(1) Copper Sulfate Is in Demand, Chemical & Engineering News Archive 1958 36 (28), 28-31(2) Removal of Copper Sulfate from Water by Ferric Floc, C. J. Brockman, Industrial & Engineering Chemistry 1934 26 (9), 924-924(3) Copper (II) Removal from Wastewaters by a New Synthesized Selective Extractant and SLM viability, Raffaele Molinari, Teresa Poerio, Roberta Cassano, Nevio Picci, and Pietro Argurio, Industrial & Engineering Chemistry Research 2004 43 (2), 623-628(4) An Introduction to Chemical Systems in the Laboratory, Department of Chemistry, University of Illinois, Urbana-Champaign, 2020.Copper/Iron StoichiometryAlexander J. ArchieBB1September 24, 2020AbstractThe goal of this lab was to determine the product of the reaction of copper (II) sulfate andiron through qualitative and analytical analysis. 5.0030 g of CuSO4 (Table 1) was reacted with 1.4862 g of Fe, and the solution created had 1.6770 g of solid copper precipitate. Based on this data, it was determined that iron was the limiting reagent, and that iron (II) sulfate was the iron sulfate product formed. The percent yield was 99.17%.IntroductionThis lab was focused around the reduction of CuSO4 to Cu using Fe. The goal was to determine what the limiting reagent was, find the percent yield, and to practice the techniques of quantitative transfer and vacuum filtration. The aforementioned reduction was identified to have two possible chemical reactions. These reactions are dependent on the oxidation state of Fe, and they are represented as such:CuSO4(aq) + Fe(s) → Cu(s) + FeSO4(aq) (1)3CuSO4(aq) + 2Fe(s) → 3Cu(s) + Fe2(SO4)3(aq) (2)In order to find which equation occurs predominantly, the limiting reagent of the reaction must be determined. By choosing a certain amount of CuSO4 and Fe to react, both the limiting reagent and predominant reaction can be discerned. Assuming that there is slightly more CuSO4 than Fe, and reaction (1) occurs, then Fe is the limiting reagent and it determines how much Cu precipitate is produced. However, if reaction (2) is favored, then CuSO4 becomes the limiting reagent. In order to ensure accurate results in this process, though, measurements must be veryprecise. This is why vacuum filtration was utilized when measuring the mass of copper precipitate, and why an analytical balance was needed for all the other measurements. The premise of this lab may seem rather benign, but contamination by copper sulfate in water supplies is an issue that has gone back many decades, due to its high solubility and many uses. Copper sulfate is used extensively as a fungicide and herbicide1, both of which can easily contaminate water supplies. The method of using “ferric floc” to decontaminate water with copper sulfate2 has been used by oil crews since at least the 1930s, as discussed by C.J. Brockman of the University of Georgia, in the September 1, 1934 edition of Industrial Engineering & Chemistry. This method cannot be used in soil and other solid media, however, and unfortunately Copper (II) contamination is still present there. In 2004, a group of Italian researchers from the University of Calabria published their findings on the feasibility of using 2-hydroxy-5-dodecylbenzaldehyde, a ligand which has an 80% recovery rate of dissolved copper (II) ions3. There is still research going on regarding this topic, because with the growing emphasis on clean industry, waste removal becomes more and more important.Materials and MethodsThis lab was performed using the methods of vacuum filtration and weighing by difference.4 1.4862 g of Fe (Table 1) was added to a solution of 5.0030 g CuSO4 in 75.0 mL of deionized water that was previously heated on a hot plate. Once the iron powder was poured in, the solution was left to cool for about 10 minutes, and during this time it was important to remainclear of the beaker due to the possible production of sulfur fumes in side reactions. Once the vacuum filtration of the copper solid is complete, it was washed with 15.0 mL of DI water, and then four portions of 15 mL acetone washes. The copper product was then left to dry for another 10 minutes. It was important not to let the copper sit out for too long, as it would start to oxidize.ResultsTo determine the limiting reagent, it was necessary to find the theoretical yield for copperin each reaction. The theoretical yield could then be compared to the experimental yield, and then both the predominant reaction and limiting reagent could be found. If reaction (1) was favored, Fe would be the limiting reagent, as shown by:0.02661 mol Fe × 1 mol Cu1mol Fe = 0.02661 mol Cu × 63.55 g1 mol Cu = 1.6911 g Cu (3)However, if reaction (2) occurred, CuSO4 becomes the limiting reagent:0.03134 mol CuSO4 × 3 mol Cu3 mol CuSO 4 = 0.03134 mol Cu × 63.55 g1 mol Cu = 1.9917 g Cu (4)The experimental mass of copper recovered was 1.6770 g, so it is clear that reaction (1) occurredpredominantly and that Fe was the limiting reagent. This could be inferred qualitatively as well, since the solution was a blue color after the reaction reached completion. Copper sulfate is noted for being a brilliant blue in solution, so if the color remains after the reaction has reached equilibrium, then CuSO4 was not the limiting reagent. Further, it was determined that the iron sulfate product that was made was FeSO4, not Fe2(SO4)3. The crucible used in filtration had a mass of 30.8091±.0002 g (Table 1), and after filtering and drying the copper, the mass of both the crucible and copper was 32.4861±.0002 g. This then means that there were 1.6770±.0003 g of copper recovered, compared to the theoreticalvalue of 1.6911 g. The percent yield was calculated to be 99.17%.Table 1: Masses of CuSO4, Fe, Crucible, and Cu ProductCuSO4(g) Fe(g) Empty Crucible(g)Crucible + Cu(g)Cu (g)5.0030±.0002 1.4862±.0002 30.8091±.0002 32.4861±.0002 1.6770±.0003Table 2: Moles of CuSO4, Fe, and CuCuSO4 (mol) Fe (mol) Cu (mol)0.03134 0.02661 0.02639DiscussionWhen reacting copper (II) sulfate and iron, there are two possible reactions. These reactions differ in the oxidation state of iron. In the first reaction, iron has an oxidation state of +2, producing FeSO4. In the second, iron has an oxidation state of +3, and it produces Fe2(SO4)3. One of these equations will be favored based on the initial conditions of the reaction, and in order to proceed with any analysis of the reaction, this must be discerned. By