Sciencemadness Discussion Board - Ferricyanide test for ferrous ions

Ferricyanide test for ferrous ions

CHRIS25 - 21-6-2014 at 01:12

Just a little confirmation here if you will.....I titrated ferric chloride this morning after adding a stoichiometric amount of peroxide to the ferrous/ferric solution.
(Moles Fe added at the beginning just for information = 0.569)

Titration results after 2 runs averaged at 0.5674 moles

Good maths I thought.

However, I then added about 50 grains of potassium ferricyanide (in a few mLs of water) to the rest of the diluted analyte solution just for confirmation that no more ferrous existed. In amongst the blue It precipitated the solid, confirming that ferrous still existed.

Taking three test tubes; one with a dissolved solution of ferricyanide, and two test tubes - to these were added in 15 mLs water and 2 drops of chloride taken from the source solution of ferric chloride. I then Added 15 mLs Peroxide to one test tube of ferric chloride only. To this was added 10 mLs of the ferricyanide solution and 10 mLs was also added to the other test tube that had no peroxide. The results are confusing. Neither test tube precipitated any blue solid matter, just the expected blue colour change. bearing in mind that the analyte solution (that was used for the titration) precipitated blue solid I do not know how to interpret things?

[Edited on 21-6-2014 by CHRIS25]

[Edited on 21-6-2014 by CHRIS25]

aga - 21-6-2014 at 02:08

Dunno if this helps :-

Just tested my Ferrous and Ferric solutions using Postassium Hexacyanoferrate(III) K₃Fe(CN)₆.

The Ferrous gives the dark blue precipitate.
The Ferric changes colour with a sight blue tint, but no blue precipitate.

Adding H₂O₂ to a sample of Ferrous, then immeditely adding the cyano gives a test tube of opaque blue liquid.

The Ferric was produced the same way, the difference being that it was boiled down after the peroxide was added.

The cyano solution was about 1 quarter of a spatula in 20mls water.

unionised - 21-6-2014 at 03:26

I think that hydrogen peroxide will reduce ferricyanide to ferrocyanide in some circumstances.
(It's odd to think of a peroxide as a reducing agent, but it happens)
http://chemistry.proteincrystallography.org/article148.html

[Edited on 21-6-14 by unionised]

blogfast25 - 21-6-2014 at 04:39

Chris and aga:

Berlin Blue is insoluble (that's why it used as a pigment), so when you obtain a precipitate it means there's still a lot of ferrous ion. But when the solution turns only slightly blue (no precipitate) that means very little ferrous ions are left. The test is very sensitive, so even traces of ferrous ions will give a bit of blue...

Unionised is right that peroxide in acid conditions can behave as a reducing agent, it's well known:

H2O2 === > O2 + 2 H+ + 2 e-

(it's powerful enough to reduce dichromate (Cr2O72-

to Cr(III), it also reduces Ce(IV) to Ce(III))

But I'm not sure it can reduce ferricyanide to ferrocyanide because they are very strong complexes. If it does, then false positives become possible because ferrocyanide also forms Berlin Blue but with ferric ions! I will try and check this later on.

To eliminate any excess peroxide try boiling the solution for a bit then cool, prior to adding the ferricyanide.

Chris: if you run the test using only a few drops of ferric chloride solution, then also only use a few drops of peroxide, 15 ml is really overkill.

[Edited on 21-6-2014 by blogfast25]

unionised - 21-6-2014 at 05:27

Quote: Originally posted by blogfast25

I'm not sure it can reduce ferricyanide to ferrocyanide
[Edited on 21-6-2014 by blogfast25]

Then read the link I posted.

blogfast25 - 21-6-2014 at 08:12

Quote: Originally posted by unionised

Then read the link I posted.

I didn't read that far down, but I have now.

"An interesting case of the oxidising and reducing action of hydrogen peroxide was discovered by Brodie. An acid solution of potassium ferrocyanide is oxidised by hydrogen peroxide to potassium ferricyanide: 2K4FeC6N6 + H2O2 = 2K3FeC6N6 + 2KOH. An alkaline solution of potassium ferricyanide, however, is reduced to potassium ferrocyanide by hydrogen peroxide, with evolution of oxygen: 2K3FeC6N6 + 2KOH + H2O2 = 2K4FeC6N6 + 2H2O + O2."

We are far away from alkaline ferricyanide here: the solutions aga and Chris use (ferric chloride) are strongly acidic.

Time for a few tests, methinks...

blogfast25 - 21-6-2014 at 09:08

Here’s what I did.

A fairly concentrated solution of K3Fe(CN)6 was split over two tubes. To one was added a flake of KOH, to the other some drops of HCl 12 M and to both the same amount (about 0.5 ml) of H2O2 32 %.

The alkaline tube immediately produced bubbles and the solution bleached, consistent with a reduction to K4Fe(CN)6.

