7_GasLaws_PostLab_Spr24

pdf

School

Arizona State University *

*We aren’t endorsed by this school

Course

113

Subject

Chemistry

Date

Apr 3, 2024

Type

pdf

Pages

Uploaded by ChiefWillpowerFalcon32

Gas Laws Michelle Lewis 02/19/2024 TOTAL: 20 points

Michelle Lewis Gas Laws Page 2 of 12 CHM 113 POST-LAB Gas Laws 1. Insert ONE picture of yourself in full PPE here ( include the thermometer ). **Remember to (1) show your full body so that we see you are wearing shoes; (2) wear your safety glasses, buttoned lab coat and gloves; (3) cover your arms and lower legs (socks are not optional, even in Arizona); (4) tie back long hair in a ponytail or a bun; (5) remove jewelry.

Michelle Lewis Gas Laws Page 3 of 12 2. Insert the picture you took of the apparatus you constructed (Part 1, Step 1) for collecting oxygen gas from the hydrogen peroxide decomposition here, making sure to show the flask, tubing, and inverted graduated cylinder. Be sure to include a figure number and legend below the picture. Insert Picture Here

Your preview ends here

Eager to read complete document? Join bartleby learn and gain access to the full version

Access to all documents
Unlimited textbook solutions
24/7 expert homework help

Michelle Lewis Gas Laws Page 4 of 12 3. Complete Table 1 with your experimental data for the decomposition of hydrogen peroxide catalyzed by yeast in Part 1. Table 1: Temperature, Pressure and Volume Data Temperature of Distilled H 2 O Room (or regional) Pressure (atm) Initial Volume of Air (mL) Final Volume of Air (after reaction) (mL) Volume of O 2 Collected (Final Volume - Initial Volume) 24.3 C 1.00 atm 10 mL 45 mL 45 mL-10 mL = 35 mL 4. In order to determine the number of moles of oxygen produced, you will need the partial pressure of the oxygen in the gas sample collected. This sample contains not just oxygen, but also a very small amount of water vapor. Using Dalton’s Law of partial pressures, we can use the following equation to find the pressure of just the oxygen (P O2 ) in the sample: ? ???𝒂? = ? ? ? + ? ?𝒂??? ?𝒂??? P total is the room pressure recorded in Table 1. Look up the water vapor pressure, P water vapor , (you may have to convert to atm) for the temperature you recorded in Part 1, Step 2 and calculate the pressure of the O 2 gas (in atm) . Show your work and include units. Hint: You can Google a vapor pressure table to find P water vapor. Temperature of room: 20 °C Pressure of water vapor ( at room temperature recorded in Part 1, Step 2 ): 0.0295 atm Partial pressure of oxygen gas (P ) = P total - P water = (1.00 - 0.0295 ) atm = 0.9705 atm

Michelle Lewis Gas Laws Page 5 of 12 5. Using the Ideal Gas Law along with your answer to Question 4 and your data in Table 1, what quantity of oxygen (in moles) was liberated from your hydrogen peroxide? Show your work and include units. 6. Using your answer from Question 5, the balanced equation below, and assuming that the decomposition of H 2 O 2 into oxygen was complete , ? 𝑯 ? ? ? (𝒂?) 𝒄𝒂?𝒂???? → ? 𝑯 ? ? (?) + ? ? (?) Ideal gas equation: PV = nRT where P = PO 2 = 0.9705 atm V is volume of Oxygen = 35 mL = 35 10 -3 T = 20 + 273 = 293 K R = 0.082 L-atm/mol.K so, number of moles of Oxygen (n) = PV / RT = 0.9705 x 35 x10 -3 / 0.082 x 293 = 0.00141

Michelle Lewis Gas Laws Page 6 of 12 a. How many moles of H 2 O 2 did you start with in solution ? Show your work and include units. b. What is the mass, in grams, of the H 2 O 2 in the solution you started with? 7. Determine the experimental percentage by mass of H 2 O 2 in your starting solution. Use your answer from Question 6 as well as the volume of the H 2 O 2 solution you used in Part 1, Step 7. Show your work and include units. Hint: Assume that the density of your H 2 O 2 solution Is 1.0 g/mL. 𝑴𝒂?? % ?? 𝑯 ? ? ? = ? 𝑯 ? ? ? ? 𝑯 ? ? ? ???????? × ???% From balanced equation: Mole ratio of H 2 O 2 to O 2 = 2: 1 Moles of H 2 O 2 started with = 2 moles of Oxygen = 2 0.00240 = 0.00480 mass of H 2 O 2 in solution = moles molar mass of H 2 O 2 = 0.00480 mol 34 g/mol = 0.163 g From Question 6 Use V from Part 1, Step 7 to find

Your preview ends here

Eager to read complete document? Join bartleby learn and gain access to the full version

Access to all documents
Unlimited textbook solutions
24/7 expert homework help

Michelle Lewis Gas Laws Page 7 of 12 8. Was the calculated (or experimental) percentage of hydrogen peroxide in Question 7 close to the theoretical percentage given on the label (i.e. 3%)? Calculate the percent error of your value. Show your work and include units. % 𝑬???? = |??????????𝒂? − ????????𝒄𝒂?| ????????𝒄𝒂? × ???% 9. Insert a picture of the flask with balloon attached to the neck (a) BEFORE HEATING , (b) AFTER HEATING , and (c) AFTER COOLING here. Be sure to include a figure number and legend below the picture. Mass % of H 2 O 2 = m H2O2 / m H2O2 solution x 100% 0.163 / 4.2 x 100% = 0.03881 % Error = experimental – theoretical / theoretical x 100% = (0.03881 – 0.03 / 0.03) x 100% = 0.2936 Yes, close to the theoretical yield of 3%

