Graham’s Law Lab
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 rapid 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 expresses as a ratio:
ra
mb
rb
ma
In this formula, ra is the rate of diffusion (or effusion) of a gas and rb is the rate of
diffusion (or effusion) of a second gas. Similarly, ma and mb are the respective molar
masses of the two gases.
In this lab, you will compare the rates of diffusion of ammonia (NH3) gas and
hydrogen chloride (HCl) gas. These two gases react chemically to form tiny white
ammonium chloride (NH4Cl) crystals that appear as an aerosol, or suspension, in the air.
You will place solutions containing these gases at different locations in a reaction tube.
By measuring where in the tube the aerosol forms, you will be able to use Graham’s law
to find the relative rates of diffusion of the gases.
Pre-lab questions:
1. What are the hazards of working with concentrated solutions of HCl and NH3, and
what precautions should you follow?
2. How will you know when gaseous NH3 and HCl molecules will have made
contact with each other?
3. Using the molar masses of the gases and Graham’s law, calculate the ratio of
diffusion rate of NH3 to the diffusion rate of HCl that you expect to find in this
investigation.
4. Predict whether the NH4Cl crystals will form closer to the source of NH3 or HCl
gas.
Safety
Wear goggles at all times in the laboratory. Gaseous ammonia (NH3) and hydrogen
chloride (HCl) molecules are irritating to skin, nasal passages, and eyes.
Hydrochloric acid solution is corrosive. Wear gloves. Do not sniff these solution.
Do not let these solutions contact your skin or clothing. Before discarding the cotton
swabs, place them in a small beaker of water to dilute the hydrochloric acid and
ammonia solutions. If these solutions contact your skin or clothing, flush these areas
with great quantities of water and notify your teacher.
Procedure
1. Put on your goggles. Splice the two straws into one tube as follows. Using
scissors cut a small slit (about 7 mm) in the end on one straw. Push the cut end
into one end of the other straw. Secure the straws together with a piece of clear
adhesive tape.
2. Place a cotton swab into each end of the straw tube. Using a marking pen, mark
the tube at the points where the swabs have penetrated. These marks represent the
starting points from which the gases will diffuse. Measure and record the total
distance between these two marks.
3. Use the marking pen to mark one of the cotton swabs. This swab will be used for
the hydrochloric acid (HCl).
4. Half fill a small beaker with water and set it aside for use later.
5. Put on latex gloves. Using a dropper, place 2 drops of concentrated hydrochloric
acid solution on the marked swab. Place two drops of concentrated ammonia
solution on the other swab. CAUTION: work under a fume hood. Concentrated
HCl and NH3 are toxic and irritating to eyes and nasal passages. HCl is
corrosive to skin and clothing. Avoid contact with the solutions and their vapors.
6. While one person holds the straw tube steady, another person should
simultaneously insert the two wet swabs in opposite ends of the tube. The swabs
should penetrate to the points you previously marked.
7. Watch for the formation of a white aerosol ring inside the tube. The ring consists
of tiny crystals of ammonium chloride (NH4Cl). Measure and record the distance
traveled by the HCl from its initial mark to the point where the ring forms. Do the
same for the NH3 gas.
8. Remove the wet swabs and place them in the beaker of water. This will minimize
the escape of HCl and NH3 gases into the laboratory. It will also dilute the
concentrated solutions and allow them to neutralize each other.
9. Rinse the beaker and swabs thoroughly with water and discard the swabs in the
garbage. Rinse the straw tube with water from the wash bottle and discard in the
garbage.
10. Wash your hands before leaving.
Data
Total distance between the marks _______
Distance traveled by the HCl gas ______
Distance traveled by the NH3 gas ______
Questions:
1. Write a balanced chemical equation to describe the reaction you observed.
Include the states of the reactants and products.
2. Calculate the ration (dNH3/dHCl) of the distances traveled by the gases. How does
your experimental ratio compare to the diffusion ration you calculated in the prelab questions?
3. Calculate the percent error of the experiment.
4. Are the results of your investigation consistent with Graham’s Law? Explain.
5. Suppose the distance between the solutions in the straw tube is 40.0 cm. Use
Graham’s law to calculate the distance the NH3 and HCl gases would travel
before they collide.
6. Consider the latex balloon described in the introduction to this investigation.
Explain why the balloon deflates more quickly if it is filled with helium than if it
were filled with air.
7. Suppose you had an unknown gas. How could you use Graham’s law to design
an experiment that would help you determine the molar mass of the unknown
gas?