Copy of 101How Much CHM 113 Oxygen Gas is Produced Worksheet RS-1 PDF

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Lab Report Worksheet: How Much Oxygen Gas is Produced? Lab Section (or Day-Time): Friday 12:30 pm

Group Number: 1

Due Date: 11/22/19

Name(s):Reagan Sanders, Lindsey Edgar, Maryama Aden

1. What is the goal of this investigation and how will it be achieved? The goal of this investigation is to measure the amount of oxygen gas produced when mixing yeast and H2O2. This will be done by creating a contraption to trap and measure the amount of oxygen released by the reaction 2. What is the specific reaction that will be taking place in the 125mL Erlenmeyer flask for this investigation? The mixture of H2O2 and yeast will emit oxygen and the pressure will push it through the tube into the graduated cylinder.

3. What

4. a. Draw the system that will be used for gas collection in the investigation, with some or all of the following equipment: ring stand with ring, 100 mL graduated cylinder, 125 mL Erlenmeyer flask, stopper with rigid tubing, flexible tubing, tape, rubber band, 600 mL beaker, and 3500mL of water in plastic container. PROVIDE LABELS.

b. Discuss how you will find how much oxygen gas will be produced in the reaction using this set-up. (This should be a brief procedure). The oxygen produced will enter the tube connecting the flask and the upside-down graduated cylinder. The oxygen will displace the water in the graduated cylinder allowing us to measure the amount of oxygen produced over a given period of time. 5. Answer the following statements as True or False. a. true . 30% Hydrogen peroxide is one of the chemicals used in today’s lab. b. true . At least 2 trials are required for this investigation. c. true . 5.00mL of a yeast solution will be used as a catalyst in the reaction. d. true . The temperature of water used in this experiment must be recorded. e. false . The atmospheric pressure of water is dependent on volume. 6. Table 1: Experimental Values and Data. Complete the following table with your experimental data for each of the two trials.

(a) (b) (c) (d) (e) (f)

(g)

T: Temperature of H2O  (K) Use Thermometer from drawer to take the temperature of the water PH2O: Pressure of water vapor for this temp (mmHg) Look for it here: h ttp://genchem.rutgers.edu/vpwater.html Patmospheric: Atmospheric pressure (mmHg) TA will provide this number from Labquest and gas pressure sensor Mass of dry yeast (g) Use balance (close balance room doors, use shields) Volume of H2 O2 store-bought  solution analyzed (mL) Use balance to take the mass of a certain volume of this solution. Plug into “h” V: Volume of gas collected (L) [Measured After Reaction] in 100mL graduated cylinder Height of water remaining in 100mL graduated cylinder above the surface level in the beaker (mmH2 O)   fter Reaction] use r uler for height of column from water [Measured A surface to gas level

Trial #1

Trial #2

295.15K

295.15K

19.827 mmHg

19.827 mmHg

729

729

0.201g

0.202g

8 mL

8 mL

55mL

64mL

101.6mm

88.9mm

7. Table 2: Experimental Calculations. Complete the following table with your experimental data and calculations for each of the two trials.

(h)

Mass of H2O2 solution  analyzed (g)

(i)

Pwater column: Pressure exerted by column of water (mmHg) Convert (g) to mmHg 1mmHg = 13.595mmH2 O  PO2: Pressure of O2 (mmHg)  PO2 = P  a tm - P  H 2O - P  w ater

(j)

Trial #1

Trial #2

8.02 g .0098 mmHg

7.99 g .0112 mmHg

709.16 mmHg

709.16 mmHg

.933 atm

.933 atm

2.12x10(^-4) moles

2.47x10(^-4) moles

8.02 g

7.99 g

3%

3%

column

(k)

= (c) - (b) - (i) PO2: Pressure of O2 (atm) Convert (j) to atm; 760 mmHg = 1 atm n: Actual moles of O2 produced  (mol) (𝑘) × (𝑓) PV=nRT ; n =

(l)

𝐿∙ 𝑎𝑡𝑚 

0.08206

× (𝑎)

𝑘∙𝑚𝑜𝑙

Mass of H2O2 in  store-bought solution (g) 2H2O2 (aq) ฀   2H2O(l) + O2(g) Use stoichiometry to convert from mol O2 (l)  to g H2O2 Mass percent of H2 O2 in  solution (%)   × 100% %m = 𝑔 𝑂2 × 100% = (𝑚)

(m)

(n)

𝑔 g 𝐻2𝑂2

(o)

(ℎ)

97.6%

Average mass percent H2O2 in  solution

Complete calculations below for Trial #1, to show all the work with correct units. 8. Calculation “e”: What is the mass of your H2O2 solution  for Trial #1? 8.02g of H2O2 9. Calculation “i”: Calculate the pressure (in mmHg) of the water column using the height of the water (measured in mm) that you found after completion of the reaction (1mmHg=13.595mmH2O).  Show your work. 1.6mm x 1mmHg/13.595mmH20= 7.47mmHg

P

O 

10. Calculation “j”: Atmospheric pressure must equal the combined pressure of the trapped gas and the water column rising above the surface level. (Pa tmospheric =  Pt rapped gas +  Pw  ater column) and Dalton’s Law of Partial Pressures dictates that the pressure of the trapped gas must equal the pressures of the water vapor and the O2 produced  in the reaction. (PH O +  PO ).  Solve for PO2 in mmHg. Show your work. = P - P - P atmospheric 

H O 

water column

PO2 =  “c” – 729-19.827-.0098=709.16

“b” –

“i”

11. Calculation “k”: Convert to mmHg to atm before continuing with the next steps. (Hint: 1atm=760mmHg). Show your work. 0.993atm x 760mmHg/1atm=755mmHg

12. Calculation “l”: Calculate the actual moles of O2 from  the ideal gas law, PV=nRT. (Equation 2). i. Show your work for the calculation.

nO2= (.933atm)(.0055L)/(.08206)(295.15k)= 2.12 x 10^-4 ii. You have collected many pressure values. What pressure will you use? What are the units? Pressure of O2 in atm iii.

