1. Rates of Reaction
“The iodide Clock”
By: Dr. Robert D. Craig, Ph.D.
For Freshman Chemistry 2
2. Measuring Reaction Rates
• In this experiment we will determine the
effect of a reactant concentration and
temperature on the rate of a chemical
reaction. A reaction is chosen which
proceeds conveniently slowly near room
temperature and which can be measured
easily by a dramatic color change.
3. Measuring Reaction Rates
• By measuring the initial rate (the rate near
reaction time zero) for a series of reactions
with varying concentrations, we can deduce
to what power the rate depends on the
concentration of each reagent.
4. • Demonstration of variations in reaction
rates:
H2O2 vs. I- vs. H30+ in water to
produce I2-
H2O2(aq) + 2I-(aq) + 2 H+(aq) ->
I2(aq) + 2 H2O(aq)
5. Measuring Reaction Rates
• this reaction proceeds, the colorless
reactants gradually develop a brown color
due to the product I2. Because of the
difficulty of timing the appearance of the I2,
we make use of another much faster
reaction in the same solution to mark the
progress of the slow reaction
7. Measuring Reaction Rates
• The reaction happens so fast that I2 produced by
reaction 1 is consumed instantaneously by the
thiosulfate (S2O3
2-
), so that the I2 color cannot
develop. Because both S2O3
2-
and S4O6
2-
are
colorless, the solution remains colorless. However,
we do not add enough thiosulfate to react with all
the I2 that will be formed from reaction 1.
8. Measuring Reaction Rates
• By this device the reaction solution stays
colorless until the instant at which all the
thiosulfate is consumed, and then free I2
begins to appear.
9. Measuring Reaction Rates
• We time the reaction from the initial mixing
until the appearance of I2.
• In order to help see this appearance we add
starch indicator which forms an intensely
colored dark complex with I2 and signals the
appearance of I2 by a dramatic color change.
10. Measuring Reaction Rates
• We know the initial concentration of S2O3
2-
and measure the time interval for this S2O3
2-
to be consumed. These quantities determine
the rate of the slow reaction 1. By repeating
the experiment with different concentrations
of H2O2 and different temperatures, we can
determine the effect of H2O2 concentration
and temperature on the rate of reaction 1.
12. Effect of Acidity
• This experiment tests the effect of different
sulfuric acid concentrations on the rate of the
iodine clock reaction. The initial concentrations
and volumes of all other reactants are the same in
all four vials. The greater the acid concentration,
the faster the reaction.
• H2SO4 Concentration Ratios (1 :0.83 : 0.67 : 0.51)
14. Effect of Iodide
• This experiment tests the effect of different iodide
ion concentrations on the rate of the iodine clock
reaction. The initial concentrations and volumes of
all other reactants are the same in all four vials.
The greater the iodide concentration, the shorter
the reaction time.
• Iodide Ion Concentration Ratios (1 : 0.75 : 0.50 :
0.375
16. Effect of Peroxide
• This experiment tests the effect of different
hydrogen peroxide concentrations on the rate of
the iodine clock reaction. The initial
concentrations and volumes of all other reactants
are the same in all four vials. The greater the
peroxide concentration, the less time it takes for
the solution to turn blue.
• H2O2 Concentration Ratios (1 : 0.80 : 0.60 : 0.40)
18. Effect of Temperature
This experiment tests the effect of
temperature on the rate of the iodine clock
reaction. The reaction rate is greater at
higher temperatures.
20. Measuring Reaction Rates
RUN 1
•
Dispense reagents for Run 1 into another set of
glassware.
• Have a clock or watch with a second hand ready.
Then in quick succession:
Add the contents of both beakers B and C into flask
A.
Swirl the mixture in A.
21. Measuring Reaction Rates
RUN 1
– Start timing as soon as beaker B is
emptied. This is time ti .
– • Place a thermometer in flask A but DO
NOT STIR WITH IT.
– • Record the time interval Δt = tf – ti in
seconds.
22. Measuring Reaction Rates
• RUN 2
– • Thoroughly rinse flask A just used, and then
dispense reagents.
– • Follow the procedure for Run 1. This one is
faster.
23. Measuring Reaction Rates
• RUN 3
– • Thoroughly rinse flask A, and then dispense
reagents.
– • Follow the procedure for Run 1. This is even
faster than 2.
24. Measuring Reaction Rates
• RUN 4
– • Thoroughly rinse flask A, and dispense reagents.
– • Place flask A and beaker C (but not B) on a hot plate
and warm these up to between 45 o
C and 50 o
C. Use a
thermometer to monitor the temperature in flask A
only.
