- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 1 (11.81 KiB) Viewed 61 times
- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 2 (42.65 KiB) Viewed 61 times
- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 3 (38.99 KiB) Viewed 61 times
- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 4 (34.78 KiB) Viewed 61 times
- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 5 (36.74 KiB) Viewed 61 times
- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 6 (34.62 KiB) Viewed 61 times
- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 7 (73.51 KiB) Viewed 61 times
- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 8 (73.51 KiB) Viewed 61 times
- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 9 (94.14 KiB) Viewed 61 times
- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 10 (101.88 KiB) Viewed 61 times
- Assume You Started The Experiment With 0 3452 G Of Copper Wire After All Is Completed And The Copper Is Dried You Have 11 (68.09 KiB) Viewed 61 times
REPORT FOR COPPER CHEMISTRY AND REDOX EXPERIMENT A. Copper Chemistry Orignal Mass of Copper, (g) Mass of Evaporating dish + recovered Copper, (g) Mass of Evaporating dish, (g) Mass of Recovered Copper, (g) Percent Recovery, (%) Observation For each step of the procedure in which a chemical reaction occurred, record your observations. Did a gas or precipitate form? Was there a color change? For each step involving a redox reaction, write a net ionic equation that has been balanced using the Va reaction method. Then write the lonie equation and chemical equation. Show your work on a separate sheet. For each step with a chemical reaction that does not involve redox, write a balanced chemical (molecular form) equation. For all of the equations, include symbols for physical states or you will lose credit! Step 2 Chemical Equation Ionic Equation Net Ionic Equation Step 4 (NOT the neutralization reaction that was provided) Observation NAME SECTION Chemical Equation 47
REPORT FOR COPPER CHEMISTRY AND REDOX EXP. (cont). Step 6 Observation Chemical Equation Step 10 Observation Chemical Equation Step 14 (Note: Two different chemical reactions occurred in this step: give the equations for both. Only one reaction involved copper.) Observation Chemical Equation A Ionic Equation A Net Ionic Equation A NAME Chemical Equation B Ionic Equation B Net lonic Equation B 48
REPORT FOR COPPER CHEMISTRY AND REDOX EXP. (cont). Step 16 (Note: This reaction does not involve copper!) Observation Chemical Equation Ionic Equation Net Ionic Equation B. Additional Oxidation - Reduction Reactions-show work on separate sheets-see lab manual procedure for products Reaction of potassium permanganate with iron(11) chloride in hydrochloric acid solution Observations NAME Chemical Equation Tonic Equation Net Ionic Equation 49
REPORT FOR COPPER CHEMISTRY AND REDOX EXP. (cont.) Reaction of sodium dichromate with potassium iodide in sulfuric acid solution. Observations Chemical Equation Ionic Equation Net Ionic Equation Reaction of nickel(II) nitrate with potassium persulfate in sodium hydroxide solution Observations Chemical Equation Ionic Equation Net Ionic Equation Reaction of aluminum with sodium bromate in sodium hydroxide solution Observations Chemical Equation NAME Jonic Equation Net Ionic Equation 50
QUESTIONS FOR COPPER CHEMISTRY AND REDOX EXP. 2. Why are CuSO4 and Cu(NO3)2 the same color? 1. What are the greatest sources of error in copper recovery in this experiment and what can be done to minimize these errors? Initial mass of copper Mass of recovered copper + dish Mass of evaporating dish NAME 3. From the following data, compute the percent recovery of copper. Assuming that the masses given are correct (no weighing errors), give a possible explanation for the fact that the percent recovery is greater than 100%. Mass of recovered copper Percent Recovery 51 0.6572 g 182.7653 8 182.1051 g 11
COPPER CHEMISTRY AND REDOX REACTIONS INTRODUCTION This laboratory exercise will examine several oxidation-reduction reactions (reactions which involve the transfer of electrons from one substance to another). In the past, the first experimental procedure was performed using pure copper pennies. However, modern pennies are not pure copper, but instead have zinc centers with copper plated only on the outer surfaces. Therefore, this experiment will use pieces of pure copper metal. The reactions that will be used to dissolve the copper metal and then recover it are outlined below: HNO, NaOH H₂SO4 Zn Cu → Cu²+ (→→ Cu(OH)2 CuO → Cu²+ (aq) → Cu 42 A The first reaction in this experiment generates nitrogen dioxide, a poisonous gas. Breathing only a small amount of this gas can result in inflammation of the lungs. In larger amounts, this gas can be fatal. In addition, later experimental procedures will generate flammable hydrogen gas. All of these potentially hazardous procedures must be performed in the fume hood. Several reactions in this experiment require the use of concentrated acids. Use these acids in the fume hoods. Do not remove the bottles from the hoods. Immediately clean up any spills using sodium bicarbonate to neutralize the acid and then wipe up the moist neutralized salts. In this laboratory exercise, if the reaction is a metathesis (or other reaction in which there is no oxidation- reduction reactions), only the chemical (molecular) equation will be written and balanced. For each step of the copper chemistry experiment in which a chemical reaction occurs, and also for each of the four redox demonstration reactions, it will be necessary to write balanced net ionic, ionic, and chemical equations.
