This simple device will allow the experimental verification of the Boyle’s law relationship between the pressure and volume of a gas as well as the Charles’s law relationship between the temperature and volume of a gas.
- Properties of gases
- Boyle’s law
- Charles’s law
Beaker tongs, insulated gloves, or a hot vessel gripping device‡
Boyle’s Law device (metal weight hanger, syringe with plunger and tubing with pinch clamp)
Set of hook weights (100, 200, 500, 1000 g)†
Support (ring) stand†
*Materials included in kit.
†Boyle’s law experiment
‡Charles’s law experiment
Wear appropriate safety eyewear. Chemical splash goggles are recommended whenever working with chemicals, heat or glasswear. Check the syringe, plunger and pinch clamp. Replace any parts that are damaged or worn. Exercise caution with the boiling water bath as hot water and steam can cause burns.
Save the Boyle’s Law Apparatus for future use.
1. Take the plunger out of the syringe. Lubricate the entire side wall of the rubber plunger with silicone grease. Re-insert the plunger into the syringe. 2. Check that the rubber tubing with the pinch clamp is securely attached to the syringe tip.
Part 1. Boyle’s Law
- Prepare a data table to record the volume of air (in the syringe) versus the total mass (of the hook weights). Note: Create enough columns to run 4 trials plus a column for the average volume.
- Open the pinch clamp on the tubing of the Boyle’s law device. Draw up the plunger to fill the syringe with air to a volume of approximately 10 cc. Close the pinch clamp.
- Attach the buret clamp near the top of a support (ring) stand.
- Use the buret clamp to hang the Boyle’s Law device so that the metal hanger and the syringe tip/rubber tubing point downward (see Figure 1). Note: The flared end of the syringe should rest on the edges of the clamp jaws. Take care not to tighten the clamp too much which would exert pressure on the walls of the syringe.
- Before adding any weights to the metal hanger, record your first data point by reading the starting volume of air with no masses on the apparatus. Note: Before reading the scale on the syringe, pull the syringe out a small bit and allow it to return to equilibrium in order to overcome the static friction between the plunger and the syringe walls.
- Hang a hook weight on the metal hanger and record the mass of the weight (in grams) and the volume of air (in cc) as indicated on the scale of the syringe. A 200-g mass is a reasonable starting mass.
- Continue adding additional hook weights on the metal hanger, recording the total mass and volume of air after each addition. Gather data for as many masses as possible. Note: This device can hold weights up to approximately 2,000 g.
- Repeat this procedure (steps 5–7), if time allows, in order to gather data for three more trials. Average the volume data for each mass and record in the data table.
- Plot the data on a sheet of graph paper. Plot the total pressure in kPa (see Tips; mass is directly proportional to pressure here) on the independent x-axis and the average volume of air on the dependent y-axis. Do not plot your first data point (0 weights). This implies that there is 0 pressure, which is not true because atmospheric pressure is present at 101 kPa.
Part 2. Charles’s Law
- Heat 300–350 mL of water in a 400-mL beaker to near-boiling. Use beaker tongs, insulated gloves or a hot vessel gripping device to handle the beaker during the experiment. Note: The water level must be high enough to cover the portion of the syringe that holds the trapped volume of air.
- Prepare a data table to record the volume of air (in the syringe) versus the water temperature (in °C). Note: Create three columns for volume—“Volume In” for the reading after pressing the plunger in, “Volume Out” for the reading after pulling the plunger out and “Volume Average” for the average volume.
- Draw 5 cc of air into the syringe and close the pinch clamp.
- Using a thermometer, measure and record the near-boiling water temperature (90–100 °C) in your data table.
- Hold the Boyle’s Law device by the metal hanger with the syringe pointing downward. Immerse the portion of the syringe containing the trapped volume of air completely in the near-boiling water (see Figure 2). Wait a minute or so for the air in the cylinder to reach equilibrium and read the volume of air shown on the syringe scale. Record the temperature and volume of air in your data table. Note: To obtain a more accurate volume reading, first quickly push the plunger down into the cylinder and release it. Record this volume as “Volume In.” Next sharply pull the plunger outward and release it. Record this larger volume as “Volume Out.” Average these two measurements to give “Volume Average,” which helps to correct for some of the friction between the plunger and the wall of the cylinder. Be very careful not to splash hot water on yourself while manipulating the syringe.
