds of the silicone remained floating on top.

Upon running electric current in the cell, I would predict the silicone to mix more.
Glacial acetic acid has a non-polar end and a polar end; once the polar end is bound up in nickel ... the other side is free to dissolve in oil. eg: a bit like soap or detergent.



[Edited on 24-11-2023 by semiconductive]

I'm happy to read old sources from online. I can even run a few experiments in water. But, I don't think they are going to have a solution to my issues.

But: Your links may be valuable to other readers, too.

----

The current in #490 ether variation, has risen to 3 [mA]. No point in running this longer since current is high enough. All the active chemicals have had an opportunity to hit the anode at least once. The glacial acetic acid is not compeletely precipitated.

In spite of the small amount of silicone oil and it's hydrophobicity, the color green in the precipitate increased. Without speculating on it, here a photo. The amount of precipitate is in the ballpark of 7 drops worth of liquid. That means about half the acetic acid, or all the sulfur compound is on the bottom.

Liquid level has dropped, so a certain amount of the chemicals have been destroyed ( ?? Kolbe reaction, maybe ??) or evaporated.

I don't know why. I'd appreciate thoughts.

No obvious signs of plating.



------

And just for completeness, I let the pure silicone oil tube I mixed up in the previous post -- sit for a few days and it became colloidal. Added a couple drops of silicone oil and it immediately cleared up again but the separation became invisible. It *now* looks like the silicone totally dissolves in the acetate+sulfur compounds.
I wonder what's up with that.

Here's a shot of how straight silicone oil on top of acetic + sulfur oxide, looks like.
Current flow is in the 100 micro-amp region,eg: not even one milliamp. But I can see the silicone oil around the anode flicker and flow due to ions moving.

I'm re-using the anode after cleaning, which is why it's pitted.





After fiddling around with the anode, I noticed there are no bubbles formed unless it is near the very bottom of the tube; and then only the bottom part of the anode bubbles. I'm pretty sure this means that the sulfer dioxde bearing molecule must be denser and sinks to the bottom. Without sulfur to affect it, acetic acid will not release hydrogen at the anode.

It's just not possible to see the liquid separation since the densities are so close and the refraction, too. Current rose to 500 micro-Amps with the anode inserted fully. It's about half the current of the previous run, using ether, at the start.



[Edited on 26-11-2023 by semiconductive]

Hi, thanks for the suggestion.

I do have propylene carbonate reagent grade, (Southern tier scientic TM) and DMSO, among others.

I only have yellow prussate no ac-nitrile, eg: I have K Fe CN_x H2On crystals. Those are safe because the iron binds the cyanide. In weak acids (eg: acetic), it can be used without release of gas.


What would you suggest trying?
Just using 21 drops PC and 7 of acetic?
Or add NaCl, CaCl2, SrCl2, or something?


I'm definitely happy to attempt propylene carbonate.

But also I'm very curious why I have current flow, yet no deposits in the last two experiments. After a certain point, I would think all the acid is exhausted and combined with nickel; so, do you have any idea why electrical reduction continues to happen at the cathode ... but no material shows up on it?

When starting in the silicone solution, I only had 500uA current; but 10 hours later, I had 1.8 [mA]. So, what-ever the reaction is ... it can change conductivity.

The question will be the same if I use Propylene Carbonate, and it doesn't plate.
What 'side' reaction can go on nearly indefinitely, but deposit nothing?


Note:
#490 -- silicone + acetic + SO2 bearing molecule; after a full 24 hours.
Some precipitate visible, much less precipitate than last experiment;
low current cathode (-) on right shows some whitish deposit near the bottom, high current cathode (-) on left does not. Current is at 500 [ uA ]; eg: Current wen't back to what it was when starting. So under silicone there is a reaction that increases conductivity, but it's mostly used up within 24 hours.

Another dud experiment. Oh well, try again.



Let me know how you want propylene carbonate mixed, and I'll set up experiment #491.


[Edited on 27-11-2023 by semiconductive]

I found this, which is a little more detailed than the wikipedia article.

Kolbe Reaction to make alkenes


The article indicates that an alkalai metal is used Na, or K, during a Kolbe decarboxylation reaction. It also indicates that hydrogen is released at the cathode.

Similarly, for acetic acid, alkanes can form:
https://www.youtube.com/watch?v=be8D0kLeOSQ

There's no hydrogen being released at the cathode in my setup; and there's no alkalai metal in the mix; just nickel. It's not proof, that Kolbe reaction isn't happening; but I don't have the normal conditions for it.

https://www.youtube.com/watch?v=-OpWbSun2d8

I added 14 drops of acetic to see if that was the limiting chemical in the reaction; no change, stayed in the microamp region. Warmed it, no change. I could see the acetic acid mix into all the liquid, uniformly.

Therefore: I Added another 7 drops SO2 bearing molecule, and immediately bubbles form on the anode. I've got a full 3 milliamps of current flowing, now.

There is also layer separation visible where the sulfur molecule has sunk; and bubbles only form below that line. I think this means that Acetic Acid will misc with silicone oil; but a sulfur dioxide bearing material will be polar enough to sink (at first).

Also, the slight silvery color on the graphite electrodes remained when I added acetic acid; but etched off immediately upon adding the polar molecule.




[Edited on 27-11-2023 by semiconductive]

While I wait for a more specific suggestion ... I tried an ester + carboxylic acid, spiked with a tiny bit of lithium carbonate.

This fizzes, and of course hydrogen bubbles form at the cathode because as lithium plates out, it will attack even the ester and release hydrogen.

After running 12 hours, the current has dropped from 38 [mA] down to 5 [mA]. But it is now plating. This is very similar to what water will do.

Notice there is a thick plating of nickel on the high current electrode (-).
But, There is only black tarnish on the low current electrode (-).



The metal is likely not adhered, well; which is fine for a first test.


[Edited on 28-11-2023 by semiconductive]

Expriment #491
Choline chloride: 136 [g/mol]
Urea 60.0 [g/mol]

Eutectic is 1:2 molar mixture. eg: 1.133 x the grams of Choline as Urea is eutectic.

Trying .500 [g] of ChCl 98%, .440 [g] of prilled urea.
Put under 1 [mL] of kerosene before melting, to keep moisture out.

Gently heat until melted, then allow to cool.
Nice liquid layer separation; add fresh nickel anode (+) and two new graphite cathodes (-).

Assembles easily , first try, I like it.

Initial charge looks like this:



Graphites are well separated, although it's hard to tell in picture.
Apply current and I get.... current limit; 37 [mA].
This solution is very conductive; lots of gas at cathodes *and* on nickel anode.
Allow it to run ... and I'll report back, later.



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

Note: In my experience, thiourea (found in Broccoli!), is less reactive than urea.
I won't bother running the full experiment, but I'll melt a tube and see if it's even a eutectic or not.

Supposedly, it is:
https://pubmed.ncbi.nlm.nih.gov/24528755/

Choline chloride 136 [g/mol]
Thiourea 76.12 [g/mol]

A 1:2 mix, then has 0.89 the mass ratio. .5g CholineCl, .561g Thiourea.
Result: I get a paste, not a liquid. I'll let it set for a couple of days, just to see; but I don't expect any changes.

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

Back to the test at hand ... it's now 2 hours into operation.
The current limiting circuit is still pegged. It's conductive enough to go above 40mA, but my supply limits it to prevent over-heating and excessive currents.

As you can see on this shot, the low current electrode has nearly stopped releasing hydrogen. The high current electrode is foaming. The bubbles are penetrating/foaming into the kerosene region above, which indicates that the ionic liquid's viscosity is increasing. I'm posting an early shot, because this went caramel color within an hour, and it's dark carmel now; I think by 12 hours, it will be opaque and difficult to see.

No clear evidence of plating on the low current electrode, yet.




Edit: It has now been 24 hours.
The solution is completely opaque and black. Removal of electrodes does not show any silvery deposit; eg: no evidence of electroplating.

The solution is still liquid, although more viscous than at first; and more conductive.
The higher viscosity is annoying, though, because it causes the foam to float up in the kerosene.

What to do next?

Obvious idea:

The voltage being applied is too high, and decomposing the urea. Possibly lowering (clamping) the voltage might stop the color change and allow an extended run. (What voltage, though, and why?)

In the previous carboxylic experiment, the low current electrode remained tarnished while the high current electrode showed reduction. Just 'lowering' the current, then, could miss a properly working reaction. The low current cathode has shown no plating, already.

Mix up another tube, yes/no ?



[Edited on 29-11-2023 by semiconductive]

Last night, when I saw the carmel color in the choline, I suspected the tube was going to go totally black (which it did);

I turned my attention to the other experiment I was already running, the carboxylic acid + ester + lithium carbonate:

I put new cathodes in and cleaned the nickel anode.
Immediately the current rose from 5.5 [mA] up to 12 [mA].
This is fairly typical in many of the experiments I've run in the past.

I decided to lower the amperage to verify that any lower current will not plate. I reduced the current limiter to 3 mA. (This also reduces whatever voltage is at the anode to make 3mA).

Then let the experiment continue to run, overnight, on a fresh pair of cathodes.
Here's the result:



At 5mA, the right cathode will coat with nickel (badly, but it does it.)
But: In this continuation (of the same solution), you can see the bright plating has stopped.

There is a film of some kind on the right (higher current) electrode. But it's dark.

Experiment #492, start.
7 drops acetic acid, 21 drops ester with multiple oxygen count, sea-salt (NaCl) recrystallized with a tiny amount of NaCO3. (<1%).

The purpose of this experiment is to see if table salt will dissolve in acetic acid + an extra oxygen containing ester. eg: I'm replacing the SO2 bearing molecule with an ester, to see if the presence of extra oxygen might bias the gas release toward the anode.



The initial current was 100 [uA]. So, I heated the tube with the soldering iron, and watched the acetic acid attack the bicarbonate in the salt; resulting in the salt crystal slumping to the bottom of the tube. This ought to expose a lot of surface area to the liquid. Current rose to 3 [mA].