– • When flask A is over 45 o
C, remove both from the hot
plate. Mix B and C into A and swirl.
– • Leave the thermometer in the reaction mixture and
record the temperature when the color turns. This one is
very fast.
26. Measuring Reaction Rates
On the data sheet, fill in the following
quantities in the table for each of the five
runs:
• Row 1: Δt is your measured time interval in
seconds.
• Row 2: Temp, oC is the temperature of the
reaction mixture.
• Row 3: [H2O2] is the [H2O2]molarity in flask
A at time t i .
• Calculate the total volume ignoring the 2
drops of starch.
27. Measuring Reaction Rates
• Row 4: [S2O3
2
-]i is the thiosulfate
molarity in flask A at time t i .
• Row 5: Δ[H2O2] is the change in
peroxide concentration during Δt.
• This calculation is illustrated below.
28. Measuring Reaction Rates
• Below is a graph plotted following the
iodine clock experiment. The graph could
be of any reaction where you can measure
the concentration of a reactant over time.
30. Measuring Reaction Rates
1. Set up an ice bath.
2. Dispense reagents into one set of
glassware A, B, and C as shown above for
Run 5. Set these in the ice bath for later
use.
31. Examples of factors affecting rate
• For each of the 5 factors give one example
from today’s demonstrations
• Nature of reactants –
• Ability of the reactants to meet –
• The concentration of the reactants –
• The temperature of the system –
• The presence of catalysts –
32. Examples of factors affecting rate
• Nature of reactants – H202 is very reactive
• Ability of the reactants to meet – Most
reactions occur in the aqueous phase
• The concentration of the reactants – How does
this effect rate
• The temperature of the system – reactions that
are exothermic are faster due to room
temperature
• The presence of catalysts – an oxidizing
reagent like MnO2 is a catalyst for this reaction
33. Measuring Reaction Rates
Answer these questions:
1. What units are associated with
concentration?
2. What units are associated with reaction rate?
3. What do the square brackets in [H202] indicate
4. Explain how the rate of a reaction is
determined
5. Plot this data (include title, axes labels):
H202 0.100 0.072 0.056 0.046 0.039 0.034 0.030 0.026
Time
(s)
0 50 100 150 200 250 300 350
34. Measuring Reaction Rates
• Rate = k[I-
]p
[H2O2]q
[H3O+
]r
• This reaction is the oxidation of iodide ion
(I-) to molecular iodine (I2) by hydrogen
peroxide (H2O2):
• It is what happens when no thiosulfate is
present “A BACK REACTION”
35. Measuring Reaction Rates
• Because we can measure the concentrations
in the rate law using the techniques
described above, the unknowns we wish to
measure are k, p, and q. One method of
directly measuring k, p, and q is called the
method of initial rates.
36. Measuring Reaction Rates
1. How does the rate at the beginning of a
reaction compare to the rate later in a
reaction?
2. Explain how the rate of reaction of
H2O2(aq) + 2I-(aq) + 2 H+(aq) -> I2(aq) + 2
H2O(aq) is determined experimentally.
37. Measuring Reaction Rates
1. Concentration: mol/L or mol•L–1
2. Rate = concentration/time: (mol/L)/s or
mol•L–1
•s–1
3. The square brackets "[ ]" is the symbol for
concentration (mol/L)
4. The rate of reaction is measured by:
Instantaneous slope; rise over run; slope of
the tangent at any point.
38. Measuring Reaction Rates
Slope (rise/run) is the reaction rate in (mol/L)/s…
H202 concentration (mol/L) 0.1 0.072 0.056 0.046 0.039 0.034 0.03
Time (s) 0 50 100 150 200 250 300
H202 concentration (mol/L) Time (s) Concentration of H2S04 (mol/L) Concentration of KI(mol/L)
0.1 0 0.1 0.1
0.072 50 0.072 0.072
0.056 100 0.056 0.056
0.046 150 0.046 0.046
0.039 200 0.039 0.039
0.034 250 0.034 0.034
0.03 300 0.03 0.03
0.026 350 0.026 0.026
40. Measuring Reaction Rates cont
1. As the reaction proceeds, the rate decreases
because reactants are being used up (recall,
concentration of reactants affects rate)
2. H202 concentration is measured indirectly by
measuring the production of I2(g) (purple) –
likely via a spectrometer
For more lessons, visit
www.chalkbored.com
41. Measuring Reaction Rates
• Through experiments and calculations, the
rate law of the Iodine Clock Reaction at
approximately 24°C was found to be…
• Rate = 1.18 X 10-4
[I-] [H2O2]
• Using the rate law, the rate of the reaction
can be found at the given temperature of
24°C with any molar concentration of I- and
H2O2.
