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Topic: Lab Procedure for dissolving Iron(II) Sulphate crystals  (Read 21229 times)

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Offline Donaldson Tan

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Lab Procedure for dissolving Iron(II) Sulphate crystals
« on: September 13, 2004, 02:25:04 AM »
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.
« Last Edit: September 13, 2004, 02:57:32 AM by geodome »
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Re:Lab Procedure for dissolving Iron(II) Sulphate crystals
« Reply #1 on: September 13, 2004, 09:24:55 AM »
Probably to prevent the formation of Fe(OH)2 which is insoluble.
« Last Edit: September 13, 2004, 09:35:48 AM by Demotivator »

Offline Donaldson Tan

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Re:Lab Procedure for dissolving Iron(II) Sulphate crystals
« Reply #2 on: September 13, 2004, 03:08:02 PM »
How does water provide sufficient hydroxide ions for the precipitation of iron(II) sulphate?
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Offline jdurg

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Re:Lab Procedure for dissolving Iron(II) Sulphate crystals
« Reply #3 on: September 13, 2004, 03:53:54 PM »
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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Re:Lab Procedure for dissolving Iron(II) Sulphate crystals
« Reply #4 on: September 13, 2004, 03:59:56 PM »
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.

Offline Donaldson Tan

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Re:Lab Procedure for dissolving Iron(II) Sulphate crystals
« Reply #5 on: September 14, 2004, 02:18:56 AM »
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?
« Last Edit: September 14, 2004, 02:21:20 AM by geodome »
"Say you're in a [chemical] plant and there's a snake on the floor. What are you going to do? Call a consultant? Get a meeting together to talk about which color is the snake? Employees should do one thing: walk over there and you step on the friggin� snake." - Jean-Pierre Garnier, CEO of Glaxosmithkline, June 2006

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Re:Lab Procedure for dissolving Iron(II) Sulphate crystals
« Reply #6 on: September 14, 2004, 10:47:15 AM »
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.
« Last Edit: September 14, 2004, 11:40:36 AM by Demotivator »

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Re:Lab Procedure for dissolving Iron(II) Sulphate crystals
« Reply #7 on: September 15, 2004, 10:50:59 AM »
You can ease your mind. I just read that iron II oxidizes much more quickly in neutral rather than acidic solution.

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