Therefore, this is an ion limited solution.

Bubbles do indeed show up on the anode, heavily. Only a very few gas bubbles on the cathodes; Both the high and low current cathode are slowly changing color with a whitish film developing on them.

This (mildly) supports the hypothesis that the presence of oxygen or sulfer. can act as a catalyst causing hydrogen loss to preferentially happen near/at the anode rather than the cathode.

I still can't rule out Kolbe reaction.

Note: 4 hours in current dropped to zero. Bubbles still happened on nickel electrode. I re-heated experiment, and bubbles switched totally from nickel anode to cathodes with current rising to 8 [mA]. Salt does not appear to have dissolved, appreciably.

I decided to redo the experiment: Fresh nickel electrode, cleaned cathodes and dried them, all chemicals the same. The only difference is I did not pre-tarnish the nickel in formic acid.

Result: Current again starts at 3 [mA] after heating; BUT -- Sigh...
No bubbles on the nickel, only a few bubbles on the high current graphite cathode.



Experiment did not repeat. Will have to retest ... again. Doh!



Experiment maintained 2 [mA] all night. Lots of precipitate this morning.
There is obvious white film on high current electrode, less film on low current (left) electrode. The film appears only on the side of the electrode nearest glass. This is not likely nickel metal.

No appreciable NaCl dissolved into solution.
The color green only appears on the Nickel anode. For some reason (either salt or ester) the solubility of Nickel Acetate is severely reduced in this experiment.

Will re-run it with no salt to see what happens on fresh nickel anode.


Edit:
7 drops glacial acetic, 21 drops ester, capped with kerosene as usual.
The attempt without salt fails. There is no conductivity even after heating. Total fail.

This means an ester with distributed oxygens does not have the same effect as a sulfer dioxide bearing molecule with concentrated oxygens. SO2 is able to cause conductivity in acetic acid and bubbles at anode, only. The original experiment #492 was possibly contaminated/anomalous. I'm not sure what was going on there. Best just to ignore it.

Will try an alternate salt.
Experiment #492 salt variation: Sodium Nitrite.

Immediately salt out-gasses, even before electrodes added to tube.
Conductivity is very low. (100 [uA] or less).
Heating raises conductivity to 2.7 [mA]. (Pretty typical.)
Everything Anode, Cathodes, and glass tube, continue to bubble slightly.
Color is changing rapidly:



Current dropped back down as liquid cooled. I added more nitrite salt and heated in order bring current back up. Current drops again as cooled. Added yet more salt ... and this time brown gas started forming in the test tube. I think I've just found another way to break Sodium Nitrate down into nitrous oxide. Experiment terminated.

Hmmm...
Will try again with potassium nitrate, to see if it reacts with the ester... then add glacial acetic acid.

Experiment #492, potassium nitrate variation.

7 drops glacial acetic, 21 drops ester, capped with kerosene as usual.
2 prill's worth ~20uL/AKA 1 drop volume worth of potassium nitrate salt.

Mixed ester and KNO3, first, heated gently until painful to the touch several times; at max heat a little gas evolves from salt, but I think that may just be boiling. There is no decomposition of the salt.

Added acetate, shook, heated, capped with kerosene.
It is conductive, but not even 3mA when hot. ~2 [mA]



In photo, clearly there is gas evolution at (+) anode.
Nothing on cathodes. Nothing from salt.

Whether this is Kolbe reaction, Acetic attacking the nickel (hopefully), or nitrate being decomposed; I don't know. I will let it run.

--- bubbles slowly but continually on nickel; conductivity has constantly risen 4 [mA ] now, at 19C. Originally solution began to color yellow, then greenish, but then suddenly formed colloid which is lighter green than the dissolved color.



No sign of plating, yet. I'll let it run a couple of days; probably centerfuge it, and see what the equilibrium chemistry is like. Might be an experiment worth studying to figure out the balance between colloid start and dissolved ions (in the future, not now.)



[Edited on 30-11-2023 by semiconductive]

I think I found another example.
Experiment #492, in order to get away from alkalai metals which can cause Kolbe reactions;
I want to try Ammonium Nitrate. This will not dissolve, though, in 7 drops of acetic acid and 21 drops of ester (or ether...,etc.)

It will dissolve in Formic acid, because of how small it is and similar to water.
I decided to mix up 7 drops of acetic, 21 drops of ester, and 1 drop of formic acid;
This ought to slowly dissolve a prill of ammonium nitrate.

Here's the experiment after running a few hours:



It has 5 [mA] flowing; which means the 7 drops of acetic acid I put in will be used up within 48 hours *if* the acid is attacking the nickel.

However, if you look closely at the cathodes; you'll notice the low current (left) cathode has a brownish coating on it; the higher current cathode, though, is clean.
There's not even a whitish film on it.

I've seen this happen several times in the past; and often (not always) this indicates a reaction that will never stop. Sometimes this kind of reaction will bubble gas for a few days, and then eventually stop gassing. But, the electricity keeps on flowing.
However, it never plates metal. But it often will show a build up of something on the low current cathode, exactly as seen in the picture.

I'm wondering if it's possible for metals to partially reduce at the cathode, say from Nickel ++, to Nickel +, and then return to the anode to be oxidized again?

Since the + electrode is always repelling positive ions; Can anyone explain the mechanism that might make this happen?

I have cellulose acetate which is permeable to +ions, and can make a separated cell; but if nickel is changing from +2 to +1 oxidation state; wouldn't it be able to pass both ways through the membrane? (Eg: a membrane wouldn't be able to force +nickel ions to stay in the cathode area and be fully reduced to metal).

Any thoughts on techniques to break this kind of cycle?

Also, I notice, the prill of ammonium nitrate continues to bubble as if dissolving; but it's not getting any smaller. This definitely puzzles me.


[Edited on 1-12-2023 by semiconductive]

Experiment #492 never plated. Fail. Continues to conduct electricity, with very slowly reducing currents.
Whatever gunk is on the low current electrode appears to be the only product.


Experiment #493.
Tackling the pH solubility question.

I know Iron, in water, is in the +3 form when pH < 4. From a couple of tests, I think it does the same thing as expriment #492. Everything turns to rust/sludge in the water, and does not plate metal.

I know nickel, in water, does not plate at pH<3, from sulfamic acid in water. But it will at higher pH.
Even though nickel is only supposed to have a +2 oxidation state for bonds, the previous paper reveals that with Chloride ion present; it attracts 3 negative chloride ligands.

So, I'm under the impression that it's possible that nickel in a waterless solution, is dragging along more vinegar molecules than it would do in water.

The normal way to solve this problem is to add a base to the solution, to buffer the pH.
Except that pure acetate doesn't ionize in the first place; but it does ionize in the presence of alkalai metals, or salts, which may be acting (artificially) as a super acidic pH since the acid molecule is in excess. Adding a metalic salt, then, would be counter productive.

It occurs to me that another common buffering agent is an extremely weak acid coupled with an alcohol.

I know that having multiple exposed oxygens (in pairs) is enough to cause anode erosion, and ion conduction.

So, it seems to me a possible choice of ionizing substance is to use Boric Acid + a poly-alcohol, and reflux the mix until a solid ester is formed. This might, hopefully, act to buffer the nickel solution to make it act *as if* it had a less acid pH.

7 drops volume of boric acid, refluxed with 7 drops poly-ol-alcohol, until very viscous/ nearly solid. Add to 1CC of pure acetic acid. Heat with soldering iron, gently, until dissolved.

Nickel electrode +, 2 graphite electrodes -.
Initial test of conductivity is good; the Boric acid's oxygen + alcohol, does cause the glacial acetic acid to become conductive.

5 mA current flows. I can see nickel haze forming slowly in the first half hour of operation.

Solution is rapidly greening, which suggests the alcohol + Boric acid might be abnormally breaking the dimmer bonds and increasing the exposure of free acetic acid to the nickel. ( It's not the boric acid, alcohol, by itself which is conducting the current; as the color is not green. )



This appears, even within three hours of starting, to be a successful plate; although it is bubbling hydrogen out at the cathodes.

There is un-evenness in the nickel plating; but that is normal when hydrogen gas is bubbling off the cathode. I'll have to wait to see if the hydrogen ever stops bubbling or not.




[Edited on 5-12-2023 by semiconductive]

It plates nicely at first, but the conductivity of the solution continues to rise; 10x.
The current density gets high enough that nickel starts plating black-ish, and brittle.

As the liquid gets near to saturation with nickel, the hydrogen evolution slows down.
But, it stops plating at that point. Current drops due to the liquid viscosity becoming extremely high; and it's opaque.

I tried adding an ester to dilute the liquid, which made it semi-transparent and the conductivity did rise again; but it didn't resume plating. This was a very interesting experiment.

#494, mixed carboxylic glacial + ether + SO2 bearing molecule.

This is the pre-plate phase, where I'm just running a titanium anode at 5-10mA.
I boiled the SO2 bearing molecule in ether, in order to drive off any water of crystallization.
Here, I'm using titanium to get rid of any remaining moisture.

But, surprisingly, it's plating titanium/rutile mixtures !!

On the low current cathode, the plating is metallic. On the high current electrode, it's blackened; also the titanium anode blackens as well.



My choice of SO2 radical bearing molecule, I've been varying to see if it makes any difference. There's solubility differences; but chemically, the same activity tends to happen when SO2 is present -- bubbling is suppressed on the cathode, and shifts to the anode.

I'm suspicious that the bubbling is "boiling" of the lightest carboxylic acid, because it refluxes on the glass of the test tube rapidly; and condcutivity increases as drops of condensate fall back into solution. Even with solution cold, the test tube fogs within a minute of starting the current. Electric ion boiling, perhaps ?