What volume will you use? What are the units?

Volume of O2 collected

iv.

What value of R will you use in the ideal gas law equation, PV=nRT? Include units.

We used 0.08206 (L x atm) (mol x K)

v.

What T will you use? What are the units?

The temp of H2O in K

13. Calculation m”: Use the stoichiometry of the balanced reaction to convert the number of moles of O2 gas  present in the collection tube into grams of H2 O2 present  in the store-bought solution. Show your work. 0.0002119 mol O2 x 2.0 mol H2O  2  / 1.0 mol O2 = 0.0004238 mol H2O  2  0.0004238 mol H2O  x 34.02 g/mol = 0.0144g H2O2 14. Calculations “n and o” (experimental value): Calculate the mass percent of H2O2 in  solution and get the average for your trials. Show your work.

7.99 g / 8.221 g = 97.5% H2O2 15. Calculate for % error by using 3% as the accepted value. Show your work.

0.03 g - 0.0144 g / 0.03 g = 52 % error

Repeat calculations for 7-14 for Trial #2. Show your work and include units, etc. 16. Calculation “e”: What is the mass of your H2O2 solution  for Trial #2?

7.99g of H2O2 17. Calculation “i”: Calculate the pressure (in mmHg) of the water column using the height of the water (measured in mm) that you found after completion of the reaction (1mmHg=13.595mmH2O).  Show your work.

1.6mm x 1mmHg/13.595mmH20= 7.47mmHg

18. Calculation “j”: Atmospheric pressure must equal the combined pressure of the trapped gas and the water column rising above the surface level. (Pa tmospheric =  Pt rapped gas +  Pw  ater column) and Dalton’s Law of Partial Pressures dictates that the pressure of the trapped gas must equal the pressures of the water vapor and the O2 produced  in the reaction. (PH O +  ).  PO  Solve for PO2 in mmHg. Show your work. 729-19.827-.0098=709.16

19. Calculation “k”: Convert to mmHg to atm before continuing with the next steps. (Hint: 1atm=760mmHg). Show your work. .993atm x 760mmHg/1atm=755mmHg

20. Calculation “l”: Calculate the actual moles of O2 from  the ideal gas law, PV=nRT. (Equation 2). i. Show your work for the calculation.

n=PV/RT : 0.933atm(0.0055 L)/(0.08206(Lxatm)(molxK)(295K)=2.12 x 10^-4 moles of O2 ii.

You have collected many pressure values. What pressure will you use? What are the units?

iii.

0.933 atm What volume will you use? What are the units?

0.0055 L iv.

What value of R will you use in the ideal gas law equation, PV=nRT? Include units.

We used 0.08206 (L x atm) (mol x K)

v.

What T will you use? What are the units? 295 K

21. Calculation “m”: Use the stoichiometry of the balanced reaction to convert the number of moles of O2 gas present in the collection tube into grams of H2O  2 present in the store-bought solution. Show your work. 0.0002119 mol O2 x 2.0 mol H2O2 / 1.0 mol O2 = 0.0004238 mol H2O2

0.0004238 mol H2O  x 34.02 g/mol = 0.0144g H2O2 22. Calculations “n and o” (experimental value): Calculate the mass percent of H2O2 in  solution and get the average for your trials. Show your work.

8.02 g / 8.221 g = 97.6%

H2O2

23. Calculate for % error by using 3% as the accepted value. Show your work. 0.03 g - 0.0144 g / 0.03 g = 52 % error

24. What evidence do you have that the gas you will collect is oxygen?

When analyzing the stoichiometry of the balanced reaction we are able to predict that when h2o2 decomposes it should produce 2 moles of water and 1 mole of oxygen gas. So when collecting the gas we also did calculations based on the information that we collected during the lab and were able to determine that the gas that we collected had to be oxygen gas. 25. Discuss possible errors and limitations in the procedures used in this investigation.

Some possible errors that we could’ve experienced throughout this lab could have been the fact that the rubber tube wasn’t sealed all the way causing some air bubbles to escape causing the error percent of our calculations to be affected. As well as the size of the flask, using a bigger flask can decrease the amount of gas collected since the stopper will not be secured enough to hold the gas. 26. How could the equipment or conditions be improved to obtain better results?

Using a much smaller gas for example the 125 mL flask that we were required to use and using some sort of sealant to reduce the amount of gas the escapes throughout the

experiment. As well as the pressure of the room, if the pressure was different depending on the conditions we may have received a much accurate result.

27. Relate the experiment to outside research conducted; i.e, how are scientists using these techniques in their experiments? How can these techniques help you in possible field research? With this experiment it may help us determine the vapor pressure of different gasses based on their boiling point and the collection of gas. We can also determine how much reactants are lost throughout an experiment and giving us the techniques that may help us determine the moles of different elements during decomposition....