2. If 1800 L of Acetylene Gas (C2H2) Is Burned, How Many Liters of CO2(G) Are Formed?

H CH 11 Homework

Stoichiometry:

1. How many liters of ammonia gas (NH3) can be produced by the reaction of 4000 L of Nitrogen(g) w excess Hydrogen(g)?

2. If 1800 L of acetylene gas (C2H2) is burned, how many liters of CO2(g) are formed?

3. Electrolysis of liquid water produced 500 L of Hydrogen gas. What volume of Oxygen gas was also produced?

4. If 200 L of carbon monoxide gas react with 400 L of Oxygen gas, how many liters of carbon dioxide gas will be formed?

5. How many liters of Oxygen gas at .78 atm and 34 C are required to react with 55 L of Hydrogen gas at the same conditions?

6. When 70 g of H2O2 (s) are decomposed, how many liters of oxygen gas are formed? 2H2O2  2H2O + O2

7. If 40 g of solid Aluminum react with 60 g of hydrochloric acid (aq), how many liters of Hydrogen gas are produced?

8. How many grams of solid potassium chlorate are required in the preparation of 90 L of Oxygen gas at 1.20 atm and 45 C?

9. How many grams of solid Sodium Peroxide must react with water to produce 35 L of Oxygen gas at 5 C and 660 mmHg? 2 Na2O2 + 2 H2O  4 NaOH + O2

10. Calculate the number of liters of Oxygen gas needed to react with 75 g of solid Aluminum.

11. If 35 g of solid Sulfur react with 80 L of Oxygen gas, how many liters of Sulfur Dioxide gas would be produced?

12. Oxygen gas was originally produced by the decomposition of solid Mercuric Oxide

(HgO). What volume of O2(g) at 15 C and .98 atm can be produced from 2.36 g of HgO (s)?

13. If 32 g of solid Sodium are added to 25 L of Chlorine gas at 45 psi and 300 K, how many grams of solid salt would be produced?

1 Collection Over Water:

1. 11 g of solid potassium chlorate is decomposed and the resulting oxygen gas is collected over water. What volume of oxygen would be collected at .93 atm and 25 °C?

2. 20 g of solid potassium bromide is mixed with 10 L of chlorine gas at 34 psi and 20 °C. What volume of liquid bromine can be collected over water at these conditions?

3. 16.7 ml of liquid water is electrolyzed. The resulting gases are sent to 2 different tanks that are initially filled with water. The hydrogen tank is at 30 °C and the oxygen tank is at 70 °C. Which tank must have the larger volume? Atmospheric pressure on this day is 710 mmHg.

Density

1. What is the density of Oxygen gas at STP? Nitrogen gas?

2. What is the density of Fluorine gas at STP and at 30 C and 725 Torr?

3. Find the density of Freon 12 (g) (CF2Cl2) at 315 K and .945 atm.

Grahams Law

1. If Hydrogen gas will diffuse at a rate of .3 m/s at certain conditions, how fast will Oxygen gas diffuse at those conditions?

2. How much faster will Carbon Monoxide gas diffuse as compared to Carbon Dioxide gas?

3. What is the relative rate of diffusion of Fluorine gas to Chlorine gas?

2 Review:

1. 47 g of solid magnesium oxide is heated and decomposes. How many liters of oxygen gas are produced at 582.3 mmHg at 300 K? What if the oxygen was collected over water?

2. What is the density of carbon monoxide gas at 27 psi and 200 °C?

3. If you combine 30 g of hydrogen gas with 15 liters of chlorine gas at 744 Torr and 73°F, how many of mL of liquid HCl (density 1.18 g/mL) would be produced?

4. Chlorine gas diffuses at .15 m/s. How fast will Ne gas diffuse at the same conditions?

3 Temperature Pressure Temperature Pressure

4 (degrees C) (mmHg) (degrees C) (mmHg)

0 4.6 50 92.5

1 4.9 51 97.2

2 5.3 52 102.1

3 5.7 53 107.2

4 6.1 54 112.5

5 6.5 55 118

6 7 56 123.8

7 7.5 57 129.8

8 8 58 136.1

9 8.6 59 142.6

10 9.2 60 149.4

11 9.8 61 156.4

12 10.5 62 163.8

13 11.2 63 171.4

14 12 64 179.3

15 12.8 65 187.5

16 13.6 66 196.1

17 14.5 67 205

18 15.5 68 214.2

19 16.5 69 223.7

20 17.5 70 233.7

21 18.7 71 243.9

22 19.8 72 254.6

23 21.1 73 265.7

24 22.4 74 277.2

25 23.8 75 289.1

5 26 25.2 76 301.4

27 26.7 77 314.1

28 28.3 78 327.3

29 30 79 341

30 31.8 80 355.1

31 33.7 81 369.7

32 35.7 82 384.9

33 37.7 83 400.6

34 39.9 84 416.8

35 42.2 85 433.6

36 44.6 86 450.9

37 47.1 87 468.7

38 49.7 88 487.1

39 52.4 89 506.1

40 55.3 90 525.8

41 58.3 91 546

42 61.5 92 567

43 64.8 93 588.6

44 68.3 94 610.9

45 71.9 95 633.9

46 75.7 96 657.6

47 79.6 97 682.1

48 83.7 98 707.3

49 88 99 733.2

100 760

6 LAB: GAS COLLECTION

1. Obtain the following: pneumatic trough #6 rubber stopper with glass bend and rubber tube and thistle tube 250 ml Erlenmeyer flask 5 - 125 ml Erlenmeyer flasks 1 piece of rubber tubing

an unknown amount of NaHCO3 A glass plate and a 50 ml graduated cylinder

2. Place the 2 pieces of the rubber tubing on the pneumatic trough. The piece of tubing on the glass bend should be connected to the bottom opening in the trough. The other piece of tubing will be connected to the top hole of the trough and the other end should be placed in the sink.