[Edited on 14-12-2023 by semiconductive]

Experiment #504. See if Acetone can complex divalent metals, such as nickel, and/or plate:

1 CC non-protonating solvent; Eg: THF, DMSO, sulfolane, polaxamer, ether, etc.
Test for conductivity before starting experiment, no-conduction indicates good purity.

Add 7 drops of acetone to solvent.
Cap mixture with kerosene to reduce moisture absorbtion.

I applied voltage. Current begins flowing from Ni(+) anode to graphite (-) cathode. ~4 [mA] (in sulfolane -- nickel anode stays clean. )
~2.5[mA] (in DMSO -- note: blackens nickel anode. )

Surprising:
Initial bubbling forms in liquid near the cathode, but not on the cathode itself. This is less apparent the lower the initial current is, eg: DMSO does it but not as strongly as sulfolane.

I can see the liquid mixing and going away from the cathode before bubbles form.
I'm not sure why/what this means. But perhaps the reduced hydrogen is initially soluble but negatively charged, and repelled by cathode? Curious!


Over night, solution turns a dingy yellow color. (Depends on the lighting, some angles look green others yellow.).
DMSO has the color but is more grey over-all; probably a side reaction.

This suggests that acetone can complex a divalent metal ion (transition metals, Ni,cu, etc. ) ( It's likely then that it will also work with trivalent Fe. Future test! )

Current level decreases to 2 [mA], over-night, and bubbling stops. Current continues, but no sign of plating.

Add 14 drops of acetone. Bubbling re-starts, but this time it is physically on the lowest part of the cathode. eg: quite normal looking.



I can see a little discoloration of the nickel anode; so perhaps there is some solid reaction product possible, too. EDIT: No, it was just pitting. The nickel is clean.

From studying acetone, I think that the alcohol/enol form of the chemical is less stable than the ketone form. I would suspect that this makes hydrogen evolution harder, since a more stable alpha hydrogen needs to be displaced before an enol can be formed. I hope this will make acetone more willing to give up complexed metal ions.

But: The hydrogen can be displaced/removed from the Enol form, and should allow acetone to chemically react with nickel: either one or two molecules per nickel ion:
I suspect: [ C3H5O- ]2 [Ni++] (I am open to suggestions, or corrections; recall, I'm an amateur at best. )

I think: Positively charged nickel ions, then, can be complexed by acetone and/or form a salt with an acetone, enol tautomer, molecule.

I find it surprising that no sign of plating occurs with this solution even after hydrogen evolution stops. I'll clean the anode, and wait until the second dose of acetone (+14 drops) stops evolving hydrogen.

This brings the acetone content up to around 50% of the solvent volume.

bubbling stopped several hours after re-starting and conductivity dropped to the micro-amp region. There's no way that all the acetone reacted, so it seems to me that the yellow/green product might be suppressing further ionization of the acetone.

Pure acetone, 1CC, will conduct about 3micro-amps for a few minutes due to moisture; and then drops to zero. Acetone by itself isn't conductive. The oxygen bearing solvents cause acetone to become conductive.


[Edited on 1-3-2024 by semiconductive]

Quote: Originally posted by semiconductive

#494, mixed carboxylic glacial + ether + SO2 bearing molecule. But, surprisingly, it's plating titanium/rutile mixtures !! [Edited on 14-12-2023 by semiconductive]

Quote: Originally posted by Twospoons

Quote: Originally posted by semiconductive   #494, mixed carboxylic glacial + ether + SO2 bearing molecule. But, surprisingly, it's plating titanium/rutile mixtures !! [Edited on 14-12-2023 by semiconductive] This is interesting - if you could get a plating of Ti4O7 that would be quite a valuable and unique accomplishment.

At my level, I'm just surprised to get any plating at all.
I don't know how to test which kind of titanium deposit I have on the cathode.

I've recently found a program molcalc, and have been drawing acetone in various configurations, to see if I can work out anything about which molecules would be stable in solution and which ones wouldn't for experiment #504.

eg: I can get the energies of vibration (etc.) and compare that against a molecule with nickel replacing the hydrogen in the enol tautomer. I was hoping I might be able to figure out a color for the ligand this way; but I am not there, yet.

The same ought to be possible with the titanium experiment, since that interests you.

I would be happy to re-run the titanium experiment in several different ways, and see what happens; but without some theoretical guidance from someone experienced, I don't have much hope of solving the technical issues of getting a particular kind of deposit.

To get Ti₄O₇ requires a deficit of oxygen.
If that's a darker version of rutile, it may already be happening.
I don't know.

I have a fiber optic spectrometer (ocean optics), basic reagents, etc. but I lack the knowledge/experience to make definitive tests about what kind of oxide is forming.

I would expect that whatever I have made (already) may be contaminated with sulfur as well. The anode blackens as well as one of the cathodes which has deposits on it.

If you are willing to make suggestions, I would certainly try your experiments and post results. Care to gamble?


[Edited on 29-2-2024 by semiconductive]

For example, I mocked up two acetone molecules in the ENOL form and joined them with a divalent atom. Nickel isn't available in MolCalc, so I used magnesium.

But, this is one possibility for what is happening to the acetone in the presence of nickel.

I see.
Well, whatever it is on the electrode; it *is* conductive and it does contain titanium.
It fails the test for sodium, and it is semi-bright metallic shine on the low-current electrode -- although I still think it's an oxide. (It's not pure titanium metal, whatever it is.)

You could replicate the experiment pretty easily using a piece of titanium wire.

I was attempting to dehydrate some reagant's water of crystallization.
Of the set I tried that day;
Sodium Saccharine is one that contains SO2 and is easy for anyone to get.

That particular experiment was a variation:
Ether eucalytis globulus (purified) or Tri-ethyl citrate as bought (both work the same), acetic acid, formic acid, as solvents.

1/2 cc TriEthyl citrate, 1/2 cc acetic, and enough Na-Saccharine to saturate it.
I boiled it with a 25 watt soldering iron in a glass test-tube until the saccharine was dry and steam stopped forming on the test tube (overnight); then I tried adding formic acid (just enough) to make half the saccharine dissolve.

This resulted in a clear fluid, which after many hours, stopped producing vapors. Which means the formic acid that was left became stable.

Then I ran electricity from a titanium anode to two graphite cathodes;
It was a very crude setup since I was just trying to get titanium to go into solultion and chemically combine with any traces of water that were left. I expected this to cause either precipitation, or to chemically bind the water to titanium which would then not affect nickel plating, since Ti is more reactive than Ni. Either way, it would do what I wanted.

60HZ , 125 VAC power, through a diode, through a 4.7K resistor into a titanium anode and a pair of graphite electrodes; one hooked directly to neutral, the other hooked through a resistor picked to reduce the current by about a factor of 10. ( I didn't write the value down. )

You're welcome to replicate it.
It's not my interest in this thread; but I'm happy to share any info, publicly as I stumble across things.


[Edited on 1-3-2024 by semiconductive]

Experiment #506.
I can't get NaCl to dissolve in acetic acid; but I notice that MgSO4 (epsom salt) is used by a few people successfully in water with vinegar.
So, I want to see if I can plate with mostly crystalline epsom salt + glacial acetic acid, diluted with an ester or ether.

I put about 1CC of "Dr. Teal's" epsom salt from the store into a test-tube. I expect this is contaminated with small amounts of copper, mangneese, and other metals. But, I can get reagant grade, later, if this works.

I added just enough glacial acetic acid to cover the MgSO4 crystals. Then I cap the experiment with 2CC's of inert kerosene.

My hope, here, is that the MgSO4 will tightly bind it's water of crystallization, but still dissolve in the acetic acid.

If it does, then during plating a certain amount of water of crystallization will be destroyed as hydrogen gas; and hopefully, whatever tiny amount remains will not want to separate from Mg into solution; eg: this might make a 'nearly waterless' electrolyte.


Nickel+ graphite-.

I added 1 drop SO2 bearing molecule to cause initial ionization.
Initial bubbling is intense as expected. After about a half hour at 10mA current on 1cm of 0.9mm diameter graphite; I can see silvery-coppery looking metal plating onto electrode. So, it does work (at least at first.)

After running several hours, I can see the epsom salt is dissolving into the glacial acetic acid. But it's becoming jelly like.
So, I add a few drops of ester or ether to dilute it.

Over-night, bubbling completely stops. Conductivity dropped to almost nothing, but it's easy to see the entire tube has become a gel again. Quite a bit of metal did plate onto the graphite. I replace graphite with a fresh one. Add 6 drops of acetic and stir.

Solution becomes liquidy but colloidal.

The new graphite appears to be plating (slowly). Most of the colloidal suspension falls to the bottom of the test tube as the solution becomes warm/hot.



There is much less bubbling this time around. I think this experiment is going to be successful. I need to figure out how much epsom salt is needed for a given amount of glacial acetic acid in order to make a final solution with minimum colloidal solids remaining. I think it's an optimization problem. Cool!

After 48Hours, pretty much the only place bubbles come from is the contact point between the graphite and the epsom salt. Bubbles are hardly forming (at all) in the liquid areas on the visible part of the graphite electrode. The current level has dropped, considerably, though. From 10mA yesterday, down to 2mA today. I'm going to allow it to run at 2mA for a full day and see if the plating properties change.

This picture a fresh electrode, that has been electroplating for 20 minutes at 2mA.



You can see: The liquid is still electroplating even in places where bubbles are not being formed, and the color of the deposit is about the same as at the start of the experiment.

It's not a normal nickel/silverly color, but is tinted salmon pink to brownish as if some copper were alloyed with it.

This is a bit puzzling. Nickel sulfate is bluish, Nickel acetate is greenish, and neither of these colors are showing as a tarnish on the plated metal;

There's likely a tiny bit of contamination metal in 1CC's worth of epsom salt. eg: But only something like 0.01%. That should have already deposited out by now!
I would expect to be getting a nickel plate at this point.

It makes me wonder if I've got some kind of magnesium-nickel alloy plating, or maybe a basic hydroxide. I'm not sure what color those would give. Google(TM) searches don't help much.


