AP+Chem+Lab+2

Properties of Gases and the Gas Laws
 * Introduction:** The properties of gases and the gas laws are important in many science and engineering applications, including physiology, meteorology, scuba diving, and even hot-air ballooning. Boyle’s law, for example, is demonstrated with every breath you take. Use this set of five activities to study the properties of gases and to investigate the relationship among the four measurable gas properties – temperature, pressure, volume, and the number of moles.

· Boyle’s Law · Kinetic- molecular theory · Atmospheric pressure · Pressure · Charles’s Law · Diffusion
 * Concepts:**


 * Background:** Pressure is defined as force divided by area. According to the kinetic molecular theory, the particles in a gas are in constant, random motion. When the gas molecules confined to any container collide with the “walls” of the container, the force of the resulting collisions causes the gas to exert a pressure against the container. The pressure of the gas is related to the total force exerted by the individual collisions divided by the area over which the collisions occur.

A. Diffusion of gas molecules B. Crush the can C. Boyles’s law in a bottle D. Effect of Charles’s law E. Volume vs pressure (class experiment)
 * Experiment Overview:** The purpose of this activities lab is to investigate the properties of gases, derive the mathematical relationships among the gas variables, and explain the behavior of gases using the kinetic molecular theory. Five lab activities are set up around the classroom. Each activity focuses on a different relationship among the gas properties and is a self contained unit.


 * Safety Precautions**: //Ammonia solution is toxic by ingestion and inhalation and is corrosive to body tissue. Phenolphthalein solution contains alcohol and is a flammable liquid. Hot objects and escaping steam can cause sever burns. Handle hot objects with beaker tongs and do not place your hands in the steam. The pressure bottle is safe if used properly. The bottle should not be inflated above 100psi. At very high pressures, the bottle might split, but it will not shatter. Do not use a thermometer as a stirring rod. Wear chemical splash goggles whenever working with chemicals, heat or glassware in the laboratory. Wash hands with soap and water before leaving the lab.//

The fact that many gases are colorless and odorless and cannot be seen may give us a misleading image of the properties of gas molecules. An accurate “molecular” picture of gases would show small particles very far away from each other, swarming about in great, rapid, and random motion, and colliding frequently with whatever “walls” the gas may be confined to. What evidence do we have for this “kinetic” picture of gas molecules and the motion of molecules?
 * //__Part 1: Diffusion of Gases__//**

· Ammonia solution, NH3, 6M, 2 ml  · Distilled water and wash bottle · Phenolphthalein indicator solution, 1 ml  · Beral-type pipets, graduated, 3 · White paper
 * Materials**

1. Working in a fume hood, place the divided Petri dish on a sheet of white paper. 2. Using a graduated, Beral-type pipet, add 2 ml of ammonia solution to one compartment of the divided Petri dish. 3. Using a clean pipet for each liquid, mix 1 ml of phenolphthalein solution and 1 ml of distilled water in a second compartment in the Petri dish. Place the cover on the divided Petri dish. Note: Do not allow any phenolphthalein to drip into the ammonia compartment. 4. Observe any changes in the color and appearance of each solution in the Petri dish. 5. Wearing gloves, rinse the contents of the Petri dish into the large waste beaker provided at the activity station. 6. Answer the questions in the Post lab section
 * Procedure:**

1. Describe the initial color and appearance of each solution and any changes that were observed when the Petri dish was covered. 2. What compound was responsible for the color change observed in the phenolphthalein solution? Assuming that none of the liquids were spilled or contacted each other in any other way, how did this compound “travel” to the indicator? 3. What is the role of the phenolphthalein “indicator” in this demonstration? Write an equation for the reaction of ammonia gas with water that explains the indicator color change. 4. What evidence does this demonstration provide that gas molecules are moving continuously about and randomly colliding with nearby walls and surfaces? 5. Describe two observations from daily life that also show us that gas molecules are able to move randomly through a “container”.
 * Post lab questions:**

Pressure- we all feel it! But what is it? In the case of the surrounding air, the pressure it exerts is a force, a surprisingly strong force. Use this “pressure-packed” activity to prove that air is a force to be reckoned with. When the water inside the can boils, it is converted to steam, which drives the air out of the can. When the can is then inverted and quickly cooled in a container of water, any steam contained in the can condenses back to a liquid. Since there are fewer air or gas molecules remaining in the can than there were originally, the gas pressure inside the can after it cools is substantially lower than its original value. The external air pressure “pressing” on the outside of the can is normal atmospheric pressure (15lb/in2).
 * //__Part 2: Crush the Can__//**

· Aluminum pie pan · Beaker tongs · Graduated cylinder, 25 ml  · Hot plate · Soda can, 12 oz, aluminum empty · Water
 * Materials:**

