Paula MoltzanSean Schaffer11-8-17Bryan Novas (L-42)Chemical Reactions of CopperIntroduction:The purpose of this lab was to complete chemical reactions showing the properties of copper and its compounds. Also, this experiment tested the law of mass conservation and our laboratory efficiency. Chemical reactions occur every second of our lives, some are more visible around us than others. Reaction may be classified by the number, physical or chemical nature theof reactants and products in the reaction, and by the arrangement of atoms in the conversion. Copper can be combined into an assortment of compound, but is usually found in the form of a sulfide. With specific reaction conditions copper can be chemically changed from one from to another. The first step was to obtain a sample of Cu that was 0.025-0.050 grams in a test tube, then find the mass of the test tube and the mass of the copper sample. Under the fume hood 10 drops of nitric acid was added to the test tube to dissolve the copper. Then, 20 drops of distilled water were added. We then returned to lab table and continuously stirred the test tube while adding sodium hydroxide until precipitate formed. Then we centrifuged the test tube for 1 minute, this separated precipitate and the supernatant. Then we put 125-mL of water in to a 250-mL beaker as well as some boiling chips. The beaker was placed on the heating plate and broughtto a boil. Once the precipitate changed colors it was removed from the heat and we let it cool. Weobtained a porcelain evaporating dish and weighed it on the scale. Then we added sulfuric acid tothe test tube until the precipitate dissolved, and transferred all the liquid into the weighed porcelain dish. Then we added a piece of magnesium turning to the solution in the evaporating dish and crushed it up with the stir stick until the reaction stopped and repeated until the solution was no longer blue. Then we added 10 drops of 6 M HCl to dissolve any remaining pieces of Mg. Next, we let the copper metal settle to the bottom of the evaporating dish, then poured the solution into a waste beaker. We then washed the copper metal with 5 mL of distilled water, allowed the metal to settle and poured the solution into the waste beaker, then repeated. Then we washed the copper metal with 5 mL of 95% ethanol and poured out the solution and repeated. Finally, we washed the copper metal with 5 mL of acetone, removed the excess solution into the waste beaker once it's settled and repeated. Then we placed the evaporating dish on top of the 250-mL beaker with the boiling chip from part c. Then we heated the water in the beaker to dry the copper metal for 5-10 minutes. When the copper looked dry, we picked up the evaporating dish with tongs and set it on the lab bench to cool. When it cooled, we wiped the bottom to remove any water and let it dry. Once the evaporating dish reached room temperature, it was weighed. Making sure both were at room temp for this last step. Safety glasses and gloves were required during this experiment. We made sure to do the first few steps described above in the hood due to the gas. Ethanol and acetone are flammable and 6M HNO3, 6M NaOH, 6M H2SO4 and 6M HCl solutions are all corrosive, so handle with care. Any remaining decanted solutions or water rinsing, as well as both the ethanol and acetone were disposed of in the correctly labeledliquid waste container. Also, any used magnesium and recovered copper was disposed of in the solid waste container.Results and Discussion: In this experiment, we started with 0.0268 grams of copper and through a series of redox reactions 0.024 grams of copper was produced. In the first step of the experiment we used nitric acid to dissolve the copper sample which caused the test tube to heat up and the copper fizzed and popped as it dissolved, the solution turned a bright turquoise color. Then water was added to the solution which diluted the color of the solution. This solution was then reacted with sodium hydroxide which formed a precipitate, which we then centrifuged to form the product copper hydroxide. After this it was heated in a hot water bath to slowly change the color of the precipitate from green blue to black, this formed the product copper oxide. Then sulfuric acid was added to the solution to dissolve the precipitate and turning the solution a blue color which formed Copper (II) ions. The magnesium was then added to the solution and stirred and crushed for a long time, allowing it to completely react with the solution. Once the solution had turned a grey color we removed any excess magnesium turnings, and allowed for the copper to settle at the bottom of the evaporating dish. We then drained the solution out only leaving the copper in the evaporating dish, we rinsed and dried the copper multiple times making sure not to lose any solid during the process. Finally, we dried the copper over the boiling water to evaporate any remaining liquid that was present in the solid. Once the evaporating dish had cooled we weighed the dish with the copper in it to find the mass of the recovered copper. This was done by subtracting the final mass of the evaporating dish and recovered copper from the mass of the empty evaporating dish (which was taken earlier in the experiment). Once the mass of the recovered copper was determined we were able to calculate the percent yield by using the formula below.Percent Yield =Actual YieldTheorectical YieldX 100The percent yield of this experiment was 89.55% this is a decent result considering it wasthe first time performing this type of experiment. A couple sources of error that could have resulted in a lowered recovered mass than originally used are in one; in step E when we were adding the magnesium turnings we could have potentially not reacted enough magnesium or when removing the excess pieces of magnesium turnings from the solution some cooper may have been attached to them. Therefore, by removing the excess magnesium it is possible to have lost small pieces of reacted copper as well. Another step in which copper could have been lost was in step F when rinsing and decanting the copper small bits of the cooper could have slipped through which would also result in a smaller recovered mass of copper then originally used. Lastly, when drying the copper in the evaporating dish we used a stir stick to check if the copper was completely dry and super tiny bits of copper would stick to the stir stick, this would also