Sciencemadness Discussion Board - Iron(III) Nitrate

Iron(III) Nitrate

Waffles SS - 31-7-2011 at 05:56

Today i tried to make Iron(III) Nitrate :
When I add 160gram of Fe2O3(brown powder) with 756gram of 50% nitric acid the color of reaction didnt changed(that is still brown,iron nitrate is white powder and soluble in water)for ensure i added more nitric acid but color didnt changed.I am really confused
Does this reaction need heat?Or i didnt maistake?
Fe2O3 + 6 HNO3 → 2 Fe(NO3)3 + 3 H2O

[Edited on 31-7-2011 by Waffles SS]

bbartlog - 31-7-2011 at 07:07

So what started brown, and what stayed brown? Did the Fe2O3 remain unchanged? Did the nitric acid start red/brown and stay that way? Depending on the source of your Fe2O3, I would guess that it might have been contaminated with lower oxides of iron, so that as a side reaction you could get some NO or NO2 gas (which, dissolved, would color your solution). Alternatively, it may be that your approach is simply unworkable and will only yield basic ferric nitrate, Fe(NO3)2OH, for whatever reason.

Waffles SS - 31-7-2011 at 07:29

Thanks my dear friend @bbartlog,
When i added iron oxide(Fe2O3) to 55% nitric acid solution color changed to brown(i think iron oxide cause color change).when i stir mixture white fumes start to come( i think that is NO) I didnt see NO2
Now in bottom of my beaker i have unreacted iron oxide.I really confused why iron oxide dont react with nitric acid?
I want to make Iron(III) nitrate for fertilizer
I think reaction of Calcium nitrate with Ferrous sulfate is easier!

Ca(NO)3 + FeSO4 = CaSO4 + Fe(NO3)2
Fe(NO3)2 -> Fe(NO3)3

blogfast25 - 31-7-2011 at 12:31

Any solution of Fe3+, no matter what anion it is matched with, is yellow to dark coloured depending on concentration and temperature. Solvated Fe3+ forms Fe(H2O)n (3+) ions which tend to hydrolyse:

[Fe(H2O)n]3+ + m H2O < === > [Fe(H2O)n-m](3-m)+ + m H3O+ with m 1 or maybe even 2.

The [Fe(H2O)n-m](3-m)+ species give iron (III) compunds their characteristic colour. In strongly acidic and dilute solution (the equilibrium is pushed back to the left), iron (III) ions are almost colourless. Temperature tends to push the equilibrium to the right: hence such solutions are also slightly thernochromic.

Your observation is correct and entirely normal, waffles.

CaSO4 is a bit soluble in water.

[Edited on 31-7-2011 by blogfast25]

hkparker - 31-7-2011 at 12:53

I have tried this before as well. At first no reaction is observed, but eventually I was left with a thick yellow/brown solution.

Waffles SS - 31-7-2011 at 19:55

Thanks dear @blogfast,
What method you suggest for prepare Iron(III)Nitrate?
I thought this is simple to prepare Iron (III) Nitrate from Nitric acid and Iron (III) Oxide
Why when i stirr mixture the white fume appear?
(For solve CaSO4 solubility i can replace Ca(NO3)2 by Ba(NO3)2)

AndersHoveland - 31-7-2011 at 21:13

Magnetite (Fe3O4) forms an extremely stable crystal structure, which is surprisingly chemically resistant.
While hydrochloric acid can dissolve it, nitric acid barely seems to attack it. The reaction rate is greatly increased at 35degC, but even at this temperature, it still takes a long time for the nitric acid to dissolve it.

Perhaps the HCl forms a more soluble FeCl4(-) complex, which could explain the significant difference in reaction rates. It should also be remembered that the Fe+3 ion is somewhat acidic, it may be that the Fe+3 ions within the crystal structure are unable to be fully hydrated until they escape, thus surface ions may likely have difficulty dissolving, as the unhydrated Fe+3 will prefer to form insoluble hydroxide, even in the presence of acid.

for more reading:
http://ccm.geoscienceworld.org/cgi/content/abstract/59/2/136

AJKOER - 31-7-2011 at 22:39

If interested, consider reacting Fe and HClO, a vigorous reaction even for dilute HClO (I will discuss methods for making HClO later). I suspect that this initially forms the unstable Ferric Hypochlorite, Fe(ClO)3 which rapidly decomposes into FeCl3 (the literature is a bit incomplete on this, but I believe I am correct). Now, react with HNO3 (be careful, I would suspect a vigorous to violent reaction so start with dilute solutions). One could also use AgNO3 in place of HNO3.