DAY 1 - A. Copper Chemistry 1. Work in groups of 3 to 4 people. Obtain a pair of beaker tongs and a piece of copper wire. Determine the mass of the copper wire to the nearest 0.0001 g. (It can be weighed directly on the balance pan.) Now place the copper wire in a clean 400 ml. beaker. Record your observations for each step in which a chemical reaction occurs. 2. Make sure the hood fans are on. In the hood, add 5.0 ml. of concentrated nitric acid to the copper in the beaker. Two of the products from the oxidation-reduction reaction that occurs are nitrogen dioxide and copper(II) ion in aqueous solution. 3. Keep the beaker at the back of the hood. Avoid unnecessary exposure to nitrogen dioxide by keeping your head out of the hood and the glass shield pulled down in front of your face. Tip and swirl the beaker occasionally to make sure the copper reacts completely. As long as the solution is green, nitrogen dioxide gas is still being formed. When the copper has completely dissolved and no mornitrogen dioxide gas is being produced, the solution will be completely blue (not green). Make sure that all of the copper wire has dissolved. Now add approximately 100 ml of cold, deionized water to the beaker, then take the beaker out of the hood and return to your bench space. The blue color of your solution is due to the aqueous copper(II) cation. 4. Add 20 mL of 6-M NaOH (stored in the hood) to your solution. The hydroxide ion will react with the copper(II) ion to form a gelatinous white copper(II) hydroxide precipitate. The precipitate appears blue because some of the copper(11) ion remains in hydrated. Of course, NaOH also neutralizes the excess nitric acid (from step 2) according to the balanced molecular equation: NaOH + HNO-NANO + H₂O 5. Obtain one large piece of broken crucible. Rinse it well and add it to your beaker containing the precipitate. It will act as a boiling chip to prevent "bumping" and splattering during heating in the next step. 6. With gentle stirring, carefully heat your precipitate mixture to just boiling (use a hot plate on a low setting). The heat will cause all of the copper ions to react and will convert the copper(II) hydroxide into black copper(II) oxide solid (water is expelled in the process). Heat until all of the blue color is gone and the solids are completely black. Use beaker tongs to transfer the hot beaker to a ceramic tile and allow the black solid to settle. 7. While the precipitate is settling, heat a large beaker of about 300 mL of deionized water to boiling. 8. Wash your evaporating dish and store it in your drawer. It needs to be clean and dry for DAY 2. It is difficult to precipitate out copper metal while NO, is present. Step 9 removes most of the nitrate. 9. After your precipitate has settled, decant the clear solution (supernatant) into a separate beaker. Try not to lose any precipitate. Using beaker tongs to handle the hot beaker, wash the precipitate with about 100 ml. of hot water. Allow the precipitate to settle. Decant the supernatant and repeat the wash procedure two more times. Be sure to keep the deionized water hot. Decant as completely as possible after the last wash. 10. Add 2 mL of 6.0-M H₂SO, to the black precipitate. Make sure all of the precipitate dissolves. If all of the precipitate does not dissolve, add more sulfuric acid 0.5 ml. at a time until it does. Decant your solution into a second clean 250 ml. beaker leaving the boiling chip in the original beaker. Rinse the original beaker containing the boiling chip with deionized water and add this rinse water to your solution in the second beaker (you are trying not lose any copper ions). 11. Return the piece of broken crucible to the container provided. 12. Cover the solution with a watch glass (curved side down) and store it in your drawer until the next lab
