How Much Oxygen Gas Is Produced?

630 words | 3 page(s)

The purpose of this experiment is to find out the amount of oxygen produced from the commercial extraction of oxygen from hydrogen peroxide. Hydrogen peroxide has both the moles of hydrogen and the moles of oxygen. To produce oxygen, the peroxide is decomposed, through a catalyst, to form water and oxygen (Harada, 55). Although hydrogen peroxide is able to decompose naturally to form water and oxygen, a catalyst is used to speed up the reaction. The following equation shows the decomposition of hydrogen peroxide to water and oxygen.

2H2O(aq) -> 2H2O(I) + O2(g)

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In the above equation, H2O2 represents hydrogen peroxide, H2O represents water, and O2 represents oxygen.

The ideal gas equation in this case is PV = nRT.

Where; P represents the pressure of the oxygen gas in the atmosphere, V represents the volume of oxygen gas collected in liters, R represents the gas constant 0.08206 L.atm/(K.mol), and T represents the temperature in Kelvin (given by the water temperature).
This gas law equation will be used in calculating the moles of oxygen produced following the decomposition of the peroxide. Following the determination of the number of moles of oxygen produced, the reaction in the balanced decomposition will be used to get the mole to mole ratios of hydrogen peroxide and oxygen produced (Schulz, 28).Finding the moles and the mass of the hydrogen peroxide will allow the accurate computation of the moles and mass of water and oxygen produced from the process.

Procedure 1
The materials used in the experiment include clear bin of shoe storage, 100 mL graduated cylinder, 125 mL Erlenmeyer flask, 600 mL beaker, stopper with unbending tubing, elastic tubing, rubber band, ring with a stand, and a tape. The first step is to Put 3,500 mL of distilled water in a clear bin of shoe storage and record the temperature in Kelvin (K). Next is to record the water’s vapor pressure, in mmHg and at the recorded water temperature. Third is to fill a 100 mL graduated cylinder with water and use a rubber band and parafilm to cover the cylinder’s opening. Fourth, using a small ring, set up the ring stand. Fifth, at the edge of the rigid tubing stopper of the Erlenmeyer flask, assemble the supple tubing steadily. This is the point where the decomposition will take place. Next, slowly invert the 100 mL graduated cylinder over the clear shoe storage bin with the help of small ring. Next, use a spatula to remove the rubber band and parafilm from the 100 mL graduated cylinder. Next, make a record of air volume in the 100 mL graduated cylinder. A very minimal amount of air should be in the cylinder, otherwise, repeat the process until a small volume is achieved. Next, following the conversion and stabilization of pressure units to mmHg by the TA, record the atmospheric pressure.

Procedure 2
Objective in this procedure is to determine the mass percentage of hydrogen peroxide solution and the amount of oxygen produced from the process (in moles). The first step is to accurately measure 10.00 mL of hydrogen peroxide and measure its mass. Next, measure 0.01g of dry yeast and combine it with the peroxide in a 125 mL Erlenmeyer flask. Next, make a record of the final air volume in the 100 mL graduated cylinder once the reaction is complete. Next, make a record of the water column height in the clear bin of shoe storage in millimeters (mm). Next, esnuer there is oxygen in the tube with the use of a burning splint in the opening of the graduated cylinder. Lastly, make sure there is no glassware in the cylinder. Proceed to calculations.

  • Harada, Hisashi. “Sonophotocatalytic decomposition of water using TiO 2 photocatalyst.” Ultrasonics sonochemistry 8.1 (2001): 55-58.
  • Schulz, Manfred, et al. “Peroxide chemistry: mechanistic and preparative aspects of oxygen transfer.” (2000).

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