- Use a few ice cubes to bring the water temperature down to ~50 °C. Stir the solution to make sure the temperature is uniform. Note: Some water may need to be removed. Exercise caution and use a hot vessel gripping devise when handling the hot beaker.
- Repeat steps 4 and 5 and record the data in your data table.
- Repeat steps 4–7 two more times. First bring the water temperature down to ~25 °C (room temperature), gather and record data. Then bring the water temperature down to near 0 °C, gather and record data.
- Plot the volume–temperature data on a separate sheet of graph paper. Plot the temperature in Kelvin on the independent x-axis and the average volume of air on the dependent y-axis.
- Grease plunger well. May need to grease before each trial, this will minimize friction in the syringe.
- The Charles’s Law experiment will give students a fairly qualitative view of the temperature–volume relationship. The volume of air in the syringe does not vary greatly in the temperature range from 0 to 100 °C; therefore, expect small volume variations of several milliliters.
- To calculate pressure in kPa:
- Convert grams to kilograms.
- Convert kilograms to Newtons (force) = mass x acceleration due to gravity, mass in kg x 9.8 m/sec2.
- Pressure = force/area and is expressed in units of Newtons/m2. Area is the area of the syringe opening (use: A = πr2, r = 7.5 x 10–3 m).
- Convert to kPa.
- Add to each, standard atmospheric pressure 101 kPa.
- To convert to Kelvin in Charles’s law experiment add 273 to °C.
Boyle’s Law Experiment
Answers to Questions
- The volume of air decreased as the mass (and thus pressure) increased.
- The graph is a curve downward, a hyperbolic curve.
- An inverse relationship exists between pressure and volume.
- See the Discussion and Tips sections.
- The volume of air decreased as the temperature decreased.
- The graph is a straight line.
- A direct or linear relationship exists between temperature and volume.
The kinetic molecular theory describes particles in a gas as far apart and in rapid, random motion. To study the properties of gases, measurements of three macroscopic variables of a given quantity of gas are needed — volume, temperature and pressure.
In 1660, Robert Boyle, a British scientist, performed an experiment that measured the volume of a trapped gas as the pressure on the gas changed, with temperature being held constant. He observed that when the temperature and the number of moles of a sample of gas are held constant, its volume is inversely proportional to the pressure applied. This is known as Boyle’s law. Volume (V) decreases with increasing pressure (P). Mathematically, this inverse proportionality may be expressed as
P • V = k (where k is a constant)
or alternately as P1
, where P1
are the initial pressure and volume of gas and P2
are the final pressure and volume. A plot of V versus P forms a curve called a hyperbola (see Figure 3), showing that the volume doubles as the pressure is halved.
After Boyle’s findings, scientists continued to study the properties of gases. In 1787, Jacques Charles, a French physicist, performed experiments measuring the effect of temperature on the volume of a fixed amount of gas (at constant pressure). He observed that when the pressure and amount of a gas are held constant, the volume of the gas is directly (or linearly) proportional to its temperature. This is known as Charles’s Law
. Volume increases with increasing temperature. Mathematically, this direct proportionality may be expressed as
or alternately as V1
(where T is absolute temperature in Kelvin, and Kelvin = degrees Celsius + 273). A plot of V versus T is a straight line (see Figure 4), showing that the volume doubles as the temperature (in Kelvin) doubles.
Herron, J. D.; Sarquis, J. L.; Schrader, C. L.; Frank, D. V.; Sarquis, M.; Kukla, D. A. Chemistry; D. C. Heath: Lexington, MA, 1996; pp 205–215.
Zumdahl, S. S. Chemistry, 3rd ed.; D. C. Heath: Lexington, MA, 1993; pp 187–190.