General Chemistry I CHEM-1030 Laboratory Experiment No. 4 (Revised 10/21/2016) Stoichiometry: The Reaction of Iron with Copper(II) Sulfate

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1 Laboratory Experiment No. 4 (Revised 10/21/2016) Introduction In this experiment, you will use stoichiometric principles to deduce the appropriate equation for the reaction between metallic iron and a solution of copper(ii) sulfate. The reaction produces metallic copper, which precipitates as a fine red powder. This type of reaction, in which a solid metal displaces a metal cation from solution, is known as a single substitution or single displacement reaction. A metal capable of displacing a metal cation from solution is considered to be more active than the displaced metal. In this experiment, iron metal is more active than copper. Iron forms two different cations, Fe 2+ and Fe 3+. You will use stoichiometric principles to determine which one of these ions is produced in the reaction between elemental iron and copper(ii) ion If Fe 2+ forms, then Equation a is correct, while Equation b is correct if Fe 3+ forms. Each reaction is represented below first in its molecular equation form and then in net ionic form. Your task is to determine whether Equation a or Equation b is consistent with the results of your experiment. Fe(s) + CuSO4(aq) FeSO4(aq) + Cu(s) (Molecular Equation a) Fe(s) + Cu 2+ (aq) Fe 2+ (aq) + Cu(s) (Net Ionic Equation a) 2Fe(s) + 3CuSO4(aq) Fe2(SO4)3(aq) + 3Cu(s) (Molecular Equation b) 2Fe(s) + 3Cu 2+ (aq) 2Fe 3+ (aq) + 3Cu(s) (Net Ionic Equation b) You will add an excess of copper(ii) sulfate solution (to ensure that all the iron reacts) to a known mass of elemental iron. After working up your reaction mixture, you will weigh the metallic copper product. By comparing the mass of copper predicted by each equation with the mass of copper actually collected, you will be able to determine which of the two equations represents the reaction that occurs. Experimental Perform the following procedures on two samples simultaneously. If for any reason, you or the instructor feel you did not perform a trial properly, do a third or even a fourth trial. Use your two best trial results for your report. 1. Pour about 35 ml of 1.0 M CuSO4 into each of two Erlenmeyer flasks. Heat the solutions almost to boiling on a hotplate. 2. Pencil your initials and the (trial) numbers 1 and 2 on two clean, dry 100 or 250 ml beakers. Weigh the empty beakers separately on a hanging pan balance. After you weigh each beaker 1

2 once, check your balance for consistency. Move the balance masses back to zero and reweigh each empty beaker at least two more times to make sure you get the same reading. Show your results to the instructor before you proceed. If your balance does not read consistently, weigh each beaker a fourth or fifth time or use another balance. Do not move or adjust your balance once you start using it. 3. For trial 1, weigh between 0.8 and 0.9 g of iron powder into one of the beakers. For trial 2, weigh between 1.1 and 1.2 g of iron into the other beaker. 4. Slowly pour 35 ml of hot CuSO4 solution onto the iron powder in each beaker. 5. Swirl each beaker to mix the reactants. On your data sheet, record all the changes you see. When visible signs of reaction cease, place the beakers back on the hotplate. Let the solutions simmer for one minute to ensure the reactions go to completion and remove the beakers from the heat source. 6. Allow the copper product to settle completely. Carefully decant the hot solution from the solid copper (pour off the liquid and leave the solid behind). Exercise care and patience here. Remove as much solution as you can but do not allow any solid copper into the waste beaker. (Ignore the thin copper sheen on the liquid surface that does not settle out.) Pour the decanted liquid into a beaker at your work station for later disposal as heavy metal hazardous waste. 7. Add about 10 ml of distilled or deionized water to the solid copper in the beaker. Break up any clumps to allow the water to contact all the solid material. Swirl to wash all iron salts and unreacted copper sulfate from the solid copper. Carefully decant the wash water and add 10 more ml of distilled water to the beaker. Swirl and decant again. Combine the liquid from the two washes with the solution in the beaker designated as heavy metal hazardous waste. 8. In a fume hood that contains no hotplate or other ignition source, add 5 ml of acetone to the wet copper metal in the beaker. (CAUTION - Acetone is very flammable. Bulk flammable liquids must never stored, poured or used near any ignition source such as a flame, a hotplate or other electrical device.) Swirl the beaker and let the copper settle. Decant as much acetone as possible into a small beaker for later disposal as organic hazardous waste. Repeat with a second 5 ml portion of acetone and pour off as much of the acetone as you can. (Acetone is miscible with water and replaces most of the water mixed with the copper. Acetone has a lower boiling point than water and is easier to remove by evaporation.) Do not leave a large amount of liquid acetone in the beaker. Discard the acetone washings in the fume hood container designated for collecting organic hazardous waste. 9. Transfer the beakers to a different fume hood equipped with a hotplate set to medium high. Place the beakers at the cooler edge of the hotplate so the acetone and residual water do not evaporate too rapidly and cause the copper metal to spatter. Gradually move the beakers toward 2