I would give you two ways of obtaining a working solution of HClO. First, mix HCl and H2O2 (yes, this commonly used etching solution behaves like HClO, and in fact is an old recipe per Watt's Dictionary of Chemistry for making HClO). Another way, especially for the home chemists, is to react vinegar and bleach (NaClO). This forms HClO and Sodium acetate. One can either distill (almost all the HClO and Cl2O is obtained when half of the solution is distilled) or use as is on iron (but note that you may now be working with Ferric Acetate instead of Ferric Chloride).

This experiment should provide a good comparison to other methods.

Waffles SS - 1-8-2011 at 03:11

Thanks @AndersHoveland and @AJOKER,
What about black iron oxide?(FeO)?Does it is stable like Iron(III)Oxide?

AJKOER - 1-8-2011 at 05:24

Per one of my references, Iron black or Iron(II,III) oxide (formula: Fe(II)Fe(III)2O4 usually written as Fe3O4), is made by the oxidation of Iron with steam below 566 C (over 566 C one gets Iron(II) oxide also written as FeO although the actual ratio of Iron to Oxygen is apparently between 0.90 to 0.95 ).

Iron black is considered to be the most stable Iron oxide. Note, Iron(III) oxide (Fe2O3) is converted into Iron black when heated in air to over 1200 C.

As to FeO, it is stable over 566 C and is considered metastable at room temperature. Source: "Concise Encyclopedia Chemistry" by deGruyter.

[Edited on 1-8-2011 by AJKOER]

Waffles SS - 1-8-2011 at 06:52

Today i Tried another route for Iron(III)Nitrate
When I added 236 gram of Calcium Nitrate tetrahydrate solution in 1lit water((solution was colorless) with 278 gram Iron(II)Sulfate heptahydrate solution in 1 lit of water (solution color was green)lot of calcium sulfate precipitated(milky color) and the color of solution turn to light green.Does the color of solution is usual?
I think solution of Iron(III)Nitrate should be colorless
Ca(NO)3 + FeSO4 = CaSO4 + Fe(NO3)2
I thought that dissolved oxygen in water cause Iron(II)Nitrate reduced to Iron(III)Nitrate.is it true?what is the color of Iron(II)Nitrate ?

[Edited on 1-8-2011 by Waffles SS]

blogfast25 - 1-8-2011 at 07:05

Quote: Originally posted by Waffles SS

What method you suggest for prepare Iron(III)Nitrate?

Precipitate Fe(OH)2 from FeSO4(aq) (or any other ferrous salt) and NaOH(aq) or ammonia solution. Wash precipitate with copious amounts of water to get rid of sulphates. The Fe(OH)2(s) will slowly oxidise to Fe(III) in air. Accelerate this by dissolving it into hot HNO3:

3 Fe2+ + HNO3 + 3 H+ === > 3 Fe3+ + NO + 2 H2O (Caution! NO will oxidise in air to toxic NO2: ventilate well or do outside!)

This is the 'standard method' (there are others, of course) of oxidising ferrous compounds.

Waffles SS - 1-8-2011 at 07:22

Thanks my dear friend @blogfast,
good method my friend but Iron(II)hydroxide dont decompose to iron oxides?
What about iron (II)carbonate?
Which one is more stable?

[Edited on 1-8-2011 by Waffles SS]

blogfast25 - 1-8-2011 at 11:44

All water insoluble ferrous compounds, including the carbonate (Siderite as mineral), eventually oxidise to 'Fe(OH)3.nH2O' in air. I write it like that because the composition of 'hydrated ferrric oxide' or 'ferric hydroxide hydrate' isn't really well known or understood.

The 'wet' forms of insoluble ferric compounds always somewhat retain their solubility in (even weaker) acids: it is heat (calcination) and driving off the water that greatly reduces reactivity and solubility in acid.

Fe2O3 (calcined, like what you're using) will dissolve in conc. H2SO4 or can be fused with bisulphate. It then enters solution as Fe2(SO4)3, after dilution with water.

not_important - 1-8-2011 at 18:05

Note that CaSO4 is slightly soluble, so there may be some calcium contamination of the ferric nitrate.

