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Waffles SS
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I dont suggest to use iron + HNO3 method because deadly,unashamed,mother f***er gas (NO2) will produce
[Edited on 4-8-2011 by Waffles SS]
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blogfast25
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If small amounts (0.1 mol or something like that) of Fe(NO3)3 is what you want to make, then just make sure you do it outside or underhood: then
dissolving iron in nitric really is the easiest route.
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Waffles SS
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What the equilibrium of reducing Fe(NO3)2 to Fe(NO3)3 by air(Or dissolved oxygen in water)?
Fe(NO3)2 + O2 =Fe(NO3)3 +?
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blogfast25
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Quote: Originally posted by Waffles SS  | What the equilibrium of reducing Fe(NO3)2 to Fe(NO3)3 by air(Or dissolved oxygen in water)?
Fe(NO3)2 + O2 =Fe(NO3)3 +? |
You meant 'of oxidising', instead of 'reducing', right?
1) Calculate the cell potential E from the half reactions (here E > 0),
2) Calculate the Gibbs Free Energy change Delta G = - n F E for the reaction (here Delta G < 0)
3) Calculate equilibrium constant Delta G = - RT ln K (here K >> 1)
But fully oxidising Fe (II) with just air is a slow boat to China. A catalyst may speed it up. And you need to bubble air through the solution...
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Waffles SS
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Yes, sorry my mean was oxidation
But I want to say what will produce at other side of this reaction(oxidation):
Fe(NO3)2 + O2 =Fe(NO3)3 + ?
*I think that will be iron oxide
[Edited on 6-8-2011 by Waffles SS]
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blogfast25
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For redox reactions, always start from the half-reactions:
Oxidation: 4 x [Fe2+ === > Fe3+ + e-]
Reduction: O2 + 4 H+ + 4e - === > 2 H2O
Add up to balance electrons:
4 Fe2+ + O2 + 4 H+ === > 4 Fe3+ + 2 H2O
If you started from ferrous nitrate and added a non oxidising acid (e.g. sulphuric - for the 4 H+ !) you would end up with a mixture of ferric nitrate
(about 2 mol) and ferric sulphate (about 1 mol).
[Edited on 6-8-2011 by blogfast25]
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Waffles SS
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Thanks my dear friend @blogfast
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White Yeti
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I'm a beginner in chemistry, but aren't iron(III) ions a reddish brown colour when dissolved in water? If you filter the mixture, is the filtrate
brown?
Why don't you just take some iron nails and dip them in nitric acid? It's much simpler and you can use iron in excess so as not to waste any of the
acid.
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blogfast25
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Concentration and temperature mainly determine what colour an Fe (III) solution is (assuming the anion is colourless).
Fe3+ solvates to probably [Fe(H2O)6]3+ which tend to hydrolyse:
[Fe(H2O)6]3+ + n H2O === > [Fe(H2O)6-n(OH)n](3-n)+ + n H3O+ ( n = 1 or 2, depending on conditions). The hydrolysed [Fe(H2O)6-n(OH)n](3-n)+ is what
give these solutions their colour. Higher temperatures also push that equilibrium to the right: solutions visibly darken (a bit) on heating. Adding
acid to an Fe3+ solution pushes back hydrolysis.
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Whirl85
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I know this is an old thread but here goes nothing,I'm trying to recrystalize ferric nitrate and the typical methods (increaseing the acidity, cook
down) are not working.
Any advice?
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Poppy
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| Quote: Originally posted by Arthur Dent | I've been reading this thread with interest. This is a compound I'd like to synthetize. Following the formula from the first post, I would figure that
the addition of an excess of iron oxide and some vigorous mixing on the stirplate would eventually yield a clear mixture with a brown oxide deposit at
the bottom once the mix settles. I think the addition of more nitric would likely contaminate the nitrate and make the mixture acidic and hard to
crystallize.
But I imagine a fresh batch of precipitated iron hydroxide and oxide (still wet) would probably be even better to synthetize
Fe(NO<sub>3</sub> <sub>3</sub> ? What about pure iron? I
know that stainless isn't affected much by nitric acid, but would plain iron wool react? Maybe i'll try this this weekend. 
Robert
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Iron hydroxides won't work. At best you can react FeSO4 with Ca(NO3)2 to produce Fe2(NO3)3 and then react this with NaClO to produce a red solution o
nearly pure Fe3+ ions in solution. Adding a little more CaCO3 will push iron2+ out, as Fe3+ is acidic it will tend to deal with it, but iron2+ won't
(you can't wait eternally btw).
Those where the results from experimental data.
Now, you can try and isolate it: evaporation? Add Na2SO4 to this final product so you can evaporate then down to crystaline form. Once done, drain off
all the water by useing mild heat (~60°) then proceed to calcination, which will ultimately drive to red iron oxide.
You can also aerate the Fe2(NO3)3 + Na2SO4 solution while slowly reducing pH 5 to produce ironIII oxide directly, filter and dry at mild heat. Fe2O3
is a very strong bond and forms following the FeO(OH) strucuture hence the Fe(OH)3 method will fail.
Good luck.
The best sample of Fe2O3 I obteined so far is a yellow powder of anhydrous Fe2(SO4)3 >90%, just because the hydrogen peroxide method is inferior
comparing to the straightforward acion of NaClO.
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