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Registered Member #69
Joined: Thu Feb 09 2006, 07:42AM
Location:
Posts: 116
This really reveals how bad my chem knowledge is. I've got a solution of ferric sulfate (nice strong red color) and I drop the pH by adding sulfuric acid to it and it gets reduced to ferrous sulfate (pale green). How does this happen? Obviously the Fe+3 is being reduced to Fe+2 but how?
Registered Member #56
Joined: Thu Feb 09 2006, 05:02AM
Location: Southern Califorina, USA
Posts: 2445
I should know how to do this, but I forgot 98% of the redox stuff from chem...
In any case, isn't it just that at low pH (like 1-2) the fe2+ is favored, where as at higher pH's the 3+ form is preferred? So by adding the extra H2SO4 you lowered the pH enough that most of the fe3+ decided to be fe2+....
I suppose I could elaborate by saying something about the the equilibrium between fe2+ and fe3+ (the half reaction illustrated as fe2+ -> fe3+ e- for simplicity's sake) is shifted far to the right at high pH's (4-6) where at low pH's (1-2) it favors the left, but that would just be plain confusing
Registered Member #69
Joined: Thu Feb 09 2006, 07:42AM
Location:
Posts: 116
Well, you can oxidize the Fe+2 completely to Fe+3 in very acidic solutions by adding H2O2, it seems quite stable. But on addition of more H2SO4 it'll convert back. At more neutral pH, Fe+2 oxidizes much more easily to Fe+3, oxygen from the air will do it spontaneously, so it would seem that what you say is true about it 'preferring' neutral pH.
I'm just lost as to where the electron is coming from upon addition of the H2SO4. It's not from the H+ or the SO4-2...
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Redox Question
