Why Does Adding Strong Acid Change the Color of Metal Ion Solutions?

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Discussion Overview

The discussion revolves around the phenomenon of color change in metal ion solutions when mixed with strong acids. Participants explore the underlying reasons for this color change, including the potential formation of complexes and the effects of dilution.

Discussion Character

  • Exploratory
  • Technical explanation
  • Debate/contested

Main Points Raised

  • One participant notes that the color of a solution containing Fe3+ changes when strong acids like HCl or H2SO4 are added, suggesting the formation of colorless complexes such as FeCl4-.
  • Another participant proposes that changes in the complexes present and the removal of OH- as a ligand due to lowered pH could shift the equilibrium and affect color.
  • A different viewpoint suggests that dilution from adding acid could be the primary reason for the fading color, referencing the Beer-Lambert law.
  • One participant questions whether the dilution effect is also influenced by the equilibrium of complexes, particularly in relation to the concentration of H+ and its effect on cation release.
  • Another participant asserts that simply lowering the concentration through dilution is sufficient to explain the fading color, dismissing the influence of complex equilibria in this context.

Areas of Agreement / Disagreement

Participants express differing views on the mechanisms behind the color change, with some emphasizing the role of dilution and others considering the effects of complex formation and equilibrium shifts. No consensus is reached regarding the primary cause of the observed phenomenon.

Contextual Notes

Participants reference the Beer-Lambert law and the role of ligands in complex formation, but the discussion remains unresolved regarding the specific contributions of these factors to the color change.

Eureka99
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Hi everybody!
Does anyone knows why when a cation is mixed with a strong acid the color of the solution (if it's colored) fades a little bit, depending on the concentration of the acid?
For example an "orangy" solution that contains Fe3+, If HCl 2 M is added the color turns yellow, and if H2SO4 6 M is added instead it turns transparent. Is it due to the formation of colorless complexes (like FeCl4-)? Is it even possible that a complex is formed adding only the acid, with an excess of the ligand (like in the first case Cl-)?
The same thing happens with, Fe 2+, Co2+, Cr 3+, Ni 2+ and Co2+ solutions ( I did it in the laboratory) , and I'm not able to explain it.
 
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Definitely there are changes in complexes present. OH- is a ligand, when you lower pH you remove it from the solution shifting the equilibrium. Cl- is a ligand, when you add it to the solution, you shift the equilibrium.

But perhaps the simplest explanation is that you are diluting the cation adding the acid?
 
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Ok. Even if it's due to the dilution, isn't the dilution also based on another equilibrium of a complex that is shifted? I know for example, that most of the time, increasing [H+] concentration, if the ligand comes from a weak acid, more cations are freed in the solution and colors can fade, as the concentration of the complex decreases. Can It also possibly be, something about that concept?
 
Eureka99 said:
isn't the dilution also based on another equilibrium of a complex

No. Do you know the Beer-Lambert law? Just lowering the concentration makes the color fade.

Neither of the acids you have listed is a weak one, so that doesn't matter here (but yes, it can be an important factor in other cases).
 
Borek said:
No. Do you know the Beer-Lambert law? Just lowering the concentration makes the color fade.

Neither of the acids you have listed is a weak one, so that doesn't matter here (but yes, it can be an important factor in other cases).

I don't know much about this law, but If I didn't get it wrong is due to optics reasons the fading color. Anyway now it's a bit clearer to me :smile:
Thank you for the help!
 

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