DOWNLOAD PDF CYCLE OF COPPER REACTIONS

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1 Chapter 1 : Copperâ chlorine cycle - Wikipedia CYCLE OF COPPER REACTIONS. PURPOSE: The goal of the experiment is to observe a series of reactions involving copper that form a cycle and calculate the percent recovery of the initial copper mass. INTRODUCTION: A schematic of the cycle of copper reactions is shown below. Illustrate variety of substances of which an element can be part: Conservation of mass and of moles: We should recover as much copper as we started with. Same amount of copper present at every stage: Experience in standard chemical techniques: Actually, the nitrate ion oxidizes the copper metal to copper II ion while itself being transformed to NO2 gas in the process; the copper II ion then binds to six water molecules. As a result, hydroxide ion can displace water from the copper II ion, yielding copper hydroxide, Cu OH 2, a blue precipitate. We have seen this reaction before in the copper chloride lab. The pieces of wire are closer to 0. Use about ml of concentrated HNO3 solution. Be careful with the nitric acid: If some copper remains undissolved by the time the production of gas is finished, then put the beaker on the hot plates in the hoods to hasten the reaction. It is important to carry out this step in the fume hood because the brown NO2 gas is an irritant. Keep the mixtures in the hood until after you add the 10 ml of distilled water after completely dissolving the copper. Add the NaOH solution dropwise to the copper solution. After a blue precipitate is formed, periodically test the acidity of the solution by dipping your stirring rod into the solution and touching it to red litmus paper. Try not to transfer the blue precipitate onto the litmus paper: The solution starts out acidic because of excess nitric acid from the previous step, so the first OH- added goes into neutralizing the acid; once the acid is neutralized, the next OH- added goes to forming the blue Cu OH 2 precipitate. Only after that is finished does added OH- hang around idle, and only at that time will it turn red litmus paper blue. We want to make sure all the copper present is turned to Cu OH 2, so we add OH- until the solution turns the litmus paper blue. Transform Cu OH 2 s to CuO s Add water to the reaction mixture obtained in the previous step, and add one or two boiling stones as well. Heat the contents of the beaker, but do not boil. Boiling makes the black CuO so fine that the filtration step is excessively long. Filter and wash the CuO as described in the procedure part C. Keep the solid on the filter paper, and discard the filtrate. Sulfuric acid solution is corrosive and will sting skin with which it comes into contact. Wash the copper metal three times with distilled water and transfer it to an evaporating dish as described in the procedure part E, and then wash it three times with 5-ml portions of isopropanol. Washing with isopropanol will reduce the time needed for the drying step. Dry the copper over a beaker of boiling water as described in the procedure E. Weigh the dry copper and record the mass. Compute the percent of copper recovered. Page 1

2 Chapter 2 : Cu Again! - A Copper Cycle Cycle of Copper Reactions Minneapolis Community and Technical College v Introduction To a beginning student of chemistry, one of the most fascinating aspects of the laboratory is the dazzling array of sights. General Chemistry Lab 1 Conservation of Mass: A Cycle of Copper Reactions Purpose. The goal of this experiment is to introduce you to several classes of chemical reactions: The reactions will be performed in a specific order and, when completed, will regenerate the starting material, pure copper. Since you will not add copper to the reaction at any point after the initial step, it should be possible to recover all of the copper at the end of the cycle, thereby illustrating the Law of Conservation of Mass. You are asked to practice your observational skills by recording the changes that occur at each step of the cycle. Color changes, the generation of gases or precipitates and changes in temperature are all examples of the types of phenomena you should be looking for when studying chemical reactions. Thousands of individual chemical reactions can be classified into a very small number of reaction types. In this experiment, several different types of reactions will be performed to transform copper into a variety of its compounds. It is important to realize that each product contains copper and that the total number of copper atoms involved in each step is the same. Therefore, at the end of the cycle, the mass of copper recovered should equal the mass that was originally used. You will calculate the percent yield of copper at the end of the cycle to determine the "efficiency" of the process. The cycle of copper reactions to be performed in this experiment The cycle of reactions to be performed is shown in Figure 1. Beginning with pure copper at the top of the figure, these are: Oxidation of metallic copper with nitric acid HNO3. The balanced equation is: It is formed in automobile engines and is one of the major contributors to urban air pollution. If you have ever witnessed the brown haze in Los Angeles or sometimes even in Seattle, the color is due to nitrogen dioxide. Precipitation of copper II ion as copper II hydroxide. In this reaction a solution of sodium hydroxide is added to the solution of copper nitrate and will cause the precipitation of the