[Edited on 5-3-2024 by semiconductive]

Experiment #507
See if conductivity can be increased by using combination of formic acid (stronger), and propionic acid (weaker), than acetic.

Start: 1/2 cc roughly of MgSO4.
Cap with Kerosene.

Boil with 25watt soldering iron for about half an hour. Attempting to dehydrate MgSO4 partially by getting rid of only excess water of crystallization / water that is extremely weakly bound. I hope this will cause experiment to take less time to run without changing final results.

Fill with Formic acid until 1/2 cc space above MgSO4 is formic acid.
Nickel+, graphite-
Run for 48 Hours.
Current is notably less than many experiments I've done in the past using formic acid alone.
Green crystals form on the graphite electrode, and have to be removed many times.
As water content drops, the amount of crystals forming on cathode decreases. Eventually it becomes a non-problem, after several cleanings.

At this point the MgSO4 is colored a bluish-green, uniformly, and tends to become thrixotropic/solid when allowed to settle.

I decanted off remaining formic acid by slowly turning test tube and allowing liquid to drain. All green material stuck well to the bottom of the test tube. This ought to be nickel-diformate and formic acid crystallized MgSO4.

Add propionic acid (which normally will not conduct electricity by itself).
Conductivity is very good (30mA), which is likely due to Formic acid dissolving in propioninc acid.
Lots of hydrogen bubbling (undesirable).

Picture after 12 hours of current applied to propioninc mixture.



Experiment hasn't been run long enough.
I think that's a crystal film on the graphite electrode, it doesn't appear to be quite metallic nickel. Maybe nickel propionate/formate mixture.

Hydrogen bubbling continues, current is still quite decent >10mA.
Will let run over the weekend, and will clean graphite electrode a few times.

Interesting how the nickel formate in the salt has become hetrogenous, with dark lines of Ni-carboxylate showing up an arbitrary distance away from the electrode. I'm not sure why that's happening.

Obviously a crystalline precipitate is sticking to the electrode (-), this March 11th,2024 morning. No signs of Nickel plating Metal.

The line of material that was greenish in color, 48 hours ago, is still green but it looks like it's darkened (concentrated) a lot. Compare the pictures. I'm puzzled.

Conductivity has dropped, It was only 100 uA, this morning. As I cleaned the electrode, scraping off a hard crystaline deposit / perhaps an oxide/hydroxide, I notice that the epsom salt continues to bubble even without the electrode present.

After cleaning, reinsertion of the electrode shows 0 uA current (immeasurably small.). Bubbles resume forming on the graphite electrode, though, so there is a current at least in the liquid if not in the attaching wire.

I'm really not sure what to think about this experiment.
It's only a small deviation from the previous experiment, but it seems a total failure.

Oh, one detail. Propionic acid wouldn't evaporate in these conditions, but the liquid film is not visible. The current loss may be due to absence of liquid solvent. I'll add a few drops (7) more propionic, and see if any change.

edit: +7 drops, still couldn't see film; it seems to mix with the kerosene a little too easily/abnormally. Added +8 drops more, then gave up. I can see diffraction swirling in the kerosene, so maybe propionic acid will settle out in a few hours.

I checked the nickel anode; it had completely dissolved below the epsom salt layer.
I replaced it with a new piece of nickel; current resumes at 400 uA.

Edit 2:
Conductivity dropped again over a period of hours, no settling of the propionic acid is visible. I took a glass stir rod, and tried to stir the epsom salt dust, but it had aggregated. It was hardest near the graphite electrode, and softest near the nickel source. I pressed down hard, and was able to break it up again. Current rose to 4mA for a while, and then dropped back down to about 2mA.

I don't know if it's thermally induced, or a chemical reaction near the graphite; but on the glass, on the graphite, anywhere near the cathode; crystals seem to grow rapidly in response to electric current flowing.

After breaking up the salt with a glass stir rod, I could see a meniscus about 2mm above the salt form. So the propionic acid is separating out now that I've homogenized the epsom salt (again.). It's hard to see in photo, but above the meniscus uniformly across the test tube is a haze of crystals in the kerosene sticking to the glass. Just below that; Where the propionic acid accumulates, the crystalline dust mostly washes off glass.





Note: I did not shake or stir the test tube between these two photos. They are with an hour or two of each other. The bubbles inside the epsom salt of the upper photo have expanded and knocked the salt around. Gas forms inside the salt, near the electrode. I suppose the earlier photos of the dark band, might be a channel where tiny gas bubbles formed and it's shape is just a matter of random luck.

The color is lightening, slowly, in the salt. That would suggest either dissolving of the darker green nickel formate from before, or a chemical reaction of some kind. The current level is staying low, so I think it's more likely a reaction than dissolution of a conductive salt. Hmmm... it's not headed in the correct direction to ever plate.


[Edited on 12-3-2024 by semiconductive]

I've waited long enough to exhaust the 7 drops of formic acid that I added. After a week of electrolysis, there is no more of the green lines forming in the powder. The powder has an aqua-marine, or jade hue. Occasionally, a very dark green powder will form on the anode; but after shaking or stirring the powder, it re-dissolves.



You'll notice that the propionic acid has a greenish hue (looks golden yellow when viewed with eyes.) So, something is now dissolving into the propionic acid which was clear in all the earlier photographs. But: The reaction/solution does not appear to be stable, as I can see color changes in the solution which will precede different kinds of plating activity.

Yesterday night 12 hours ago, there was a silvery black deposit forming on the electrode. It was partially metallic. But, then the solution's color changed and it rapidly began forming crystals of the epsom salt/nickel formate. It's as if the solution turned slightly colloidal, and electroplating causes particles floating in the solution to adhere/precipitate on the cathode and (also) to the glass of the test-tube near the cathode. It's electrically induced crystallization.

This is a very interesting experiment. The conductivity does not drop to zero when the acid charge is basically depleted. Each time I add an acid charge, there is more conductivity after the formic acid is used up. The first time conductivity dropped to 500uA; The second time conductivity dropped to 2mA; and after this last charge of 7 formic acid drops -- the current drops to 3 to 4mA. (3 with crystals on electrode, 4 when electrode is clean.)

I'll see how much nickel I can get to dissolve into the solution and epsom salt, will post results below. I think nickel salts (double salts, perhaps?) are increasing the conductivity. Propioninc acid isn't conductive by itself, but it does appear to dissolve nickel formate / nickel sulfate / whatever I have in there.


[Edited on 17-3-2024 by semiconductive]

#507 is a very frustrating (and long running) experiment. The color of the salt slowly darkens (greenish) with time.
Often the solution will show color change from yellowish to greenish, and then suddenly will plate a black coating.
After that the current drops and plating stops for a while (just when your hopes are up that plating has begun!).

I've figured out that currents higher than 8mA are vaporizing the acids. To get maximum nickel dissolution, it's necessary to keep the current below 8ma on 1 inch or more of nickel.

I'm letting the experiment sit, and seeing if dissolution happens slowly from the salt. Perhaps I can get it to plate black a few times in a row, by just letting it sit for a few days between attempts.

On the bright side ... I've found amidobutyric acid can be dissolved in small amounts of propionic acid & many different esters / ethers.

Today's picture is aluminum going into solution and trying to get rid of water before attempting nickel. As the reaction proceeds, the color of the liquid becomes less colloidal and more clear/golden.



I see a very thin but noticable aluminum / perhaps oxide film on the graphite; although no metallic buildup. At higher current rates, the film is more obvious in the kerosene area ... but etches back off with heating of the solution.


[Edited on 29-3-2024 by semiconductive]

Well lookie there!
It's aluminum plating near the tip (only) of the electrode in a 4-aminobutyric acid (not amido-sorry about the typo in previous post) solution. I haven't even added a nickel anode or bromine, yet. This is still the moisture destroying phase of the experiment before I actually intended to run a metal plating test. I've never seen this happen before!

Usually I centerfuge out the aluminum hydroxide, and then put a nickel anode in the tube.

The plating been getting better, but only at a certain depth in the solution.
Notice the very bottom of the electrode is not plating, and nor is the region close to the kerosene.

(You can see the hydroxide, so I think maybe a significant amount of moisture or oxygen is still penetrating the kerosene layer which is meant to reduce moisture, sigh.)

I bought a compressed gas tank to supply an inert shielding gas. I'm going to try drying CO2 gas with a dessicant made specifically for that task, and then slow flowing the gas to prevent atmospheric oxygen getting into the test tube as air or water. Hopefully, I'll have that working in a week or so.



[Edited on 2-4-2024 by semiconductive]

I'm still waiting April 2024, for CO2 tank parts. When I get the correct safety rupture disks, and a stainless mesh to hang sorbead blue dessicant inside the tank with; I'll take photos. Anyone work with sorbead before? The stuff looks like silica gel with blue dye in it. I just got Jupiter Pneumatics (TM) stuff off ebay.
Hopefully, I can bake it in an oven hang it inside the CO2 tank, and it will wick up traces of moisture in the CO2 that's put into the tank.

Today's shot is again of the preparation phase. I threw out the last test tube, and setup for a retry. This time I'm using butyl buterate as the solvent (which, unfortunately miscs significantly with the kerosene), 4-Gaba & and aliaphatic hydrocarbon plus a prill of solid iodine for ion carriers.

Iodine dissolves in the aliphatic hydrocarbon, which dissolves in the butyl buterate, quite well. Normally, literature i've seen online suggests using Bromine to plate aluminum. But considering I got it to plate a little bit with no halogen at all, I decided to try a safer/cheaper iodine and see if it improves the plating. Normally, by the 24 hour mark, iodine color would have lightened and visibility in the tube would improve. This isn't normal.

But the answer, is *yes*. Iodine and/or aliaphatic hydrocarbon, and/or butyl buterate imrpoves the aluminum plating.



The original aluminum anode is most visible all the way to the right, against the glass, and is darkened with iodine. The bright silvery thing, just to the left of it, is an electroplated graphite electrode.

