Studying the Kinetics of a Reaction
In this experiment we will study the kinetics of a chemical reaction. The reaction is called a
"clock" reaction because of the means of observing the reaction rate. The reaction involves the
oxidation of iodine ion by the bromate ion in the presence of an acid:
6 I1- (aq) + BrO31-(aq) + 6 H1+(aq) ----> 3 I2 (aq) + 3 H20 (l) + Br1- (aq) (1)
The reaction is somewhat slow at room temperature. Its rate depends on the concentration of
the reactants and on the temperature. The rate law is a mathematical expression that relates the
reaction rate to the concentrations of reactants. If we express the rate of reaction as the rate of
decrease in concentration of bromate ion, the rate law has the form:
a b c
Rate = - ∆[BrO31-] /∆t = k[I1-] [BrO31-] [H1+]
where the square brackets refer to the molar concentration of the indicated species. The rate is
equal to the change in concentration of the bromate ion, divided by the change in time for the
reaction to occur. The term "k" is the rate constant for the equation, and changes as temperature
changes. The exponents a, b, and c are called the "orders" of the reaction with respect to the
indicated substance and show how the concentration of each substance affects the rate of
reaction. Adding the exponents gives the order of the overall reaction.
One purpose of the experiment is to determine the total rate law for the process. To do this we
must measure the rate, evaluate the rate constant, k, and determine the order of the reaction for
each reactant, the values of a, b, and c. Also, we will see the effect a catalyst has on the rate.
To find the rate of the reaction we need some way of measuring the rate at which one of the
reactants is used up, or the rate at which one of the products is formed. The method that we will
use is based on the rate at which iodine forms. If thiosulfate ions are added to the solution they
react with iodine as it forms in this way:
I2(aq) + 2 S2O32-(aq) ----> 2 I1-(aq) + S4O62-(aq) (2)
Reaction (1) is somewhat slow. Reaction (2) proceeds extremely rapidly, so that as quickly as
iodine is produced in reaction (1), it is consumed in reaction (2). Reaction (2) continues until all
of the thiosulfate is used up. After that, iodine begins to increase in concentration in solution. If
some starch is present, iodine will react with the starch to form a deep blue-colored complex that
is rapidly apparent.
Carrying out reaction (1) in the presence of thiosulfate ion and starch produces a chemical
"clock." When the thiosulfate is consumed, the solution turns blue.
In all of our reactions we will used the same quantity of thiosulfate ion. The blue color appears
when all the thiosulfate is used up. An examination of equations (1) and (2) shows that 6 moles
of S2O32- are needed to react with the I2 formed from 1 mole of BrO31-. Knowing the
amount of thiosulfate used allows the calculation of the amount of I2 that is formed, and also the
amount of BrO31-that has reacted at the time of the color change. The reaction rate is expressed
as the decrease in concentration of BrO31-ion divided by the time it takes for the blue color to
The experiment is designed so that the amounts of the reactants that are consumed are small in
comparison with the total quantities present. This means that the concentration of reactants is
almost unchanged during the reaction, and therefore the reaction rate is almost constant during
The experiment is designed using a microscale procedure. Only 12 drops of reactants delivered
from capillary droppers will be used for each measurement. The steps involved are as follows:
1. Determine the volume of a drop of solution. This must be done so that the number of
moles of thiosulfate ion can be found, and so the amount of bromate ions that react can be
2. Fine the order of the reaction for each of the reactants, and determine the rate constant.
We will do this by carrying out an experiment at specific concentrations of each of the
reactants and measuring the reaction rate. Then we will change the concentration of one
reactant and observe how the reaction rate changes. This will be repeated for each reactant.
This data allows the calculation of the order of each reactant. Once the orders are known, the
value of the rate constant can be calculated.
3. In the last part of the experiment we will observe the effect of a catalyst on the rate.
Potassium iodide, KI, 0.010 M Sodium thiosulfate, Na2S2O3, 0.0010 M
Potassium bromate, KBrO3, 0.040 M Hydrochloric acid, HCl, 0.10 M
Starch solution, 2% Copper (II) nitrate, Cu(NO3)2, 0.1 M
Baking Soda, NaHCO3, Solid
Beral pipets with capillary tips Trough for hot and cold water bath
Thermometer Microplate, 12-well (well plate)
Beaker, 10-mL or 50-mL Toothpicks for stirring
Sensitive balance Cotton swabs
Safety Alert - Hydrochloric acid is hazardous to skin and eyes. Wash off spills with
lots of water. Neutralize spills on counter top with baking soda. Wear Chemical Splash
Goggles and a Chemical-Resistant Apron.
