Lab 20: Ideal Gas Law In this experiment, we will use yeast to accelerate the decomposition of the hydrogen peroxide into water and O2 gas. Yeast contains the enzyme catalase, which is a catalyst for this reaction. You will add yeast activated in warm water to a known amount of hydrogen peroxide and quickly seal off the system so that the O2 gas formed is collected in a graduated cylinder. After measuring the total volume of gas produced, its temperature, and the atmospheric pressure, the ideal gas law can then be used to calculate how many moles of O2 gas is formed. We can do this by solving the ideal gas law equation for n. n= PV RT Once the number of moles of O2 gas is calculated, the percent of H2O2 present in the solution can be determined. To do this, you first need to calculate the theoretical number of moles of O2 there would be if the solution was 100% hydrogen peroxide. This can be found by using the following equation: Theoretical moles O2 = H2O2 used × H2O2 density × 1 mol H2O2 1 mol O2 × 34.0 g H2O2 2 mol H2O2 For this experiment:  mL H2O2 used is the volume of H2O2 you actually use (approximately 5 mL).  H2O2 density is 1.02 g/mL  1 mol H2O2 / 34.0 g H2O2 is the reciprocal (inverted fraction). of the molar mass of H2O2 . The molar mass of H2O2 is 34.0 g /mol, so this is equal to 1 mol H2O2 / 34.0 g H2O2.  1 mol O2 / 2 mol H2O2 is used since the decomposition produces 1 mole O2 from 2 moles of H2O2 .  The units in the entire equation cancel to give moles of O2. The percent hydrogen peroxide can now be found. To do this, divide (n), the actual number of moles you calculated, by the theoretical moles of O2 there would be if the hydrogen peroxide were 100%. This number is then multiplied by 100%. % H2 O2 = Figure 3: Carbonated beverages contain dissolved CO2 at high pressure. When the container is opened, this pressure can create a powerful burst, such as with this sparkling wine bottle, or when your soda “explodes.” Actual moles O2 (n)  100 Theoretical moles O 2 This value can now be compared to the 3% hydrogen peroxide shown on the label to see if any decomposition has occurred. 212 Lab 20: Ideal Gas Law Pre‐Lab Questions 1. What is it in yeast that aids in the decomposition of hydrogen peroxide? 2. List the ideal gas law and define each term with units. 3. How many moles of O2 were produced in a decomposition reaction of H2O2 if the barometric pressure was 0.980 atm, the temperature was 298 K and the volume of O2 gas collected was 0.0500 L? 4. If you decomposed 10.00 mL of 100% H2O2, how many moles of O2 could you theoretically obtain? 213 Lab 20: Ideal Gas Law Experiment: Finding Percent H 2 O 2 with Yeast Materials Safety Equipment: Safety goggles, gloves Yeast 10 mL Hydrogen peroxide 10 and 100 mL Graduated cylinders Erlenmeyer flask Stopper with hole Rigid plastic tubing (3 in.) *You must provide Rubber band Flexible tubing (18 in.) 2 Droppers (pipettes) 250 mL Beaker Stir rod 600 mL Beaker Thermometer Stopwatch Warm water* Ring stand* Distilled water* Large ring* *Optional Materials (not provided) Graduated cylinder Collected gas Rigid tubing Flexible tubing 600 mL Beaker Stopper Ring Stand Rubber Band Erlenmeyer flask Procedure 214 Figure 3: Gas Collection Apparatus (not to exact scale) 1. Prepare the materials for the apparatus as shown in Figure 1. Insert the smaller rigid tubing into one end of the larger, flexible tubing. Insert the free end of the rigid tubing securely into the rubber stopper hole. 2. Bend the free end of the flexible tubing into a U shape, and use a rubber band to hold this shape in place. This will allow you to more easily insert this end of the flexible tubing into the inverted graduated cylinder. Make sure the tubing is not pinched and that gas can flow freely through it. 2. Fill the 600 mL beaker with 400 mL distilled water. 3. Fill the 100 mL graduated cylinder with distilled water slightly over the 100 mL mark. Lab 20: Ideal Gas Law 4. Take the temperature of the water in the 600 mL beaker, and record it in the Data section. Also, determine the barometric pressure in the room, and record it in the Data section. HINT: The pressure in your region may be found online—if necessary, convert this value to mm Hg. 5. Mix 100 mL of warm water (45°C) and 1 packet of baker’s yeast in a 250 mL beaker. This will activate the yeast from the dormant (dry) state. Be sure to mix well with a stir rod until the yeast is completely dissolved. 6. Use a 10 mL graduated cylinder and pipette to measure out 5.00 mL of hydrogen peroxide. Pour this hydrogen peroxide into the Erlenmeyer flask, and place the stopper with stopper tube over the top. 7. Clean the 10 mL graduated cylinder by rinsing it at least three times with distilled water. Dispose of the rinse down the drain. 8. Cover the opening of the graduated cylinder with two or three fingers and quickly turn it upside down into the 600 mL beaker already containing 400 mL of water. DO NOT remove your fingers from the opening until the graduated cylinder is fully submerged under the water. If the amount of trapped air exceeds 10 mL, refill the cylinder and try again. 9. Insert the U shaped side of the flexible tubing into the beaker, and carefully snake it into the submerged opening of the graduated cylinder. You want as little air as possible to be in the graduated cylinder. 