Required Reading: Week 1: Chapter 5, sections 5.1 & 5.2 Week 2: Chapter 5, sections ; Chapter 7, sections

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1 Name: Chem 121 Lab Clark College Partner: Instructor initials Week 1: Experiment 6: Reaction of Magnesium Metal with Hydrochloric Acid Adapted from: Chemistry, Experimental Foundations, 4th Ed. Laboratory Manual, by Merrill, Parry & Bassow. Content Goals: Write and balance chemical equations. Determine which gas law(s) are necessary in problem-solving situations. Use gas laws and experimental data to solve experimental problems. Process Goals: Read and follow complex instructions. Construct and organize a data table. Accurately measure volumes. Solve problems that combine stoichiometry, gas laws, and experimental data. Identify and explain experimental error. Required Reading: Week 1: Chapter 5, sections 5.1 & 5.2 Week 2: Chapter 5, sections ; Chapter 7, sections Instructor initials Week 2: Pre-Lab Questions - Do this with your lab partner before starting the lab! 1. Write the balanced equation for the reaction between magnesium metal and aqueous hydrochloric acid (HCl) to form hydrogen gas and aqueous magnesium chloride. 2. Read through this experiment carefully. Underline every sentence that instructs you to make a measurement. Neatly construct a data table on p. 3. Part A, Data Collection 1. Obtain a piece of magnesium ribbon approximately 4 cm long. Measure the length of the ribbon carefully to the nearest 0.01 cm. Your instructor will give you the mass of 1 m of ribbon. Since it is uniform in thickness, you can calculate the mass of the magnesium used. 2. Fold the magnesium ribbon so that it can be encased in a small spiral cage made of fine copper wire (Figure 1). Let enough copper wire serve as a handle so that the cage can rest at the 50 ml mark of the gas measuring tube. Mg/HCl S11 SB/KB Page 1 of 6

2 3. Set up a ring stand and utility clamp in a position to hold a 50 ml gas measuring tube. Place a 400 ml beaker about two-thirds full of tap water near the ring stand. 4. Tilt the gas measuring tube slightly and carefully fill it with with 3M hydrochloric acid (HCl) to about the 15 ml mark. 5. With the tube still tilted, slowly fill it with distilled water from a beaker. While pouring, rinse any acid that may be on the side of the tube. The final liquid in the top of the tube should contain very little acid, so try to avoid stirring up the acid layer. The tube Figure 1 should be brimming full of water. 6. Use the handle to insert the copper cage into the tube until is is positioned at about the 50 ml mark. Hook the wire over the edge of the tube and secure it in place with the stopper. The stopper hole(s) should be full of water and you should have no air bubbles in your tube. 7. Cover the hole(s) in the stopper with your finger and invert the tube into the 400 ml beaker of water. Clamp the tube in place and watch the reaction occur. 8. After the reaction stops, wait about 5 minutes to allow the tube and contents to come to room temperature. (Although you may not be able to feel it with your hand, this is a slightly exothermic reaction, and this waiting period is important.) Dislodge any bubbles by gently tapping on the side of the tube. 9. Again, cover the hole in the stopper with your finger and transfer the tube to a large cylinder of room temperature water (provided by your instructor). Raise or lower the tube in the water until the level of liquid inside the tube is the same as the level of water outside the tube. At this position, measure and record the volume of gas in the tube. Raise your hand to have your instructor check and initial your measurement. (See Figure 2) 10. Remove the gas measuring tube from the cylinder, and dispose of its contents in the disposal hood. Rinse your tube with tap water and replace it and your stopper on the prep bench. Pour the water from your 400 ml beaker down the sink. 11. Record the room temperature and the atmospheric pressure. Figure 2: Measuring the volume of gas at room pressure. a) P room<p gas b) P room=p gas c) P room>p gas Mg/HCl S11 SB/KB Page 2 of 6

3 Data table: Results and Calculations, show all work. 1. Compute the mass of the magnesium ribbon that you used. 2. Compute the number of moles of magnesium ribbon that you used. 3. The hydrogen gas (H 2 ) in this experiment was collected over water. This means that the gas in the collection tube contained both H 2 gas and water vapor. When you equalized the level of liquid in the two containers, the pressure of the hydrogen gas and water vapor equaled the atomospehric pressure: Patm = PH2 + PH2O Appendix A at the back of the lab has P H2O values listed for given temperatures. Use Appendix A to fin the vapor pressure of water for your experiment and compute the P H2. When your done, convert the P H2 from mm Hg (or Torr) to atm for your final answer. Mg/HCl S11 SB/KB Page 3 of 6

4 4. Use the ideal gas law to compute the volume 1.00 mole of H 2 gas occupies at STP. This will be the predicted volume of H 2 gas for this experiment. 5. The predicted volume above is based on 1.00 mole of H 2 gas. Based on the balanced chemical equation, how many moles of gas should have formed in your reaction? How many moles of Mg did you use in the reaction? Use this number to compute how many moles of H 2 gas should have been produced in your experiment. 6. Use your volume of H 2 gas and number of moles of H 2 gas formed in your reaction (#5 above) to compute the volume (in L) that would be occupied by 1.00 mole of hydrogen gas in your reaction. 7. To compare your result with predicted volume, use the combined gas law and your data to solve for the volume of H 2 gas (in L) your gas sample would have occupied if the reaction had been performed at STP. (Note: P 1 is your result from #3, V 1 is from #6, moles are constant in this calculation!) Mg/HCl S11 SB/KB Page 4 of 6

5 Follow-up Questions 7. Compute the % difference between your value and the predicted value. 8. List at least three reasons why your experimental value might have differed from the actual value. Measurement errors (such as we had trouble measuring the level of liquid in the tube) are NOT valid! Extension questions: 1. Through complex metabolic processes, our bodies oxidize glucose (C 6 H 12 O 6 ) to form carbon dioxide gas and liquid water. ( Oxidize, in this context, means that one of the reactants is oxygen gas.) a) Write a balanced equation for this process: b) A can of Dr. Pepper contains 40.0g of sugar. Assuming that all of this sugar is glucose (it isn t, but it s close enough!), how many moles of glucose are in the soda? c) How many moles of carbon dioxide are produced when all of the glucose in the soda is completely metabolized? Mg/HCl S11 SB/KB Page 5 of 6

6 d) What volume is occupied by that carbon dioxide at body temperature (37 C) and at atm? 2. The average volume of air exhaled in a normal breath for an average-size male is 5.0 x 10 2 ml. Roughly 5% of that volume is carbon dioxide. What mass of glucose would have been metabolized to produce that amount of carbon dioxide at body temperature and 1.00 atm pressure? (Show all your work and feel free to continue on the following page.) Mg/HCl S11 SB/KB Page 6 of 6

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