Chemical Forums
Chemistry Forums for Students => Analytical Chemistry Forum => Topic started by: Donaldson Tan on September 13, 2004, 02:25:04 AM
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Often in the lab, diluted sulphuric acid (H2SO4) is used for dissolving Iron(II) Sulphate (FeSO4.7H2O) cystals. Since this Iron(II) salt is water soluble, why isn't distilled water preferred?
(1) O2 + 4H+ + 2e <-> 2 H2O E=+1.23V
(2) O2 + 2H2O + 4e <-> 4OH- E=+0.40V
With reference to the equations above, an acidic medium would favour the forward reactions of both equilibria. Hence, Iron(II) should have a higher tendency to be oxidised to Iron(III) in an acidic medium. I would think that an acidic envionment will contribute to the relative instability of iron(ii) ions in solution. Hence, dilute sulphuric acid should not be employed.
Someone please enlighten. Thank you alot.
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Probably to prevent the formation of Fe(OH)2 which is insoluble.
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How does water provide sufficient hydroxide ions for the precipitation of iron(II) sulphate?
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The only thing I can think of is that there's the natural dissociation of water, and that the SO4-2 ion would be prone to taking an H+ from water and form the bisulfate ion. That's really the only explanation I can think of.
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My guess is that some, not all, precipitation of iron hydroxide might occur. Neutral water has OH conc 10^-7 M. Solubility product of FeOH2 is 8x10^-16. So there's enough hydroxide for some iron to precipitate out.
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If an acidic medium encourages the oxidation of iron(II), why should I even use it to dissolve iron(II) sulphate? I intend to maintain a solution of iron(II) for use..
2FeSO4 (s) + 2H2O (l) -> Fe(HSO4)2 (aq) + Fe(OH)2 (s)
Are u all suggesting that the above reaction occurs when Iron(II) Sulphate dissolve in water? And that an acidic medium will neutralise the Iron(II) Hydroxide formed to retain Fe2+(aq) in solution?
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I may have overstated the case. Yeah, that's the equation but only a tiny fraction of material will react.
The SO4 ion hydrolysis in water is very small and OH conc is very small. Perhaps an acidic medium is a cautionary measure against hydroxide build up over time.
The equations you have for acidic medium oxidation are not complete.
You ignored the other half reaction: Fe++ -> fe+++ + e
E = -.77V (unfavorable).
Also, there is a concentration dependence of emf according to the nernst equation. The emf potentials as given are for reactants at standard concentrations of 1 molar and 1 atm for gases. Dilute acid and oxygen would reduce that potential. The net result is probably a small enough potential as not to be worrisome reaction rate-wise. Besides that, the bottle is capped so how much dissolved oxygen can possibly exist ?
Solutions may degrade somewhat over a long period of time, however.
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You can ease your mind. I just read that iron II oxidizes much more quickly in neutral rather than acidic solution.