Materials Included In Kit

Additional Materials Required

Prelab Preparation

Activity D. Charles’s Law—Effect of Temperature on the Volume of a Gas

Water Baths: Prepare water baths at different temperatures as follows:

1. Add crushed ice and a small amount of water (total volume 259 mL) to a 400-mL beaker, followed by about 10 scoopfuls of salt to prepare a bath between –10 to 15 ºC.
2. Add crushed ice and water to a 400-mL beaker for a bath at about 0 ºC.
3. Add 250 mL of water to a 400-mL beaker and heat it on a hot plate on a low setting to prepare a hot water bath at around 60 to 65 ºC.
4. Instruct students to add hot water or ice as needed during the course of the experiment to maintain the temperature of each bath within ±5 ºC of the desired temperature.

Safety Precautions

Ammonium hydroxide solution is toxic by ingestion and inhalation and is corrosive to body tissue. Phenolphthalein is an alcohol-based solution and is flammable. Hot objects and escaping steam can cause severe 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 60–80 psi. At very high pressures, the bottle might split, but will not shatter. Wear chemical splash goggles whenever working with chemicals, heat or glassware in the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information. Remind students to wash hands with soap and water before leaving the lab.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. Excess ammonia solution may be neutralized for disposal according to Flinn Suggested Disposal Method #10. Excess phenolphthalein may be disposed of according to Flinn Suggested Disposal Method #18b.

Lab Hints

• For best results, set up two stations for each activity throughout the lab. This will allow eight groups of students to rotate through four activity stations in a 50-minute lab period, if needed. A double lab period (two 50-minute class periods) will allow time both for a review of basic gas laws before lab and for a collaborative class discussion after lab.
• The activities may be completed in any order. Also, since each activity is a self-contained unit, the experiment may be set up with as many or as few of the activities as the teacher desires. Students should need only 7–8 minutes per station—keep the pace fairly brisk to avoid dawdling. Questions in the Calculations and Analysis section may be answered during downtime between stations.
• Prelab preparation is an essential component of lab safety, and it is also critical for success in the lab. (Standing in front of the lab station is not a good time for students to be reading the activity for the first time.) Having students complete the written prelab assignment for this lab will help teachers ensure that students are prepared for and can work safely in the lab.
• In Activity A, the indicator color change provides a visual reminder how quickly toxic gases can travel. Since many gases cannot be smelled or seen, this activity reinforces two important safety practices—work in a well ventilated room, and carry out any reactions that may produce toxic fumes in a fume hood.
• In Activity C, if the bottle leaks air around the cap when pressurized, remove the cap assembly and put additional petroleum jelly around the inside seal of the cap.
• The best results in the Boyle’s Law activity are obtained using a bicycle pump with an attached pressure gauge. If one cannot be obtained with student help, the experiment can be done using common automobile tire gauges. Tire gauges come in many shapes and sizes. The best gauges for this activity are those with an attached pressure dial or digital readout rather than a “pop-out” sliding scale. The units shown on some pressure gauges may be psig (gauge pressure per square inch).
• Students may need instruction in how to use a tire gauge. The pressure scale on a tire gauge is marked in units of pounds per square inch (psi). The scale starts at zero when the gauge is exposed to the surrounding air. This means that the total pressure is equal to the gauge pressure plus the pressure of the surrounding air.
• For best results, use a barometer to measure the local barometric pressure. The National Weather Service website reports corrected sea-level air pressures. Note that these are not actual barometric pressure readings. Meteorologists convert station pressure values to what they would be had they been taken at sea-level. The following equation can be used to recalculate the barometric pressure (in inches Hg) from the reported sea-level pressure (in inches Hg). Elevation, which must be in meters, can be obtained from geological survey sites.

Barometric pressure = sea-level pressure – (elevation/312 m)

On the syringes, the black rubber seals have two “seal lines.” Make sure students are consistent in where they measure the volume!

• In Activity D, teachers who have access to computer-based graphing programs may want to schedule additional time for students to graph and analyze their data. The graphs will be more precise than hand-drawn graphs and allow students to obtain more accurate estimates of absolute zero. The actual temperatures of the water baths are not important as long as students measure volumes over a wide enough temperature range. The difference between any two baths should be at least 15–20 ºC.

