Can anyone tell me what the reaction is between Iron (III) Nitrate and Sodium Thiosulfate solutions?

Thanks

JWP

Thanks.
Will excess Fe(3+) oxidise the S4O6(2-) to SO4(2-)?

Sulfuric acid decompose thiosulfate.
Thisulfate itself forms a complex with Fe 2+.

Deep purple sounds to me like thiocyanate Fe(III) complex (no idea why the color disappears, no idea what the color of thiosulfate complex is). Solution of Fe²⁺ almost always contains at least traces of Fe³⁺, as it gets oxidized with air oxygen very easily.

Deep purple? Fe(SCN)₆^3- is "bloody red"

The deep purple color definitely was from thiosulfate, not thiocyanate.  I am familiar with the latter blood/wine/bromine red color, the purple color is different.  One speculation I had was that the purple color is a Fe3+ complex with thiosulfate, then reduction of Fe3+ to Fe2+ by the thiosulfate resulted in either a colorless Fe2+/thiosulfate complex, or no complex at all.

I have no idea if that is close or not.

Repeating this experiment, making measurements of the amounts involved...

Note:  I usually calculate solution strength in moles / kg solution, not moles / liter since I can more accurately and conveniently measure mass than volume.  This is similar to molality, but not quite (I count total mass of solution, not mass of solvent.)  I'm not sure of the right name for my unit...

I have previously measured 1 drop of H2O as 0.055 ml, for the iron solutions involved I assumed that 1 drop then was 0.055 g of solution, this may be inaccurate.

Fe3+ / thiosulfate:
To 0.75 g of 0.26 moles/kg sodium thiosulfate (0.195 mmol), I added 1 drop of 0.247 ferric ammonium sulfate (in 1.781 moles/kg H2SO4), this is approx. 0.0135 mmol Fe3+.
I observed a purple color which faded quickly, after c. 10 sec it was mostly gone, after c. 30 sec completely gone.  After a little more time, I observe a yellowish/white cloudy suspension, assumed to be precipitated sulfur.

Before this, I did the same thing with twice as much thiosulfate, I saw still the purple color but not as intense.

Fe2+ / thiosulfate:
Now I repeat this using ferrous ammonium sulfate, nominally 0.487 moles/kg, with 0.162 moles/kg of H2SO4.  Since I had observed a slight purple color earlier with this solution (see earlier posts) I had added some iron filings to this (I heard somewhere this improves Fe2+ solution keeping properties.)  This was a day or so ago, immediately after adding the Fe filings, I got lots of bubbling (assuming this to be H2), by now this has mostly settled down, but my solution is now presumably overstrength.

Anyways...  To 0.65 g of the same 0.26 moles/kg thiosulfate (0.169 mmol thiosulfate), I added 4 drops of the ferrous solution.  (At least 0.107 mmol Fe2+, probably somewhat more.)
I observed no color change at all.  The resulting solution was very light green.  After 15 minutes or so, there is a slight cloudiness (possibly sulfur, not sure).

Just for fun, I then combined approximately equimolar amounts of Fe3+ and thiocyanate (as KSCN), and observed the familiar blood red color, which does not fade.  (At leat in 30 sec or so, I just did this.)

Ahh, molinity then.  (Now I need to find the abbreviation for it!)

Is molality used much?  The ratio of moles of solute to kg of solvent, instead of kg of solution, seems strange to me.  Moles per kg of total solution makes more sense to me.  Besides, there are cases where the identity of the solute vs solvent is not clear - e.g. 50% of EtOH in H2O.

Molinity would have many of the same advantages as molality, namely independence of temperature, and convenience/accuracy of not having to measure volumes.

Back to thiosulfate...

According to J Chem Ed, Vol 67 No 2, Feb 1990, p 146-149, "In Praise of Thiosulfate", the reaction between Fe3+ and S2O3- proceeds as follows...
First, a deep violet complex (Fe(S2O3)2)- forms almost immediately.  It is quite stable and decomposes very slowly, with a fading away of the violet color:
(Fe(S2O3)2)- + Fe3+ -> 2Fe2+ + S4O6--
The presence of Cu++ catalyses the decomposition, and then the violet color fades much more rapidly.

Perhaps I have Cu++ contamination?  There shouldn't be anything other than trace amounts but I cannot be sure.

I also noticed this purple complex between iron (III) and thiosulfate. I used reagent grade ammonium iron (III) sulfate and photo grade sodium thiosulfate. I also had the deep purple color (definitely different from the blood red color of iron (III) and thiocyanate). The color faded within a minute or 2 at a temperature of 15 C.

In fact, thiosulfate forms complexes with quite some metals. It also coordinates to silver and to copper (a strange yellow complex, which very slowly decomposes, giving CuS). As I understood, the complex formed with iron is [Fe(S2O3)2(H2O)2](-).

Iron (II) does not form a highly colored complex with thiosulfate.