So I got myself a large Erlenmeyer flask, some steel wool and the bleach from the bathroom, mixing them under safe conditions in small quantities to see the sparky reaction I'd heard about. Trouble is, I didn't know the bleach we had was non-chlorine!

EDIT: My silly mind assumed that no reaction meant I had non-chlorine bleach, but staring me in the face is hypoCHLORITE. My apologies, I'd not meant to mislead. The brand was Shaw's at the time, or "Signature Selections" which is their store brand. I got no sparks and no visible fumes nor immediate discoloration.

Disappointed, I left the mixture alone and uhm.. forgot about it.

Going back to check on it, I found the steel wool dissolved into a powder at the bottom, and the rest of the fluid was purple.

Curious, I used a giant coffee filter and strained the rusty cake-mix residue out of the rest of it, which was now a lighter purple.

I didn't know what I'd made and thought it best to simply leave it be for a while again. I kept it in the basement, visible when the door was open and kept an eye on it. Slowly the purple went away and the fluid decreased leaving a white powder and bits of clear-quartz like appearing crystals. After several months the crystals were larger, and the white powder has climbed the entire container and spread down the sides. It's now been about 18 months since the initial experiment, and the white powder (which is similar to efflorescence in appearance) has climbed about halfway down the outside of the container, and the inside crystals have grown slightly.

What in the world have I made?

  • $\begingroup$ Well, purple was most probably ferrate, but you got colorless stuff later? That's weird... $\endgroup$
    – Mithoron
    Oct 25 '19 at 20:31
  • $\begingroup$ Interesting observation and reminds me of my own childhood experiments. I think that steel wool has a small amount of other transition metals which are causing this purplish color. One could have formed ferrate or even permanganate. This can happen with chlorine bleaches only. On the other hand, ferrates cannot form with oxygen bleaches. Ferrates can form when you react iron salts with hypochlorites (while gently heating). I rather suspect chromium alum type compounds (purple) or perhaps even a trace of vanadium. It is certainly not iron. What is the brand of your oxygen bleach? $\endgroup$
    – M. Farooq
    Oct 25 '19 at 22:56
  • $\begingroup$ [It changed my number for some reason when I created the account/this post, but I am the OP) Shaw's store brand (Signature Selection) Also I should clarify I never got sparks. I checked the container today, the crystals are still there and the white powder is still up the side, though it has peeled away a bit (about a half inch) and is cracked down to the bottom. There is still clear fluid under the crystal at the bottom, and the fluid comes through when tilted, so it is not solid. The crystal looks somewhat like clear ice, with bits of snow on it. $\endgroup$ Oct 28 '19 at 3:06
  • $\begingroup$ I sent a request to merge your accounts. In the meantime, this looks to me more like an edit/update to the question, not an answer. $\endgroup$
    – andselisk
    Oct 28 '19 at 3:51
  • $\begingroup$ It might help to tell us what the bleach actually was (read the ingredients on the label). $\endgroup$
    – matt_black
    Oct 28 '19 at 11:49

After the edits by the original post, the OP clarified that they used a hypochlorite (not an oxygen bleach as stated earlier) and left a piece of steel wool in it. Due to corrosive nature of chloride/hypochlorite, the steel wool crumbled to iron (III) oxides. When you gently heat rust in highly alkaline medium in the presence of hypochlorite ion, it forms a purplish sodium ferrate. This has iron in the (+6) oxidation state. Over time the purple color disappeared because ferrate slowly decomposes back to lower iron oxidation states in water solution, releasing its excess oxygen. Ferrates are more stable in alkaline solution than in neutral or acidic ones, so absorption of carbon dioxide over time, making the solution less alkaline, could have promoted this decomposition.

This decomposition of ferrates by acids, carbonic acids, and acidic salts was known as early as 1890s as noted in a review by Moser in Journal für Praktische Chemie 1897, 56, 425 (see relevant comments). Ferrates were known in 1700s.

Also note that hydrogen peroxide and ferrate decompose each other immediately, just like permanganate ion in aqueous solution.

The quartz like crystals are nothing but sodium chloride with perhaps some sodium carbonate.

Image from: https://www.youtube.com/watch?v=5u8LOfTrO6w


  • 1
    $\begingroup$ Added a brief discussion of tge ferrate disappearing. Please attribute the photo if not your own. $\endgroup$ Oct 31 '19 at 0:06
  • 1
    $\begingroup$ My mistake, I added the link in image description which was not visible. The Youtube video is in German & English from where the image was taken. $\endgroup$
    – M. Farooq
    Oct 31 '19 at 1:17
  • $\begingroup$ What is it about the alkaline conditions that stabilises ferrate? $\endgroup$
    – Gimelist
    Oct 31 '19 at 1:35
  • $\begingroup$ I think there is no simple answer as to why alkalis stabilize it. The half life of ferrate is less than hour in neutral pH. It can oxidize water. The why part of science is often the most difficult part of a phenomenon. $\endgroup$
    – M. Farooq
    Oct 31 '19 at 1:51
  • $\begingroup$ @M.Farooq that’s interesting. I’m at a high pressure and temperature lab, and we noticed that alkalis and primarily sodium tend to stabilise higher oxidation states of transition metals in silicate melts and glasses. We have some explanation based on the structure of the silicate framework, and it’s really interesting that these things also happen at STP in aqueous chemistry. $\endgroup$
    – Gimelist
    Nov 1 '19 at 22:07

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.