The Ideal Gas Law. R = x 10-2 Latm. **Note the complex units for the gas constant, R. These units are part of the gas constant, R.

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1 Chem 110 Lab The Ideal Gas Law Clark College Name: Partner s Name: LEARNING OBJECTIVES After completing this experiment, you should feel comfortable with: Working with glassware. Proper waste disposal of hazardous waste. Observing phase changes and recording temperature Using the ideal gas law to calculate molar mass PART I. A. Introduction The behavior of gases is described by a number of laws. The three gas laws listed below show the relationship between gas volume (V) and the absolute temperature (K); gas volume and pressure (P); and gas volume and quantity of gas (n). The symbol means directly proportional to V T Charles Law V ; T (n, P constant) V 1/P Boyles Law V ; P (n, T constant) V n Avogadro s Postulate (n = # moles) V ; n (P, T constant) Combining these laws, we come up with: The ideal gas law: PV = nrt It was determined, for 1.00 mole of any gas, at K, that the value of R, gas law constant is equal to R = x 10-2 Latm molek **Note the complex units for the gas constant, R. These units are part of the gas constant, R. In this experiment, you will use the ideal gas law (PV=nRT) to find the number of moles of gas present in a 250 ml erlenmeyer flask. You will assume a vaporized sample will act ideally, driving out all the air and totally filling the erlenmeyer flask with the unknown gas. You will measure the internal volume of the flask (thus the volume the unknown gas occupies), the temperature at which all the gas is vaporized, and the pressure inside the flask (which will be equal to atmospheric pressure). Those will be the values for V, T and P. Since you know R, the gas constant, you will be able to find n, the moles of the unknown gas. The flask of gas is then cooled and the gas condenses to a liquid. You will measure the mass of the condensate so that you can find the molar mass (g/mol) of the unknown. During the experiment handle the erlenmeyer flask with your fingers as little as possible. (Some moisture/oils on your fingers could be transferred to the flask, increasing its mass). The Ideal Gas Law F09 AEM Page 1 of 5

2 B. Sources of Error 1. When the gas is condensed to measure its mass, you are neglecting the fact that not all of the displaced air can re-enter the flask. This results in a slightly low mass for the gas. 2. If you continue to heat the container after all the visible liquid has vaporized, lower mass air will gradually displace the gas leading to a low mass determination. 3. Insufficient heating fails to vaporize all the liquid to gas leading to a high mass determination. 4. Failing to completely remove water from the glass exterior will give a high mass determination. C. Experimental Procedure 1. Fill a clean 1000 or 800 ml beaker about half full with tap water. Add ~4 boiling stones. 2. Using a hot plate, heat the water in the beaker until it boils. (Continue with steps 3 & 4 while you wait). 3. Obtain a clean, dry 250-mL Erlenmeyer flask. Mass the flask with a square piece of aluminum foil to the nearest 0.01 gram, and record the mass on the Data Sheet. 4. Use a graduated cylinder to measure approximately 4 ml of the unknown liquid. Add the unknown to the flask and cover the top of the flask with the aluminum foil. Crimp the foil around the top of the flask to make a small cap. Make a small pin hole in the top of the foil to allow excess vapor to escape. See Figure After the water reaches a full boil, carefully remove the beaker from the hot plate. Use tongs to lower the Erlenmeyer flask into the beaker very carefully so that the hot water comes as far up the neck of the flask as possible (do not get water into the foil.). When lowering the flask, tip it to pool the unknown liquid on one side (see Figure 2). Watch the liquid closely & continuously. Figure 2 6. While holding the flask mostly submerged, carefully place a thermometer in the beaker. Continue watching the unknown! The Ideal Gas Law Revised F09 AEM Page 2 of 5

3 7. Leave the flask immersed in the water until all the liquid in the flask has vaporized (1-5 minutes). Watch carefully! At the moment when all the liquid has turned to vapor, one person should read the thermometer and record the temperature (as precisely as possible sig figs!) while the other removes the flask from the water bath. o T = 8. Cool the lower part of the flask under running tap water for one minute. As the flask cools, the gas that fills the flask will condense back into liquid. 9. Wipe the outside of the flask completely dry. Weigh the flask to the nearest 0.01g. Record the mass in the space provided on the data sheet. **Waste Disposal: Pour the unknown liquid that remains in your Erlenmeyer flask into the Organic Liquid Waste container in the hood. 10. Determine the volume of the Erlenmeyer flask: Fill the Erlenmeyer flask to the brim with tap water. Accurately determine the volume by pouring the water (approximately half at a time) into a large, empty (250-mL) graduated cylinder. Record the two volumes and add them together to find the total volume. 11. Using the barometer value given to you from the laboratory wall, record the pressure on the space provided on the Data Sheet. Convert the pressure into atmospheres. TABLE I. DATA SHEET 1. Mass of clean, dry and empty flask and foil g 2. Final mass of flask, foil, and unknown liquid (after cooling) g 3. Mass of unknown liquid (#2 - #1) g 4. Temperature of hot water (at the moment the unknown vaporized) Conversion from C to K: C K 5. Barometric Pressure mmhg atm Conversion from mmhg to atm: 6. Volume of flask ml L Conversion from ml to L: The Ideal Gas Law Revised F09 AEM Page 3 of 5

4 D. Experimental Calculations (always show your work) 1. Moles of gas at vaporization temperature moles Equation and Gas Constant: x 10 Latm R = molek PV = nrt 2. Molar mass of the unknown grams/mole Mass divided by moles. PART II. A. Calculation of Experimental Error If the gas is C 6H 14, what is the % error between your experimental value for molar mass and the true molar mass? Solution: To calculate the % error: % error = ( your value - accepted value) x 100% accepted value 1. Calculate the molar mass of C 6H 14. This is the accepted value. 2. Subtract the accepted molar mass of C 6H 14 (#1 above) from your calculated molar mass (#2 experimental calculations). 3. Divide this value (#2 above) by the calculated molar mass of C 6H 14 (#1 above). 4. Multiply the value in #3 (above) by 100. The Ideal Gas Law Revised F09 AEM Page 4 of 5

5 B. Experimental Questions and Problem Solving 1. What factors might cause your experimentally determined molar mass value to be: a. too high? b. too low? 2. What is the pressure, in atm, on the inside of a jar that contains 88.0 grams CO 2? The temperature is 20.0 o C and the volume of the jar is 1.00 liter. The Ideal Gas Law Revised F09 AEM Page 5 of 5

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