1

I don't need to necessarily separate pure iron, but from a ferrous sulfate solution, I am trying to form a precipitate with iron. By getting a precipitation of an iron compound, I can find out the amount of iron in the solution. Any help would be appreciated. I'm completely new and tried finding information online but all of it was beyond my comprehension. I apologise for my ignorance, but help would really be great.

tripleee
  • 129
  • 1
  • 7
hayne newell
  • 33
  • 3
  • Are you a high school student, first semester, who asked you to do that? Please give some background to get a useful answer. – Karl Jun 23 '20 at 07:12
  • 2
    Anyway, ask your prof/teacher for a textbook on quantitative inorganic analysis, specifically gravimetry. You won´t get good results unless you do this by the book. The standard technique involves precipitation as hydoxide, filtration, and calcination to get the oxide, which you then weigh. – Karl Jun 23 '20 at 07:45
  • @Karl thankyou, yes I am a highschool student in Australia. My teacher doesn't know how to do it, but I'll look into the textbooks. I haven't been able to find any standard techniques anywhere yet but I'll keep trying. – hayne newell Jun 23 '20 at 10:40
  • @Karl and also, it's a depth study, so we chose our own experiment. I starting looking into doing something like extracting a substance from a product. Then came the idea of extracting (or finding the amount) of iron in ferrous sulfate supplements. I don't know enough about chemistry yet to know what works and what won't. The main aim of the assignment is to discuss percentage yield. I would like to learn why each step is done, but at the moment im just trying to find out how to actually do the experiment. Thankyou for the general procedure, I'll look into how to do both of them – hayne newell Jun 23 '20 at 10:45
  • Did I understand correctly: You would like to quantify the iron content of a ferrous sulfate solution? – Martin - マーチン Jun 23 '20 at 10:51
  • Yes, thankyou for that experiment you linked to. I'll try it out next lab and see how it works – hayne newell Jun 23 '20 at 11:18
  • 2
    As an alternative to precipitation, you should look into titration. Iron forms very stable complexes with EDTA. – Martin - マーチン Jun 23 '20 at 12:01
  • @Karl I was just wondering what the calcination was for? Once the hydroxide precipitate forms, can it be weighed and from there find the iron contents? – hayne newell Jun 24 '20 at 03:38
  • 2
    No, hydroxides are notoriously unstochiometric, i.e. full of water. You need to transform it into the pure oxide. – Karl Jun 24 '20 at 20:18
  • What a cute teacher. Instead of teaching, he leaves it to the pupils to find a suitable (university-level) textbook. wtf – Karl Jun 24 '20 at 20:20
  • @Karl thankyou, he is helpful with what he knows. I wasn't sure if calcination was needed. I did a rough run with just the hydroxides and got 74% yield (I understand this is likely innacurate due to not calcinating, and unreliable since I didn't do it thoroughly). You've been a great help in making it simple enough for me to understand – hayne newell Jun 25 '20 at 01:40
  • @Karl my teacher said that not doing the calcination and just discussing the unreliability of the hydroxides is enough. I would have liked to do the calcination to make it accurate, but I think this is the best I can. Is there a way I can contact you again for more resources/guidance. Thank you again for your help – hayne newell Jun 25 '20 at 01:45
  • If I were you I would have suggested a displacement reaction like putting a Mg rod in salt solution and finding out weight gain on the Mg rod, weigh the rod before and after reaction. It should be OK. – Harikrishnan M Jan 24 '24 at 08:30

2 Answers2

2

Decomposing ferrous sulfate to get iron(III) oxide should work:

$$\ce{2FeSO4 ->[\Delta] Fe2O3 + SO2 ^ + SO3 ^}$$

Iron content can be measured from iron(III) oxide. Commercially, iron is recovered from waste ferrous sulfate in the form of magnetite ($\ce{Fe2O3}$) by co-precipitation and magnetic separation$\ce{^{[1,2]}}$. Iron content is measured from magnetite. It was observed that after processing the magnetite, the grade of iron in magnetite concentrate increased from 62.05% to 65.58% and the recovery rate of iron decreased from 85.35% to 80.35%.

References

  1. Wang YU, Ying-lin PENG, Ya-jie ZHENG Recovery of iron from waste ferrous sulphate by co-precipitation and magnetic separation, Transactions of Nonferrous Metals Society of China, Volume 27, Issue 1, January 2017, Pages 211-219, DOI: https://doi.org/10.1016/S1003-6326(17)60024-4

  2. Yu, W., Peng, Y. & Zheng, Y. Effect of purification pretreatment on the recovery of magnetite from waste ferrous sulfate. Int J Miner Metall Mater 23, 891–897 (2016), DOI: https://doi.org/10.1007/s12613-016-1304-2

Nilay Ghosh
  • 25,509
  • 26
  • 91
  • 197
  • 2
    You don't really need step two if you only want to determine the iron content. This lab manual (pdf) oxidises first: $\ce{2Fe2+ + H2O2 + 2H+ -> 2Fe3+ + 2H2O}$ Then precipitates it $\ce{Fe3+(aq) + 3 OH- (aq) + x H2O(l) -> Fe(OH)3.xH2O (s)}$ Then probably heating it thoroughly: $\ce{Fe(OH)3.xH2O (s) ->[\Delta] Fe2O3(s)}$. – Martin - マーチン Jun 23 '20 at 10:59
  • @Martin-マーチン This case is also possible. I edited the answer to address the issue. Shouldn't need pure iron to quantify. – Nilay Ghosh Jun 23 '20 at 11:23
  • Last week magnetite was Fe3O4, and hematite was Fe2O3. When did it change? – Oscar Lanzi Jan 24 '24 at 11:46
1

I would do the following to get pure iron powder without contamination:

  1. Evaporate to dryness.

  2. Roast the $\ce{FeSO4}$ in an oxygen enriched environment to form $\ce{Fe2O3}$. You will notice the change in color. This has to happen at high temperatures (minimum: $\pu{400°C}$ ideal: $\pu{600-700°C}$) and can take quite a while. The best thing to use is a quartz tube as you can see the change in color.

  3. Hydrogen reduce the $\ce{Fe2O3}$ in an hydrogen atmosphere at about $\pu{400°C}$ to get the pure iron powder.

I have done this procedure - or at least a variation of this procedure with other metals and it works.

But be aware that this procedure will yield toxic gases and you have the risk of explosion when you work with hydrogen - especially when you do not know exactly what you do.

Nilay Ghosh
  • 25,509
  • 26
  • 91
  • 197
Andreas
  • 426
  • 3
  • 12
  • 1
    This is a very unorthodox procedure. Would you mind telling why exactly you did that? – Karl Jun 24 '20 at 20:23
  • The use of hydrogen $\ce{H2}$ is the best way for producing metallic iron out of iron compounds. @ Karl. I doubt Andreas will give you an answer more than three years after his remark. – Maurice Jan 24 '24 at 10:55