Rate of natural and catalyzed decomposition of hydrogen peroxide

Enzymes are proteins that catalyze biochemical reactions by take downing activation energy. Enzymes are specific to one reaction due to their construction ; they bind with a specific substrate at an active site to bring on a conformational alteration that allows chemical bonds to be broken easier-thus cut downing the sum of energy needed to originate a reaction. Catalysts, and in this enzymes, are neither reactants nor merchandises in chemical reactions. Therefore, the same enzyme can be used to catalyse legion reactions.1 Several conditions affect the activity of an enzyme: salt concentration, temperature, pH, and other activators/inhibitors. Enzyme action is characterized by peak conditions, so a great addition or lessening in certain conditions could do the enzymes to denature and lose the ability to increase the rate of reactions. For illustration, enzyme activity typically increases in correlativity to an addition in temperature due to more kinetic energy, but when temperature reaches a certain point ( about 40-50EsC for most enzymes ) the protein is denatured and can no longer map right. 2

Natural decomposition of H peroxide experiences fluctuation due to alterations in temperature and concentration.3 These variables are eliminated in this experiment by the changeless room temperature and changeless concentration of H peroxide. This allows the lone variable in the experiment to be the sum of clip that the catalase reacts with the H peroxide.

This experiment tests the consequence of an enzyme, catalase, on the dislocation of H peroxide ( H2O2 ) . The decomposition of H2O2 into H2O and O occurs spontaneously, but at an highly slow rate. A basal line will be established to mensurate the initial sum of H2O2 in the solution. After the base line is established, the dislocation of H peroxide will be catalyzed by catalase. After a certain sum of clip, the add-on of sulphuric acid ( H2SO4 ) will halt the activity of the enzyme, holding the decomposition of H peroxide. The staying sum of H2O2 will be titrated with aid of K permanganate ( KMnO4 ) . The K permanganate will respond with the extra H peroxide and the sulphuric acid. Once all the H peroxide has been consumed, the add-on of K permanganate will for good dye the solution a pinkish or chocolate-brown colour. This full reaction is shown in the equation: . 2 It is likely that the catalase will ease the dislocation of increasing sums of H2O2 as more clip elapses during the experiment. Because of the slow rate of natural dislocation of H peroxide, it is likely that an un-catalyzed decomposition reaction that sits out for 24 hours will incorporate about the same concentration of H2O2 as the base line check.

Materials and Methods:

Experiment 1:

1. Establishing a Base Line2
2. 10 milliliter of 1.5 % solution of H2O2
3. 1 milliliter of H2O
4. 10 milliliter of H2SO4
5. 5 milliliter of KMnO4
6. 10 mL syringe ( 2 )
7. 5 milliliters syringe
8. 2 beakers
9. 1 milliliter Pipette

White paper

10 milliliter of 1.5 % H2O2 was put into a clean beaker. The pipette was used to add 1 milliliter of H2O to the solution. A clean 10 milliliter syringe was used to add 10 milliliter of H2SO4 and the full solution was assorted good. A 5 mL sample of the solution was removed utilizing the 5 milliliter syringe ; it was placed into another clean beaker. A clean 5 milliliter syringe was used to mensurate 5 milliliter of KMnO4. The KMnO4 was added to the 5 milliliter sample one bead at a clip, and the solution was assorted good after each bead. The solution was compared to the white paper so that it would be clearly evident when a lasting colour alteration occurred. When the solution was for good dyed a pinkish or chocolate-brown colour, the degree of solution left in the syringe was measured. The sum of KMnO4 used was measured by deducting the concluding reading of the syringe from the initial 5 milliliter. This base line value was used in computations subsequently in the experiment. 2

Experiment 2:

The Uncatalyzed Rate of H2O2 Decomposition2

1. 10 milliliter of 1.5 % solution of H2O2
2. 1 milliliter of H2O
3. 10 milliliter of H2SO4
4. 5 milliliter of KMnO4
5. 10 mL syringe ( 2 )
6. 5 mL syringe ( 2 )
7. 2 beakers
8. 1 milliliter Pipette

White paper

The same process for Experiment 1 was followed, except the beaker incorporating merely H peroxide was stored at room temperature for 24 hours before the anything was added to it. The observations and computations of this experiment were used to cipher the natural rate of decomposition of H2O2. 2

Experiment 3:

Enzyme Catalyzed Rate of H2O2 Decomposition2

1. 100 milliliter of 1.5 % solution of H2O2
2. 10 milliliter of catalase
3. ice
4. 100 milliliter of H2SO4
5. 50 milliliter of KMnO4
6. 10 mL syringe ( 2 )
7. 5 mL syringe ( 2 )
8. 14 beakers
9. 1 milliliter Pipette

White paper

Stopwatch

Catalase was kept on ice until needed in the experiment. 10 milliliter of H2O2 were added to a clean beaker. The 1 milliliter pipette was used to add 1 milliliter of catalase to the H2O2 and the solution was assorted for 10 seconds. After 10 seconds, a clean 10 milliliter syringe was used to add 10 milliliter of H2SO4. After the solution was assorted, a 5 milliliter sample was removed and placed into a clean beaker. A clean 5 milliliter syringe was used to mensurate 5 milliliter of KMnO4. The KMnO4 was added to the 5 milliliter sample one bead at a clip, and the solution was assorted good after each bead. The solution was compared to the white paper so that it would be clearly evident when a lasting colour alteration occurred. When the solution was for good dyed a pinkish or chocolate-brown colour, the degree of solution left in the syringe was measured. The initial and concluding readings of the syringe incorporating the KMnO4 was observed and recorded. These stairss were repeated at clip intervals of 30, 60, 90, 120, 180, and 360 seconds ( measured from when the catalase was added to when the H2SO4 was added ) .