Chem& 161
General Chemistry w/Lab I

Experiment: Determination of the Gas Law Constant

In this experiment you will calculate the gas law constant, R, by collecting a known quantity of hydrogen gas and measuring the temperature, pressure and volume of the gas collected.

From the ideal gas law, PV = nRT, you can see that it is possible to determine a value for R if you can isolate a sample of gas for which P, V, T and n are all known. In this experiment you will accomplish this by collecting hydrogen gas formed in the reaction of magnesium metal with hydrochloric acid. The equation is:

Mg (s) + 2 HCl (aq) ——> MgCl2 (aq) + H2 (g)

When you collect the hydrogen gas you will also measure the temperature, pressure and volume of the gas collected. From this data a value for R can be calculated.

Since the hydrogen will be collected over water in a gas collection tube, it will be saturated with water vapor. According to Dalton’s Law, the total pressure of the gas mixture is the sum of the partial pressure of H2 plus the partial pressure of the water vapor. The partial pressure of the water vapor can be looked up in tables in the CRC, and subtracted from the total pressure to find the pressure of the H2. The volume of the gas will be measured directly from the gas collection tube.

The necessary calculations are as follows:

Note: In this laboratory you will be working with a lab partner.


  1. Obtain a gas collection tube and a plastic bucket from the cart. You will also need a clamp for mounting the tube on a ringstand.
  2. Get some copper wire, a strip of magnesium, and a one hole stopper from the cart.
  3. Using the fact that one mole of gas occupies 22.4 L at STP, calculate the amount of Mg metal that will produce between 30 and 40 mL of H2 gas. Weigh out this amount of Mg. Note that this is only an approximation that allows you to calculate an appropriate amount of Mg to use.
  4. Roll the Mg into a loose ball, and wrap the Mg with the Cu wire to form a cage around the Mg. Leave a 2.5 cm length of Cu wire coming off of the cage.
  5. Obtain about 25 mL of 6 M HCl. Add approximately 20 mL to the gas collection tube. Then carefully add deionized water to completely fill the tube, being careful to not disturb the HCl as the water is added. Pouring the water slowly down the side of the tube will help.
  6. Run the length of wire from the copper cage through the one hole stopper, bend the wire over the top of the stopper, and insert the assemblage into the gas collection tube so that the cage of Cu holding the Mg is inside the tube. Make sure that the tube is completely filled with water and that no air bubbles remain.
  7. With one finger covering the hole in the stopper, invert the tube into the bucket (or beaker) of water. Mount the tube to a ringstand, being careful to not let the end of the tube come out of the water in the process. As the HCl flows down the tube it will react with the Mg, generating H2 gas which will rise to the top of the tube and displace water out of the bottom.
  8. When the reaction is complete attempt to adjust the height of the gas collection tube so that the liquid levels inside and out are the same. This will assure that the pressure inside the tube is the same as that outside. Record the volume of gas in the gas collection tube.
  9. Record the temperature of the gas collected, assuming its temperature is the same as that of the water in the bucket and gas collection tube. Also record the barometric pressure. The instructor will demonstrate how to read the barometer.
  10. Repeat the experiment a second time.

Analysis and Calculations

  1. Assuming that the magnesium reacted completely, calculate the moles of hydrogen formed from the mass of magnesium you actually used. As part of your calculations, confirm that the Mg was the limiting reagent.
  2. Calculate the pressure of the dry H2 from the barometer reading and the vapor pressure of water. (Look in the CRC index under vapor pressure, aqueous vapor, below 100°C.)
  3. Using the experimental values for pressure, temperature, volume, and moles of the gas, calculate a value for R. Repeat the calculations for your second trial. If your two trials agree report an average.
  4. Calculate a percent error.

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