Name
Period
Date
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
2H2O2  2H2O  O2
(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:

Δ[O2 ]
1 Δ[H2O2 ]

2
Δt
Δt
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:
H 2 O2  I  OI  H2O
2OI  O2  2I
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:

Δ[H2O2 ]
 k1[H2O2 ]m [I ]n
Δt
(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
*
This is an AP Chemistry course recommended experiment.
1
Determination of the Rate of the Decomposition of Hydrogen Peroxide
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,

ΔnO2
1 ΔnH2O2

2 Δt
Δt
While we cannot measure the change of number of moles of O 2, using the ideal gas law we can
calculate it from the change of pressure, which we can measure:

1 ΔnH2O2
V ΔP

2 Δt
RT Δt
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 H 2O2 as well:
Vs [H2O2 ]  nH2O2
Vs  Δ[H2O2 ]  ΔnH2O2

Vs Δ[H2O2 ]
V ΔP

2
Δt
RT Δt
where
Vs = volume of the solution
We can rearrange the formula to get the rate of the reaction:
Δ[H2O2 ]
2V ΔP
 
Δt
Vs RT Δt
Substituting this into Equation 2, we can calculate the rate constant:
2V Δp
 k1[H2O2 ]m [I ]n
Vs RT Δt
k1 
2
2V
ΔP
[H2O2 ] [I ] Vs RT Δt
m
(3)
 n
PS-2897B
Student Inquiry Worksheet
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:
Rate1  k1[H2O2 ]m [I ]n


Rate2  k1[H2O2 ]m 2[I ]
n
The ratio of the two rates yields


k1 [H2O2 ]m 2[I ]
Rate2

Rate1
k1 [H2O2 ]m [I ]n
n
Rate2
 2n
Rate1
 Rate2 
n
ln 
  ln 2
Rate
1 

 Rate2 
ln 
  n ln 2
 Rate1 
 Rate2 
ln 

Rate1 

n 
ln 2
(4)
Using the same argument to derive m from reactions 1 and 3:
 Rate3 
ln 

 Rate1 
m 
ln 2
(5)
The values of the order of the reactants, m and n, then will be used to determine k1, using
Equation 3.
3
Determination of the Rate of the Decomposition of Hydrogen Peroxide
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:
. Next click on
“temperature” and the display button:
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
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.
4
PS-2897B
Student Inquiry Worksheet
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:
f.
Continue to record the data for three minutes. Hold the stopper firmly during
the entire experiment.
g. Stop recording data.
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:
Then select:
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.
5
Determination of the Rate of the Decomposition of Hydrogen Peroxide
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 H 2O2.
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–]
(M)
∆P/∆t
(kPa/s)
1
2
3
Average k1:
6
PS-2897B
n
m
k1
(M–1s–1)