Michelle Lewis Gas Laws Page 8 of 12 10. Answer the following: a. What happened to the volume of the air in the balloon when the balloon was exposed to changing temperature conditions? Hot air ________ the balloon slightly inflated ____________ Cold Air _______ the ballon deflated ____________ b. Using your understanding of the relationship between temperature and pressure of a gas, explain what might happen if an inflated Mylar balloon (the shiny silver non-stretchy kind) were left inside a car on a hot July day in Phoenix (for those who don’t live in Phoenix, it gets HOT!). (Answer should be 2 – 3 sentences) Hint: Think about the characteristics of Mylar.

Michelle Lewis Gas Laws Page 9 of 12 The balloon's air volume increases in hot air exposure and decreases in cold air exposure. An inflated Mylar balloon could explode on a sweltering July day in Phoenix due to the confluence of rising pressure and high temperatures. 11. Complete the following table with your experimental data for the temperature and volume data from the syringe experiment in Part 3 . Table 2: Temperature vs. Volume of Gas Data Temperature Conditions Temperature (°C) Volume (mL) Room Temperature 23.4 3 Hot Water 45.8 5 Ice Water 8.3 2

Your preview ends here

Eager to read complete document? Join bartleby learn and gain access to the full version

Access to all documents
Unlimited textbook solutions
24/7 expert homework help

Michelle Lewis Gas Laws Page 10 of 12 12. Plot the volume of gas as a function of the temperature of gas using the volume and temperature data from Table 2 in Question 11 (do NOT connect your points). Your graph will consist of just three data points. Draw in your plot a trend line (or straight line of best fit) through your data points, and then extrapolate the line until it intersects the (negative) x-axis. Include the equation of your trend line and the R 2 value in the figure legend below the graph. Hint: To extrapolate the trend line, change the minimum values on the x-axis and y-axis to negative values until the line intersects with the x-axis. Use graphing software to create an x-y scatter plot. On your graph, the x-axis should be the temperature (in °C), and the y-axis should be the volume (in mL) at that temperature. Do NOT connect your points; but include a trend line and m ake sure to: • Label the axes including the units in parentheses. • Include a figure legend (text below the graph) explaining the data in your graph that includes the equation of your trend line and the R 2 value . • Remove any default title, legend, and gridlines. You will not receive credit if you draw your graph by hand. Excel or Google Sheets are good choices that you can learn how to use quickly if you don’t already have a favorite graphing program. 0 1 2 3 4 5 6 0 10 20 30 40 50 Volume (mL) Temperature (C) Figure 1.1: Shows the temperature and volume of Room Temperature, Hot water and Ice water

Michelle Lewis Gas Laws Page 11 of 12 See other general tips for making graphs in the How to Make a Graph in Excel document located in the Introductory Materials for this lab. 13. Putting it all together! Using the graph in Question 12 and your knowledge of absolute temperature, answer the questions below. a. At what temperature does your line intersect the x-axis? What volume corresponds to this temperature? Temperature ______23.4__________ Volume 3 mL_________ b. Using your graph, predict the volume of the air in the syringe at -15 o C? Volume at -15 o C _______2.5 mL______________ c. Would it be possible to cool a real gas down to zero volume? Why or why not? What do you think would happen before that volume was reached? (Answer should be 2 – 3 sentences) No, the actual gas should either react like a solid and be so dense that no air could escape the syringe into the experiment, or it should transform into a solid at 0 volume. d. Is your measurement of absolute zero close to the theoretical value (i.e. -273 o C)? Calculate a percent error to indicate how close (or far) your experimental value (from Q13a) is to absolute zero. Show your work and include units. % 𝑬???? = | ??????????𝒂? − ????????𝒄𝒂? ????????𝒄𝒂? | × ???% -0.03881 / 0.2936 = -0.13218 C Error = (-0.13218) - (-273) / -273 x 100% = -0.99951582 The measured number is not near to the actual value because of the extremely high error percentage.

Michelle Lewis Gas Laws Page 12 of 12 e. Discuss TWO sources of error and ONE way to change the experiment to get closer to the theoretical value (“Human error” is NOT an acceptable source of error—be specific!). (Answer should be 3 – 5 sentences) Error 1: U sing extremely hot and extremely cold water temperatures, which would have badly affected or interrupted the experiment. Both elements might influence the lab's findings. This mistake would cause you to have to redo the experiment. Error 2: Errors in the experiment, such as too much or not enough air entering the chamber. Not knowing when to let up on the pistol in order to adjust the syringe's volume. These two could have an impact on the lab's findings. Proposed Change: Determine the proper water temperature for the lab to run smoothly. ensuring that the equipment is configured appropriately to prevent mistakes in the outcomes. Both of the suggested adjustments can lessen the effects of a failed lab.

Your preview ends here

Eager to read complete document? Join bartleby learn and gain access to the full version