If you notice the tip, it's not plated the same as the whole body. This is the same issue that happened in the last run. It looks like a hexagonal set of dendrites on the tip with some graphite exposed.

The adherence of the aluminum in the last experiment to the graphite was not very good. I could slide an aluminum oxide ring off the graphite. In this one, there is a thick enough plate that it should be more difficult to get off. But, it still might be contaminated with organics heavily (don't know.) There is *much* less hydroxides/precipitates in this experiment than the last one. Current is on the order of 3 mA at the start, and 800 micro-amps this morning. Still, for a second try, I couldn't ask for more.

Until I get Argon or CO2 shielding gas, I don't really have a way to completely stop oxide formation.


[Edited on 8-4-2024 by semiconductive]

Experiment #510
A nice separation,

I've been trying to make alkoxides (strong) to attempt plating in.
In general, I've been attempting to use K-OH and electrolysis, to make a progressively stronger and dryer alkoxide.

But; most solutions turn super dark brown or black during the process.
I tried 1,3 propane-diol, glycerol, and di-propyl-diol (food grade, O bond in center.). All of them turn brown (or reddish-brown) during electrolysis.

I've also noticed that formic acid tends to have a brown color after exposure to air/light for a while. I'm not sure if this is oxidized chemical, or if it's an impurity. But, since it makes the experiments opaque; I decided to look for ways to remove the color (if possible).

In this experiment, I used ~1 cc K-OH crystals, then I added enough 1,3 propanediol to fill all pores and cover the crystals.

Then I added 1 CC kerosene, as an inert cover layer in the test-tube.

Since dissolving KOH is a problem, and tends to leave chunks and solidify during electrolysis; I decided to use a few drops of formic acid, to bring the bulk of the potassium into a liquid solution. I figure that the water formed, and possibly the K+formate, can be removed by electrolysis.

I heated and stirred, and slowly added just enough formic acid (drop by drop) to cause almost all the KOH to dissolve, but leaving a trace of it.

Then I added a few drops of poloxamer-124. This is much less dense than the potassium/diol/formate solution and slightly soluble in kerosene. I was hoping it would mostly float between the kerosene and alkoxide mix; it does (very well!)

Then I placed a graphite cathode, with an insulating sleeve except for the very tip, into contact with the diol/KOH solution; and I placed an anode *above* the poloxamer 124 layer by about 0.5 mm. This is close enough that ions in the poloxamer will penetrate the kerosene and hit the anode. The anions then electrolyse to release oxygen (from water). I think formate ought to turn into some kind of gas (An alkene or CO2, Kolbe electrolysis; this is not an acidic electrolyte, but it's CO2 when acid as shown by Woelen: https://woelen.homescience.net/science/chem/exps/precision_e... ).

This result appears to be 4x (or more) gas bubbles at the cathode compared to the gas at the anode.

As the reaction proceeds, the slight brownish color in the KOH/diol layer migrates *up* into the poloxamer 124 layer.

This does clarify the KOH/diol layer. Heating the glass test-tube with a soldering iron speeds up the color migration and clarification a lot. It doesn't cause re-browning, either.

I've run it for 24 hours, and the KOH/diol layer is very nearly clear; with the poloxamer layer increased in width (about double original) and turning medium brown. It's a very successful separation. This should (I'm hoping), allow a build up of alkoxide at the bottom of the test tube.



I'll allow this to continue to run; and see how it proceeds.
The current is resistor limited to roughly 30mA or less.



[Edited on 24-5-2024 by semiconductive]

Experiment #510; postlog.

The use of poloxamer to separate our darkening electrolysis impurities during the addition of KOH, works reasonably well.

Electroysis of the resulting mixture for about a week with a graphite anonde, caused the solution to become a very light amber color which could not be removed by more poloxamer. Still, the solution is clear enough to see the effects going on, which was my goal. The graphite electrode was oxidized at the tip, and I suspect this is the source of the amber darkening that poloxamer could not remove.

I switched to a nickel anode and attempted plating for a week. The solution turned green, slowly, which suggests that Ni + a di propyl alcohol, or potassium makes this color.
No metal plating was observed (at all). But, there was a continuous whitish crystalline formation on the electrode. Conductivity slowly dropped over a period of a week.

I wonder if this could be K2O, since I'm destroying water; and I would have expected nickel to remain green. Anyway, it's run nearly to exhaustion, and the green is strong enough that it's getting difficult to see the electrode; I could add ester and centerfuge in order to make it transparent enough to watch again; but the conductivity would be very low and there's no sign of metal deposit which I would have expected by now; so I'll give up.

Quote: Originally posted by semiconductive

But; most solutions turn super dark brown or black during the process. I tried 1,3 propane-diol, glycerol, and di-propyl-diol (food grade, O bond in center.). All of them turn brown (or reddish-brown) during electrolysis.

Yes, indeed. I have both urea and sulfourea, and happen to be running an experiment since Wednesday on both of them. xUrea+Formic acid (1 drop), in an ester.

I haven't gotten metal plating from a solution containing either of them; but they are quite conductive.

Urea also works well for increasing solubility of various salts both in molten form, and not.

My experince so far: Urea has a low electrochemical resistance window, and tends to be a little too easily broken down into hydrogen, CO2, and ammonia. Sulfourea is less reactive. I think it may be more reducing.
But that causes precipitation of sludge in a lot of experiments.

I tried urea + CholineChloride a few months ago at ?Draconic Acids? request; It ended up turning black over-night due to electrolysis and plated sludge, only. I'm not sure how to tame it ...

Idea are welcome. I am an experimentalist. I like to test out ideas.

Quote: Originally posted by semiconductive

Looking at the paper, the first thing that comes to my attention is these are hydrated salts (quad-hydrate.)

Quote: Originally posted by semiconductive

Is there any particular reason you found this paper interesting; or is it just the fact that it's molten at very low temperatures?

Quote: Originally posted by semiconductive

In your copper plating solution with Urea, what source did you use to obtain the copper oxide? I have copper sulfate on hand, which is easy to convert to copper carbonate; but I'm not sure if I hit the sulfate with a torch, if it would decompose to an oxide easily.

Quote: Originally posted by semiconductive

If it had enough conductivity to plate, it's an experiment I'd like to try replicating even though it's not nickel based. I am able to plate copper in water, very successfully, but I've not tried it in ionic liquids yet. If you were able to get clean copper to plate out of the solution, I'd be amazed.

Quote: Originally posted by semiconductive

Hmm ... I'll have to try this again, later, with double the sodium carbonate and see if the conversion efficiency goes up. I got more than half converted, so it seems at least some Sodium Hydrogen sulfate was formed.

Quote: Originally posted by semiconductive

About how much urea, do you think, would be appropriate for a first attempt?

Hmm.
The copper reduced to a fine oxide powder in the test tube; kind of greenish-grey/black. ( Slightly less than 400C cooking temperature, roughly 35 watt ).

It's a bit of a pain to work with this small amount of copper oxide (I'm mildly disabled. )

So, I'm going to just put about 1/2cc worth of urea prills on top of it, cover with kerosene (which just acts as an air barrier), and melt the urea into the powder. This ought to make bubbling from 'fizzing' as you called it, visible.

I can make more copper oxide, later, if this shows any promise at all.
Picture of test tube just before capping with kerosene.

Quote: Originally posted by semiconductive

So, I'm going to just put about 1/2cc worth of urea prills on top of it, cover with kerosene (which just acts as an air barrier), and melt the urea into the powder. This ought to make bubbling from 'fizzing' as you called it, visible.

It will be very visible: it will foam a lot. I did today as you said, putting urea on top of copper oxide but without kerosene. I don't know if it would stop the foam from rising up the test tube (probably, yes). When solidified, it looks like the picture below.

A couple of pictures of the copper complex (tetraammine, or so it seems) dissolved in molten urea from yesterday.


Edit: Typo.

[Edited on 7-7-2024 by bnull]

I have a lot of things to report. I took a bunch of pictures, and can post images for whatever is of interest; but I'd rather not post 100 shots of nothing...

I put kerosene on top; and heated the kerosene rather than the urea. This caused a slow melting process of the urea and gave me good control. There was significant gas coming off the urea as it turned blue. It did not foam, at all, in my experiment.
It just released gas, some of which I think was steam.



Conductivity rose before the urea had melted. I used a copper anode and a graphite cathode. It rapidly rose in conductivity as it started to melt, in fact it was too conductive. I had to put a resistor limiter in to reduce the electric heating. I've already got a 30W soldering iron touching the test tube from behind (the philips screw on it is just visible in the picture. ) This is enough heat.

Overcurrent, like I had, creates side products. I'm not sure how much it might have affected the rest of the experiment.

It's pretty obvious that the copper is extremely reactive with urea. And from the odor, I think copper oxide encourages transformation of urea into ammonia when water is present. The slight ammonia odor went away as the reaction progressed.

For my next attempt, I won't bother to oxidize the copper; because I'd like to test if urea is reactive enough to decompose carbonate; and if the ammonia smell will be avoided.

A couple of times, I got it too hot; but quickly reduced temperature. If significant gas came off the anode, I would cool it.

As the urea finally melted, I could see a thin layer of copper form on the graphite electrode. It re-dissoved shortly thereafter. The color reminds me of ammonia-copper etching color, sort of off-color salmon pink/brown. Once the urea was finally melted, I removed the graphite electrode; cleaned it, and replaced it in the liquid to make sure there was nothing on the electrode interfering with plating.

No furthur plating after returning it to service for two hours. The current level was easily adjustable from 3mA to 45mA, and gas in proportion to the current was forming at the electrode.

The color was so dark blue that it was difficult to see the electrodes. So I added about another 1/2 CC of urea prills on top. I watched the blue liquid wick up onto them as they melted. It's quite obvious that the melting point of the blue liquid was *much* lower than that of the urea prills. And it's also obvious that adding extra urea was raising the melting point of the whole mixture (not desirable.)