The acid tube showed no visible change. Frustratingly, when I added a drop of what I believed to be strong FeCl3 (no ferrous iron at all) I still got a blue reaction. Whether that was due to some actual ferrous iron in the ferric chloride solution or due to some reduction of the ferricyanide to ferrocyanide, I do not know.

Looks like we could be with a decent testing protocol for ferrous ions by ferricyanide…

CHRIS25 - 22-6-2014 at 07:03

Quote: Originally posted by blogfast25

The acid tube showed no visible change. Frustratingly, when I added a drop of what I believed to be strong FeCl3 (no ferrous iron at all) I still got a blue reaction. Whether that was due to some actual ferrous iron in the ferric chloride solution or due to some reduction of the ferricyanide to ferrocyanide, I do not know.

Hi gert, all I can say is that the ferric ion in ferric ammonium citrate is mixed with the ferric ion in ferric potassium cyanide, when this is brushed on paper it is lime green, (from the citrate); when exposed to sunlight (or UV) the lime green becomes brown as a result of the ferric becoming ferrous in the light sensitive citrate (not blue as is written in some chemistry papers), the ferrous mixes with the ferric at the moment it hits water and it is at this point that the blue starts to appear. This is probably of no help at all, but it appears that when water is involved any ferrous ions in solution with a cyanide molecule become ferric again, ferric cyanide, (with the potassium salt absent). Now my knowledge is incomplete but I wonder if there is something here that might trigger a thought of for you? (from my days making cyanotyopes by the way).

blogfast25 - 22-6-2014 at 08:55

Interesting Chris, I'll have to look into that.

But I think a more established testing procedure (better than what we've been doing) should solve our woes easily. I'm looking out for such a procedure now.

It does appear that obtaining 100 % ferrous or 100 % ferric iron is not as easy as it may seem. Some Mohr's salt ((NH4)Fe(SO4)2.6H2O - an ammonium ferrous sulphate double salt) I prepared carefully some years ago still tested weakly positive for ferric ions.

Similarly, carefully prepared ammonium ferric alum tested weakly positive for ferrous ions.

CHRIS25 - 22-6-2014 at 09:30

Co-incidence here is that I am preparing a second batch of ferrous sulphate ready for a mohr's salt crystal. That it could contain ferric is something I certainly did not expect to read. How is that possible? I will do some research.

blogfast25 - 22-6-2014 at 10:00

It is as always due to inadvertent oxidation of a bit of the ferrous to ferric ions by air oxygen. slow as that process is it does happen.

Because Mohr's salt is prepared in acidic conditions that oxidation is suppressed but not eliminated altogether.

I read somewhere that adding some blank steel nails to the mixture of ammonium sulphate and ferrous sulphate during mixing and crystallisation can prevent it because the iron in the steel reduces any formed ferric ions back to ferrous ions:

Fe + 2 Fe3+ === > 3 Fe2+

I've no idea whether this is true.

Let me know how your Mohr's salt prep goes: mine have always had surprisingly low yields (50 %) for such a simple preparation.

[Edited on 22-6-2014 by blogfast25]

blogfast25 - 22-6-2014 at 12:40

From the electrochemical series:

Fe === > Fe2+ + 2 e | Eoxidation = +0.447 V

2 Fe3+ + 2 e === > 2 Fe2+ | Ereduction = +0.771 V

Total is Fe2+ + 2 Fe3+ === > 3 Fe2+ | Eoxidation + Ereduction = +1.218 V

Since as Eoxidation + Ereduction > 0, this means this redox reaction can proceed, at least from a thermodynamical point of view (Delta G < 0). So this little trick of adding some mails, a bit of steel wool or some steel plate to a preparation of Mohr's salt should really work to keep things ferric free!

I remember someone using this reaction to reduce a ferric solution to a ferrous one, using steel wool. The end product was ferrofluid...

[Edited on 22-6-2014 by blogfast25]

[Edited on 22-6-2014 by blogfast25]

MrHomeScientist - 23-6-2014 at 09:32

That someone was me!

Iron(II) and (III) chlorides - http://www.youtube.com/watch?v=iWpfHkWr5DY
Using these to produce ferrofluid - http://www.youtube.com/watch?v=LlQw9dfexBQ

I did use steel wool to reduce iron(III) back to iron(II), which seems to work (based only on visual inspection). The ferrofluid I ended up with was pretty mediocre, though.

aga - 23-6-2014 at 11:10

Quick note : today i have Ammonium thiocyanate, N₄SCN available, so tested my ferric and ferrous solutions with it.
They both give the same result - a blood red colouring.

chemrox - 23-6-2014 at 15:57

check bipyridine solution

blogfast25 - 24-6-2014 at 04:12

Quote: Originally posted by MrHomeScientist

That someone was me!

No, I got it from another non-video ferrofluid instructable.

aga - 24-6-2014 at 11:29

Quote: Originally posted by blogfast25

Quote: Originally posted by MrHomeScientist

That someone was me!

No, I got it from another non-video ferrofluid instructable.

This sounds suspiciously like a follow up to your doctor telling you to contact any sexual partners you have had in the past two years ...