insoluble copper II hydroxide. Thermal decomposition of copper II hydroxide to copper II oxide. Compounds that are stable under standard conditions often become unstable at elevated temperatures. Many times this can result in the loss of gases, for example, the loss of carbon dioxide from calcium carbonate: Reaction of copper II oxide with sulfuric acid. Reduction of copper II with zinc. To accomplish this, the mixture is treated with hydrochloric acid. At this point the solid, pure copper can be isolated, dried and weighed. Toxic NO2 is produced. Keep your apparatus away from open flames. Get a piece of copper foil and cut an approximately 0. If it is not bright and shiny, clean it with a piece of steel wool, rinse with water and dry with a paper towel. Then get an accurate mass measurement using an analytical balance. Place in a ml beaker and add about 4 ml of concentrated nitric acid slowly and carefully. Record in your notebook a description of what you see. After the copper has dissolved, add 10 ml of deionized water to dilute the sample for step 2. Record your observations in your notebook. Dilute the solution with deionized water to about ml in preparation for step 3. Add a magnetic stir bar and place on a heatable stir plate. Boil gently while stirring for about 4 minutes. Record any changes that occur in your notebook. Remove the beaker from the heat and allow to cool. Add 40 ml of deionized water into a second clean beaker and begin heating. Prepare a filter paper and funnel to filter the copper II oxide. Use a ml beaker to collect the filtrate, supporting the funnel with a funnel support or iron ring on a ringstand. Filter the copper II oxide. The filtrate should be colorless and free of any solids. Transfer the last traces of solid material from the beaker using a stream of deionized water from a wash bottle. Use the deionized water that you have been heating to wash the solid collected in the filter paper. Pour about 5 ml of the hot water through the filter. Leave the filtration apparatus in place for the next step. In your notebook, describe the appearance of the collected solid. Dissolve the CuO by carefully pouring about 15 ml of 3 M sulfuric acid directly through the residue on the filter into a ml beaker. Record any changes that you see. If the solid is not completely dissolved the first time, replace the collection beaker with a clean new one and pour the acid in the first beaker through the filter again. Pour very carefully so as not to lose any of the liquid. Repeat this procedure as often as necessary to dissolve all of the solid. It may take four or five times. Once the solid is dissolved, you need to collect all of the copper II Page 2

3 containing solutions in the same beaker. Rinse down the walls of the collection beaker that is not in position below the filter with deionized water from a wash bottle and pour the rinse water through the filter into the other beaker. Wash the empty filter paper with three or four 5 ml portions of cold deionized water and collect the washings in the beaker containing the acid solution. Add about 2 g of zinc metal to the copper II solution and stir rapidly. The reaction between the zinc and copper II ion will be complete when the blue color of the copper solution is gone. If any blue color remains after the zinc has been consumed, add approximately 0. Record this in your notebook. When the copper has been completely reduced, decant most of the solution, even if the zinc is still generating gas. Add 25 ml of 3 M HCl to speed up the rate of zinc oxidation. When no more bubbles are seen, proceed to the epilogue. Allow the copper metal to settle to the bottom of the beaker. Carefully decant the supernatant. Wash the copper metal precipitate with three 50 ml portions of water, removing each portion by decantation. Weigh a clean, dry evaporating dish and record the mass. Let the solid settle in the evaporating dish and carefully decant most of the liquid. Remove from the oven and allow to cool to room temperature place on a paper towel with your name on it. Weigh the evaporating dish and calculate the mass of copper recovered and your percent yield using the following equation: Page 3

4 Chapter 3 : Chemistry Lab Report (Copper Cycle) â Sarah Jackson Dr. Caddell A Cycle of Copper Reactions Chemistry Weigh out about 3 grams of zinc metal. Using the shears provided cut the zinc into thin strips. Add this zinc to your ml beaker. Place a magnetic stir bar in the beaker and place on a magnetic stirrer. The Nature of a Chemical Reaction Introduction Chemical reactions are often accompanied by formation of a precipitate, evolution of gas, change in color, or pronounced temperature change. In this activity, you will observe these characteristics of chemical reactions. Go to Top Purpose To recognize that change of state, change in color, formation of a precipitate, or the evolution of heat are associated with a chemical change; to study reactions of copper. Go to Top Safety Considerations Wear protective glasses and an apron at all times. Avoid skin contact with solids and solutions. Dispose of all solutions in the containers provided by your teacher. Wash your hands before leaving the laboratory. Procedure First, prepare a data table. Record what you see throughout the laboratory activity. Your teacher will demonstrate for the class the reaction of metallic copper with concentrated nitric acid, HNO3. This must be done in the fume hood. Avoid breathing poisonous gases. Avoid contact of skin with