AJKOER2

My experience with heating a copper penny was disappointing. As I was using a methane flame, I eventually realized that inserting the hot CuO / Cu2O in CH4 is like heating the oxide in hydrogen, with a reduction of the hot copper oxide back to Cu/Cu2O by CH4! If one somehow heats copper powder in air till it is black, is that just a CuO coating on possibly Cu/Cu2O? Not exactly a preparation of pure CuO it seems to me. Also, no heat needed. Just pour chlorine bleach (NaOCl) onto room temperature copper powder, and at least the surface looks like CuO! Here is a reference: https://www.hunker.com/13401530/what-does-bleach-do-to-copper .

Perhaps a more meaning insight can be gathered from employing a non-parametric correlation measure. For example, look at the rank correlations, the so called Spearman Rank correlation statistic. If they are computed here to be 1.000 for both, there is no meaningful difference. One can also look up the approximate standard deviation measure for R2, and, I would guess, with even a very large sample size , no statistical meaningful difference here as well.

Interesting to me would be the following reaction: Fe(NO3)3 + 3 [Cu(NH3)4(H2O)2]OH --> Fe(OH)3 (s) + 3 [Cu(NH3)4(H2O)2]NO3 Or: Fe(NO3)3 + 3 [Cu(NH3)2(H2O)4]OH --> Fe(OH)3 (s) + 3 [Cu(NH3)2(H2O)4]NO3 I have previously reacted the basic [Cu(NH3)4(H2O)2](OH)2 with Mg salts resulting in a precipitate of Mg(OH)2 and the corresponding tetra amine complex. Note: [Cu(NH3)x(H2O)y](OH)2 where x + y = 6 can be formed based on ammonia concentration used in the salt preparation.

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?

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.

Save your CuSO4 and try dissolving a small amount of Cu powder (copper sources include electric wires, copper pipe,..) into a mix of excess ammonia water, O2 (from an air pump or H2O2), optionally NH4HCO3, and a required small amount of sea salt to serve as an electrolyte. Reactions and source: 2 Cu + 4 NH3 + 1/2 O2 (or H2O2) + H2O --> 2 [Cu(NH3)2]OH (see https://onlinelibrary.wiley.com/doi/abs/10.1002/bbpc.19630670412 ) 2 [Cu(NH3)2]OH + 4 NH3 (aq) + 1/2 O2 + H2O --> 2 [Cu(NH3)4](OH)2 Cu + [Cu(NH3)4](OH)2 <---> 2 [Cu(NH3)2]OH CO2 + H2O = H2CO3 = H+ + HCO3- [Cu(NH3)4](OH)2 + H+ + HCO3- --> [Cu(NH3)4]CO3 + 2 H2O If employing H2O2 + NH3 + HCO3- + electrolyte, one can jump start the reaction for a minute in a microwave in an open vessel, which I usually place in a plastic bag with expansion room to control ammonia fumes. In a few hours, the solution will approach royal blue in color, at which point, remove unreacted copper metal. Then, boil down the solution, in a fume hood or outdoors, releasing steam, NH3, CO2 creating a precipitate of copper oxides, predominantly CuO if an excess of ammonia /oxygen source was used. See comments at US Patent 5,492,681 https://patents.google.com/patent/US5492681 especially, to quote: "(IVb) Cu(NH3)4 (OH)2 + HEAT → CuO + 4 NH3 + H2O " Note, an advantage of this method is that one can pass the above formed NH3/CO2 back into fresh water (or dilute H2O2) to make another batch upon adding copper. Also, the suggested path outlined above differs from the patent in the expressed use of sea salt or plain NaCl as an electrolyte given the underlying electrochemical aspects of the dissolution reaction of copper with ammonia and oxygen. The cited patent preference for the addition of ammonium sulfate may stem for its added benefit in acting as an electrolyte, in my opinion.

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.