# Please Return! Gas Laws: LAB Directions

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1 Gas Laws: LAB Directions Please Return! Background Applications of the gas laws are important in physiology, meteorology, scuba diving, tire manufacturing, even hot-air ballooning. Robert Boyle built a simple apparatus to measure the relationship between the pressure and volume of air. The apparatus consisted of a J-shaped tube that was sealed at one end and open to the atmosphere at the other end. A sample of air was trapped in the sealed end by pouring mercury into the tube. In the beginning of the experiment, the height of the mercury column was equal in the two sides of the tube. The pressure of the air trapped in the sealed end was equal to that of the surrounding air, equivalent to 29.9 inches (760mm) of mercury. When Boyle added more mercury to the open end of the rube, the air trapped in the sealed end was compressed into a smaller volume. The difference in height of the two columns of mercury (Δh) was due to the additional pressure exerted by the compressed air compared to the surrounding air. Boyle found that when the volume of trapped air was reduced to one-half its original volume, the additional height of the column of mercury in the open end of the tube measured 29.9 inches. The pressure exerted by the compressed air was twice as great as atmospheric pressure. The mathematical relationship between the volume of the air and the pressure it exerts was confirmed through as series of measurements. This became known as Boyle s law: P 1 V 1 =P 2 V 2 Charles s law, which describes how the volume of a gas changes as it is heated or cooled, was inspired by one of these applications, the first sensational flight in the history of hot-air ballooning. The world s first manned hot-air balloon flight took place in France in A few months later, Jacques Alexandre César Charles, professor of physics at the Sorbonne University in France, made the first flight in a lighter- than-air (hydrogen) balloon, climbing to a height of 9000 meters. Motivated by twin interests in science and ballooning, Charles not only made more flights, he also studied the factors that influence the flight of a hot air balloon. A hot air balloon works on the principle that the volume of a gas expands when heated. When the gas inside a balloon is heated, the gas expands and the hot air balloon becomes less dense than the air it displaces, causing the balloon to rise and float in the atmosphere. This became known as Charles law: Charles investigated this principle in the laboratory by measuring the increase in volume of a fixed amount of air when it was heated at a constant pressure. Charles s unpublished work was taken up by another French scientist and balloon enthusiast, Joseph-Louis Gay-Lussac, who studied the expansion of oxygen, hydrogen, nitrogen, carbon dioxide, nitrous oxide, and ammonia. In 1802 Gay-Lussac concluded that all of these gases expanded equally when heated from 0 to 100 C the volume of a gas is proportional to its temperature and does not depend on the nature of the gas. Evetually, this was developed in Gay-Lussac s law: This conclusion led to the development of more precise air thermometers for measuring temperature. Although very precise, air thermometers were still arbitrary there was no absolute basis for any of the numbers on the scale. In 1848 the British mathematician William Thomson (knighted Lord Kelvin in 1866) proposed an absolute temperature scale based on the assumption that there must be a lower limit to temperature. He wrote: Infinite cold must correspond to a finite number of degrees of the air thermometer below zero, since if we push the strict principle of graduation sufficiently far, we should arrive at a point corresponding to the volume of air being reduced to nothing Thomson estimated a value for the infinite cold temperature on the Celsius scale. This temperature is known today as absolute zero and the temperature scale is called the Kelvin or absolute temperature scale. The units on this scale are equal in magnitude to the degree units on the Celsius scale and are called kelvins (abbreviated K), in honor of Lord Kelvin. 1

4 Gas Laws Lab Report Purpose: Summaries the purpose of this lab: Names: Period: Row(s): Hypothesis: A: B: C: D: Data and Results: Part A: 1. Observe and record what you see in your lab report for Part A, including what happens to the balloon in both environments. 2. Circle the variables you are dealing with: temperature pressure volume 3. Complete this: As increases, increases. 4. The relationship between these variables is (direct, inverse) Part B: 1. When I push the syringe, the marshmallow. 2. When I pull the syringe, the marshmallow. 3. Circle the variables you are dealing with: temperature pressure volume 4. Complete this: As increases, decreases. 5. The relationship between these variables is (direct, inverse) Part C: 1. Observe and record what you see. (include what happened to the can.) 2. Circle the variables you are dealing with: temperature pressure volume 3. Complete this: As decreased, decreased. 4. The relationship between these variables is (direct, inverse) 4

5 Part D: Data tables: include units! Circumference of balloon: cm Mass of system before releasing gas: Mass of system after puncturing balloon: Mass of gas (subtraction!) Room temperature (in Kelvin): Atmospheric Pressure (mm Hg): Pressure of water vapor at room temp Calculations for Part D: 1. The pressure inside the balloon is equal to atmospheric pressure, which includes both CO 2 and water vapor. Subtract the value for water vapor at this temperature from atmospheric pressure (Dalton s law of partial pressure P total of atmosphere = P H2O + P CO2 ). You now have the pressure of your gas. 2. Convert this to atmospheres (1atm = 760mmHg): 3. Using the circumference of the balloon, calculate the radius and then volume of the gas produced. Circumference = r Volume of a sphere = (4/3) л r 3 Convert cm 3 to ml then liters. (2л) 3. You now have all the values needed to solve for n with the ideal gas law. The value for R is liter-atm/mol K. (PV = nrt) Solve for n. Don t forget to convert to correct units as needed (see background information for units) Don t forget units in your answer!!! 4. The gas is in the balloon is CO 2. Using your value for n and the mass of the gas you generated, what is your value for the molar mass of CO 2? (Molar mass = g/mole) Analysis: 1. Which law do you think was demonstrated in a. Part A? b. Part B? c. Part C? d. Part D? 5

6 2. Explain, for Part A, why the balloon did what it did. (Relating to the variables and the particles of air). 3. Explain, for Part B, why, the marshmallow did what it did. 4. Explain, for Part C, why the can did what it did. 5. For part D, find the actual molar mass of CO 2 using the periodic table. 6. What is your percent error for your value of the molar mass of CO 2? your value periodic table value X 100 = periodic table value 7. What do you think are the biggest source of errors in this lab? 8. Look at your answer to the question above. Would these errors make your value come out too high or too low? Explain. 9. CO 2 is quite soluble in water. Would this fact make your value for molar mass too low or too high? Why? 10. If the balloon full of gas that you produced were brought out into a cold day (-20 C) and put under 3.5 atm of pressure, what would the new volume be? (show work!) 6

7 11. What is the volume of a sample of oxygen gas that contains 2.36 moles at 28 C and 12.0 atm pressure? g of solid CO 2 is put in an empty sealed 4.00 L container at a temperature of 300 K. When all the solid CO 2 becomes gas, what will be the pressure in the container? (Hint: use the molar mass of CO 2 to first convert grams to moles, then use the ideal gas law.) 13. A flask with a volume of 32.5L has a mass of 536.8g. At STP, the flask is filled with a gas and the mass of the flask is now g. What is the gas that replaced the air? (Hint: find moles, then molar mass of the gas) Conclusion: Write a conclusion including: what you learned, were your hypothesis correct of not, and likes, dislikes or improvements. (can write conclusion on this page) 7

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