28. Determination of the Rate of the Decomposition of ...



28. Determination of the Rate of the Decomposition of Hydrogen Peroxide*

Driving Question

How does changing the concentration of the reactants affect the value of k, the rate law constant?

Background

Hydrogen peroxide (H2O2) in aqueous solution decomposes very slowly under ordinary conditions. The equation for the decomposition is

[pic] (1)

A catalyst such as potassium iodide, manganese dioxide, or catalase enzyme may be used to increase the rate of reaction. Conducting a catalyzed decomposition of H2O2 in a closed vessel enables the determination of the reaction rate based on the pressure increase from the production of oxygen gas. Every 2 H2O2 molecules yields one O2 molecule; therefore, the rate at which H2O2 disappears is half the rate at which O2 is formed:

[pic]

Because the concentration of oxygen is proportional to its pressure, we can calculate the rate at which H2O2 decomposes by monitoring the rate of increase of the pressure due to the formation of oxygen. By varying the initial molar concentration of H2O2 solution, the rate law for the reaction can be determined.

There are two steps involved in the decomposition of hydrogen peroxide with potassium iodide as the catalyst:

[pic]

The first reaction determines the rate, that is, it goes much slower than the second reaction. The rate of the rate-determining reaction is calculated as follows:

[pic] (2)

where

k1 = the rate constant of the first reaction

m = the order of I– in the first reaction

n = the order of H2O2 in the first reaction

In this experiment, we determine n and m, as well as k1.

As O2 is a gas, it makes more sense to work with the number of moles than with concentration to obtain the concentration of H2O2. For Equation 1,

[pic]

While we cannot measure the change of number of moles of O2, using the ideal gas law we can calculate it from the change of pressure, which we can measure:

[pic]

where

V = the volume that the O2 can occupy (L)

R = the gas constant (L kPa/mol K)

T = the temperature inside the flask (K)

P = the partial pressure of the O2 generated by the reactions (kPa)

Now we can return to calculating the change of concentration for H2O2 as well:

[pic]

where

Vs = volume of the solution

We can rearrange the formula to get the rate of the reaction:

[pic]

Substituting this into Equation 2, we can calculate the rate constant:

[pic]

[pic] (3)

To determine the order of the reactants, n and m, the reactions are performed according to the following table.

Table1: Concentration ratios between the two reactants

|Reaction | Conc. of H2O2 | Conc. of I– |

|1 |[H2O2] |[I–] |

|2 |[H2O2] |2[I–] |

|3 |2[H2O2] |[I–] |

Determining the rate for reactions 1 and 2:

[pic]

The ratio of the two rates yields

[pic] (4)

Using the same argument to derive m from reactions 1 and 3:

[pic] (5)

The values of the order of the reactants, m and n, then will be used to determine k1, using Equation 3.

Materials and Equipment

For each student or group:

|Data collection system |Stopper with two holes for the Erlenmeyer flask |

|Absolute pressure sensor with quick-release |Beaker, 50-mL |

|connectors and plastic tubing |Glycerin, several drops |

|Stainless steel temperature sensor |0.1000 M Potassium iodide (KI), 60 mL |

|Sensor extension cable |3% Hydrogen peroxide (H2O2), 40 mL |

|Beaker (3), 100-mL |Water, deionized, 100 mL |

|Erlenmeyer flask, 250-mL |Electrical tape, 60 in. (optional) |

|Graduated pipet (3), 25-mL with rubber bulb | |

Safety

Follow all standard laboratory procedures.

Procedure

After you complete a step (or answer a question), place a check mark in the box (() next to that step.

Set Up

1. ( Connect to the USB port that contains the stopper with the pressure sensor and thermometer connected to it. Click on “Sparkvue” and then on “Build”. Click on “Pressure” and “Time” (in that order) and then the graph button: [pic]. Next click on “temperature” and the display button: [pic] and “ok”. Pressure should appear on the y-axis and time should be on the x-axis. There should be a box where temperature will be displayed. Set time interval [pic]to 2 seconds.

Collect Data

2. ( Find the mass of a dry 250-mL Erlenmeyer flask. Record. You will need to make a data table to record this value along with other values in future steps of the procedure.

Table 3: Reactant amounts to use for the three reactions

| Reaction |0.1 M KI (mL) |Water (mL) |3% H2O2 (mL) |

|1 |15.00 |35.00 |10.00 |

|2 |15.00 |25.00 |20.00 |

|3 |30.00 |20.00 |10.00 |

3. ( Perform each of the three reactions, using the measured amounts given in Table 3, according the steps listed below.

a. With graduated pipets, measure and transfer the water and potassium iodide solution into the 250-mL Erlenmeyer flask.

b. Place a Teflon stirrer into the flask and place the entire flask onto the magnetic stirrer. Turn on the stirrer about 1/3 of the way.

c. With a graduated pipet, measure and transfer the prescribed amount of H2O2 solution into the 50-mL beaker.

d. Pour the H2O2 solution into the Erlenmeyer flask and immediately insert the rubber stopper into the flask.

Important: Make sure that the stopper is sitting firmly in the flask. Pressure is building in the flask and a loose stopper might pop out. If that happens, you will need to repeat the experiment.

e. Quickly start recording data by clicking on the green arrow: [pic]

f. Continue to record the data for three minutes. Hold the stopper firmly during the entire experiment.

g. Stop recording data. [pic]

Note: The initial portion of the pressure versus time graph is not straight, which is attributed to the fact that the reaction does not begin immediately.

4. ( Record the temperature at which the reactions took place.

5. ( Display the three data runs on a graph. To do so make sure the 3 trials are checked in the legend in the upper right hand corner.

6. ( Save your experiment. First click on the snapshot button: [pic] Then select: [pic] and “save as…” Then print your graph. Click on “journal, then finally “print journal”. Print to the printer in the classroom (S21) and be sure to print one for each member of your lab group.

7. ( Find the slope of each line by:

Record these values in the table below as (P/(t.

8. ( To determine the volume of the flask: mark where the bottom of the stopper is in the neck of the flask with a sharpie. Fill the flask up with water to that mark. Determine the mass of the flask with the water in it. Record.

Data Analysis

1. ( Convert the percent by mass concentration to molarity for the different volumes of H2O2. Record the values in Table 4. Also calculate the molarity of I-. Keep in mind the volume of the entire solution once the KI, H2O2 and water are combined. The H2O2 is a 3% solution by mass. You can assume it has a density of 1 g/mL.

2. ( Calculate the order of the two reactants from the respective rates using Equations 4 and 5.

3. ( Calculate the volume of the flask by using the density of water. Also calculate the volume occupied by the oxygen gas. Finally, calculate the rate constants, using Equation 3, by substituting your respective experimental data. Record the values in Table 4 and determine the average value of k1.

Table 4: Experimental data and data analysis

|Reaction  |[H2O2] (M) |[I–] |(∆P/∆t (kPa/s) |n |m |k1 |

| | |(M) | | | |(M–1s–1) |

|1 | | | | | | |

|2 | | | | | | |

|3 | | | | | | |

|Average k1: | |

* This is an AP Chemistry course recommended experiment.

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