The liquid slowly darkened again, and I thought I had added too much heat. Since the tube was messy, I decided to change the kerosine and the test tube. After letting it solidify, I decanted; broke up the solid urea and transferred it to a clean test tube. I could see the solid chunks were slightly greenish in color, except for where the air could reach it, where it became pure blue again.

This makes me suspect that the blue color is actually a hydrated ammonia ion / ligand. But: Under kerosene, (hypothesis), the mass turns green slowly as it drys; and it's melting termperature rises, cauing solidification.

After quite a while, it would no longer melt. I raised the heat, and am not sure if I caused decomposition or not. But, it still conducts electricity even when semi-solid. It is not until the temperature becomes almost room temperature that conductivity stops.

At this point, I decided to add an ester that is liquid. I was hoping it would lower the melting point of the mixture. It did; although most of the mass remained solid. However, after about an hour and a half I could see the ester starting to brown; which means the heat was well above the melting point of normal urea.

On the other hand, in the presence of ester; copper began plating on the electrode.
It's dingy, not bright; but it still was plating out; This is at 2 [mA] current.

Shall we try again?
Suggestions / questions ?






[Edited on 7-7-2024 by semiconductive]

Quote: Originally posted by semiconductive

I put kerosene on top; and heated the kerosene rather than the urea. This caused a slow melting process of the urea and gave me good control. There was significant gas coming off the urea as it turned blue. It did not foam, at all, in my experiment. It just released gas, some of which I think was steam.

Quote: Originally posted by semiconductive

It's pretty obvious that the copper is extremely reactive with urea. And from the odor, I think copper oxide encourages transformation of urea into ammonia when water is present. The slight ammonia odor went away as the reaction progressed.

Quote: Originally posted by semiconductive

As the urea finally melted, I could see a thin layer of copper form on the graphite electrode. It re-dissoved shortly thereafter. The color reminds me of ammonia-copper etching color, sort of off-color salmon pink/brown.

Quote: Originally posted by semiconductive

I could see the solid chunks were slightly greenish in color, except for where the air could reach it, where it became pure blue again. This makes me suspect that the blue color is actually a hydrated ammonia cation. Under kerosene, the mass turns green slowly; and it solidifies.

Quote: Originally posted by semiconductive

After quite a while, it would no longer melt. I raised the heat, and am not sure if I caused decomposition or not. But, it still conducts electricity even when semi-solid. It is not until the temperature becomes almost room temperature that conductivity stops. At this point, I decided to add an ester that is liquid. I was hoping it would lower the melting point of the mixture. It did; although most of the mass remained solid. However, after about an hour and a half I could see the ester starting to brown; which means the heat was well above the melting point of normal urea. On the other hand, in the presence of ester; copper began plating on the electrode. It's dingy, not bright; but it still was plating out; This is at 2 [mA] current.

Quote: Originally posted by semiconductive

Shall we try again?

Quote: Originally posted by semiconductive

Suggestions / questions ?

Don't apologize! I write books unintentionally for posts, darn it!

In the presence of water, I have observed esters breaking down into their constituents whenever acid or strong base is present. But, I didn't add any ester until after the solution had become nearly solid and green under kerosene.

On the other hand, I think there are too many steps and possible side reaction variables for me to really guess whether it could or could not. I'll need to do additional tests to figure that out.

The ester I used this time was citrate based, because citric acid is resistant to oxidation. I don't have enough experience to know for sure whether it would react since this is my first attempt at copper; BUT: It generally *doesn't* react with nickel ions present.

I don't know how the ethanol I esterified it with could get displaced when the urea has already been cooking long enough to drive out any free water and is solidifying even at high tempertature ?

I didn't see bubbling when adding the ester; so I don't think significant ethanol was released as it was hot enough to boil. eg: Adding an ester did not make a visible reaction before electrodes were re-inserted. The picutres are bubble free.

If it did react with electrodes, the obvious product would be copper citrate dissolved in molten urea. I could make copper citrate either in water, or perhaps (to avoid water) in methanol, DMSO, etc. ( Or suggest something simple! ) I have the citric acid; would that answer the question?

The article you linked was interesting. I have a working visible Jaz spectrometer, which could characterize the green or blue color spectrum; but nothing that can take and compare against the FTIR spectra listed in the document. I have very little way of figuring out what kind of ligand I have in solution, and my chemistry experience from is from college is 35+ years ago, and was my worst subject. I'm a honors BSEE, but not even an undergrad chemist, here.

I'm pretty sure, though, that Molten urea won't be fully spectrum tested by me; the square optical vials I have are disposable plastic and would melt. I could take surface reflectance spectra of molten urea with a surface reflectometer in a glass test-tube, but it would have to be manually done since the drivers for Jaz spectrometers are written in java and crash on linux systems. . I'd love to be able to monitor ligand concentration by color, but I'm limited to manual spectroscopy at the moment unless I can find an open source Linux driver that *works* with ocean insight spectrometers. Their support is not very good.

Wax; yes, I have that and occasionally use it. But it's a pain to clean out of test tubes. I have silicone oil, as well, which is less flammable; but I've never had a flame problem without chlorates or oxide-nitrates involved; so I'm not too concerned yet. The quantity of kerosene I'm using is so small, (Those pictures are through a microscope lens) even if it goes up it just is like a match striking.

But: I generally see smoke gathering in the tube slowly before ignition; and then there's usually not enough oxygen because of the smoke filling the test tube.

Squirrel knocked over the CuCarbonate I set outside today; will have to re-precipitate a new batch.
Look for post below ... eventually.

Quote: Originally posted by semiconductive

I'd love to be able to monitor ligand concentration by color, but I'm limited to manual spectroscopy at the moment unless I can find an open source Linux driver that *works* with ocean insight spectrometers. Their support is not very good.

I prefer something that actually works on a Raspberry PI™ which is the educational platform, open source, (de-facto) used throughout the world; and using it without charging me industrial prices for a hobby.

I don't mind writing a python script to control soldering iron temperature in response to SeaBreeze feedback.
I'll attempt to use a USB RS232 dongle, to turn on and off a 110VAC outlet so I can temperature regulate the test tube via soldering iron and monitor the color change vs. time.


[Edited on 9-7-2024 by semiconductive]

This bug has been known since 2017, fix is to set compiler to --std=c++03
https://sourceforge.net/p/seabreeze/tickets/34/

But still no joy, new error:

re-running make now complains about a missing USB.h, that is *upper* case; from the same line!

I know that file exists in SeaBreeze as /include/native/USB.h
What educated i***t made gcc report an upper case file as lower case, before...
Oh well, I've shown how to update Rasperry pi ™ as a bonus.

Editing, src/native/usb/linux/NativeUSBLinux.c
commenting out line 35, since <USB.h> is a c++ header, and this is a *C* file.

And now I have a bunch of undefined linux functions. Replacing line 35 with: #include <llibusb-1.0/libusb.h> // fixes nothing.

Checking kernel headers, and Linux USB native does not have the missing functions .

Checking the SeaBreeze readme.txt, I see it! They want libusb version 0.1 for Linux.
I'm not doing that. I'd end up breaking my raspberry pi which uses version 1.0 already.

so, time to manually upgrade seabreeze to usblib-1.0.
Small headache! I've got to read a bunch of api names, and changes.


More to come *if* I'm able to debug SeaBreeze device driver....


[Edited on 10-7-2024 by semiconductive]

I successfully ported SeaBreeze's native linux USB to use a modern libusb-1.0
SeaBreeze's USB interface will now will compile on any Raspberry PI.

I see the jazUSB being built, which is my spectrometer.
Therfore, I'm getting excited This might actually work, and not be a waste of time.

The package is almost completely built, but ...

In file src/common/Log.cpp, false indentation had to be deleted in two places.

And after that, I got this weird error which took a while to figure out:

And, finally, that's the last of the bugs.
Make finishes building SeaBreeze with no errors.
Therefore: I have all of SeaBreeze and test programs installed on my Raspberry PI ™.

Now to actually find a USB cable, plug in my spectrometer ... and see if it works.

EDIT: Yes !!!!!!!!!! it DOES! This is AWESOME.

There is a single bug that I noted during testing; Linux does not call USB close when when signals are caught. If your seabreeze application is killed by a unix signal; the USB interface will be left *claimed* as if open. This results in the spectrometer not being openable again, until a USB reset or power-cycling happens.

I have a Jaz, so I reset my spectrometerer with vendorID (Ocean Optics), 0x2457, and product Jaz 0x2000.

Copper sulfate, pentahydrate 255 [mg] in 45 ml R.O. water.
+ 170 [mg] NaHCO3 in 5 ml water.

Extra metallic support on 5 sides, makes a partial Faraday cage.
My experiment ought to be safe even if lightning strikes.

Scheduled to change water 3x rinsing precipitate thoroughly, through Saturday 7/13.

There was a slight loss of material during transfer; and I forgot to cover the beaker with a watch glass.
So, this might be 1 or 2 mg in error.
The squirrels didn't invade this time, though. Success is good.

The precipitation gave 110 [ mg ] of basic carbonate from 255 [mg] of pentahydrate.

I have 1.02₁ [ m · mol ] of copper

Cu=63.546 [ g / mol ] CO₃=60.01 [ g / mol ] OH- = 17.008 [ g / mol ]
Cu₂ · (OH)₂ · CO₃ = 221.₁ [ g / mol ]

Assuming perfect conversion to basic carbonate, I ought to have: 112.₉ [mg] of product.
110/112.₉ ≈ 97.4₃ % conversion.

I've gotten pretty close to theoretical maximum yield.

My spectrometer is also giving me data samples, although the data is coming into my laptop raw. It doesn't give the data after the spectrometer processed it. A bit annoying. The data also has noise spikes.

I'm writing some software to filter the data, statistically, since I don't have an operational ocean optics software package to process it. I'll get a good color scan of the basic copper carbonate, and post it (eventually) below. I'm slow, but I am getting there!

