Beginning of the lab
The copper reacting with the nitric acid.
Final product
At the end of the experiment. The copper weighed 3.85g
Purpose
The purpose of the lab is to see what percentage of copper can be recovered after performing various reactions to a sample of copper with a known mass. The Law of Conservation of mass is being tested. The experimenter is trying to get a 100% for the percent yield.
Qualitative data
ReactionCu + HNO3
Cu(NO3)2 + NaOH Cu(OH)2 CuO + H2SO4 CuSO4 + Zn Zn + H2SO4 |
ObservationsOrange gas forming, the copper is dissolving, and the solution is turning blue
A precipitate is forming and gas is being released The precipitate is black fizzing, turning green/clear, the beaker is warm, turning teal/blue fizzing, cloudy, smells, gas being released, beaker is warm, burgundy/orange precipitate. the beaker is warm, murky, small orange pellets |
Quantitative data
Initial mass of copper
Mass of evaporating dish Mass of evaporating dish and copper Mass of recovered copper |
2g
54.1g 57.95g 3.85g |
Net ionic equations
Calculations
Final mass of copper
Final mass of copper = Mass of evaporating dish and copper (57.95g) - Mass of evaporating dish = 3.85g
Final mass of copper = 3.85g
Final moles of copper
1 mol of Cu
3.85g of Cu X -------------- = .061 moles of Cu
63.55g of Cu
Initial moles of copper
1 mol of Cu
2g of Cu X ------------- = .031 moles of Cu
63.55g of Cu
Percent yield
3.85g of Cu
------------- X 100% = 193%
2g of Cu
Final mass of copper = Mass of evaporating dish and copper (57.95g) - Mass of evaporating dish = 3.85g
Final mass of copper = 3.85g
Final moles of copper
1 mol of Cu
3.85g of Cu X -------------- = .061 moles of Cu
63.55g of Cu
Initial moles of copper
1 mol of Cu
2g of Cu X ------------- = .031 moles of Cu
63.55g of Cu
Percent yield
3.85g of Cu
------------- X 100% = 193%
2g of Cu
Conclusion
The final mass of copper was 3.85 grams. There was an excess of copper at the end of the experiment due to various sources of error. The initial amount of copper was 2g. 1.85 grams of copper and other substances added mass to the initial amount. The copper reacted with nitric acid to obtain copper nitrate. Sodium hydroxide was added to the copper nitrate to obtain copper hydroxide. The copper oxide was set aside to settle and the water was decanted from the precipitate. Sulfuric acid was added to the copper oxide to form copper sulfate. There was an excess of zinc in the solution so am additional 50mL of 3.0M sulfuric acid was added the solution. The 2g of copper was used to create various reactions and by the end of the experiment we ended with 3.85g.
Discussion of Theory
The lab was conducted to test the Law of Conservation of Mass. Various reactions were performed on 2g of copper to see if the experimenters ended with 2g of copper, as stated by the Law of Conservation of Mass. However, at the end of the lab the final amount of copper weighed was 3.85 grams. There were many possible areas for sources of error to have occurred. Stoichiometry was used to calculate the number of moles yielded from the experiment. The percent yield of copper was 193%.
At the beginning of the experiment copper started in the solid state. Nitric acid was added to copper to form a blue, aqueous solution. The reaction produced copper nitrate, nitrogen dioxide, and water. Nitric acid splits into ions because it is a strong acid. Strong acids complete ionize and therefore we get hydrogen and nitrate ions. By the end of the reaction copper was not in a solid state, it was in the form of copper ions in the solution mixture. The products were placed in an ice bath to keep the temperature of the beaker controlled because it is an exothermic reaction. The cooler temperature of the water will absorb more heat than the room temperature water. Exothermic reactions are favored in nature because they have a lower activation energy than endothermic reactions. The reaction released heat because the potential energy stored in chemical bonds is converted to thermal energy.
