Gas Laws. E k = ½ (mass)(speed) 2. v101613_10am

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1 Gas Laws v101613_10am Objective: In this lab you will become familiar with the Ideal Gas Law and Dalton s Law of Partial Pressures. You will be able to use the information collected along with stoichiometry to determine the percent of hydrogen peroxide in a commercially available product. Background: An ideal gas is one that follows the Kinetic Molecular Theory. (Chapter 5, Section 8 of Tro 2 nd edition) The Kinetic Molecular Theory states: 1. A gas is composed of molecules that are much smaller than the space between the molecules. 2. Gas molecules move randomly at various speeds and in every possible direction. 3. Forces of attraction and repulsion between gas molecules are negligible. 4. When collisions between molecules occur, they are elastic. 5. The average kinetic energy of gas molecules is proportional to the absolute temperature. E k = ½ (mass)(speed) 2 The ideal gas law equation (Chapter 5, Sect. 4 of Tro 2 nd edition) allows for the calculation of one variable needed to describe a gas under ideal conditions, given knowledge of the other three variables. The Ideal Gas Law is: PV = nrt In this equation, P stands for the pressure and V for volume. Pressure can have units of torr, mmhg, inhg, atm, Pa, etc. Volume may have units of L, ml, cm 3, m 3, etc. The unit used for pressure and volume in this equation will depend upon the pressure and volume units associated with R, which is the ideal gas constant. Two ideal gas constants commonly used include: R = atm. L / (mol. K) or torr. L / (mol. K). n stands for moles and therefore its units are mols, and T stands for temperature. When dealing with gasses, the only acceptable temperature unit is Kelvin, K. In this lab, you will determine the concentration of aqueous hydrogen peroxide by analyzing the amount of oxygen generated during the decomposition of the peroxide to form oxygen gas and liquid water. The decomposition of hydrogen peroxide is a naturally occurring process that you will be speeding up by use of a catalyst (peroxidase) present in yeast. The method used in this laboratory involves collecting oxygen gas generated by the decomposition of hydrogen peroxide above water (Chapter 5, Section 6 of Tro 2 nd edition) in a closed container using the apparatus shown below: Gas Laws - 1

2 In order to determine the amount of oxygen collected, you will be using the ideal gas law to solve for the number of moles of oxygen generated. The volume and temperature, associated with the gas, will be simple to determine. To determine the volume of the gas you will adjust the water level within the beaker so that it matches the water level within the graduated cylinder. You will then measure the volume directly from the 100-mL graduated cylinder. This method of measurement is acceptable since gases spread out to fill the space available. The temperature of the gas will be the same as the temperature of the water that the gas bubbled through. The pressure is the most complicated of the variables to determine. At all temperatures the surface molecules of a liquid can absorb sufficient energy to break free of the remaining liquid molecules and enter into the gas phase. The amount of liquid molecules entering the gas phase will depend upon the amount of energy absorbed by the liquid. The amount of molecules that break free and move from liquid phase into gas phase will cause a change in the pressure applied by that substance to the walls of its container. In this experiment, the liquid is water. The amount of water that will enter into the gas phase will be determined by the temperature of the water. The higher the water s temperature, the greater the number of water molecules that will break free from the liquid phase to enter into the gas phase. So the gas collected in the 100- ml graduated cylinder will be a combination of water vapor and oxygen. According to Dalton s Law of Partial Pressure, the total pressure of a mixture of gases is equal to the sum of the partial pressures of each gas present. In this experiment, therefore, the total pressure of the gas is the same as the pressure within the room (barometric pressure, when the water level inside the cylinder = water level in the beaker), is the sum of the pressure of the water vapor plus the pressure of the oxygen gas. P total = P water + P oxygen To solve for the partial pressure of the oxygen gas, the equation can be rearranged to give: P total P water = P oxygen Gas Laws - 2

3 Below is a list of water vapor pressures at various temperatures: Temperature in o Water Vapor Pressure in C Torr Once the pressure, volume, and temperature of the oxygen have been determined, the ideal gas law can be used to calculate the number of moles of oxygen generated by the acid. Important pressure conversion units: 760 torr = 1 atm = 760 mmhg Gas Laws - 3