nitric acid. It burns skin and clothing. One student will be asked to feel the side of the demonstration beaker and report to the group. A solution of copper II nitrate was prepared earlier. One of your thin-stem pipets contains this solution. Transfer all of the solution to your labeled test tube. Another thin stem pipet contains sodium hydroxide solution, NaOH aq. Avoid contact with sodium hydroxide. Add this slowly to the test tube. This reaction may give off heat. In other words, it may be exothermic. Tap the tube firmly to mix or use a stirring rod. Tap or stir to mix. Place your test tube in a hot water bath. Remove the test tube when you see no further change occurring. This product is copper II oxide, CuO s. Run cold water over the outside of the test tube to cool it. Allow the material to settle or use a centrifuge to spin the solution. Decant or use the pipet labeled "waste" to remove the clear liquid supernatant liquid above the copper II oxide. Do not remove any of the solid. Discard the liquid and wash precipitate by adding about half a pipet of distilled water. Allow the solid to settle. Then remove and discard the liquid, or centrifuge and decant. To the precipitate, add hydrochloric acid, HCl, from the final thin stem pipet. Avoid contact of both skin and clothing with hydrochloric acid. This new product is copper II chloride solution, CuCl2 aq. Add the precut piece of aluminum wire to the test tube. The wire should be bent like a fish hook to hook over the top of the test tube. Place the test tube in a cold water bath. This reaction is very fast! Instead, your teacher may tell you to run cold water over the outside of the test tube. Be sure to use a test tube holder to hold the test tube since the reaction gives off considerable heat. The products forming in this step are hydrogen gas, H2 g, aluminum chloride solution, AlCl3 aq, and metallic copper, Cu s. When the reaction is finished, remove and discard the liquid. Wash the solid with half a pipet of distilled water. Wash the solid again with half a pipet of fresh distilled water, and filter. Use tweezers to remove any leftover pieces of aluminum wire. Observe the final product. Take the filter paper with your product to the fume hood. Your teacher will test a small sample of your product with concentrated nitric acid. Clean pipets, test tube, and funnel. Wash hands thoroughly before leaving the laboratory. Go to Top Data Analysis, Concept Use your observations to complete a drawing similar to the one above. Alongside each arrow, write in the chemicals used. Also near each arrow, write in key words to convey what you saw. As chemical changes occurred, what observations did you note? List the observations that alerted you that some type of chemical change was occurring. Did each reaction show all of the kinds of changes you listed above? Use your observations to justify your answer. What happened to the hydrogen gas generated in the last reaction? Add hydrogen gas to the cycle. What happened to the aluminum chloride? Show this on the cycle. What did you enjoy most about this activity? Go to Top Imply, Apply Reread the title of this laboratory activity. Why is this series of reactions often called the "copper cycle"? How did the last part of the procedure Step 16 complete the cycle? Recycling is one important way to conserve precious natural resources. What common metals have you observed being collected for recycling? Think about all types of materials that you know are recycled. Compose a list of these materials. What types of materials might prove too difficult or costly to recycle? Page 4

5 Chapter 4 : What is the copper cycle and why is it important? Yahoo Answers Types of Reaction: Copper Cycle, Teacher's Guide 2 Types of Reactions: The Copper Cycle In this laboratory experiment, students will perform a series of reactions known as the copper cycle. Exothermic and produced nitrogen dioxide gas Exothermic, precipitate formed, solution became thicker, and gas was produced. Black Copper Oxide precipitate formed and water formed. Black precipitate is denser than water and sinks to the bottom of the beaker. Copper Sulfate and water formed. Zinc sulfate is formed and is aqueous. Hydrogen gas is also formed and released. Solid copper is found on the bottom of the beaker. Due to the Law of Conservation of Mass, one would predict that after any number of chemical reactions and phase changes, the final mass and number of moles of copper would remain the same as the initial amounts. After the lab was finished, the final mass of the copper was 6. Clearly these values are not the same as hypothesized, due to a variety of sources of error throughout the lab. The percent yield of copper calculated was Discussion of Theory The main chemical concept that this lab involves is the Law of Conservation of Mass. This states that in a chemical reaction, matter can neither be created or destroyed, or more simply, the mass of the products must equal the mass of the reactants. In this lab, we tested this by putting copper through five different reactions and many phase changes. We started our experiment with 2. In our lab, this was not the case, and we ended up with approximately three and one half times the amount of copper we started with, According to the law previously stated, this is impossible. So how did this happen? After going back and analyzing the reactions in the lab, we concluded that he majority of the sources of error in the lab would cause a loss of copper and a decrease