Teacher Tips

• Many students have bicycle pumps or tire gauges they are willing to loan to the classroom for a short time in the interest of science (and of course, extra credit). Ask your students for assistance. Electric air pumps are common items in vocational education or mechanical arts classrooms. If using an electric pump, the teacher should inflate the pressure bottles ahead of time for each group. The students may then use tire gauges to measure the pressure.
• Activity A is a modification of the Graham’s law demonstration that uses hydrochloric acid and ammonium hydroxide to produce a visible cloud of ammonium chloride (NH₄Cl) particles. This activity uses one gas and a stationary indicator. Phenolphthalein is an acid–base indicator and turns red when exposed to a base. “Ammonium hydroxide” is a concentrated solution of ammonia gas and water. Ammonia is a base and very volatile.
• What does a pressure of 14.7 psi feel like? The “Atmosphere Bar” available from Flinn Scientific (AP5882) is a 52-inch steel bar that weighs 14.7 lbs., with a base of one inch square.
• In Activity C, do not exceed the maximum pressure recommended in the Procedure section. It was found that the graph of P versus 1/V became non-linear as the pressure bottle was pressurized above about 70 psi (total pressure 85 psi). This can be used as a teaching point—deviations from ideal gas behavior are more important at higher pressures. Many textbooks show graphs of real versus ideal gas behavior as a function of pressure. For many gases, deviations from ideal behavior become significant at pressures greater than about 200 psi. Even at modest pressures, however, small deviations are common in the P x V “constant.” Some of this deviation may also be due to a change in temperature. Compressing the gas will increase the temperature of the gas.
• Jacques Charles was at least partially inspired by his interest in hot-air ballooning to study the properties of gases. If it worked for him, it may work for your students as well! Inspire your students to learn more about the properties of gases with the hot-air balloon activity kit “Up, Up and Away” (Flinn Catalog No. AP6310). The kit contains enough materials for 15 pairs of students to construct and launch their own giant hot-air balloon. Students will learn how hot-air balloons lift from the ground, stay aloft, and eventually descend. The volume of a gas is always proportional to temperature. It is directly proportional, however, only to the absolute temperature. Thus, in gas law calculations temperature must always be reported in kelvins.

Answers to Prelab Questions

In Activity A, the ability of gas molecules to diffuse will be studied by observing the reaction of ammonia with phenolphthalein, an acid–base indicator.

In Activity B, the force exerted by the atmosphere will be studied by observing its effect on a partially evacuated soda can.

In Activity C, the relationship between the volume and pressure of a gas will be determined by performing a new version of Boyle’s experiment, using a syringe and a pressurized soda bottle. As the pressure in the bottle is varied, the changes in the volume of the air in the syringe will be measured. Graphing of these two variables will yield their relationship.

In Activity D, the relationship between the volume and temperature of a gas (Charles’ law) will be discovered by measuring the volume of air in a sealed syringe at four different temperatures. The syringe will be placed in different water baths at –15 °C, 0 °C and 55 °C. As the syringe is placed in each bath, the changes in the volume of the air in the syringe will be measured. Graphing of these two variables will yield their relationship.

Sample Data

Activity A. Diffusion of Gases

Observations and Analysis

1. Describe the initial color and appearance of each solution and any changes that were observed when the Petri dish was covered.

   Initially, both solutions are clear. When the Petri dish is covered, the phenolphthalein solution turns a hot pink or bright red within minutes.

2. What compound was responsible for the color change observed in the phenolphthalein solution? Assuming that none of the liquids was spilled or contacted each other in any other way, how did this compound “travel” to the indicator?

   Ammonia, NH₃, a weak base, is responsible for the color change. Ammonia volatilizes from the ammonium hydroxide solution and subsequently dissolves in the phenolphthalein solution.

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.

   The indicator color change shows visually how quickly gases can travel.

   NH₃(g) + H₂O(l) → NH₄⁺(aq) + OH^–(aq)

4. What evidence does this demonstration provide that gas molecules are moving continuously about and randomly colliding with nearby walls and surfaces?

   Since the two solutions are physically separated from each other, the only way for the gas molecules to dissolve in the phenolphthalein solution is for the gas molecules to move all through the space above both liquids.

5. Describe two observations from daily life that also show us that gas molecules are able to move randomly through a “container.”

   Smells coming from the kitchen. Hot air moving into a cold room.

1. Describe your observations; be specific. What happened when the can was heated? When it was plunged into the water bath?

   The pop can was filled with water, then heated for several minutes to boil the water and produce steam. Once steam flowed out of the can, the can was quickly inverted and plunged into a pan of cold water. The sides of the can started to collapse inward.

2. What “force” caused the can to collapse inward on itself?

   The greater air pressure surrounding the can acted as a force on the outside of the can.

3. What “drove” the air out of the can as it was heated?

   The steam from the boiling water inside the can.

4. Why was there less air pressure inside the can after it was quickly cooled in the water “bath”?

   With the air in the can driven out, the can was filled with mostly water vapor. Plunging the can into cold water caused the water vapor to condense. This drastically reduced the pressure inside the can, allowing the much higher air pressure outside to crush the can.

*See Post-Lab Calculation 2.
†See Post-Lab Calculation 5.

Answers to Questions

Activity C. Boyle’s Law in a Bottle

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 = 760 mm Hg = 29.92 in Hg = 14.7 psi.
   {12545_Data_Equation_1}
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 in psi inside the pressure bottle. Record the total pressure in the table.

   Sample calculation for trial 1: Total pressure = 42 psi + 15 psi = 57 psi
   Refer to the Data and Results Table for the results of other calculations.

3. 1. Identify the independent and the dependent variables in this experiment.

      The independent variable is pressure and the dependent variable is volume.

   2. 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 units) and to give the graph a title.
      {12545_Data_Figure_3}
   3. 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.

