Sciencemadness Discussion Board

Ferric nitrate and sodium thiosulphate

CHRIS25 - 31-5-2012 at 03:52

There is little agreement on various sites about the way this reaction should be written, even trying to balance it with two different sites yields the comment: This is an impossible reaction? Well they are up the creek because it is most definitely a reaction!

Anyway I can't balance it, so I need help: Fe(No3)3 + Na2S2O3.5H2O = NaNo3 + S2 + 2FeO3. I have to say that I know this equation is ridiculously wrong but I am not a chemist. I also notice that when I add more thiosulpate gradually the Iron nitrate colour dissappears magically. If I have a base foundation to work from, instead of guessing, then I can build different reactions from these two products that I have made. (A variety of Great results for my copper are happening, but I need more control)

Appreciations as usual.

[Edited on 31-5-2012 by CHRIS25]

ScienceSquirrel - 31-5-2012 at 04:16

Have a look here;

http://www.chemicalforums.com/index.php?topic=36416.0

CHRIS25 - 31-5-2012 at 07:37

Quote: Originally posted by ScienceSquirrel  
Have a look here;

http://www.chemicalforums.com/index.php?topic=36416.0


Thanks, so I see that thiosulphate is an oxidiser and reduces FE3 to FE2. However I am not using any catalysts, but the info here is quite scant - I guess by eye then, have done some experimenting with excellent results at different ratios but really can not understand what is happening other than guessing that the thiosulphate oxidizes the Iron 3 in solution to Iron 2 which then adheres to the copper surface. Whereas heating gently the same solution on the copper itself turns it to exactly the same product (FE2) Iron oxide, which of course is black. However in solution the copper turns shades of deep red-purple with a 2:1 ratio of Fe nitrate and thiosulphate and it turns metallic blue with my original stoichemetry balancing at an 8:1 ratio.

Nicodem - 31-5-2012 at 23:28

Quote: Originally posted by CHRIS25  
Thanks, so I see that thiosulphate is an oxidiser and reduces FE3 to FE2.

This sentence is oxymoronic. Oxidizers oxidize, reductors reduce. In any case, thiosulfate is a reducing reagent in this case. It is the Fe(III) that is the oxidizer. Check the flow of electrons when in doubt.
Quote:
However I am not using any catalysts, but the info here is quite scant

What catalyst? The reaction should be rapid without any catalysis. This should be one of those aqueous redox reactions that complete in matter of seconds or utmost minutes.
Quote:
I guess by eye then, have done some experimenting with excellent results at different ratios but really can not understand what is happening other than guessing that the thiosulphate oxidizes the Iron 3 in solution to Iron 2 which then adheres to the copper surface. Whereas heating gently the same solution on the copper itself turns it to exactly the same product (FE2) Iron oxide, which of course is black. However in solution the copper turns shades of deep red-purple with a 2:1 ratio of Fe nitrate and thiosulphate and it turns metallic blue with my original stoichemetry balancing at an 8:1 ratio.

You lost me here. Where did copper come from. There is no copper in the equation.

CHRIS25 - 1-6-2012 at 03:15

Hi Nicodem, thanks for replying. Yes you're right a very oxymoronic sentence - unfortunately I was not using the term "Reduced" in its chemical context - merely as a grammatical construction, should have avoided it then. Not being a chemist or studying chemistry professionally I will make these mistakes. I only mentioned about the catalyst because the site you pointed me to used a catalyst.

Ok, the copper, I am using the reaction and applying it as a boiled solution to copper. Hence my original post about the different colours.

I have recently come to understand that this reaction is a complex interaction between Iron and the thiosulphate, even more that it forms complex whatevers with metals such as copper. I do not understand ion exchanges although I am familiar and have read about them, I am unable to apply it to my own experiments. I am a metal worker and instead of simply performing exercises and experiments and reading formulas I want to understand and learn the chemistry behind what I do. this gives me an edge because so far in my early stages I have been able to produce brand new mixtures and achieved some nice results - simply by applying some chemistry and forethought.

My ultimate goal is Colour through chemistry, not chemistry for its own sake. I don't just learn what I need I try to understand so as to take things much further.

