Any suggestions of how to dissolve Fe2O3 in solution WITHOUT using HCl or HF?
Dear Sir. Concerning your issue about how to dissolve iron(III) oxide Fe2O3 in solution WITHOUT using HCl or HF. Information on the “solubility of Fe2O3 with different procedures and pH levels” also does not confirm to chemical logic and is self-contradictory. According to Mehra and Jackson (1960), the precipitation of FeS occured between pH 6.30 and 6.40 and also at pH 7.20 but not at pH 6.80 to 12.16. However, according to the previous observations of Aguilera and Jackson (1953), the precipitation of FeS did not occur in the presence of citrate in a neutral or alkaline pH. These two reports are, therefore, contradictory. Moreover, FeS is stable at higher pH and dissolves in acidic solutions; therefore, the formation of FeS, if any, should be more prominent in procedures employed at highly alkaline solution (pH > 12) rather than in the solutions with pH < 7. It may be mentioned here that Aguilera and Jackson (1953) did not observe the formation of FeS in their systems (pH 7.3) whereas Mehra and Jackson (1960) reported the formation of FeS in a similar system. Mehra and Jackson (1960) have subsequently attempted to justify the choice of a reaction pH of 7.3, based on the intersection of the oxidation potential curve for dithionite with the solubility curve of Fe2O3. According to the authors, the “oxidation potential increases sharply up to pH 8 and then levels off.” It is not known how the authors came to this conclusion because according to the reaction mechanism, the value of the oxidation potential (E) increases linearly with increase in OH- concentration (eqn. 1); it reaches a value of 1.12 V under standard conditions (as mentioned earlier). Hence the observation by Mehra and Jackson (1960) that the oxidation potential levels off at about 0.7 V at a pH 8.0 and above, is difficult to rationalise. The solubility of Fe2O3 with two procedures carried out by Mehra and Jackson (1960) do not match at all. Thus, by utilising one procedure, the solubility of Fe2O3 at pH 7.4 is only 65.1%; while with another procedure, it is reported that 100% dissolution can be obtained at about the same pH. Normally, solubility versus pH studies should have been done independent of procedures employed; this has not been done. Therefore, both the set of data, viz., oxidation poten. I think the following below links may help you in your analysis:
http://www.rcais.res.in/oxide_dissolution.pdf
https://pubchem.ncbi.nlm.nih.gov/compound/Iron_III__oxide#section=Experimental-Properties
http://www.atmos-chem-phys.net/12/11125/2012/acp-12-11125-2012.pdf
Thanks
You may also want to check the following somewhat related discussion: https://www.researchgate.net/post/How_can_I_remove_iron_oxide_from_teflon_paddle_stirrer
Several hydrates of Iron(III) oxide exists. When alkali is added to solutions of soluble Fe(III) salts, a red-brown gelatinous precipitate forms. This is not Fe(OH)3, but Fe2O3·H2O (also written as Fe(O)OH). Several forms of the hydrated oxide of Fe(III) exist as well. The red lepidocrocite γ-Fe(O)OH, occurs on the outside of rusticles, and the orange goethite, which occurs internally in rusticles. When Fe2O3·H2O is heated, it loses its water of hydration. Further heating at 1670 K converts Fe2O3 to black Fe3O4 (FeIIFeIII2O4), which is known as the mineral magnetite. Fe(O)OH is soluble in acids, giving [Fe(H2O)6]3+. In concentrated aqueous alkali, Fe2O3 gives [Fe(OH)6]3−.
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