      The graph is curved. The volume decreases as the pressure increases. At first, there is a sharp reduction in the volume as the pressure increases. The decrease in volume then becomes more gradual and the volume appears to level off as the pressure increases further. Mathematically, the shape of the curve is described as hyperbolic. A hyperbolic curve of this type is obtained when there is an inverse relationship between two variables (e.g., y = 7/x). See the graph for the best-fit curved line through the data.

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 ∝ 1/V. Calculate the value of 1/V for each volume measurement and enter the results in the table.

   Sample calculation for trial 1: V = 2.0 mL
                                                1/V = 0.50 mL^–1
   Refer to the Sample Data and Results Table for the results of the other calculations.

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.
   {12545_Data_Figure_4}
6. Another way of expressing an inverse relationship between two variables (P ∝ 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” and the average deviation. (Rounded to two significant figures.)
   {12545_Data_Table_3}

1. 1. Identify the independent and the dependent variables in this experiment.

      Temperature is the independent variable and volume is the dependent variable.

   2. 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.
      {12545_Answers_Figure_5}
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.

   A straight line graph implies that the volume of a gas is directly proportional to its temperature.

3. For each of the three temperatures in this experiment, calculate the value of the volume/temperature (in °C) ratio. How do these ratios compare with one another?

   Calculation for trial 1: V = 13.6 mL, T = –14 ºC
                                      V/T (ºC) = –0.97 mL/ºC
   Refer to the Sample Data and Results Table for the results of the other calculations.
   Numbers vary widely. At 0 ºC no number can be determined.

4. 1. Convert each of the temperature readings in this experiment to absolute temperature (kelvins, K).

      Calculation for trial 1: T ºC = –14 ºC
                                         T(K) = T (ºC) + 273.15 = 259 K
      Refer to the Sample Data and Results Table for the results of the other calculations.

   2. 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?

      Calculation for trial 1: V = 13.6 mL, T = 259 K
                                        V/T (ºC) = 5.3 x 10^–2 mL/K
      Refer to the Sample Data and Results Table for the results of the other calculations.
      The ratios are basically constant.

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?

   Volume is only directly proportional to temperature in kelvins.

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? Or, if the amount of gas in the syringe was different? Explain in terms of the KMT and the amount of empty space in gas.

   Since most of the volume of a gas is empty space, the relative size (molecular weight or volume) of a gas particle does not have an appreciable effect on the overall volume of the gas sample. The results using a different gas would be identical.

References

Special thanks to Patricia Mason (retired) Delphi Community H.S., Delphi, IN, and to Kathy Kitzmann, Mercy H.S. Farmington Hills, MI, for providing Flinn Scientific with general ideas and specific activities for “activity station” lab kits.

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 four “mini-lab” activities to study the properties of gases and to investigate the relationships among the four measurable gas properties—temperature, pressure, volume and the number of moles.

Concepts

• Boyle’s law
• Atmospheric pressure
• Charles’s law
• Kinetic-molecular theory
• Pressure
• Diffusion
• Temperature

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.

Activity A. Diffusion of Gases
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?

Activity B. Crush the Can
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 (15 lb/in²).

Activity C. Boyle’s Law in a Bottle
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 a safe and environmentally friendly manner!

Activity D. Charles’s Law—Effect of Temperature on the Volume of a Gas
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.

Experiment Overview

The purpose of this “activity-stations 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. Four mini-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.

Activity A. Diffusion of Gases
Activity B. Crush the Can 
Activity C. Boyle’s Law in a Bottle
Activity D. Charles’s Law—Effect of Temperature on the Volume of a Gas

Materials

Prelab Questions

Read the Background material and Procedure for each activity A–D. Write a brief, 1- to 2-sentence description of each experiment. For example: In activity A, the ability of gas molecules to diffuse will be studied by observing the reaction of ammonia with an acid–base indicator.

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 severe 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 60–80 psi. 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.

Procedure

Activity A. Diffusion of Gases

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 ammonium hydroxide 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 phenolphthlein 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 observations and analysis questions.

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 inches of water. Set the pan next to the hot plate. Warning: Notice the “HOT” sign in front of 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 trash. Do not move the “HOT” sign!

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 1-L pressure bottle.
4. Run a small bead of petroleum jelly around the rim of the 1-L bottle and in the inside of the special tire valve cap.
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. Warning: Do NOT exceed 60–80 psi. Using a manual tire pump is a safety feature—it is very difficult to pump more than about 70 psi into the 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, as shown in Figure 1.
   {12545_Procedure_Figure_1}
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 values 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.

1. Remove the tip cap, if necessary, from the 30-mL 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 (see Activity C).
5. Place the syringe in the saltwater–ice bath (–15 to –20 °C) and submerge the syringe just to the bottom of the plunger as shown in Figure 2. Measure and record the temperature of the bath in the Data and Results Table.
   {12545_Procedure_Figure_2}
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 (0–5 °C) as shown in Figure 2. 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) as shown in Figure 2. 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.

Student Worksheet PDF

12545_Student1.pdf