So once again I am having trouble writing this reaction correctly to show what is happening. And then to understand how it oxidises the copper is naturally what I am also trying to learn. In many other situations such as copper acetate and ammonia, or NACL and Ammonia, or potassium polysulphide, or the carbonates or the ammonium acetate or the ferric solutions or the oxidising power behind the K.permangenates I can understand where the colours are coming from. But this Iron nitrate and sodium thiosulphate reaction is very complex indeed, both cold and hot producing different colours, and more importantly the varied ratio of thiosulphate concentration inside the iron nitrate and water producing two separate results, the white prcepitate that appears when you cross in a very minute small way the amount of water in the thiosulphate and iron nitrate solution. So these are my questions. Searching on the internet for help did not yield anything. So to keep things simple I thought that I would ask a basic question - a starting point if you like from where I could build some understanding about this reaction.

Kind regards

[Edited on 1-6-2012 by CHRIS25]

weiming1998 - 1-6-2012 at 04:53

Quote: Originally posted by CHRIS25  
Quote: Originally posted by ScienceSquirrel  
Have a look here;

http://www.chemicalforums.com/index.php?topic=36416.0


Thanks, so I see that thiosulphate is an oxidiser and reduces FE3 to FE2. However I am not using any catalysts, but the info here is quite scant - I guess by eye then, have done some experimenting with excellent results at different ratios but really can not understand what is happening other than guessing that the thiosulphate oxidizes the Iron 3 in solution to Iron 2 which then adheres to the copper surface. Whereas heating gently the same solution on the copper itself turns it to exactly the same product (FE2) Iron oxide, which of course is black. However in solution the copper turns shades of deep red-purple with a 2:1 ratio of Fe nitrate and thiosulphate and it turns metallic blue with my original stoichemetry balancing at an 8:1 ratio.


I do not think iron (II) oxide is formed in such a reaction. If you don't boil down the mix, the Fe2+ ions and the tetrathionate ions, S4O6-, will not react. I speculate that it will not react because Fe2+ is a fairly weak Lewis acid, at least compared to Al3+, etc. This is because FeS, unlike Al2S3, etc, is stable in water and will not hydrolyze easily, indicating that only a very, very tiny amount of it disassociates to H+ ions and complexes, so it will not react easily with the tetrathionate ion to form tetrathionic acid, which is a Lewis adduct of SO3 and H2S itself, decomposing to SO2, S, and other complex products easily. Read http://en.wikipedia.org/wiki/Lewis_acids_and_bases
for more information.

So at the end of your reaction, Fe2+ and S4O6-2 ions will be left after the colour change. I do not think heating such solutions with copper will generate anything, as elemental copper is fairly inert and reacts with quite strong oxidizing agents only (most of the time!), and such a solution is a reducing agent. Over time, O2 might induce pitting corrosions in the copper (http://en.wikipedia.org/wiki/Pitting_corrosion), but it will not happen in the timeframe of an experiment. Even CuO wouldn't dissolve in the solution, as the pH is too high for that.

If you heat this solution, eventually you will get Fe(OH)2, which simultaneously dehydrates/oxidizes in air to form Fe2O3. It even disproportionates spontaneously below 575 degrees celsius to form Fe metal and Fe3O4. http://en.wikipedia.org/wiki/Iron(II)_oxide

The only possible situation the copper can dissolve in the solution, as far as I know, is excess nitric acid acidifying the solution. The pH will drop down, causing CuO to dissolve, then the oxidizing power of acidified NO3- ions will slowly dissolve the copper, forming Cu2+ ions, which is blue.

This is what I think, but looking at this site (http://www.crscientific.com/article-redox3.html) the reaction could very well be much more complicated.

If I have posted a lack of reliable sources, I apologise.

CHRIS25 - 1-6-2012 at 05:28

Hi Welming, thankyou very much for that last website pointer. Your explanation is exhaustive - many appreciations. Firstly, not a criticism ok, I am not trying to dissolve copper, just dipping it into the boiling solution, that is how I get the colours on the copper, the oxidation.

Now however if I understand you generally, what is happening is that the solution is becoming Ferric Hydroxide with thiosulphate complexes? This is coated over the copper surface, withdrawn from solution and a purple to fire red to metallic blue is there depending upon length of time submerged, (10 seconds to 50 seconds). This leads to a Ferric oxide being formed on the copper surface when it is exposed to the air.

I notice this purple when the thiosulphate immediately begins to dissolve in the diluted ferric nitrate, you can see what appears to be smokes of purple leaving the thiosulphate crystals and immediately dissappearing into the solution(before being heated this is).

