The Ideal Gas?
Dr. Roger Sandwick and Thomas M. Moffett Jr., SUNY Plattsburgh, 2004
In this experiment you will observe a reaction that produces a gas in one section, and determine
the molar mass of another gas in the second part.
In both gases the gases will behave as ideal gases, which allows for the use of the Ideal Gas Law:
PV=nRT
(1)
Both experiments are occurring in open systems, which means that the pressure of the gas is
equal to atmospheric pressure.
Part A – Decomposition of Hydrogen Peroxide
Hydrogen peroxide is a chemical used in the chemical industry and research for its oxidative
properties and in medicine as an antiseptic. (You may have a bottle of 3 % hydrogen peroxide in
your medicine cabinet.) As it name implies, hydrogen peroxide has one hydrogen per every
oxygen; its actual molecular formula is H2O2.
At room temperature, a water solution of hydrogen peroxide (H2O2) will undergo a slow reaction
to form molecular oxygen gas (O2) and liquid water (H2O). Write this reaction as a balanced
equation in your pre-lab.
This experiment will look at the decomposition of H2O2 in the presence of a catalyst (yeast). It is
important to remember that a catalyst doesn’t really factor into the chemical equation; it simply
increases the rate of the reaction.
You will calculate the amount of oxygen that should be evolved under the conditions in the
room, measure the amount of oxygen actually produced, and determine your percent error.
% error =
actual - theoretical value x 100
theoretical value
(2)
Part B – Molar Mass of Butane
Butane is a fuel used in ....duh ... butane lighters. One of the reasons it is preferred over other
hydrocarbons in this application has to do with its boiling point. Under a little pressure, butane
can be converted at room temperature to a liquid. Since liquids take up much less space than
gases, the manufacturer can then place a reasonable amount of fuel in a small unit. Once you
push down on the lever, the butane comes out and, at normal pressures, converts immediately to
a gas. It is this gas form which is ignited.
Here, you will perform three trials to experimentally determine the molecular mass of butane.
Remember that molecular mass is in the units grams/mole.
mass
moles
MM =
(3)
By rearranging equation 3 to solve for moles and then substituting into equation 1, we get the
following equation which can be used to calculate molar mass.
MM =
(mass)RT
PV
(4)
Procedure – Part A:
Set up the experimental system as directed by your lab instructor. Fill a 100 mL graduated
cylinder with water. Cover the graduate with your fingers and inverted it into a large beaker
filled three-quarters with water. Record the starting volume of the air in the graduated cylinder.
If the cylinder is too full of water and you can not get a reading (i.e. less than 5 mL), let some air
into the cylinder until you can accurately record the starting level. One end of the tubing should
now be fed into the graduate, the other end is connected to the side arm flask.
Record the temperature of the water in the beaker.
Add 5 mL of H2O2 to the side arm flask. Add 0.2 g of catalyst (yeast) into the flask, quickly
secure the rubber stopper to the top of the flask. Wait until reaction is complete and determine
the amount of gas produced. .
Do a second trial.
Calculation:
Calculate the amount of O2 which will theoretically evolve when 5 mL of a 3.0 % solution of
H2O2 completely decomposes. A 3.0 % solution in this case is by percent mass, i.e.
3 g per 100 g of solution. (The rest is water.) How does this compare to the actual yield?
Calculate % error.
Procedure – Part B:
1. Invert a 50 mL graduated cylinder filled with water into a 400 mL beaker filled three quarters
with water. Record the initial starting level of air.
2. Weigh a butane lighter to the nearest thousandth of a gram.
3. Hook a piece of small Teflon tubing onto the gas exit port of the butane lighter. Place the
other end under the inverted cylinder. Release gas by pushing down on the lever and let gas
replace the water until about 30 to 40 mL has been displaced. Take a final reading.
4. Reweigh the butane lighter to compute how much butane was used.
*** Do two trials.
Setup for part A:
* For part B, the setup is essentially the same, with the lighter replacing the filter flask and using
a different sized beaker, graduate, and tube.
Sample Data Sheet
Part A. (Show all calculations for full credit)
Temperature (oC)
Temperature (K)
Pressure (torr)
Theoretical moles of oxygen evolved
Theoretical Volume of oxygen evolved (L)
Final Volume in Graduate
Initial Volume in Graduate
Actual Volume of oxygen evolved (L)
%-error
Average %-error
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Part B. (Show all calculations for full credit)
Temperature (oC)
Temperature (K)
Pressure (torr)
Initial mass of lighter
Final mass of lighter
Mass of butane released
Final volume in graduate
Initial volume in graduate
Volume of butane released
Molar Mass of butane
Average Molar Mass
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Post Lab Questions
1.) In this experiment the temperature of the gas evolved is equal to the temperature of the water
in the beaker (which ideally should be the same as the air temperature). Explain how the
volume of oxygen evolved would change if you used ice water instead of room temperature
water. How would it change if you used boiling water?
2.) Based on your calculation of the molar mass of butane, which of the following is the
molecular formula of butane?
a.) C2H6
b.) C3H8
c.) C4H10
d.) C5H12
3.) Balance the following reaction:
____ N2 (g) + ____ H2 (g) → ____ NH3 (g)
4.) Using the balanced equation above, determine the mass of ammonia produced when 36.0 L
of hydrogen reacts at STP.
Pre-Lab Questions
1.) What is the purpose of the yeast in part A?
2.) Write the balanced reaction for the decomposition of hydrogen peroxide. Be sure to include
all information for full credit.
3.) How much space does 25.0 g of argon gas occupy at STP?
4.) A 48.5 g sample of methane (CH4) occupies a 32.5 L container at 15 oC. Determine the
pressure inside the container.