Chem 131 lab report

Lab 10 Ideal Gas Law

Finding the Percent H2O2 with Catalase from Yeast

Objectives:

• Use Dalton’s law of partial pressure to determine the pressure of a dry gas. • Use the ideal gas law to determine the moles of a dry gas. • Determine the percentage of hydrogen peroxide in a commercially available

hydrogen peroxide solution

• Observe how a catalyst affects a reaction

Introduction:

In this lab we will catalyze the decomposition of hydrogen peroxide with the enzyme catalase

from yeast as shown in the reaction below. The catalase is neither a product nor a reaction; it

merely speeds up the reaction.

Imagine you were given the problem “How many grams of H2O2 would have reacted if 0.500g

of O2 were produced?” This is a familiar stoichiometry problem that follows the pattern

gA→molA→molB→gB as shown below:

0.500 g O2 1 mol O2 2 mol H2O2 34.0 g H2O2 = 1.06 g H2O2 32.0 g O2 1 mol O2 1 mol H2O2

But what if we changed the problem just a little so it read, “How many grams of H2O2 would

have reacted if 0.375 L of O2 gas were produced.” The difference here is that we have to start

our problem with the volume of oxygen instead of the grams of oxygen. In this case our

problem follows the pattern

Liters A → mol A → mole B → gB

We need a way to convert liters into moles. In this lab we will use the ideal gas law to convert

volume of a gas into moles of a gas and more generally to check the concentration of a

hydrogen peroxide solution bought in a store.

The ideal gas law is very valuable when dealing with gases since it establishes a relationship

between temperature, pressure, volume, and amount of a gas.

PV = nRT

In this equation:

• P is the gas pressure in atmospheres

• V is the volume of the gas in liters

• n is the number of moles of the gas

• R is the constant value of 0.0821 L·atm/mol·K

• T is for the temperature of the gas in Kelvin.

Safety/disposal:

Wear safety goggles/safety glasses and always follow laboratory safety rules while performing

experiments.

Disposal/wastes:

All solutions can be disposed of down the drain.

Materials:

• Baker’s yeast

• 600 mL beaker

• Erlenmeyer flask 125 mL

• #5 rubber stopper with rubber hose

• 12.5 cm long copper wire

• 2000mL graduated cylinders

• eudiometer

• 3% Hydrogen peroxide

Figure 1: The water filled eudiometer is inverted and placed in the beaker over the hose as shown.

Figure 2: Oxygen gas generated by the decomposition of hydrogen peroxide flows through the hose and bubbles into the inverted eudiometer where it collects at the top.

Procedure:

1. Fill the 600 mL beaker with tap water to about an inch from the top. 2. Using the 150mL beaker, fill the eudiometer to the top with tap water. 3. Place your finger over the opening of the eudiometer. Flip the eudiometer with your finger

over the opening so water does not run out.

4. Place the eudiometer in the 600 mL beaker. Remove finger. Air should not be allowed to enter the eudiometer.

5. Secure the water filled eudiometer with the clamp as shown in figure 1.

6. Insert the 12.5 cm copper wire into the rubber hose. Bend the hose to create a narrow j-bend – bottom of “J” should be 1 cm wide or less. Place the end of the j-bend in the hose so it is

inside the eudiometer as in figure 1.

7. Record the temperature of the water in the 600 mL beaker. The temperature of the gas collected will be the same as the water temperature. (SEE VIDEO FOR TEMPERATURE)

8. Record the barometric pressure in the room. The pressure of the gas you collect will equal the pressure in the room when the levels of the water inside and outside the eudiometer are

equal at the end of the experiment as shown in figure 3. (SEE VIDEO FOR PRESSURE)

9. Pipette 5.00 mL of hydrogen peroxide solution into the Erlenmeyer flask. 10. Measure 10.0 mL of yeast solution in a 10 mL graduated cylinder. 11. Pour the yeast solution into the Erlenmeyer flask and stopper immediately. 12. Start the stopwatch. Swirl the two solutions together. Hold tube so it doesn’t come out of

eudiometer while swirling.