3 the hotter center of the hotplate as the copper dries. Heat the beakers strongly for an additional 15 min after all visible traces of moisture are gone from the upper walls of the beakers. 10. Remove the dry beakers from the hotplate and let them cool to room temperature at your work station. Weigh the beakers and their copper contents at least three times on the same hanging pan balance you used previously to weigh the beaker and iron. 11. Scrape as much possible of the copper metal out of the beakers and discard the copper in the heavy metal waste collection container. Safety Wear chemical splash goggles and a waterproof apron at all times during this and all chemistry experiments, from the very beginning to the very end of the time you spend in the laboratory. Acetone is mildly toxic and extremely flammable. Use tap water to rinse off any acetone you get on your skin. Do not remove the acetone supply containers from the fume hood. Never expose acetone to any ignition source. Do not transfer any bulk acetone to a hood where a hotplate or other heat source is in use. Waste Disposal Discard all aqueous solutions and rinse water containing copper ion as heavy metal hazarous waste. Rinse all traces of copper sulfate from your glassware and discard the rinse water as heavy metal hazardous waste. Do not use an unnecessarily large amount of water. Excessive dilution of hazardous waste only serves to increase the expense of hazardous waste disposal. Discard your bulk copper metal product as heavy metal waste. All acetone must be discarded in the fume hood in the container designated for organic hazardous waste. Cleanup Wipe any residual copper out of the beakers with a damp paper towel and discard the towel as ordinary trash. Clean all your work surfaces with a damp sponge before you leave the laboratory 3

4 Laboratory Experiment No. 4 Data Page Trial Number Mass of Empty Beaker (Three or more separate weighings to check for balance consistancy) Average Mass of Empty Beaker Mass of Beaker and Powdered Iron Metal (Three or more separate weighings to check for balance consistancy) Average Mass of Beaker and Powdered Iron Metal Mass of Iron (By Subtraction) Observations on Mixing Iron Powder and Copper Sulfate Solution: Trial Number Mass of Beaker and Copper (Three or more separate weighings to check for balance consistancy) Average Mass of Beaker and Copper Mass of Copper Formed (By Subtraction) 4

5 Laboratory Experiment No. 4 Report 1. From the mass of elemental iron weighed out, calculate the mass of copper expected to form in each procedure. Use the two possible ways that cupric ion can react with iron metal. For full credit, show clear, concise stoichiometric setups and calculations. Express all answers with the correct units and with the correct number of significant figures. a) For your two best trials, calculate the masses of copper metal expected to form assuming that iron metal is oxidized to iron(ii) as shown in Equation a. (4 points) b) For your two best trials, calculate the masses of copper metal expected to form assuming that iron metal is oxidized to iron(iii) as shown in Equation b. (4 points) 2. What two visible signs did you observe indicating that the changes that occurred were chemical rather than physical changes? (1 points) 5

6 3. Compare your theoretical calculations in Question 1 to your experimental results and state whether Equation a or Equation b represents the reaction that actually occurred. Clearly state the reasons for your conclusion. (2 points) 4. The chemical change you carried out is a redox or oxidation-reduction reaction. Which reactant is oxidized and which one is reduced? Which reactant is the oxidizing agent and which is the reducing agent? (2 points) 5. What was the percent yield of one of your trials? In other words, what percent of the theoretical or expected yield of copper did you recover? Show a clear setup for the calculation. (1 point) mass of product recovered percent yield = x 100% theoretical mass of product 6. Should the stoiciometry of the reactions you carried out be affected by the differing masses of iron in the two trials? Yes or no? Give a clear explanation for your answer. (2 points) 7. Calculate the number of iron atoms you weighed out in trial 1. (2 points) 8. Calculate the number of copper atoms you recovered in trial 1. (2 points) 6

7 Laboratory Experiment No. 4 Prestudy Copper forms two different cations, cuprous (Cu + ) and cupric (Cu 2+ ). When solid copper reacts with silver nitrate solution, two reactions are possible, as shown in equations a and b. a) Molecular Equation: Cu(s) + AgNO3(aq) CuNO3(aq) + Ag(s) Net Ionic Equation: Cu(s) + Ag + (aq) Cu + (aq) + Ag(s) b) Molecular Equation: Cu(s) + 2AgNO3(aq) Cu(NO3)2(aq) + 2Ag(s) Net Ionic Equation: Cu(s) + 2Ag + (aq) Cu 2+ (aq) + 2Ag(s) 1. What mass of metallic silver can form from g of copper metal according to Equation a? (3 points) 2. What mass of metallic silver can form from g of copper metal according to Equation b? (3 points) 3. A student reacts g of copper with a solution containing excess silver nitrate and recovers g of silver metal. Which reaction, a) or b), most likely occurred in the experiment? Explain your answer completely. (2 points) 4. When copper metal and silver ion react, which reactant is oxidized and which is reduced? (1 point) 5. When copper metal and silver ion react, which is the more active metal? (1 point) 7