DAY 2 - A. Copper Chemistry (cont.) 13. Calculate the mass of zinc metal needed to completely react with the mass of copper with which you started the experiment and all of the sulfuric acid added in step 10. 14. In the hood, add approximately the mass of 30-mesh zine that you calculated to your solution (a small amount at a time) with gentle stirring. Two different reactions are occurring. Zinc metal is donating electrons to the copper(II) ion, precipitating copper metal and forming zinc ion in solution. In addition, zinc metal is reacting with the excess hydrogen ion from step 10 on day 1 to produce zinc ion and hydrogen gas (which is flammable). 15. Continue to stir the mixture. Every few minutes, stop stirring and look for hydrogen gas formation. After the gas production has slowed down, look at the supernatant. It should be colorless (or have a slightly gray color). If it is still blue, add more zinc, a small amount (0.1 to 0.2 g) at a time until the blue color is gone. Do not add too much zinc. Note: it is now necessary to dissolve any excess zinc metal that may be present. Zinc will react with hydrochloric acid and dissolve but copper metal will not react. 16. When you no longer see gas formation in the beaker and the copper precipitate has settled, carefully decant and discard the supernatant. While still in the hood, first add 5.0 ml of distilled water and then 10.0 ml of concentrated (12-M) hydrochloric acid to the precipitate. Rinse your graduated cylinder immediately. When gas evolution in the beaker has stopped, remove the beaker from the hood. 17. Allow the copper to settle while you weigh the evaporating dish (which was cleaned on DAY 1) to the nearest 0.0001 g. (The dish must be clean and dry.) 18. Decant the supernatant into another beaker. Pour the supernatant into one of the large beakers located by the sinks in the lab. Fill the beaker with water and pour the resulting solution down the drain in the sink. Rinse both the large beaker and your beaker. Transfer the copper to the weighed evaporating dish. Use your wash bottle to rinse any solid remaining in the beaker into the evaporating dish. Washi the copper in the dish three times with about 30 mL of deionized water each time. Drain as completely as possible after each washing. 19. In the hood, wash the copper (still in the dish) with 10 ml of methanol. Decant the methanol into a beaker and discard the used methanol in the waste methanol container. Repeat with a second 10 ml. sample of Methanol. 20. In the hood, wash the copper with 10 ml of acetone, decant, and discard the used acetone in the waste acetone container. Repeat with a second 10 ml. sample of acetone. 21. Gently warm the copper in the evaporating dish on a hot water bath. Stir gently until the copper is dry and granular. Do not overheat or the copper will oxidize (react with the oxygen in the air). Note: Do not put your copper away wet. It will oxidize!! 22. Store the evaporating dish containing the copper uncovered in your drawer until the next laboratory period. DAY 3-A. Copper Chemistry (cont.) 23. Weigh your copper and evaporating dish. Determine your percent recovery. g of recovered Copper g of original Copper Percent Recovery 24. Discard copper in the container provided. *100
25. For each step of the Copper Chemistry experiment in which a chemical reaction occurred, write a chemical equation. B. Additional Oxidation - Reduction Reactions. 1. Observe your professor's demonstration of four oxidation-reduction reactions. 2. Record your observations (color changes, phase changes, appearance of reactants and products). 3. Follow the procedure outlined in the previous sections and write a balanced net ionic equation for each reaction. Derive the balanced ionic and chemical equations for each reaction. The reactions are: a. Potassium permanganate reacts with iron(II) chloride in hydrochloric acid solution. Two of the products of the reaction are manganese(II) ion and iron(III) ion. b. Sodium dichromate reacts with potassium iodide in sulfuric acid solution. Two of the products of the reaction are chromium(III) ion and molecular iodine. c. Nickel(II) nitrate reacts with potassium persulfate, K.S.Os, in sodium hydroxide solution. Two of the products of the reaction are nickel(IV) oxide (a black precipitate) and sulfate ion. d. Aluminum metal reacts with sodium bromate, NaBrO,, in sodium hydroxide solution. Three of the products of the reaction are bromide ion, hydrogen gas, and the Al(OH), ion. In this reaction there are two reductions and they can be combined in the same half-reaction. However, the ratio of bromide to hydrogen must be determined experimentally. It has been found that one mole of hydrogen gas is produced for every mole of bromide ion produced.