iron sulphate

[Suollo George]

I learnt that copper sulphate changes from blue to a green colour in the presence of iron. please i would like to know the stoichiometry of this reaction. Any help? thank you

[hypervalent_iodine]

Is this homework? You might like to start by writing the equation for the reaction out.

[Suollo George]

Thank you dear

Please I am trying to use an understanding of such a reaction to help myself in a project work. The work is in Histopathology though. I do not know much about the reaction. I read it somewhere and tried it at the hostel by using copper sulphate from the blood bank and an iron nail. it worked but my problem is when I use it in the project work it doesn't yield positive results. i am hoping to gain a deeper understanding of the reaction in the areas of the concentration of the copper sulphate needed, valency and the form of the iron ion.

[hypervalent_iodine]

I would typically recommend that you don't do chemistry if you lack a rudimentary understanding of what you are doing, for safety reasons.

The reaction involved is simply between CuSO₄ and elemental Fe. The ion it produces is likely the (III) oxidation state, I think. You should be able to Google it.

What sort of project is this? 

[Suollo George]

Thank you dear

[John Cuthber]

10 hours ago, hypervalent_iodine said:

The ion it produces is likely the (III) oxidation state, I think

I'd bet against that.

[hypervalent_iodine]

2 hours ago, John Cuthber said:

I'd bet against that.

Upon thinking about it, I see your point. Do you know why, or is it simply because its more energetically favoured not having to oxidise all the way to Fe3+ and undergoing a single displacement instead? 

[John Cuthber]

Iron (metal) will rapidly reduce Fe(III) ions in solution.

The strongest oxidant among the starting materials is Cu(II).

The process for etching metallic copper with ferric chloride (by a redox reaction giving Fe(II) and Cu(II) )shows that Fe(III) is a stinger oxidant that Cu(II).

You can't make a stronger oxidant by using a weaker one.

 

[AJKOER2]

On 2/28/2018 at 4:32 PM, John Cuthber said:

Iron (metal) will rapidly reduce Fe(III) ions in solution.

 

An extract of this article: “Kinetics of FeIII EDTA complex reduction with iron powder under aerobic conditions”, by Feiqiang He, et al, link: http://pubs.rsc.org/en/content/articlelanding/2016/ra/c6ra05222c/unauth#!divAbstract  suggests some qualifications to the above statement.To quote from the abstract:

“Reduction of FeIII EDTA is the core process in a wet flue gas simultaneous desulfurization and denitrification system by FeII EDTA solution. Metal powders, such as aluminum, tin, and zinc, have been proposed to reduce FeIII EDTA. In this paper, iron powder was chosen as a reductant to regenerate the absorption solution.”

So, more correctly and generally, the following equilibrium reaction is valid when expressed as follows:

 Fe(0) + 2 Fe(lll)-ComplexA = 3 Fe(ll)-ComplexB  

most likely in an oxygen free conditions (as in the presence of H+ and O2 the Fe(ll) is converted back to Fe(lll) ).

As such relative to the thread question of the action of iron on aqueous CuSO4, the claimed reductive powers of Fe on any formed acidic solution of iron(lll) sulfate, creating Fe(ll), in the presence of any dissolved oxygen, assuming sulfate serves in the role of a complexing agent, is questionable, in my opinion.

Edited March 30, 2018 by AJKOER2
Back to thread 's question, balance eq , conclusion pH

[John Cuthber]

9 hours ago, AJKOER2 said:

An extract of this article: “Kinetics of FeIII EDTA complex reduction with iron powder under aerobic conditions”, by Feiqiang He, et al, link: http://pubs.rsc.org/en/content/articlelanding/2016/ra/c6ra05222c/unauth#!divAbstract  suggests some qualifications to the above statement.To quote from the abstract:

“Reduction of FeIII EDTA is the core process in a wet flue gas simultaneous desulfurization and denitrification system by FeII EDTA solution. Metal powders, such as aluminum, tin, and zinc, have been proposed to reduce FeIII EDTA. In this paper, iron powder was chosen as a reductant to regenerate the absorption solution.”

So, more correctly and generally, the following equilibrium reaction is valid when expressed as follows:

 Fe(0) + 2 Fe(lll)-ComplexA = 3 Fe(ll)-ComplexB  

most likely in an oxygen free conditions (as in the presence of H+ and O2 the Fe(ll) is converted back to Fe(lll) ).

As such relative to the thread question of the action of iron on aqueous CuSO4, the claimed reductive powers of Fe on any formed acidic solution of iron(lll) sulfate, creating Fe(ll), in the presence of any dissolved oxygen, assuming sulfate serves in the role of a complexing agent, is questionable, in my opinion.

Two points;

Firstly, your  equation " Fe(0) + 2 Fe(lll)-ComplexA = 3 Fe(ll)-ComplexB  "

is the same as my assertion "Iron (metal) will rapidly reduce Fe(III) ions in solution."

So, your paper actually confirms what I said and

 

Secondly I have done the reaction, and reality doesn't agree with your bizarre opinion.

(Incidentally, it's at least as likely to be an aquo complex as a sulphato one -though I have also done the reaction with ferric chloride- where chloro complexes probabaly dominate.)

 

[AJKOER2]

So, John can we agree that not citing the need for a soluble form of Fe(lll) (and not rust, for example) may be educationally misleading with respect to the reductive powers of iron metal?

Further, to explain the significance of oxygen and pH, first:

Fe(ll) + O2 = Fe(lll) + O2.-

which is the so called metal auto-oxidation reaction, where oxygen acting on ferrous, produces ferric and the superoxide radical anion. Importantly, the reaction is reversible. But, if we add +H to both sides:

Fe(ll) + O2 + H+ = Fe(lll) + (H+ + O2.-)

And, at pH < 4.88, the removal of the superoxide radical anion:

H+ + O2.- → .HO2

So, not surprisingly, upon adding sufficient acid to lower pH, one can move the metal auto-oxidation reaction to the right, converting ferrous to ferric.

But, if the solution pH is above 4.88, no significant oxidation of Fe(ll) to Fe(lll).

As such, it is a question of pH also, hence one cannot be definitive, my point. Note, to quote a source,

“Since most Ferrous Sulfate solutions have a pH of approximately 2”  (source: http://www.qccorporation.com/solutions/ )

such a pH could be problematic in the presence of air.

If a purpose of this forum is to educate, than citing driving reaction factors should not be assumed as obvious in my opinion.

Edited March 30, 2018 by AJKOER2
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[John Cuthber]

When you have finished, metallic iron will still reduce ferric to ferrous.

Also, re

1 hour ago, AJKOER2 said:

So, not surprisingly, upon adding sufficient acid to lower pH, one can move the metal auto-oxidation reaction to the right, converting ferrous to ferric.

In the real world, iron (II) is oxidised to iron(III) much more quickly in alkaline conditions- If you want to stabilise Fe(II) you add acid.

You seem to be trying to say it's the other way round.

[AJKOER2]

John, your comment "When you have finished, metallic iron will still reduce ferric to ferrous" is not that immediately obvious to me.

Some of my thinking, first, as long as there is H+ and O2, the following reaction continues:

Fe(ll) + O2 + H+  = Fe(lll) + .HO2 (or,  H+ + O2.-  when pH > 4.88)

Also, I would argue that the oxidation of iron to ferrous is slowed down in neutral conditions, as is the reduction of ferric.

Interestingly, the salt Fe3O4 (actually a mixed salt, FeO.Fe2O3) is created in neutral conditions absent oxygen.

So why is the answer not some equilibrium mix?

 

Edited March 30, 2018 by AJKOER2
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