Finally you said: ===="If you don't boil down the mix, the Fe2+ ions and the tetrathionate ions, S4O6-, will not react. "====This actually is very interesting because if I put the solution cold onto the copper I get quite an ugly deposit and very inconsistent multi colours of yuck! So the fact that it is being heated means that the colours are actually coming from the Fe and tetrathionate mixing. If this is so how could I write an equation for this last reaction. This where I do have to beg for spoon feeding on this one point, it might help me to understand the different reactions that are taking place within minutes of each other. It is a bit like there are multi-reactions- should I say consecutive reactions therefore equations in the copper chloride 2 and copper chloride 1 solutions that I make and work with.

Thanks Welming

[Edited on 1-6-2012 by CHRIS25]

watson.fawkes - 1-6-2012 at 06:16

I see two questions here. The first is what's happening in preparing your solution. The second is what's happening on the Cu surface that's giving rise to color.

For the first question. I've seen the thread where you prepare the iron nitrate solution. You're writing that as Fe3+, but there may be significant Fe2+ content as well. Excess metallic iron Fe0 can reduce Fe3+ to Fe2+, analogous to the reaction for copper etching. On the other hand, nitrate acts as an oxidizer, reversing this, but if nitrate is exhausted and you've got excess iron, there could be Fe2+. This matters because if Fe2+ is present it could be acting as an electron donor when you add thiosulfate.

When you add the thiosulfate solution, is there any evolution of gas?While writing this, I see you've replied that this does seem to be happening. Thiosulfate decomposes in acidic solutions, liberating SO2. The ordinarily-state decomposing also has elemental sulfur as a product; it forms small particles that scatter light and would change the visual texture of the solution. One the other hand, with other ion species present, there could be thiosulfate decomposition without liberating sulfur. Also if the nitrate is acting as an oxidizer, it would release NO or NO2. One or both of these could be happening.

All that said, you might be getting tetrathionate out of as a product, given that some of the color you report are similar to tetrathionate complexes.

All this is to say that I'm very uncertain what's happening when you prepare your solution, because I'm not certain what the starting iron nitrate solution is. As I recall, there's excess nitrate, but there were also cases where there was excess iron remaining when the solution stopped reacting. So it's conceivable to me that you have a mix of Fe3+ and Fe2+ in an acidic solution with excess NO3-.

So, a perhaps-simple question. What's the pH of your starting iron nitrate solution? This addresses excess nitrate. Also, if you add H2O2 to that solution, do you get any qualitative change? If so, that would indicate Fe2+.

One way to get a consistent iron nitrate solution would be to crystallize iron(III) nitrate out of your original solution and redissolve it in distilled water. This would get rid of excess nitrate. It might also not work for your coloring mixture, if the excess acid is necessary to make it work.

watson.fawkes - 1-6-2012 at 06:33

Now for the surface chemistry. Similar to the last answer, there are a lot of possibilities. One thing to realize is that iron can form mixed-valence compounds. The most well-known is magnetite Fe3O4, which has one Fe2+ and two Fe3+ in its crystal structure. Since you're using heat, there's a number of intermetallic compounds that you could be getting. Copper itself has two oxidation states. The possible combinations are large, and because the reaction isn't driven by chemical potentials but by heat, reaction that are ordinarily most likely may not be the ones that are happening.

Iron(III) chloride is a copper etchant, and so there's reasonable belief that you're getting iron reduction and copper oxidation. But you've also got any number of nitrogen and sulfur oxoanions running around. Given that your colors change significantly depending on application time, the resulting color, particular for the short immersion times, could be transitory product that would otherwise be an intermediary in a reaction that went to "completion".

One of the difficulties in assaying the chemical composition of the color is that you are possibly also getting thin film interference colors. Films that are consistently thin near the wavelength of light will preferentially reinforce certain frequencies. Your resulting films may be a combination of the chemical color and the thin film color. You'd need a spectrometer to differentiate these.

CHRIS25 - 1-6-2012 at 08:49

Hallo Watson.Fawkes. Well,you have given me quite a lot to research about and digest - I look forward now to some starting point. I can however say that I have made two more nitrate solutions, not that the first onewas successful, rather I have done this three times now. My last solution was as follows: 8.8grams rusty barbed wire (The rust was intentional for the iron oxide) + 100mils 30% nitric acid. Heated vigorously untill only steam was emitted, and then heated for a further 30 minutes until that 100mils was reduced to 50 mils precisely. Rich red clear colour filtered. This was my Ferric nitrate, my personal hope and belief was that there couild not be any FE2 ions present at all, there was black powdery bits left over from both reactions (The last two) but not much, I assume this to be no indication of excess Iron because when this happened the second time I simply added more acid and it remained there.

