Kinetic Molecular Theory Gas Laws - Oakland Schools

Real-World Connections
Why do balloons filled with helium deflate more quickly than balloons filled with air?
Graham’s law was used in WWII to extract uranium-235 for use in nuclear reactions
Why is nitrogen added to “high-performance” tires?
Sources
Adapted and modified from Prentice Hall Lab Manual Chemistry: Connections to Our
Changing World, 1996
Procedure/Description of Lesson
Introduction
Have you ever noticed how quickly a helium balloon deflates? A common latex balloon
filled with helium will lose much of its gas overnight, yet the same balloon filled with air
will remain inflated for several days. Why is this so? The kinetic-molecular theory
states that gases consist of tiny particles in constant random motion. These particles
have mass, and they frequently make elastic collisions with each other and the walls of
their container. Different gases, however, differ in the rate at which they are able to
move among each other (diffusion) or through tiny openings (effusion), such as a hole
in a balloon.
Thomas Graham recognized that the different rates of movement of gas particles at
constant temperature are related to the molar masses of the gases. Graham’s law
compares the rates of diffusion or effusion of any two gases as follows: Under constant
temperature and pressure, the rate of diffusion or effusion of two gases is inversely
proportional to the square roots of their molar masses. Mathematically, Graham’s law
may be expressed as a ratio:
(m 1 / m 2 ) 1/2 = r 2 / r 1
In this formula r 1 is he rate of diffusion (or effusion) of a gas and r 2 is the rate of
diffusion (or effusion) of a second gas. Similarly, m 1 and m 2 are the respective molar
masses of the two gases.
In this laboratory investigation, you will compare the rates of diffusion of ammonia and
hydrochloric acid, and sodium hydroxide and hydrochloric acid. These gases react to
form a white salt that appears as a ring in the reaction tube. By measuring where the
salt ring forms, you will be able to use Graham’s law to find the relative rates of
diffusion of the gases.
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