Chemical Reaction Lab Report

The wire reacted with the solution to yield copper; any remaining aluminum was dissolved with a few drops of 6M HNO3. Following the isolation of the copper, it was separated from the solution with a Büchner funnel and flask. A piece of filter paper was weighed, and its mass was recorded. Then, the filter paper was placed on the vacuum flask, which was turned on, and the copper solution was poured over it. The copper was also rinsed with 95% ethanol solution before being set on a paper towel to dry. Finally, one day after the previous step, the copper was weighed, and its mass recorded. The recorded masses were then used to determine the mass of copper, chlorine, and water in the unknown copper chloride hydrate. The resulting masses were used to find the formula of the unknown copper chloride hydrate according to CuxCly•nH2O.

Although copper was washed with water and isopropyl alcohol in order to get rid of the zinc chloride formed, some might have remained in the dish that was not visible to the eye. As a result, when zinc chloride was heated with copper in the evaporating dish and then weighed, the mass of the copper would have increased, thereby increasing the ratio of moles of copper to moles of oxygen in the empirical formula. Another error that might have occurred during the lab was the amount of hydrochloric acid added to copper oxide. Should too little be added into the beaker, the copper would not be able to completely dissolve, decreasing the amount of copper formed in the reaction. On the other hand, if too much hydrochloric acid was added, then the excess would have reacted with zinc as it was added into the reaction, producing zinc chloride. And as mentioned before, if zinc chloride was not completely removed from the dish, the mass would have increased when it was weighed. The “guess” method used in adding hydrochloric acid during the lab would undoubtedly affect the mass of copper and thus, result in differences in empirical

Procedure: First you will need to get all the lab equipment and material that is needed to do this experiment. After that fill the 150ml beaker to ¼ full of water. Than you will be provided by a teacher or an instructor who will give copper with an scoopula. Than pour the copper that was provided into the beaker, before you stir

Determine the mole ratio-the ratio of the number of moles of silver to the number of moles of copper. Note: Round the results to the nearest whole number

The purpose of this lab was to calculate the percent of water mass the copper (II)

According to the Law of Conservation of Mass, the mass of the products must equal the mass of the reactants, so logically one would expect that if 2 g of copper was used to start the lab, the lab would result in 2 g of copper. Unfortunately, that was not the case in this lab, and the final mass of copper exceeded the initial mass by 4.841g. There were many different sources of error throughout this lab, and I believe that this was the reason for such a significant difference between the initial and final masses and moles of copper that were calculated. The first step of the lab was to measure 2 g of copper and place it in a beaker. Error may have occurred at this step if the balance that was used to weight the copper was not calibrated correctly, or if the amount of

In this lab there was a performance of a single replacement chemical reaction. The reactants of the reaction was solid iron metal, and copper (II) sulfate solution. This was a gravimetric lab and therefore it was crucial to begin by massing the reactants individually. The copper (II) sulfate was initially a solid. A procedure was conducted to dissolve the copper (II) solid by heating it in water to create the necessary solution. Once that was complete the following action was to react the solid iron and the copper (II) sulfate solution. The single replacement chemical reaction took place and there was a replacement of copper with iron. The products of the reaction was solid copper metal and an iron (II) sulfate solution. The solid copper was

The main purpose in doing this experiment was to figure out how much of the percentage of Copper’s mass is in a brass alloy using a spectrophotometer that is hooked up to the laptop. Another purpose in doing this experiment is to learn how to use the Beer’s Law plots to learn how to use the equation and get the correct concentrations. This is important in everyday context because it is important for people who work doing many different types of jobs. Safety Information: In this lab, we used nitric acid to reactant with copper.

