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Can Fe2+ ions exist in an acidic solution or will they always pick up available protons and become Fe3+? My hunch is they can't exist but I'm not really sure.

If the latter, does that mean that dissolving iron in an acid will raise the pH since it ties up some of the H+? Or will the additional protons still be just as free to float around and interact with other stuff as they are before iron comes to the party?

If they can exist, however, is there some kind of ferric:ferrous equilibrium based on... something?

I ask because I have a solution of oxalic acid and water, with iron dissolved in it, and it is bright yellow-green (like a green or yellow highlighter), and the internet says Fe3+ ions are yellow-green. Also I'm pretty sure there can't be actual iron oxalate particles (which are also green) suspended in the solution (conclusion from this lesson). I'm trying to work out the actual contents of the solution on my own.

So I'm pretty sure I'm looking at ferrIC ions -- and it doesn't seem to make sense that ferrous ions could exist here -- but I just want a sanity check on that guess, since I'm still trying to wrap my head around this.


As an aside: Now after a little more reading, I'm wondering if maybe my oxalic acid solution contains ferrioxalate ions too (or instead), although that's specific to my oxalic acid experiments rather than acids in general. They're also bright green, it seems. Apparently those are light sensitive though, so maybe my UV laser will do something if the wavelength (405nm) is right and so I can test for them? That's an experiment for tomorrow.

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    $\begingroup$ Ferrous ions are easily oxidized by oxygen in the air: chemistry.stackexchange.com/q/31243/79678. Other questions and answers here may be helpful. $\endgroup$
    – Ed V
    Nov 8, 2021 at 3:07
  • $\begingroup$ Thanks for digging that up! Sort of OT but does that mean iron ions will readily precipitate as iron oxides in an alkaline solution? Cause that is related to a question I was actually going to post later, to confirm my suspicions about why the acidic solution turns reddish brown as it presumably becomes saturated with iron (which is incidentally the real reason I asked about iron raising the pH in this question), and why adding more acid turns it bright green again. Also why it turns reddish brown when I mix it with sodium bicarbonate. (No spoilers please lol) $\endgroup$
    – Jason C
    Nov 8, 2021 at 3:16
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    $\begingroup$ I think the answer to the first of your two questions is yes, but I am not an expert on iron chemistry and it is a bit complicated anyway. As for adding baking soda, someone else may have an answer or hunch. $\endgroup$
    – Ed V
    Nov 8, 2021 at 3:20
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    $\begingroup$ @JasonC - Over several pH ranges you will get iron(III) precipitates. As a geochemist, I am most familiar with naturally-occurring goethite and limonite in areas I have studied. Have a look at the iron(III) oxide-hydroxides for more info, including precipitation of iron(III) oxyhydroxide from solution at pH ranging from 6.5 to 8 by addition of NaOH to a solution of ferric chloride/nitrate. $\endgroup$ Nov 8, 2021 at 3:55
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    $\begingroup$ Consider bidental complexes of Fe(II/III) and oxalate have probably different colours than respective ion aquacomplexes. $\endgroup$
    – Poutnik
    Nov 8, 2021 at 10:47

1 Answer 1

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Can Fe2+ ions exist in an acidic solution or will they always pick up available protons and become Fe3+? My hunch is they can't exist but I'm not really sure.

I was intrigued by your query because there is a hidden misconception. There are no available free floating protons in solution. If you want $\ce{Fe^2+}$ accept a proton, you are asking to cause a nuclear reaction. Of course this is far from reality. Notationally, even if you wanted $\ce{Fe^2+}$ to accept a proton (=from an acid), like the way organic chemists talk about it, it would be written as $\ce{[HFe]^3+}$. You don't want that either, because such a species does not exist.

In short, if you have an acidic solution, free from dissolved oxygen and kept in an inert atmosphere, $\ce{Fe^2+}$, will sit swim happily for a long time.

When we talk about the color of the transition metal ions, usually we are talking about the free ion surrounded by water, for example, $\ce{Fe^2+}$, which is a light green solution, chemists are then actually talking about $\ce{[Fe(H2O)_6]^2+}$.

Iron(II) oxalate, as you can see from Wikipedia is yellow... just like the dry yellow highlighter ink. When the metal ions are complexed (=surrounded) by something which can hold the metal ion like a crab's claw (chelate), you cannot predict the color beforehand.

Another example is that iron(III) thiocyanate is blood red. This has nothing to do with the yellow color of Fe(III).

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    $\begingroup$ Perhaps more importantly, even in an acidic solution with available protons (though they are not really free), $\ce{Fe^2+}$ is not going to become $\ce{Fe^3+}$, because, er, well, like you said, this is simply not a thing that can happen. $\endgroup$ Nov 8, 2021 at 7:16
  • $\begingroup$ Now I have to do a lot more thinking to figure out the difference between the bright highlighter greens vs deeper yellow solutions I see. I ordered an electronic pH tester, hopefully that'll help me eliminate some guesses. $\endgroup$
    – Jason C
    Nov 8, 2021 at 15:08
  • $\begingroup$ @JasonC, What are you trying to do with highlighter inks? $\endgroup$
    – AChem
    Nov 8, 2021 at 18:32
  • $\begingroup$ Nothing... it's just the words I used to describe the colors I've observed. :) $\endgroup$
    – Jason C
    Nov 8, 2021 at 19:22
  • $\begingroup$ Check it: semanticscholar.org/paper/… $\endgroup$
    – Jason C
    Nov 11, 2021 at 5:14

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