I've also ordered some bluetooth outlets which will allow me to robotically control various apparatus during experiments and do closed loop temperature control.

And I bought two Owon BT41 Digital multimeters, that can monotor temperature and current, during the experiment; but just realized I should have bought a third one to monitor voltage. Oh, well, for this second experiment we'll just have a current profile and temperature record.

I tested my DMM's out today, and am able to control them perfectly and record data on my Rasperry Pi.
So, I ought to be able to build a precision automated test bench, relatively soon.


The seller of the Raman spectrometer head hasn't responded to my inquiries. So, I don't think I will risk buying it. Boy, I'd love to have that, though!

Quote: Originally posted by semiconductive

I'm writing some software to filter the data, statistically, since I don't have an operational ocean optics software package to process it.

I'm winning the battle. I got a background noise fingerprint.

My Jaz spectrometer sensor looks like this in the dark, with thermal drift happening over 8 hours.



Each line is 4096 samples, at 1 second apart; so roughly an hour's worth of data.

The noise is fairly repeatable, like a fingerprint, but I see a few of the sensor noise peak's change visibly in the plot ... which is not good.

I'll probably let this run for another 24 hours, use the data to see if any of my pixels are problems; Then see about writing a filter program to reduce errors by injecting anti-correlated noise.

After that, I just need to calibrate the intensity scale against the built in Tungsten lamp. eg: It's time to look up blackbody radiation curves ... and refresh my memory. Power per frequency, not power frequency is what I recall the formula being; so I need to figure out how to convert power density into total energy at a wavelength.

I bought some TPLink™ KASA™ plugs, since those are on the market and can be controlled by a Raspberry PI™ using python; therefore qualified people with safety equipment can duplicate what I've done. This is just for educational purposes, only and I'm not recommending it. The KASA plugs use reed-relays, which means they will wear out; But, it's good enough for running a soldering iron. I'll report any failures if they occur and lifespan.


This is a picture of urea under kerosene with the soldering iron in the background.



There is about 3cc's of kerosene in the test tube, and I have metallic shielding in case of fire as I don't trust KASA yet to run un-attended. There is roughly 1/2cc's worth of Urea Prills in the tube with a 30W soldering iron behind it. This is wired solidly to a chemistry clamp stand.

My multimeter says the tube is at 132.1 [ °C ], with the KASA program pulse width modulating the iron at 72% of full power via 1.5 second long pulses. The tube's been heating all afternoon and stays within ±1 [ °C ] stably.

Notice: no prills are melted, yet. I expect the Urea is now quite dry, with excess moisture driven off.

At room temperature, my ohm meter's thermal probe registered 3 degrees celsius higher than the wall thermostat does. So, it's possible the temperatures I report might need adjusting downward by up to 3 [ °C ]. I haven't calibrated the probe more precisely than that.

The wall barometer, shows 29.85 inches of water at my elevation. But, kerosine at these temperatures is not very volatile and will remain in the test tube for several days without issue.

A quick web search shows the lowest listed melting point of Urea as, 132.7 [ °C ]
I don't know how accurate that is.

https://www.accessscience.com/content/article/a722900

I will continue to raise the temperature this afternoon at 1 degree C every two hours, and report what temperature the urea prills start to melt at. I'll be checking for decomposition smells, and stability of the liquid also.

No odors, yet, except for slight oily feel of warm kerosene.

The following program is what I'm using; and if it gets interrupted during an on cycle -- it can cause problems by leaving the switch permanently on.
Fix it, or use at your own risk.

https://pypi.org/project/python-kasa/#description

[Edited on 26-7-2024 by semiconductive]

Good morning.
I'm not sure what to say ... the repeat experiment did NOT go as planned.
The Urea, unadulterated, did not melt under Kerosene.
The prills are still quite visible. Although the Urea does appear to have partially dissolved into the Kerosene; and re-crystallized after cooling.



I've ordered two more Owon BT41 meters, to double check the thermocouples.
It sort of looks like a design flaw in the meter; but since this is one of the few meters that actually works with Linux and windows, both, over bluetooth,
I'd really like to be able to use it with my Raspberry PI.

I also notice that it registers 0 degrees celsius about once in 100 readings; so the meter's sampling algorithm is flawed, too. ( Something I can write software to detect and reject, though. )

I've got a difference between the two meters that I have of about 6 degrees celsius. It's consistent. They're really inaccurate. I'm not sure what the temperature actually were, though, during the experiment. Argh!

I had an enclosed K type thermocouple in a steel sleeve for a different meter, and plugged it into the Owon only to get readings which were much worse.

I'm dipping one of the thermocouples into silicone, to see if I can at least make a temporary shield for it. But, this is not a really good idea, because silicone is chemically reactive.

I need to either figure out how to make a teflon sleeve, or get my inert gas setup finished so I can make a true glass sleeve. Stainless steel thermometers aren't really a good idea. Setbacks ...


[Edited on 26-7-2024 by semiconductive]

Note: heating rates are way off in heating script of last post as I changed s formula after comments, but I can't edit to fix the comments any more.
It's closer to power+=.0045 makes a 25% change in 10 minutes, not in an hour.
eg: .0009 is 5% change in 10 minutes.

You'll need to experiment to figure out the other rates.

I calibrated one Owon thermometer after dipping it in orange sensor-safe silicone from permatex.℠. It reads 5 degrees C high; but this is consistent at all temperatures so I wrote a program to adjust the temperature.

Reheating, again, using power+=.0009 # 5% in 10 minutes, 40% in an hour @ 19% start.

The thermocouple is at the front of the tube, where it should be coolest.
Most of the test tube de-fogged by 30 [°C]. But: The dust around the prills did not evaporate until about 108 [°C]. Picture is 110 [°C]. The prills are not melting.



130 [ °C ] @ 72.9% of 30W power.
131 [ °C ] @ 73.3% ...
132 [ °C ] @ 74.0% ...
133 [ °C ] @ 74.6% ...


134 [ °C ] @ 75.0% ...


135 [ °C ] @ 75.8% ...


137 [ °C ] @ 77.3% ...


139 [ °C ] @ 78.6% ...


141 [ °C ] @ %80.0 ...


143 [ °C ] @ %81.3 ...
144 [ °C ] @ %82.0 ...
145 [ °C ] @ %83.0 ...
146 [ °C ] @ %83.8 ...
147 [ °C ] @ %85.4 ...


148 [ °C ] @ %87.2 ...
149 [ °C ] @ %88.0 ...
150 [ °C ] @ %88.75 ...


I don't get it.

This is hot. Burns my finger, hot. Not even the tops of the urea prills have melted.
Does driving off water slowly before melting raise the melting point???

151 [ °C ] @ 89.5%
152 [ °C ] @ 90.5%
153 [ °C ] @ 91.3%
154 [ °C ] @ 92.2%
155 [ °C ] @ 93.2%

I give up (for today). Stumper.

The silicone part that was submerged in hot kerosene, softened. Silicone's weakness is gasoline; and Kerosene is not far from it. But, the silicone dip survived intact; it's good enough, I can use silicone for a while.

( AND YES, those really are different pictures taken a few minutes apart. )


[Edited on 27-7-2024 by semiconductive]

OK, I'm going to change one variable; and retry the melting experiment.
The variables I can think of are pre-drying by baking, kerosene as a reactant, urea brand (fertilizer) has contaminant, air may be necessary to melt urea normally; for it may interact either with oxygen or nitrogen during the melting process.
Did I overlook anything? ( suggestions welcome. )

I don't think kerosene ought to react with urea, but ...

Instead of kerosene, I'm going to try silicone oil, cosmetic grade.
I'll do the same thing, wire the soldering iron to the tube and make good contact; then bake about 1/2CC of urea at just above the boiling point of water for several hours to dry it; finally do a temperature sweep up to 155 [°C] and find the melting point of dry urea (if any.) I'll keep the picture count low, if I don't see a reaction and I'll just sumarize. If you want more pictures, just ask...

Setting the soldering iron to 50% power, I get a nice 105 [°C] temperature.
I don't see any steam or bubbling... but only a slow swirling of the silicone due to convection of heat. This suggests that the urea may be dry, already, in the bag. I am probably only driving off surface moisture, then, from the air.

This is how the drying begins:


I baked it dry for four hours.
Then, I did a temperature sweep all the way up to 100% power over 3 hours.
Urea did not melt under silicone, either!.



1CST silicone oil, on the data sheet says it boils at 151 C.
Picutre says 154C, with rolling boil and a significant amount (10 to 20%) of silicone refluxing. Little bubbles began happening around 143C.

So, my thermometer is giving reasonable values.

I don't see any discoloration, which I expect silicone would cause if it chemically reacted. I'm more inclined to believe urea might react with air when melting, or perhaps I have an impurity which strangely raises the melting point rather than lowers it.

Not sure what to try next.... but this is repeatable. So something unusual is definitely going on. I've ordered reagant grade urea prills, from Loud Wolf™ to check for contamination issues.

Edit: Note, allowed last experiment to reflux silicone at 154 [ °C ] for 8 hours. No melting or fusing was observed. After cooling, there was no residue on the walls of the test-tube. Urea either doesn't dissolve in silicone, or what little does dissolve is temperature independent. However, a slight yellow color showed up on the urea prills during cooling.



I'm not sure if it's just inconsistency in the lighting. The color went back to white, slowly, while at room temperature; 21 [ °C ].


[Edited on 30-7-2024 by semiconductive]

[Edited on 30-7-2024 by semiconductive]

I tried the experiment again. It's consistent.

I didn't misread a label and pull chemicals from the wrong bag. Whether the product is really urea ...

The bag is clearly labeled "the SEEP plant, UREA 46-0-0 granular fertilizer."
I have to wait until I get reagent grade before I can rule out the seller giving me fake product.