1. Rinse out an empty, 12 oz aluminum soda can, and then using the graduated cylinder, add about 15-20 ml of water to the can. 2. Fill the pie pan with 2-3 inces of water. Set the pan next to the hot plate. 3. Holding the aluminum can with a pair of beaker tongs, heat the can on the hot plate at a high setting until the water comes to a boil and steam is observed coming out of the can. 4. After steam has steadily come out of the can for 30-60 seconds, remove the can from the hot plate. Immediately turn the can upside down and plunge the open end of the can into the pan filled with water. 5. After the can is cool to the touch discard the crushed can in the recycling.
 * Procedure:**

1. Describe your observations; be specific. What happened when the can was heated? When it was plunged into the water bath? 2. What “force” caused the can to collapse inward on itself? 3. What “drove” the air out of the can as it was heated? 4. Why was there less air pressure inside the can after it was quickly cooled in the water “bath”?
 * Post Lab Questions:**

More than 350 years ago, Robert Boyle used air trapped in a glass tube above a column of mercury to study the relationship between the volume and pressure of air. The purpose of this activity is to carry out a modern version of Boyle’s classic experiment, using only a syringe and a special, “pressurized” soda bottle. Discover Boyle’s law in an unsafe and unenvironmentally friendly manner!
 * //__Part 3: Boyle’s Law in a Bottle__//**

· Barometer · Bicycle pump with pressure gauge · Graph paper · Petroleum jelly, small amount · Pressure bottle, 1L with tire valve cap · Syringe, 10 ml, with syringe tip
 * Materials**

1. Using a barometer, measure and record the value of the local air pressure in the Data Table. 2. Remove the tip cap from the syringe and pull on the plunger to draw about 9 ml of air into the syringe. Replace the tip cap to seal the air inside the syringe. 3. Place the sealed syringe inside the 1L bottle. 4. Run a small bead of petroleum jelly around the rim of the 1L bottle. 5. Cap the bottle with the special tire-valve cap assembly. Tighten the cap securely. 6. Connect the tire valve to a bicycle pump. Pump air into the pressure bottle to obtain a pressure reading of 50-60 psi on the attached pressure gauge. Record in the Data Table the initial gauge pressure and volume in the syringe. Do NOT exceed 100psi!! Note: Using a manual tire pump is a safety feature – it is very difficult to pump more than about 70 psi into a pressure bottle by hand. 7. Loosen the connection between the valve and the pump to release a very small amount of air from the bottle – as soon as you see the plunger in the syringe begin to move, immediately retighten the connection between the tire-valve cap and the pump. 8. Measure the pressure using the attached pressure gauge and record the pressure to within +/- 1 psi in the Data and Results Table. 9. Measure the volume of air trapped in the syringe inside the pressure bottle and record the volume in the Data Table. Note: Measure the volume at the black insert rubber seal not at the inverted V-like projection. 10. Loosen the connection between the pressure bottle, tire valve, and the pump to release some air from the pressure bottle and reduce the gauge pressure by about 10 psi. Immediately retighten the connection between the tire valve and the pump. 11. Measure the new gauge pressure and the resulting volume of air inside the syringe and record both value in the Data Table. 12. Repeat steps 10 and 11 to measure the volume of gas at several different gauge pressures down to about 15 psi. It should be possible to obtain at least 5-6 pressure and volume measurements in this range. 13. Remove the tire valve from the pump and press down on the brass pin to release the excess pressure in the pressure bottle. Measure and record the final volume of air contained in the syringe at atmospheric pressure. Note: The gauge pressure is equal to zero at atmospheric pressure.
 * Procedure:**


 * Data and Results Table**
 * Barometric Pressure ||  ||   ||   ||
 * Gauge Pressure ||  Volume of Air in syringe  ||  Total Pressure  ||  1/V  ||  PxV  ||

1. Convert the local barometric pressure to psi units and enter the value to the nearest psi in the Data and Results Table. (1 atm = 760mmHg = 29.92 in Hg = 14.7 psi) 2. The tire pressure gauge measures the relative pressure in psi above atmospheric pressure. For each pressure reading in the Data and Results Table, add the local barometric pressure, in psi, to the gauge pressure to determine the total pressure of air inside the pressure bottle. Record the total pressure in the table. 3. a. Identify the independent and the dependent variable in this experiment. b. Plot a graph of the dependent variable on the y-axis versus the independent variable on the x-axis. Choose a suitable scale for each axis so that the data points fill the graph as completely as possible. Remember to label each axis. (including the units) and give the graph a title. c. Describe the shape of the graph. Draw a best-fit straight line or curve, whichever seems appropriate, to illustrate how the volume of a gas changes as the pressure is varied. 4. The relationship between pressure and volume is called an inverse relationship – the volume of air trapped inside the syringe decreases as the pressure increases. This relationship may be expressed mathematically as P µ 1/V. Calculate the value of 1/V for each volume measurement and enter the results in the table. 5. Plot a graph of pressure on the y-axis versus 1/V on the x-axis and draw a best-fit straight line through the data points. Choose a suitable scale for each axis. Remember to label each axis and to give the graph a title. 6. Another way of expressing an inverse relationship between two variables (P µ 1/V) is to say that the mathematical product of the two variables is a constant. (P x V = constant). Multiply the total pressure times the volume for each set of data points. Calculate the average value of the P x V “constant”.
 * Calculations and Analysis**