13. You will begin to see bubbles coming up into the eudiometer as in figure 2. If you do not see bubbles, make sure that the stopper is on tight and the tube is not leaking. You will need to

start over if you have leaks.

14. Continue to swirl the Erlenmeyer flask until no more bubbles form to assure the reaction has gone to completion. Record the time when the reaction is finished.

15. Place your hand in the 600 mL beaker and seal the opening of the eudiometer with your finger. Without letting air in or water out transfer the eudiometer to the 2000 mL graduated cylinder in

your sink. Raise or lower the eudiometer so the level of water in the eudiometer is the same as

the graduated cylinder. Record the volume of gas in the eudiometer (line 6 on data table). (SEE

VIDEO FOR FINAL GAS VOLUME)

16. Pour all liquids down the drain and clean the labware.

Using Dalton’s law

When two or more gases are mixed the total pressure of the gases can the thought of as the sum

of the pressures of the individual gases. This is known as Dalton’s Law and is written as

PT = P1 + P2 + . . . + Pn

where PT is the total pressure, P1 is the pressure of the first gas, P2 is the pressure of the

second gas and Pn is the pressure of the nth gas. The diagram below shows how Dalton’s

law works for a mixture of argon and water vapor. The pressure of the argon plus the vapor

pressure of the water equals the total pressure.

Sample Problems as they might appear in lab exams and lecture exams

Problem 1: What is the vapor pressure of water when the temperature is 18.0oC?

Solution: Find a table that lists the vapor pressure of water at various temperatures such as the

table below. According to the table the vapor pressure of water is 15.5 torr when the

temperature is 18.0oC. Since a torr and a mmHg are the same thing we can also say the vapor

pressure of water is also 15.5 mmHg at 18.0oC.

Temperature in oC

18.0 19.0 20.0 21.0 22.0 23.0 24.0 25.0

Pressure in torr

15.5 16.5 17.5 18.6 19.8 21.1 22.4 23.8

Problem 2: A student collected hydrogen gas over water where the pressure was 755.8 mmHg

and the temperature is 25.1oC. What is the partial pressure of the hydrogen gas in atm?

Solution: We know the total pressure of the hydrogen gas and water vapor is 755.8 mmHg.

The vapor pressure of the water is about 23.8 torr = 23.8 mmHg. Therefore, the partial

pressure of the hydrogen is 755.8 mmHg – 23.8 mmHg = 732.0 mmHg. Since 760 mmHg = 1

atm we can convert 732.0 mmHg to 0.9632 atm.

Name:_________________________

Lab 10 Finding the Percent H2O2 with Catalase from Yeast

Data Table (SEE VIDEO TO COMPLETE THE DATA TABLE)

1. Temperature of the water oC

2. Barometric pressure in Hg

3. Initial eudiometer volume

4. Reaction start time

5. Reaction end time

6. Final eudiometer volume

7. Volume of O2 gas generated in mL mL

8. Volume of O2 gas generated in L L

1. Convert the 5.00 mL of 3% H2O2 solution into grams using the density of 3% H2O2

which is 1.02 g/mL. This is the mass of the solution in grams.

2. Convert the temperature from oC to K.

3. Convert the pressure in inches of Hg to atm using the fact that 1 atm = 29.92 in Hg.

This is the Ptotal .

4. From the table 1 find the vapor pressure of water, Pwater in torr and convert to atm.

Temperature in oC 18.0 19.0 20.0 21.0 22.0 23.0 24.0 25.0 Pressure in torr 15.5 16.5 17.5 18.6 19.8 21.1 22.4 23.8

5. Use Dalton’s law of partial pressures to solve for the pressure of oxygen, Poxygen.

Ptotal = Poxygen + Pwater

6. Find the moles of O2 using the ideal gas equation, PV = nRT where R = 0.0821 L atm /

mol K. Use Poxygen for P in the equation. This is the experimental yield in moles of O2.

7. Convert the experimental yield in moles of O2 into grams of H2O2 using stoichiometry. This the experimental grams of H2O2 in the solution.

8. Calculate the percent H2O2 in the sample.

% H2O2 = (experimental grams of H2O2)/(grams of solution) x 100