in the final mass. So then how could the final mass possibly be larger than the initial mass? After more thought, we came to the conclusion that the final amount of copper that was weighed to calculate the final mass contains impurities introduced by the reactions in the lab, and these impurities altered the mass of the final product. In the final reaction, solid zinc was added to aqueous copper sulfate. If there was any solid zinc that did not react with the copper sulfate and was left over with the copper precipitate, zinc particles may have significantly changed the mass of the final product. As a result, the Law of Conservation of Mass was still held true, and the same amount of copper that was used initially was also found in the final product of the reaction. Sources of Error 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. 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 copper weighed was not exactly the 2 g required. Even a small deviation, such as that of 0. There were many possibly sources of error in the lab that could have been due to incorrect measurement from systemic or random error. Throughout the entirety of the lab, there were many steps that presented possible sources of error such as adding too much or too little of a compound to the copper solution, or the loss of copper during transport or by being left on the stirring rod. Another situation which was a source of error was during step 5 when the copper solution and water was heated so that the contents of the beaker would boil. The error in this step was caused by the variable heating of the solution to prevent bubbling of the solution. The length of heating was another source of error because the mixture may have needed to be heated less or more than the 5 minutes stated in the direction to get to complete the step. During step 5, stirring of the mixture was required, so another source of error was that copper from the beaker may have been left on the stirring rod. During the next step, error may have occurred if there was copper precipitate left on the side of the beaker instead of washed with water. Decanting also proposed a source of error because copper may have been accidentally lost, or not enough water may have been decanted from the beaker. When the final mass of copper was weighed, the balance may have not been calibrated correctly, and the copper in the evaporating dish may have contained impurities, such as un-reacted zinc, that changed the Page 5

6 mass that was measured. All in all, there was an abundance of sources of error in this lab. The reaction between copper and nitric acid is an exothermic reaction. The function of using the ice bath in this step is to control the temperature of the reaction, because the reaction may become too violent if nothing is used to cool it down. Ice water is used as an absorber for the heat released in this reaction. A double replacement reaction took place in step 4 between the copper nitrate and sodium hydroxide. The reaction that occurred in step 7 was a dehydration reaction. The reaction that occurred in step 9 was a single replacement reaction. The reaction between the excess zinc and sulfuric acid in step 9 is essential to ensure that there is so un-reacted solid zinc mixed in with the copper, and that any excess zinc is reacted to form zinc sulfate than is aqueous and can be poured out with water. If this was an incomplete reaction, there would be leftover solid zinc that would be weighed with the copper precipitate, and the zinc would increase the measured final mass of the copper and cause a significant difference between the initial and final masses of the copper. The reaction between zinc and sulfuric acid is as follows: When the CuO was washed, the excess hydroxide ions that could have remained from the previous reaction were removed. In the final step of the lab when the copper precipitate was washed, zinc ions were removed. The previous reaction that took place involved aqueous copper ii sulfate and solid zinc. The products of the reaction were solid copper and aqueous zinc sulfate. After finding the net ionic equation of the reaction, one determined that the products of the reaction were solid copper and aqueous zinc ions. After this knowledge was obtained, one can conclude that the zinc ions were washed out with the water during the final step. Powered by Create your own unique website with customizable templates. Chapter 5 : Transformation of Copper: A Sequence of Chemical Reactions of reactions and then recovered as copper. After which, relevant errors will be identified and explained. 2) Hypothesis: a) Reaction A: If copper metal is added to nitric acid, then a single displacement reaction might occur because the copper the nitric acid will displace the hydrogen while creating oxygen. Chapter 6 : Copper Lab - AP Chemistry Lab Reports In an oxidation/reduction reaction, electrons are transferred from one reactant to the other. In the simplest form of these reactions, single-displacement reactions (also called single-replacement reactions), metal ions react with pure metals. If the reaction proceeds, the pure metal gives electrons to the metal cation. Chapter 7 : General Chemistry - Summer - Lab 1 Dissolving a Solid to Obtain a Soluble Copper Product If the reaction starts with a solid copper compound and produces a soluble copper product, then you want to completely react the solid copper reactant. A good clue that the solid has completely reacted is the disappearance of that solid. Page 6