Charles’s law describes the relationship between the temperature of a gas and its volume. In order to understand this relationship, we must imagine what happens to the particles in a gas when the gas is heated or cooled. The temperature of a gas measures the average kinetic energy of the moving gas particles – how fast they are moving. When a gas is heated, the average kinetic energy of the particles increases and they move faster. When a gas is cooled, the average kinetic energy of the particles decreases and they move slower.
 * //__Part 4: Charles’s Law- Effect of Temperature on the Volume of a Gas__//**

· Temperature · Charles’s Law · Kinetic-Molecular Theory
 * Concepts:**

· Hot water · Ice · Graph paper · Paper towels · Petroleum jelly · Stirring rods, 3 · Syringe, 30 ml with syringe tip cap · Thermometers · Wood splint · Temperature baths o Salt-ice water, -15 to -20 ° C  o Ice water, 0 to 5 ° C   o Hot water, 60 to 65 ° C
 * Materials:**

1. Remove the tip cap, if necessary, from the 30ml syringe, and take the plunger out of the syringe. Place a very small dab of petroleum jelly on the black rubber gasket, and spread the petroleum jelly out in a thin layer on the surface of the gasket using a wood splint. 2. Place the plunger back in the syringe and draw the syringe to about one-half full with air. Seal the syringe with the syringe tip cap. 3. Place the syringe on the lab table and measure the ambient air temperature around the syringe. Let the thermometer equilibrate in air for 1-2 minutes before measuring the temperature. Record the ambient “room temperature” reading in the Data and Results Table. 4. Measure and record the precise volume of air in the syringe at room temperature. 5. Place the syringe in the saltwater-ice bath (-15 to -20 ° ) and submerge the syringe just to the bottom of the plunger. Measure and record the temperature of the bath in the Data and Results Table. 6. Hold the syringe in the saltwater-ice bath for at least 2 minutes, then quickly push down on the plunger once and release. Measure and record the precise volume of air in the syringe when the syringe stops moving. 7. Remove the syringe from the saltwater-ice bath and place the syringe in the ice-water bath (0 to -5 ° C). Measure and record the temperature of the ice-water bath. 8. After two minutes, measure and record the volume of air in the syringe. 9. Remove the syringe from the ice-water bath and place the syringe in the hot water bath (60 -65 ° C). Measure and record the temperature of the hot water bath. 10. After two minutes, measure and record the precise volume of air in the syringe. 11. Remove the syringe from the hot water bath and remove the tip cap and plunger. Wipe the plunger and gasket with a paper towel.
 * Procedure:** There are three water baths set up at this activity station – notice the temperature range specified for each bath in the Materials section. Add hot water or ice and stir as needed during the course of the activity to maintain the average temperature of each bath in the desired range. Note: Do not use thermometers as stirring rods.

(ml/C) || Absolute Temp., K || Volume/T (ml/K) ||
 * Data and Results Table**
 * Water Bath || Temperature || Volume of Air in Syringe, ml || Volume/T
 * Saltwater-ice ||  ||   ||   ||   ||   ||
 * Ice Water ||  ||   ||   ||   ||   ||
 * Room Temp ||  ||   ||   ||   ||   ||
 * Hot Water ||  ||   ||   ||   ||   ||

1. a. Identify the independent and the dependent variable in this experiment b. Plot a graph of the dependent variable on the y-axis versus the independent variable on the x-axis. Choose a suitable scale for each axis so that the data points fill the graph as completely as possible. Remember to label each axis, including the units and to give the graph a title. 2. Draw a best fit straight line through the data points on the graph. Describe the mathematical relationship between the temperature and volume of a gas. 3. For each of the four temperatures in this experiment, calculate the value of the volume/temperature (in C) ratio. How do these ratios compare with one another? 4. a. Convert each of the temperature measurements in this experiment to absolute temperature (K). b. Calculate the value of the volume/temperature (in K) ratio for each of the four temperatures in this experiment. How do these ratios compare with one another? 5. Which volume/temperature ratio (in C or K) appears to be more constant? Saying that the ratio of two variables is a constant is to say that the two variables are directly proportional to each other. Why is it important to specify absolute temperature (in K) when stating Charles’s Law. 6. According to the kinetic-molecular theory, the volume of the gas particles is extremely small compared to the volume the gas occupies – most of the volume of gas is “empty space.” Based on this theory, does Charles’s law depend on the identity of the gas? Would the results in this experiment have been different if different gases had been used in the syringe? On the amount of gas in the syringe? Explain in terms of the KMT and the amount of empty space in gas.
 * Calculations and Analysis**