4 Materials: 10 ml graduated cylinder Large test tube Small test tube Large Erlenmeyer flask Rubber tubing with stopper and glass tubing 100 ml graduated cylinder 1000 ml beaker Tub of water Thermometer Hydrogen peroxide solution Yeast Procedure: 1. Lower a 1000-mL beaker and a 100-mL graduated cylinder into a filled tub of water. Make sure there are no bubbles in each. 2. While the 1000-mL beaker and the 100-mL graduated cylinder are submerged under water, invert the cylinder into the beaker. 3. Remove the pair simultaneously, and dump some of the water from the beaker. 4. Obtain approximately an eraser size amount of yeast and place it inside the smaller test tube. Wipe the outside of the tube clean. 5. Obtain a clean, dry 10-mL graduated cylinder. Determine the mass of the empty graduated cylinder. 6. Pour about 8 ml of aqueous hydrogen peroxide into the cylinder and determine the mass with the H 2 O 2 solution. 7. Pour the peroxide into the large test tube. 8. Gently slide the small test tube into the large test tube. Make sure that none of the peroxide splashes into the tube with the yeast or the reaction will begin prematurely and the remaining measurements will be inaccurate. 9. Stand the large test tube in an Erlenmeyer flask. 10. Put the stopper end of the tubing into the large test tube so that the seal is tight. 11. Place the other end of the tubing with the curved glass gently under the lip of the 100-mL graduated cylinder. Be careful not to let any air into the cylinder during this step. 12. Carefully mix the peroxide and yeast by gently tilting the large test tube. Avoid sloshing the mixture too vigorously to keep any of it from splashing up into the rubber tubing. Bubbling will occur. 13. Continue mixing until the bubbling has stopped. 14. Alter the water level in the beaker until the water level in the beaker is equal to the water level in the graduated cylinder. 15. Read and record the volume of the gas in the graduated cylinder. 16. Measure and record the temperature of the water in the beaker. 17. Carefully pour the resulting solution down the drain. 18. Repeat procedure. Gas Laws - 4

5 Data: Trial 1 Trial 2 Amount of 3% Peroxide used (ml) Mass of peroxide solution used (g) Volume of Oxygen Collected (ml) Temperature of Water ( o C) Barometric Pressure (atm) (provided by instructor) Water vapor pressure from table (torr) Water vapor pressure (atm) Results: Trial 1 Trial 2 Pressure of oxygen (atm) Volume of oxygen collected (L) Temperature of oxygen (K) Moles of oxygen collected (mol) Moles of H 2 O 2 in sample (mol) Mass of peroxide in sample (g) % hydrogen peroxide in solution (%) % error ( accepted value measured value / accepted value) *100 Gas Laws - 5

6 Post Laboratory Questions 1. If you ignored the partial pressure of water vapor in your collected gas sample, would your experimental percent peroxide in the sample have been too high, too low, or not effected? Explain. 2. Why did the procedure specify to wipe the outside of the test tube containing the yeast before placing that test tube into the larger test tube containing the hydrogen peroxide? Would your experimental peroxide in the sample be too high, too low, or not effected? Explain. 3. Why did the procedure specify to gently slide the smaller test tube into the larger test tube to avoid accidentally splashing the hydrogen peroxide? Would your experimental percent peroxide in the sample have been too high, too low, or not effected? Explain. 4. What difference would it make to your results if you placed the bent tubing into the 100-mL graduated cylinder prior to putting the stopper into large test tube? Explain. 5. Suggest reasons at least two reasons that the percent hydrogen peroxide obtained from experiment compared to that stated on the bottle were likely different. Gas Laws - 6

7 Pre-Laboratory Exercise 1. Write the balanced chemical equation for the decomposition of aqueous hydrogen peroxide, H 2 O Research details about hydrogen peroxide to answer the following question. Why is commercially available hydrogen peroxide stored in brown bottles 3. A student performed an experiment in which they collected hydrogen gas from the reaction of magnesium with hydrochloric acid using exactly the same apparatus that you will be using in today s lab. (See p. 201 of Tro 2 nd ed.). The reaction is: The data they collected are shown below: Mg(s) + 2 HCl(aq) H 2 (g) + MgCl 2 (aq) Mass of Mg Volume of HCl Volume of H 2 collected Temperature of H 2 collected Barometric pressure g 10 ml 57.5 ml 22 o C inhg a. Barometric pressure in atmospheres? b. Water vapor pressure in atmospheres? c. Pressure of the hydrogen in atmospheres? d. Volume of hydrogen in liters? Gas Laws - 7

8 e. Temperature in Kelvin? f. Moles of hydrogen collected? g. Moles of hydrochloric acid reacted? h. Molarity of hydrochloric acid? i. If the bottle claimed that the molarity was M, what is the percent error? Gas Laws - 8

9 Optional Laboratory Exercise Using the procedure and data sheet of this lab as an example and using the % hydrogen peroxide determined experimentally, develop a procedure and data sheet to determine the Gas Law Constant, R. Now perform the experiment twice. Procedure: Gas Laws - 9

10 Data Sheet: Gas Laws - 10

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