Lesson 9:

Concepts conveyed | Materials | Procedure | Benefits | Resource

Concepts conveyed:

The purpose of this demonstration is to teach students basic concepts about gases such as compression and expansion, and how they respond to changes in volume, temperature, moles, and pressure. This includes the Ideal Gas Law, Charles’ Law, and Boyle’s Law.

Materials:

(for one group of students)

• enough rubber balloons (~ 10 in. & round) for every student or small group of students plus five for the actual demonstration
• 2 small-mouth glass bottles of equal volume (~ 12 oz.)
• liquid nitrogen
• insulated gloves to protect hands from liquid nitrogen
• water (about 12 oz)
• 1 metal pan for liquid nitrogen (~ 2 quarts or more)
• 3 to 4 large marshmallows
• 1 mechanical vacuum pump with flexible vacuum hose for hookup to the Erlenmeyer flask (below)
• 1 Erlenmeyer flask with a rubber stopper with a glass tube through it for hookup to the vacuum pump.
• 1 helium-filled balloon

Procedure:

Part 1: Properties of Gases

A. Compression and Expansion.

Ask for two volunteers and ask each student to take one of the small mouth glass bottles. Give each student a balloon. Ask the students to hold onto the top of the balloon and push the bottom of it inside the bottle. Now, they must stretch the top of the balloon over the mouth of the bottle. When they finish, tell the students to try to inflate the balloons by blowing into them. The balloon expands only slightly because the air in the bottle is being compressed (the molecules of gas are getting closer to one another). This illustrates two properties of gases, that they may be compressed and that they expand to fill their containers uniformly.

B. Relative Densities.
Heavier than air. Remove the balloons from the bottles and fill one bottle with water and leave the other one filled with air. Since density equals mass divided by volume,

density = mass / volume

and the volumes of the two bottles are equal, the density is dependent on the mass. Obviously, the mass of the bottle filled with water is more than that filled with air. Thus, illustrating another property of gases, which is that all gases have low density (the density of water is 770 times greater than the density of air).

Lighter than air. Next, show the balloon filled with helium. Demonstrate the difference in the density of air and helium by showing that the helium balloon floats and by breathing in the helium from the balloon. The helium in the balloon is less dense than the air in the lungs. Therefore, the voice will have a much higher pitch than normal.

Part 2: Charles’ Law

The equation for Charles’ Law is:

V₁/V₂ = T₁/T₂, when P is constant
where V₁ = initial volume, V₂ = final volume,
T₁ = initial temperature, T₂ = final temperature

Perform a demonstration with balloons and liquid nitrogen. Ask one student to blow up a balloon. After the student blows up the balloon and ties it, place it into the liquid nitrogen. The liquid nitrogen is colder than the temperature of air. The temperature of liquid nitrogen is -195.8 degrees Celsius (-320 degrees Fahrenheit). As a result of the decrease in temperature, the volume of the balloon will also decrease. Remove the balloon from the liquid nitrogen, place it on a desk or table, and watch it expand again. As the temperature increases, the volume increases. The volume and temperature vary directly proportionally to each other. The pressure, which is approximately the same as the atmospheric pressure, remains constant. When you finish, give each student a balloon to take home to remember the concepts they have learned. Caution: When working with liquid nitrogen, use insulating gloves so as not to freeze your fingers!

Part 3: Boyle’s Law

The formula for Boyle’s Law is:

P₁V₁ = P₂V₂, when T is constant
where P₁ = initial pressure, P₂ = final pressure,
V₁ = initial volume, V₂ = final volume

Drop three or four large marshmallows into an Erlenmeyer flask, seal it with the rubber stopper and hook up the vacuum pump. Lower the pressure in the flask with the vacuum. The marshmallows will expand. This can be explained by considering Boyle’s Law. Pressure is inversely proportional to volume. As pressure decreases, the volume of the air pockets in the marshmallows increases. Next, allow the pressure to be restored in the flask by disconnecting the vacuum line. The marshmallows will return to their normal size. As the pressure increases, the volume decreases.

After all three parts of this demonstration have been completed, discuss with the class the theory behind what they observed. The discussion should include the explanation of the Ideal Gas Law:

PV = nRT
where P = pressure (atmospheres), V = volume (liters),
n = number of moles of the gas, R= Ideal Gas constant = 0.08206 (L•atm)/(mol•K),
and T = temperature (Kelvin).

The equation can be derived using the various gas laws. (Let “µ” represent proportionality.)

Boyle’s Law: V µ 1/P (constant n, T)
Charles’ Law: V µ T (constant n, P)
Avogadro’s Law: V µ n (constant P, T)

Note that this equation does not require knowledge of the mass (m), the molecular weight (MW), or density (D) of the gas. However, if we wish to relate these parameters in mass terms, we can divide both sides of the Ideal Gas Law equation with mass, then we have a relationship which involves the molecular weight and density:

Benefits:

• Students can take home a balloon to help them remember the concepts they were taught.
• Demonstrating the properties of gases in class allows instructors to assess immediately the students’ understanding of the material.
• P, V, T, and n for gases are all interrelated. These different exercises are intended to help students to visually observe the generality of this relationship.

Resource:

• VanCleave, J. P. In Chemistry for Every Kid; Sobel, D., Ed.; John Wiley & Sons: New York, 1989. p 20.
  
• Peters, E. I. and Kowerski, R. C. Introduction of Chemical Principles. Saunders College Publishing: Fort Worth, TX, 1994; 102.
  
• Chemistry the Central Science; Brown, T. L., LeMay, H. E., and Bursten, B. E., Eds.; Prentice-Hall: Englewood Cliffs, NJ, 1991; p 337.