I've been experimenting with electrolysis for rust removal, electroplating, de-plating, etc. Right now I have a rust removal setup with carbon steel electrodes and a solution of oxalic acid in distilled water as the electrolyte. I've never used oxalic acid with electrolysis before (I've never used any acidic electrolysis process for de-rusting before). It seems to be working well, and quickly.

The setup is:

  • Half a teaspoon (4-5g if I did the math right) of oxalic acid crystals dissolved in about 350 mL of distilled water. That's ~10% of the concentration recommended on the box. It seems pretty dilute compared to what I'm used to (for bleaching wood).
  • Carbon steel anode and cathode. Cathode initially had some rust (brown) on it.
  • Running at 6V (which yields ~300mA) to keep heat output low.
  • About 2cm distance between electrodes.

When I turned it on, a sufficient amount of current began to flow, which I was hoping for but wasn't sure (see below). After a while, the solution turned a sort of orange-brown color which I'm assuming is iron oxalate, or at least it looks like a similar color. When the process is turned off, it settles very quickly. The cathode is relatively clean compared to a sodium bicarbonate process.

But, I'm not really sure if I'm accomplishing anything beyond splitting up water and whatever happens when I soak rusty metal in oxalic acid without electrolysis.

So my question is: What is happening in the electrolytic process compared to soaking rusty iron in oxalic acid without electricity, and what's the mechanism behind it? Or am I just wasting electricity?

From here down isn't really important.

The path I took to get to this process (for context) was:

  1. I've only been using alkaline processes for rust removal, usually with sodium bicarbonate + water as the electrolyte.
  2. I came across a video of somebody instead using acetic acid + sodium chloride + water as the electrolyte. The main points in the video:
    • Described as using electrolysis to "boost the performance" of vinegar (which seems to be the case).
    • Table salt as an ion source (from my own research I gathered this is because acetic acid is a weak electrolyte).
    • Stated that advantage of acidic process is stuff tends to stay in solution or at least not stick to the electrodes as much (no idea why, but I'm assuming it's for the same reason that e.g. copper plating fails when the pH is too low).
  3. I decided to try it, but I was out of vinegar. I decided to use oxalic acid, based on:
    • I had a box of crystals on my shelf.
    • It didn't really introduce any new elements to complicate things further.
    • I looked it up and found it dissolves into oxalate ions, then I just sort of guessed that at some point "iron oxalate" would exist. I looked that up and patted myself on the back because it actually turned out to be a real chemical (lol); but more importantly, it wasn't going to instantly kill me.
    • The internet said it was a strong electrolyte so I thought maybe I wouldn't need the sodium chloride (ended up being correct).
    • If I did need to add salt, I didn't know what the reactions would be but I figured (uneducated guessing) it'd probably involve some combination of sodium oxalate, HCl, maybe ferric chloride (?), maybe some chlorine gas, and I have the appropriate setup to deal with (and dispose of) all of those safely.
    • If any zinc got in the mix (it happens) I figured it'd end up as zinc oxalate, zinc chloride (?), neither being a cause for concern. This was also an uneducated guess, though.
  4. I set everything up, turned it on, and to my pleasant surprise, there was indeed good current flow without having to add any sodium chloride (a relief since I didn't have to test my guesses). But I'm not really sure if I'm accomplishing anything, so I headed here.
  • 1
    $\begingroup$ Comments are not for extended discussion; this conversation has been moved to chat. (Feel free to carry on there.) $\endgroup$ Commented Nov 12, 2021 at 0:08