Then sodium hydroxide was added to the copper nitrate to produce a solid (copper (II) hydroxide) and sodium nitrate. The solid was copper hydroxide. The copper hydroxide was left to settle so that we could decant the excess H2O produced, sodium, nitrogen, and oxygen ions away from the precipitate. Decanting is the process of separating the solid and the liquid a solution. Majority of the aqueous solution was removed and sulfuric acid was added to the copper oxide. The reaction produced copper(II) sulfate and water. The solid precipitate of copper oxide is now in an aqueous state. The copper ions are floating in the solution along with water and sulfate ions. Zinc was added to the copper sulfate to produce zinc sulfate and copper. Copper changed from its aqueous state to a solid state. The excess zinc was removed by adding more sulfuric acid to react with the zinc. The higher mass of copper at the end of the lab was primarily due to the incomplete reaction of zinc and sulfuric acid. The solid zinc that did not react with sulfuric acid was included in the final mass of copper.
In Step 4, a double replacement reaction occurred that formed copper hydroxide (solid). This was a precipitation reaction because a precipitate formed. Step 7 was a double replacement reaction that resulted in an aqueous solution of copper sulfate and water. This was a acid-base reaction. The sulfuric acid is the acid in the reaction and the copper (II) oxide is the base. In Step 9, a single replacement reaction took place. This was a reduction reaction. The sulfuric acid was added to zinc to produce zinc sulfate and hydrogen gas.
At the beginning of the experiment copper started in the solid state. Nitric acid was added to copper to form a blue, aqueous solution. The reaction produced copper nitrate, nitrogen dioxide, and water. Nitric acid splits into ions because it is a strong acid. Strong acids complete ionize and therefore we get hydrogen and nitrate ions. By the end of the reaction copper was not in a solid state, it was in the form of copper ions in the solution mixture. The products were placed in an ice bath to keep the temperature of the beaker controlled because it is an exothermic reaction. The cooler temperature of the water will absorb more heat than the room temperature water. Exothermic reactions are favored in nature because they have a lower activation energy than endothermic reactions. The reaction released heat because the potential energy stored in chemical bonds is converted to thermal energy.
Then sodium hydroxide was added to the copper nitrate to produce a solid (copper (II) hydroxide) and sodium nitrate. The solid was copper hydroxide. The copper hydroxide was left to settle so that we could decant the excess H2O produced, sodium, nitrogen, and oxygen ions away from the precipitate. Decanting is the process of separating the solid and the liquid a solution. Majority of the aqueous solution was removed and sulfuric acid was added to the copper oxide. The reaction produced copper(II) sulfate and water. The solid precipitate of copper oxide is now in an aqueous state. The copper ions are floating in the solution along with water and sulfate ions. Zinc was added to the copper sulfate to produce zinc sulfate and copper. Copper changed from its aqueous state to a solid state. The excess zinc was removed by adding more sulfuric acid to react with the zinc. The higher mass of copper at the end of the lab was primarily due to the incomplete reaction of zinc and sulfuric acid. The solid zinc that did not react with sulfuric acid was included in the final mass of copper.
In Step 4, a double replacement reaction occurred that formed copper hydroxide (solid). This was a precipitation reaction because a precipitate formed. Step 7 was a double replacement reaction that resulted in an aqueous solution of copper sulfate and water. This was a acid-base reaction. The sulfuric acid is the acid in the reaction and the copper (II) oxide is the base. In Step 9, a single replacement reaction took place. This was a reduction reaction. The sulfuric acid was added to zinc to produce zinc sulfate and hydrogen gas.