When I add the thiosulphate to the Ferric nitrate and water mixture of 8g Ferric nitrate 1g thiosulphate and 50mils water there is no smell at all and no gas and the red ferric colour changes to a nice yellow. However when I add 8g ferric nitrate and 4g thiosulphate and 50 mils water there is a clear distinct smell of sulphur and the solution turns milky. After a while a white densely packed precipitate lies on the bottom of this solution after about two hours.

Very much appreciated your response above,



[Edited on 1-6-2012 by CHRIS25]
Have just finished another Ferric nitrate preparation and am pleased because I have a clear yellow solution, my last solution which is the one I used with the thiosulphate is a bright wine red solution. Now I know the yellow one is more purer than the red one but have not understood these differences. I really am not aware that I did anything different EXCEPT boil that yellow one at least three times longer.

[Edited on 1-6-2012 by CHRIS25]

weiming1998 - 1-6-2012 at 18:00

Quote: Originally posted by CHRIS25  
Hi Welming, thankyou very much for that last website pointer. Your explanation is exhaustive - many appreciations. Firstly, not a criticism ok, I am not trying to dissolve copper, just dipping it into the boiling solution, that is how I get the colours on the copper, the oxidation.

Now however if I understand you generally, what is happening is that the solution is becoming Ferric Hydroxide with thiosulphate complexes? This is coated over the copper surface, withdrawn from solution and a purple to fire red to metallic blue is there depending upon length of time submerged, (10 seconds to 50 seconds). This leads to a Ferric oxide being formed on the copper surface when it is exposed to the air.

I notice this purple when the thiosulphate immediately begins to dissolve in the diluted ferric nitrate, you can see what appears to be smokes of purple leaving the thiosulphate crystals and immediately dissappearing into the solution(before being heated this is).

Finally you said: ===="If you don't boil down the mix, the Fe2+ ions and the tetrathionate ions, S4O6-, will not react. "====This actually is very interesting because if I put the solution cold onto the copper I get quite an ugly deposit and very inconsistent multi colours of yuck! So the fact that it is being heated means that the colours are actually coming from the Fe and tetrathionate mixing. If this is so how could I write an equation for this last reaction. This where I do have to beg for spoon feeding on this one point, it might help me to understand the different reactions that are taking place within minutes of each other. It is a bit like there are multi-reactions- should I say consecutive reactions therefore equations in the copper chloride 2 and copper chloride 1 solutions that I make and work with.

Thanks Welming

[Edited on 1-6-2012 by CHRIS25]


Sorry, I overlooked the fact that there might be excess Fe3+ ions in your solution. If that is the case, you have got a mix of Fe3+ ions, which are brown, Fe2+ ions, which are green, Cu2+ ions, which are blue, and thiosulfate complexes of copper and maybe even small amounts of CuO/CuS. The leftover solution, especially when heated, probably is a mix of numerous different compounds that is difficult to tell apart. The colours could all be complexes of a certain kind reacting and decomposing. Pouring the solution onto copper could mean that insoluble Cu+ compounds, as well as maybe even CuO, etc, can be formed, making the ugly deposit.

Writing a general equation for such a reaction is impossibly hard, as impurities might also play a role in the colour changes observed, and the fact that this is a complex series of reactions.

Finally, H2O2 will not work in determining the Fe2+ concentration in the solution. What it will do instead is making the H2O2 oxidize the Fe2+ to Fe3+, but then reduce the Fe3+ to Fe2+, making reactive species of oxygen in the process. See http://en.wikipedia.org/wiki/Fenton's_reagent
for more detail.

CHRIS25 - 2-6-2012 at 01:18

Hi welming, If I had excess Fe2 ions I would have a greener looking solution. My previous solution (as mentioned to watsonfawkes is red, and my last solution prepared yesterday is yellow - brown and fairly clear, this latter indicates a thoroughly Fe3 concentrated sol? I also thought that adding H2O2 to the solution would ensure oxidation of Fe2 to Fe3 and consume a little bit of Hno3? But I did not think that this process would reverse itself in the same solution?

On to the thiosulphate additions - I see that yes it is quite a complex reaction and could not be written, as you have said. What do you mean by "reactive species of oxygen"? But the main point you raise is about the Fe2 ions which I should not have at all. I am referring of course to the Fe Nitrate solution before anything is added.