3. Place the rubber tubing in the 250 ml flask so that the thistle tube is just above the bottom of the flask.

4. Fill the trough to just below the overflow outlet (top hole).

5. Since we are going to be determining the volume of gas produced by collecting it in the 125 ml flask, we need to calibrate the flask using the graduated cylinder so that we can determine the exact volume of the flask.

6. Fill your 5 - 125 ml flasks with water and, by placing a glass plate over the top of the flask, invert the flasks and place them in the trough. Place one of the flasks over the opening in the bottom of the trough.

7. Add your unknown amount of NaHCO3 to the 250 ml flask (NOT through the thistle tube). Write the letter of your unknown in the data table.

8. Measure out about 20 ml of 3M acetic acid. Place the lid on the 250 ml flask and add no more than 5 ml of acetic acid through the thistle tube at a time. Make sure to swirl the acid around to make sure that it is all reacted. Once all the acid has reacted, add more.

9. Collect the CO2 in the trough. When the 125 ml flask is almost full, you will need to remove it and place the next flask over the opening in the trough.

10. Once all of the acetic acid has reacted (ie: no more bubbling when you add acid) measure the volume of the CO2 produced. Also determine the temperature of the water in the trough.

Calculations:

1. Calculate the amount of NaHCO3 that you started with. It was either 1.00g, 1.50g, 2.00g.

Correct the volume of your gas to standard conditions REMEMBER: Reactions do not produce 100% yield. /20

7 Experiment: Determination of the Gas Law Constant Purpose In this experiment you will calculate the gas law constant, R, by collecting a known quantity of hydrogen gas and measuring the temperature, pressure and volume of the gas collected. Introduction From the ideal gas law, PV = nRT, you can see that it is possible to determine a value for R if you can isolate a sample of gas for which P, V, T and n are all known. In this experiment you will accomplish this by collecting hydrogen gas formed in the reaction of magnesium metal with hydrochloric acid. When you collect the hydrogen gas you will also measure the temperature, pressure and volume of the gas collected. From this data a value for R can be calculated. Since the hydrogen will be collected over water in a gas collection tube, it will be saturated with water vapor. According to Dalton’s Law, the total pressure of the gas mixture is the sum of the partial pressure of H2 plus the partial pressure of the water vapor. The partial pressure of the water vapor can be looked up in tables in the CRC, and subtracted from the total pressure to find the pressure of the H2. The volume of the gas will be measured directly from the gas collection tube. The necessary calculations are as follows:

 Determine the number of moles of H2 produced from the experimental mass of Mg used in the reaction.

 Record the temperature of the gas.

 Measure the volume of gas collected in the tube.

 Determine the total pressure of the gas in the tube using a barometer.

 Determine the partial pressure of the H2 gas by subtracting the partial pressure of the water vapor from the total pressure.

 Calculate a value for R.

Procedure 1. Obtain a gas collection tube and a plastic beaker. You will also need a clamp for mounting the tube on a ringstand.

8 2. Get 10 cm copper wire, a 1.00 cm strip of magnesium, and a “00” one hole stopper. 3. Mass the magnesium. 4. Roll the Mg into a loose ball, and wrap the Mg with the Cu wire to form a cage around the Mg. Leave at least 2.5 cm length of Cu wire coming off of the cage. 5. Obtain about 10 mL of 6 M HCl. Add the acid to the gas collection tube. Then carefully add water to completely fill the tube, being careful to not disturb the HCl as the water is added. Pouring the water slowly down the side of the tube will help. 6. Run the length of wire from the copper cage through the one hole stopper, bend the wire over the top of the stopper, and insert the assemblage into the gas collection tube so that the cage of Cu holding the Mg is inside the tube. Make sure that the tube is completely filled with water and that no air bubbles remain. 7. With one finger covering the hole in the stopper, invert the tube into the beaker of water. Mount the tube to a ring stand, being careful to not let the end of the tube come out of the water in the process. As the HCl flows down the tube it will react with the Mg, generating H2 gas which will rise to the top of the tube and displace water out of the bottom. 8. When the reaction is complete (you will see a dramatic change in the amount of hydrogen being given off) record the volume of gas in the gas collection tube. 9. Record the temperature of the gas collected, assuming its temperature is the same as that of the water in the beaker and gas collection tube. Also record the barometric pressure. 10. Repeat the experiment a second time.

Analysis and Calculations

1. Assuming that the magnesium reacted completely, calculate the moles of hydrogen formed from the mass of magnesium you actually used.

2. Calculate the pressure of the dry H2 from the barometer reading and the vapor pressure of water. (Look in the CRC index under vapor pressure, aqueous vapor, below 100°C.)

9 3. Using the experimental values for pressure, temperature, volume, and moles of the gas, calculate a value for R. Repeat the calculations for your second trial. If your two trials agree report an average.

4. Compute a percent error.

LAB: GRAHAM'S LAW

1. Obtain a glass tube and test tube clamp.

2. Clamp the tube to the ring stand so that the tube is level.

3. Obtain a cotton ball and using forceps, dip the cotton ball into the hydrochloric acid. Do the same with another cotton ball and the unknown.

10 4. Place the cotton balls in each end of the tube and gently place a rubber stopper in each end of the tube to close it.

5. Watch for a precipitate to form inside the tube.

6. Measure the distance from each end of the tube to the precipitate and determine which of the following gasses is your unknown:

NH3 NH4OH (NH4)2SO4 NH4Br NH4C2H3O2

7. Turn in all calculations and data.