However, after decantation and heating the solid copper a mass of 0.0698 g was recovered and the percentage yield of the solid copper resulted in 117.9%. A majority of sources of error in the lab would cause loss of copper and the decrease in the final mass however in this case the final mass was higher than the initial mass. The main reason behind this could have possibly been due to impurities present in the solid copper which has altered the mass of the final product. A source of error may have been in the final reaction with zinc, if the solid zinc did not react with copper(II) sulfate completely and the left over zinc particles in the copper precipitate may have have increased the mass of the final product. Another source of error could have been not enough water being decanted from the beaker, leaving copper with liquid. Heating the copper at the end may have not dried off enough to eliminate all the the liquid which may have added more weight to the solid copper being

First we have to find the Balanced Equation: Fe2O3 + 3CO yields 2Fe + 3CO2 Then you would need to find your number of moles for Fe, which is 0.07890 Then you would need to find your number of moles for CO, which is 0.34451 So that also helps us because it tells us that 1 singluar mole Fe behaves with 3 mole CO. 0.07890^3 mole of Fe which gives us a total of 0.2367 mole 1 singular mole of Fe behaves with 2 mole of Iron, you would then do 0.07890^2 to get: 0.1578 mol of Iron. Then you would take 55.845 (iron's singular mole) and multiply it by the 0.1578 and that would be the theoretical yield which is 8.812341

In order to calculate the concentration of iron in an unknown substance, many steps had to be taken. First, a stock iron solution was prepared by measuring exactly 0.0135g of ferrous ammonium sulfate [Fe(NH4)2(SO4)2·6H2O], mixing it with approximately 5.2mL of 2.0 M H2SO4, and diluting the solution to 200mL. Following this, a diluted 1,10-phenanthroline solution was prepared by obtaining approximately 50.0mL of 1,10-phenanthroline monohydrate in ethanol solution and diluting it to 200mL with deionized (DI) water. Then 9 200-mL volumetric flasks were cleaned and labeled A-I. Flasks B-F were filled with exactly 2.00, 5.00, 10.00, 25.00, and 50.00 mL of the stock iron solution respectively. Flask A was used as a control and was instead filled with 100mL of DI water. Flasks G,H, and I were filled with exactly 25.00mL of Unknown Iron Solution E0717. Then all flasks A-I received approximately 2.0mL of hydroxylammonium chloride solution,

The change in colour of the copper wire from bronze to black signifies that the composition of the copper wire have changed during the heating process. The copper wire, Cu had reacted with oxygen, O2 during heating and forms a new substance that is copper oxide, CuO. As the composition of the copper changes with the presence of oxygen, the chemical properties have also changed and thus this causes the resulting substance, copper oxide, CuO to differ physically and chemically with the original substance, Copper, C. Thus, it can be concluded that the heating of copper wire causes a chemical change to occur. Meanwhile, the change in physical state and colour of the iodine crystals, I2 during heating does not signify that a chemical change has taken place.

In relation to this experiment, the per cent recovery, or “how much of the original pure copper mass recovered at the end of the experiment”ii iii iv v should be 100% because all of the copper should return to elemental form. However, the per cent recovery from this experiment was calculated to be 56.5% recovery, much lower than the anticipated value, and therefore does not appear to support the Law of the Conservation of Mass. One explanation for the per cent recovery being too low would be that some copper was lost while decanting the solution after reaction three. The time constraints also should be taken into consideration, as if the reaction resulting in the precipitation of copper by the aluminium wire was allowed to run longer, more of the copper could have been precipitated out of solution, closing the gap toward the 100% recovery. The product of the reaction was a dark reddish brown colour.

Millimoles of Fe (II) in each solution = (10.00 mL) x (MB). = 10.00 x 0.000694 = 0.00694 millimoles. Mole ratio ‘n’ = (millimoles of ortho-phenanthroline at intersection)/(millimoles Fe^(2+) in each solution ) =0.0196/(0.00694 )=

In “Part 3: Preparing Copper (II) solutions of Different Concentrations”, 20.0 mL of a 0.500 M solution was from CuSO4 * 5H2 O was made. The solid was weighted and was added to the 100 mL beaker. About 15 mL of water was added to the same beaker. The mixture was transferred back and forth between the graduated cylinder and a beaker to dissolve all the crystals. The solution was transferred back to the graduated cylinder. The beaker was rinsed well with a small amount of distilled water and the rinsed solution was transferred to the same cylinder containing the dissolved solid. After adding the solution, the total volume in the cylinder was about less than 20 mL. With the use of the Pasteur pipette, the water was added until the volume in the cylinder was 20 mL. This was the first solution.