But, It did have a slight ammonia smell from it when melting with copper oxide.
The odor went away after getting above the boiling point of water for a while.
It's only pure urea prills under non-polar liquids that are refusing to melt.

Placing urea in a test tube, exposed to air, and heating tube with same soldering iron at 100% from the start; result:

Urea nearest the soldering iron began melting within seconds. Liquid wicks into remaining urea and causes fusion of the surfaces.
Mild ammonia smell is present.

But: I could only smell the urea and oils, no ammonia, when it was under kerosene or silicone oil for a while.



The thermal contact area is limited. I have to add more urea prills to get enough liquid. But it melts fine when there's enough prills to get the liquid level up to where the iron is. In fact, I need to turn down the heat; the liquid is bubbling and smells strongly of ammonia, now.






[Edited on 31-7-2024 by semiconductive]

Quote: Originally posted by bnull

Quote: But it melts fine when there's enough prills to get the liquid level up to where the iron is. Hold on. Is the soldering iron touching the tube at the same spot as that day?

Yes.

You can see it in the picture, eg: the seam line of the soldering iron's barrel.
The soldering iron is aimed downward at 75 to 80 degree angle behind the tube.
This causes the hot iron barrel to make solid contact with the test tube right where the glass changes from cylindrical to spherical.

The most intense heat always comes in at the point where the test tube changes shape.

There's generally about a 15 degree Celsius difference from the outside-back glass to the front of the tube. I am now putting the thermometer closer to the front, so it reads cold side. I have been trying to set my experiments up to be as repeatable as possible (for me, anyway.)

The power is computer controlled and changed very slowly over a period of hours.
For example, the last experiment is still running while I write this, many hours after my last post, and the power has only increased by about two watts.



The temperature of melted urea is about 130 degrees C. Which agrees fairly closely with the literature I found. Although it stays melted even at lower temperatures for a while. The urea also is slow bubbling ammonia gas out, at the point of contact with the soldering iron. There is liquid urea all the way to the bottom of the test tube, but then it goes solid as the front of the test tube (farthest away from the iron) is approached.

I'm not sure.
I suspect it would have to be more than just the capacities.

The long time periods of my experiments allow thermal equilibrium to be approached closely. You'll notice in the picture that the urea is melted everywhere except the front round bottom where the most surface area exists for the tube. There are no visible changes going on, as the heat input is very slowly changing.

The location of the heat input is why I always make sure to add enough chemical to go above the rounded bottom. Theres an artificially cool spot there.

But, I don't know:
Is urea much more heat conductive than silicone oil and kerosene, both?

For, the silicone oil was at 155 degrees in the coldest part of the tube. That's 25 degrees overheat in the coldest part of the tube.

In order for urea to not melt, at all, when in contact with 155 degree liquid at least at the top ( for hours ) ... the Urea would have to be conducting the heat away. I could insulate the bottom of the test tube, I suppose, to test for that. I've got fumed silica, pearlite, and plaster; so I could make an insulation cup. But it seems strange that none of the urea melted, not even the stuff on top.


Ohh! fully melted, finally!




[Edited on 1-8-2024 by semiconductive]
I need to improve the heating program a bit. I'd like to make detailed comments, but can't since there's more than one source of possible error. I told the program to maintain power at 62.25% permanently before heading to bed.

There was some urea condensing as solid, cotton candy like, on the test tube walls. I decided to place the melt under silicone to see if it stopped bubbling or triggered re-solidification; it didn't re-solidify for at least 4 hours; but Ammonia gas output reduced rapidly as temperature began to rise slowly even though power input was programmed not to change, anymore. Then some brown flecks formed in the urea. The temperature dropped (not sure when/why, need to modify program), and found tube as follows in the morning.



I've told the program to increase power slowly this morning, rather than maintain it; we'll see if the urea under silicone remelts or not.

[Edited on 1-8-2024 by semiconductive]

The answer is a definite "no.". It absorbed the silicone oil slowly, and then started blackening where the iron touched the tube. There was some kind of chemical reaction. ugh.

Maybe it's forming biuret, or triuret ; although the temperature is not high enough according to what I see online.

I'll try it again with a new program that compensates for line voltage changes, to guarantee my heat quantity is known and doesn't vary;

Right after I do a simple test dropping a urea prill into pre-heated silicone at 131 [ °C ]. I'm going to guess it melts, since it was exposed to air before hitting something hot; and hopefully I'll get to see see how long it can stay at 131 before turning brown or solid white.


[Edited on 1-8-2024 by semiconductive]

Ok, I've put two CC's (measured) of silicone oil in the test tube.
Nothing else is in there except the temperature probe.



I figured out about how much energy it took to get the oil to 131 [ °C ], 72% or 20.05 [ W ] of power last night.

I programmed the KASA plug to report the voltage and current being supplied to the soldering iron so I could graph it and calibrate my temperature set-point.

Note:
This is a graph demonstrating the scientific concept of systematic error; even though I've tried to compensate for voltage changes on the power line. Apparently, I did it a little wrong. I've got to fix this before I can get really reliable data for science experiments.



The spot where the temperature goes 'up' is in the hours between midnight and 8AM, local time. This is likely where all my neighbors have turned off their lights and gone to bed. But it might also be a time where the climate control in the lab (basement), changes a bit. I'll have to set up a second thermal monitoring ohm-meters in order to be able to compare what's inside the testube vs. the air around the test tube.

But, anyhow:
The temperature of the soldering iron bath rose 3 degrees during this dark of night, and I see increased power output to the soldering iron (on average) during this time.

So, power line regulation issues are a likely cause why the temperature rose on my last experiment, instead of being evidence for a chemical reaction.

I'll let this program run for another day or so, and get a long-term plot of the error and make sure it repeats as it hypothetically should, every night.

Then I can try to reprogram my heater code to *properly* compensate for line voltage variations.

Well, I'm winning. Finding a source of error is half the battle to getting rid of it.

There is definite climate control change on top of line voltage fluctuations. After several days experimenting, I've found that the the KASA smart plug's current sensing is too coarse (around 100mW steps) to be accurate. The voltage sensing is OK, though, and matches what my multimeter says my power line is doing.
I've written a power regulation script that is closed loop, based on the line voltage and a manually measured resistance. ( Will post a link, later. )

Regarding the Owon temperature probes:
There are definite temperature calibration and drift problems; at least 1 [°C] scale drift/error, and 6 degree offset error.

I'll be attempting to solve those, here:
Suggestions, welcome.

Owon BT41T+ multimeter calibration

Quote: Originally posted by bnull

]Let's see how Planck's function behaves under the assumptions

I'm back. I have some progress to report!

The Planck integral, even after a year, is not easily solved. Two professionals, paid, could not come up with a better answer than I had already. Very annoying.

I've found a way around the ceramic cup plugging issue. I can buy water filters with 1 micron pores made out of polypropylene. P.P. (USA plastic recyling #5) has chemical resistance to acetone and many acids. When I dissolve cellulose acetate into acetone to make a very weak solution, I can get it to diffuse into the pores of the filter and transform the filter into a very low cost semi-permeable membrane.

I've got the inkbird thermometer device measuring temperatures. I can now temperature regulate. Inkbird has very satisfactory temperature accuracy out of the box; but it's difficult to get to work reliably with my Raspberry PI.

https://www.sciencemadness.org/whisper/viewthread.php?tid=16...

I will be trying a battery eliminator to see if I can make my battery based equipment more accurate.

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

Now for a "warm up" test. Urea is more stable than thio-urea, so I'm going to start with the more likely to reduce thio-urea first. I want to see if I can get it to dissolve and form complexes in acetone.

I put thio-urea, boric acid, and lithium citrate into a test tube. I ground them together with a glass stir rod, and then added acetone at about triple the volume of the powder. I capped it with kerosene (as usual -- to keep out air, and slow acetone evaporation ), and put in two steel nails as electrodes. The boric acid is just to make the dissolution of iron faster, although it also tends to make a little bit of water (unfortunately).

I made the lithium citrate the day before in water using Li₂CO₃ + citric acid; and dried it a 101 [°C] for an hour. SO the only water in the test tube is from boric acid reacting with rust.

The acetone solution is not initially conductive. I added a tiny bit of oxalic acid. My theory here is that citrate has three protinated bonds and one of them will be weaker than the two oxalic acid protinated bonds, so that the oxalic acid will cause ionization of some of the citric acid's lithium.

The lithium is highly soluble, and I hope it will replace hydrogen over time and form double salts with iron and/or nickel complexes. Acetone is also able to complex iron and nickel, so there are multiple chances to form ions.

Initially the solution bubbles furiously at 7 [mA] of current with 1[mA] of DC bias. The idea here is to dissolve the electrodes in the free acid as quickly as possible without a lot of side reaction products. AC current reverses reactions multiple times per second...

Note:I've done this same experiment in the past using alcohol rather than thio-urea, and I can get iron and/or nickel to plate out a purplish-grey color depending on how much oxalic acid I add. (More tends to be better!) But, the alcohol (1,3 propanol) slowly turns black or glycerine tends to carmalize.

Initially, the plating is always black or mixed with organic crystalizing dusts; but over time the gas evolved at the electrodes reduces and the color of the plating at the electrodes lightens.

When I dip in a glass stir rod, and pull it out, there is a coating upon drying. This does not happen in the alcohol mixture, so I think it shows a small amount of thiourea did dissove into the acetone.

I've set the temperature of my experiment to 132 [°F], 55 [°C].
This puts it close to the boiling point of pure acetone.

With a small amount of oxalic acid (50mg or less), I get a pale yellow complex forming in the test tube. The lithium citrate powder also slowly turns bright yellow and swells in size. Swelling is indicative of solubility increasing, so I think I'm on the right track.

After two days, The bubbling reduces so much that I can double the current to [14 mA] A.C. to restore the initial rate. But, the color of the plating is still black oxide. ( or perhaps sulfide, if thiourea is decomposing. )




[Edited on 17-11-2025 by semiconductive]