>Green potassium ferrioxalate solution exposed to light will turn yellow, and then brown, releasing carbon dioxide in the process. After keeping them in the dark for some time, they become green again. Adding small amounts of oxalic acid greatly helps to speed up the process.
I understand why they would turn yellow, but how can they turn back to green when keeping them in the dark? They absorb carbon dioxide passively somehow?
I believe that the reason it turns green again in the dark is because the oxygen of the air oxidizes the iron back to the iron (III) form. This proceeds easily if there is excess oxalic acid. In the absence of excess oxalic acid, this could still happen via a disproportionation sort of reaction that turns some of the iron (II) oxalate into insoluble iron crud (mixed oxides/hydroxides) and frees its complexed oxalate to reform the iron (III) complex.
You could test to see if it's oxygen from the air by exposing some of the solution to sunlight until it turns yellow, then storing half in a capped container with no head space for air, and the other half in an open topped container. Does the solution in the open topped container turn green faster? Does the solution in the closed container turn green again at all?
I'm pretty sure it's oxygen. The same thing happens with solarized (sun-bleached) cyanotypes, according to the Cyanomicon, except that there the colors are white and blue rather than yellow and green. I think Chase is right that with oxalate this requires excess oxalic acid, so you can't repeat it indefinitely, while the ferricyanide/ferrocyanide redox cycle is reversible.
Light reduces the iron center of the ferrioxalate ion, and oxidizes the oxalate ion into carbon dioxide. Since the green color of the solution is due to ferrioxalate ions, a loss of them will make the solution browner, which is the color of iron.
If there is excess oxalic acid present, then the extra oxalate ions can form ferrioxalate ions in the absence of light, causing the solution to revert to green.
Probably different chemistry but made me think that you could also pin an insect, leave it for a hundred-plus years, and get some [slightly smaller crystals](https://www.researchgate.net/figure/A-carpenter-moth-Cossida...), a well known curatorial issue in entomology collections.
Many of them suffer from desiccation and I wonder if they can be "sealed" in some manner. Or if the UV-sensitive crystals can be protected by some kind of coating.
It would also be interesting to see what would happen using the crystals grown in one solution as seeds in a chemically different solution. The right sequence might not be reactive, but you could layer them ...
They are hardy enough to be handled, but will chip if scratched, and break if dropped.
I store my crystals in a sealed container, and they are fine after a month. People who have left them outside do not report dehydration, but the surface still turns white due to exposure to light.
Some crystals can be successfully sealed in resin to prevent dehydration. Preventing light from photodecomposing potassium ferrioxalate crystals would be very tricky. I do not know whether it is possible.
My copper sulfate crystal, about the size of a half a banana and grown from information on the article’s site, is doing well submerged in mineral oil in a sealed glass jar.
If you do this, get it fairly dry first because surface moisture will form a thin layer between the oil and the crystal making a shiny surface from different refractive indices.
It would also be interesting to see what would happen using the crystals grown in one solution as seeds in a chemically different solution. The right sequence might not be reactive, but you could layer them ...
Ooh, epitaxial growth! Could try that with a series of alums - Al(iii), Cr(iii) and Fe(iii) is all readily available.
Sorry about that. I tried to keep it short, as the guide is long enough already. Since some of you have commented about that, I'll make some edits to it. Thanks for the feedback.
I heated it at 200C for 1 hour on a hot plate. You might need to heat a bigger batch for longer. Generally, I do double the time for it to get completely dry, to make sure all the iron hydroxide is converted to iron oxide.
The hardware stores here sell potassium bioxalate (salt of lemon) as a rust remover, which I suspect produces precisely this material. (Someone here on HN helpfully corrected me last year when I expressed skepticism that oxalic acid would remove rust. I was wrong!)
Sometimes I wonder that too. Have you reached the point where sometimes you imagine people responding to your invisible posts? Regardless, Happy New Year!
>To carry out electrolysis, I first prepared salt solution as the electrolyte. Then, I used a piece of iron as the positive electrode and a steel wire as the negative electrode.
Wouldn't you risk producing chlorine gas due to the salt in the solution?
"Some of you might have concerns that chlorine gas will develop, but because chlorine is so reactive, and because the salt solution is dilute, it’s safe for this setup."
Yes, as the text in the part I quoted says "dilute" now as well. Which is good as there are many "bucket chemists" out there, people who think, why add a pinch as more will be better and quicker mindsets kick in.
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I understand why they would turn yellow, but how can they turn back to green when keeping them in the dark? They absorb carbon dioxide passively somehow?