10. Secure the graduated cylinder to the ring stand by sliding a ring under the submerged cylinder, then attach‐ ing the ring to the stand. OPTIONAL PROCEDURE: If your kit does not include a ring stand, you will hold the graduated cylinder in place while gas is collected. Make sure to keep the open end of the cylinder completely submerged to pre‐ vent additional gas from entering. Rest the graduated cylinder against the side of the beaker during experi‐ mental setup. 11. With the cylinder vertical, record the volume of air inside (the line at which the water reaches in the cylinder) in the Data section in Table 1. 12. Using the pipette, measure out 5.00 mL of yeast solution into the rinsed 10 mL graduated cylinder. NOTE: Do not immediately pour the yeast solution into the Erlenmeyer flask. 13. Prepare to place the stopper (still connected to the hose) on the Erlenmeyer flask. Reset the stopwatch. 14. Quickly pour the 5.0 mL of yeast solution into the Erlenmeyer flask. Immediately place the stopper securely in the opening of the Erlenmeyer flask by twisting it down into the flask gently. 15. Start timing the reaction with the stopwatch. 16. Swirl the Erlenmeyer flask to mix the two solutions together. 17. You will begin to see bubbles coming up into the 100 mL graduated cylinder. HINT: If gas bubbles are not immediately visible, make sure the stopper is on tight enough and the tubing is not leaking. You will need to start over after correcting any problems. 18. Continue to swirl the Erlenmeyer flask and let the reaction run until no more bubbles form to assure the reaction has gone to completion. This should take approximately 6‐10 minutes. HINT: Catalase works best around the temperature of the human body. You can speed the reaction up by warming the Erlenmeyer flask with your hands. 19. Record the time when the reaction is finished in Table 2 of the Data section, along with the final volume of air in Table 1. Remember to read it at eye‐level and measure from the bottom of the meniscus. 20. Pour all other liquids down the drain and clean the labware. 215 Lab 20: Ideal Gas Law Data Water temperature: ⁰C Barometric Pressure: mm Hg Table 1: Volume data Initial volume of air (mL) Final volume of air after reac‐ tion (mL) Volume of O2 collected (Final volume ‐ initial volume) Table 2: Reaction time data Time reaction started Time reaction ended Reaction time (s) Calculations The goal is to find the percentage of hydrogen peroxide in the solution! This can be found by working through the following steps. 1. Convert the temperature of the water from ⁰C to Kelvin (K). Use the equation K = ⁰C + 273. This will be your value for absolute T or the temperature in Kelvin. T = 2. ⁰C + 273 = K If necessary, convert the barometric pressure in the room from mm Hg to atmospheres (atm). Divide the measured pressure from the Data section by 760 mm Hg. This will give you pressure (P) in atmospheres. P = 216 mm Hg * 1 atm 760 mm Hg = atm Lab 20: Ideal Gas Law 3. Convert the volume of oxygen from mL to liters (L). V = mL * 1L 1000 mL = L 4. Rearrange the ideal gas law to solve for n. 5. You are now ready to solve for the number of moles of O2. Be sure the units cancel so that you end up with only the moles of O2 left. Use the value for the constant R given: R = 0.0821 L∙atm/mol∙K Actual number of moles of O2 (n) = moles 217 Lab 20: Ideal Gas Law 6. Calculate the theoretical number of moles of O2 there would be if the hydrogen peroxide were 100%, and not an aqueous solution. mol H2O2 g H2O2 Theoretical moles of O2 = H2O2 volume * H2O2 density * * 1 mol O2 2 mol H2O2 To use the above equation, calculate the following: — H2O2 volume is the volume (mL) of hydrogen peroxide used: — H2O2 density is known: — mol H2O2 g H2O2 Volume = mL H2O2 Density = 1.02 g/mL is the reciprocal of the molar mass of H2O2. First write the molar mass of H2O2 then find the reciprocal. Molar mass of H2 O2 = g H2O2/1 mol H2O2 Molar mass of H2 O2 reciprocal = Now you have all of the information needed to solve the equation for the theoretical moles of O2. All you need to do is fill in the blanks and do the calculations. Theoretical moles of O2 = Theoretical moles of O2 = 218 * * mol * Lab 20: Ideal Gas Law 7. Find the percent hydrogen peroxide. % H2O2 = 8. Actual moles O2 Theoretical moles O2 * 100% = % You can also easily determine the reaction rate. To do this, divide the total volume of oxygen collected by the total time of the reaction. Reaction rate = Volume O2 (mL) Reaction time (s) = mL/sec Post‐Lab Questions 1. Was the calculated percentage of hydrogen peroxide close to the same as the percentage on the label? 2. Considering that catalysts are not consumed in a reaction, how do you think increasing the amount of catalyst would affect the reaction rate for the decomposition of hydrogen peroxide? 219 All the observable changes you see when matter changes phase—the bubbles during boiling, the crystallization during freezing, and the disappearing of matter during evaporation—can can be explained at the particle level. Use your understanding of kinetic molecular theory and phase changes to complete the following drawing assignment. Use the whiteboard or paper, pencil, and a scanner to create particle-level diagrams showing the particle's relationship to other particles according to the phase they are in for two of the following three situations: • bubbles of gas form as pasta water begins to boil on the stove • a refreshing drink contains both ice cubes and water • freeze-drying technology dehydrates foods through sublimation Clearly label which two situations you are depicting and include a paragraph (three to five sentences long) explanation to accompany each drawing.
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