Source of Error
At the end of the experiment, there was 1.85g more copper than the initial amount. There was an excessive amount of copper at the end of the experiment possibly because of contamination, incomplete reactions, and loss of reactants. If contamination occurred not the reactants may not be able to completely react to form the full amount of products. If all of the reactants did not completely react then the amount of products are less than the estimated amount and therefore, the proceeding reactions will be short of reactants because the products from the previous reaction are the reactants in the next reactions. In step 5 the solution did not boil properly and had to be left over night for the copper oxide to settle. The copper hydroxide could have not completely settled and we decanted the copper oxide that was in the solution. In step 7 not enough of the coper oxide could have been in the reaction to react with the sulfuric acid. In the reaction did not completely complete , in step 9, not enough sulfuric acid could have been added to react with the zinc and the zinc could have been weigh with the copper.
Questions
1. Why is the product of the reaction between copper and nitric acid in step 2 placed on ice?
The product of the reaction between copper and nitric acid in step 2 was placed on ice to keep the temperature of the reaction down. The reaction was an exothermic reaction and therefore, released heat. The ice bath will absorb more heat than the room temperature water, which will allow for it to keep the reaction temperature controlled.
2. What type of reaction (synthesis, decomposition, single rep. double rep., redox, or dehydration) occurred in steps 4, 7, & 9?
Step 4 was a double replacement reaction that formed copper hydroxide (solid). Step 7 was a double replacement reaction that resulted in an aqueous solution of copper sulfate and water. This was a acid-base reaction. Step 9 was a single replacement reaction. An additional 50mL of sulfuric acid was added to react with the excess zinc. The reaction resulted in zinc sulfate and hydrogen gas.
3. The reaction of excess zinc with sulfuric acid is critical step in this investigation. Write the balanced equation for this reaction. What problems would arise from an incomplete reaction?
The balanced equation for this reaction is Zn (s) + H2SO4 (aq) ---> ZnSO4 (aq) + H2 (g).
An incomplete reaction would result in a higher mass than the initial amount because the zinc that did not completely react will be weighed with the copper.
4. What ions did you remove when you washed the CuO?
When the copper oxide was washed nitrogen, oxygen, and sodium ions were removed.
5. What form of copper is present in the beaker after you added the H2SO4?
The aqueous form of copper is present in the beaker after sulfuric acid was added. The copper oxide and sulfuric acid produce copper sulfate, which is aqueous. The copper, sulfur, and oxygen ions are floating in the solution.
6. What ions did you remove when you washed the precipitated copper?
When the precipitated copper was washed zinc and sulfate ions were removed.
The product of the reaction between copper and nitric acid in step 2 was placed on ice to keep the temperature of the reaction down. The reaction was an exothermic reaction and therefore, released heat. The ice bath will absorb more heat than the room temperature water, which will allow for it to keep the reaction temperature controlled.
2. What type of reaction (synthesis, decomposition, single rep. double rep., redox, or dehydration) occurred in steps 4, 7, & 9?
Step 4 was a double replacement reaction that formed copper hydroxide (solid). Step 7 was a double replacement reaction that resulted in an aqueous solution of copper sulfate and water. This was a acid-base reaction. Step 9 was a single replacement reaction. An additional 50mL of sulfuric acid was added to react with the excess zinc. The reaction resulted in zinc sulfate and hydrogen gas.
3. The reaction of excess zinc with sulfuric acid is critical step in this investigation. Write the balanced equation for this reaction. What problems would arise from an incomplete reaction?
The balanced equation for this reaction is Zn (s) + H2SO4 (aq) ---> ZnSO4 (aq) + H2 (g).
An incomplete reaction would result in a higher mass than the initial amount because the zinc that did not completely react will be weighed with the copper.
4. What ions did you remove when you washed the CuO?
When the copper oxide was washed nitrogen, oxygen, and sodium ions were removed.
5. What form of copper is present in the beaker after you added the H2SO4?
The aqueous form of copper is present in the beaker after sulfuric acid was added. The copper oxide and sulfuric acid produce copper sulfate, which is aqueous. The copper, sulfur, and oxygen ions are floating in the solution.
6. What ions did you remove when you washed the precipitated copper?
When the precipitated copper was washed zinc and sulfate ions were removed.