Yesterday I added a small amount of Peroxide (6%- strength - is all I can get), I added this to a 5ml Ferric nitrate solution in a test tube. There was absolutely no change in colour and no visible bubbling or reaction, I understand from this that it means that this is a good Fe3 solution since nothing needed to be oxidized meaning that there was no appreciable amount of Fe2 in solution?

weiming1998 - 2-6-2012 at 05:22

Quote: Originally posted by CHRIS25  
Hi welming, If I had excess Fe2 ions I would have a greener looking solution. My previous solution (as mentioned to watsonfawkes is red, and my last solution prepared yesterday is yellow - brown and fairly clear, this latter indicates a thoroughly Fe3 concentrated sol? I also thought that adding H2O2 to the solution would ensure oxidation of Fe2 to Fe3 and consume a little bit of Hno3? But I did not think that this process would reverse itself in the same solution?

On to the thiosulphate additions - I see that yes it is quite a complex reaction and could not be written, as you have said. What do you mean by "reactive species of oxygen"? But the main point you raise is about the Fe2 ions which I should not have at all. I am referring of course to the Fe Nitrate solution before anything is added.

Yesterday I added a small amount of Peroxide (6%- strength - is all I can get), I added this to a 5ml Ferric nitrate solution in a test tube. There was absolutely no change in colour and no visible bubbling or reaction, I understand from this that it means that this is a good Fe3 solution since nothing needed to be oxidized meaning that there was no appreciable amount of Fe2 in solution?


Adding in H2O2 doesn't do anything, as the H2O2 will simply reduce the Fe3+ to Fe2+, then simultaneously oxidize it. Taken from the Wikipedia article, the two redox reactions are as:(1) Fe2+ + H2O2 → Fe3+ + OH· + OH−
(2) Fe3+ + H2O2 → Fe2+ + OOH· + H+
The reactive species of oxygen are the OH. and OOH., each having short half-lives (though not few seconds, instead being around for half an hour or so). They might react with each other to form O2 and H2O, or simply oxidize any Fe2+ to Fe3+. But telling the purity of a solution from colour change is not very accurate.

Anyway, a way to oxidize the solution is to simply leave it outside. Over time, Fe2+ solutions will slowly oxidize to Fe3+. You can also try the H2O2, but that requires large amounts of 6% is what you have.

Finally, there is Fe2+ contamination in your ferric nitrate because even nitric acid cannot oxidize all Fe2+ to Fe3+, and some ferrous nitrate will no doubt be formed.

barley81 - 2-6-2012 at 13:33

Maybe permanganate titration would be a good way to determine Fe2+ content, but I don't think there is much ferrous nitrate in solution. Ferrous nitrate in (even dilute) nitric acid is unstable and easily decomposes to nitric oxide and basic ferric nitrate, which should form ferric nitrate if there is excess acid present.
Quote:

Ferrous nitrate is prepared by dissolving iron in cold dilute nitric acid (1.10 sp. gr.). But a considerable amount of ammonium nitrate is also formed in the solution, according to the reaction:-

4 Fe + 10 HN03 =4 Fe(N03)2 +NH4NO3 + 3 H20.

This solution is very unstable and decomposes when heated even slightly, forming basic ferric nitrate and liberating nitric oxide. To prepare a pure ferrous nitrate, decomposition of a ferrous sulphate solution by barium or lead nitrate is employed:-



Read more: http://www.lenntech.com/chemistry/nitrates.htm#ixzz1wfpZYjyj

CHRIS25 - 2-6-2012 at 13:37

Ok, but leaving outside would take forever. But thanks for the elucidation about the FE 2 and 3. and for clarifying the reaction a bit better. But I am sure that I read somewhere that adding H2O2 to a sample of your nitrate was a reliable indicator about whether there was free FE2 still in solution, trouble is i can not source that info anymore - it was just something that I remembered which is why I did it. But if you say so....

watson.fawkes - 4-6-2012 at 05:51

Quote: Originally posted by weiming1998  
Adding in H2O2 doesn't do anything, as the H2O2 will simply reduce the Fe3+ to Fe2+, then simultaneously oxidize it.
This is missing the other part of the reaction, which is the formation of free radicals. In the Wikipedia page you reference, these free radicals attack organic compounds. This is the applications of Fenton's reagent for wastewater treatment, where it is an ordinary industrial process. In the present system, lacking easy-to-chew organics, the radicals aren't going to sit around doing nothing. I won't deny that the iron ions participate as you outline to split up the hydrogen peroxide, but the Fe2+ will also react with the radicals so produced and oxidize.