Every prep I've seen for this compound, as well as ferrous nitrate, states that there should be excess HNO3 so that the solution is distinctly acid.

Waffles SS - 1-8-2011 at 21:14

Thanks dear @blogfast and my friend @Not_important(i am very happy to see your post),
I find out the solution of Iron(II)Nitrate is light green and solution of Iron(III)Nitrate is yellow
After adding calcium nitrate solution to Iron(II)sulfate i got light green solution and still that is green .I think the dissolved oxygen in solution should reduce Fe(NO3)2 to Fe(NO3)3 but it seems after 2 day my solution color didnt change.

MrHomeScientist - 2-8-2011 at 05:28

Quote: Originally posted by Waffles SS

I think the dissolved oxygen in solution should reduce Fe(NO3)2 to Fe(NO3)3 but it seems after 2 day my solution color didnt change.

You could test that by taking a small amount of your solution and adding some hydrogen peroxide to it. It should change color as Fe(II) is oxidized to Fe(III). I've done this before with iron(II) chloride (green) to make iron(III) chloride (brown).

Waffles SS - 2-8-2011 at 05:54

Thanks @MrHomeScientist ,
I wonder for my today experiment
I added 20gram Iron(III)Oxide to 200ml of 55%sulfuric acid but it seems it didnt react
I can see small bubble release from bottom of my beaker(from iron oxide layer) very slowly(~2-3 bubble every 2 second)when i heat solution to 90c
My solution color is brown and there is unreacted iron (III)oxide in bottom of beaker

[Edited on 2-8-2011 by Waffles SS]

blogfast25 - 2-8-2011 at 06:18

Depending on degree of calcination, 55 % is too weak. You need to use an excess of 95 % (at least) to dissolve most commercial grades of ferric oxide. 55 % will dissolve 'a bit' but far from all. Nota also that some Fe2O3 contain and acid insoluble portion: usually silicates (see e.g. so called 'pottery grades')

Treatment with conc. H2SO4 will not give you an actual solution but more like solid Fe2(SO4)3. Allow to cool down, then leach with 20 % (or so) H2SO4 to get your ferric sulphate back in solution.

[Edited on 2-8-2011 by blogfast25]

Waffles SS - 3-8-2011 at 03:02

Does adding more calcium nitrate to iron(II)sulfate solution cause another iron compund precipitate?
When i add more calcium nitrate to solution the brown compound precipitate

blogfast25 - 3-8-2011 at 04:31

Can you test whether the solution still contains significant amounts of Fe3+ after that?

Waffles SS - 3-8-2011 at 06:18

How can i test it?
I just can say that the color of my solution after adding calcium nitrate turn brighter now that is very very bright green/yellow
of course i should say my iron (II)sulfate consist heptahydrate and monohydrate crystal also i decide to add 1ml hydrogen peroxide 35% in my solution tomorow and i dont know what will happen(maybe it change solution color to colorless

)

[Edited on 3-8-2011 by Waffles SS]

blogfast25 - 3-8-2011 at 07:45

Add some NaOH solution or ammonia solution to the liquid. If plenty reddish/brown precipitate forms, there is still Fe3+ in solution. A black precipitate would point to Fe3O4 and Fe (II) still present too.

Or use KSCN solution: it forms red FeSCN2+ with Fe3+...

Waffles SS - 4-8-2011 at 02:32

Thanks dear blogfast,
I tried to evaporate water and heat solution(~50c) for 3 hour and now i got lot of brown precipitate( that is like ferric oxide)
I find out that iron nitrate solution decompose on contact of air(decomposition increase at higher temp)

[Edited on 4-8-2011 by Waffles SS]

Arthur Dent - 4-8-2011 at 04:01

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(NO3

3 ? 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

Waffles SS - 4-8-2011 at 05:53

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]

blogfast25 - 4-8-2011 at 07:09

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.

Waffles SS - 6-8-2011 at 02:09

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 +?

blogfast25 - 6-8-2011 at 04:33

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...

Waffles SS - 6-8-2011 at 04:56

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]

blogfast25 - 6-8-2011 at 06:32

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]

Waffles SS - 6-8-2011 at 11:16

Thanks my dear friend @blogfast

White Yeti - 17-8-2011 at 11:53

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.

blogfast25 - 17-8-2011 at 12:04

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.

Whirl85 - 26-9-2013 at 04:56

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?

Poppy - 30-9-2013 at 16:34

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(NO3

3 ? 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

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.