1. Find the Volume of a Drop of Solution.
Obtain micro tip Beral pipets. Then place a small beaker on a sensitive balance and find it's
mass. Holding the dropper vertically, deliver 5 drops of water into the beaker, and find the total
mass. Add an additional 5 drops of water, again determine the mass. Deliver 5 more drops and
again find the mass. Record your data in a table in your notebook. See the Data and
Calculations section for help in setting up your notebook.
Assume that the density of each of the dilute solutions that will be used is the same as that of
water, 1.00 g/mL. Calculate the volume of 1 drop of each solution.
2. Determine the Reaction rate and Calculate the Rate Law.
The table that follows shows the reagent quantities to be used in carrying out the reactions
Because we don't want the reaction to start until we are ready, be sure the KBrO3 solution is the
last solution added. It is important to use care in measuring out the solutions. Since the total
solution volume is quite small, even one extra drop can cause a substantial change in
It is necessary to use consistently good technique to obtain reproductive data. Hold droppers
vertically and be sure no air bubbles are introduced. Since such small quantities of reagents are
used, it is very easy to repeat measurements. Practice your technique by carrying out the first
experiment at least three times (more, if necessary) until your values are reproducible.
Calculations of the orders of reactants are all based on the values obtained for the first
experiment, so be sure to get reproducible data from the beginning. All other experiments should
be carried out at least twice.
A study of Table 1 shows that all experiments contain the same total number of drops of
solution. Only one drop of sodium thiosulfate, Na2S2O3, and one drop of starch solution are
added to each well. In experiments 1, 2, and 3, the concentration of potassium iodide, KI, is
gradually increased while all other volumes remain constant. Experiments 1, 4 and 5 have an
increasing concentration of potassium bromate, KBrO3. Experiments 1, 6 and 7 show an
increase in the concentration of hydrochloric acid, HCl. Experiment 8 will be a test to see if
calculated orders of reactants agree with experimental values.
Table 1. Reagent Quantities for Experiments
Experiment KI H2O HCl Starch Na2S2O3 KBrO3
Number 0.010M 0.10M 2% 0.0010M 0.040M
1 2 drops 4 drops 2 drops 1 drop 1 drop 2 drops
2 4 drops 2 drops 2 drops 1 drop 1 drop 2 drops
3 6 drops 0 drops 2 drops 1 drop 1 drop 2 drops
4 2 drops 2 drops 2 drops 1 drop 1 drop 4 drops
5 2 drops 0 drops 2 drops 1 drop 1 drop 6 drops
6 2 drops 2 drops 4 drops 1 drop 1 drop 2 drops
7 2 drops 0 drops 6 drops 1 drop 1 drop 2 drops
8 3 drops 1 drops 3 drops 1 drop 1 drop 3 drops
Measure out the drops of solutions required for experiment 1 in one of the wells of a 12-well
strip. Be sure to add KBrO3 last. Stir the mixture thoroughly with a toothpick. This is very
important, because it is impossible to achieve good mixing in the small well without stirring.
Begin timing the reaction as soon as the KBrO3 is added. Record the time required for the first
blue color to appear. Repeat the measurements until consistently reproducible values are
obtained. Record the room temperature as the temperature of these reactions.
Empty the well plate, rinse with water and shake to dry the wells. Use detergent and a cotton
swab, if necessary, to be sure the wells are clean and dry for each experiment. Carry out the
experiments with solution volumes described in Experiments 2 through 8.
3. Observe the Effect of a Catalyst on the Rate
Repeat the procedure of Experiment 1, but this time add 1 drop of 0.10 M copper(II) nitrate
solution, Cu(NO3)2, and only 3 drops of water to the mixture. The total volume will still be 12
drops. Record the reaction time.
Disposal and Cleanup
The solutions can be washed down the drain with a 20-fold excess of water. Clean the well
plates with detergent and water using a cotton swab.
Data and Calculations
1. Find the volume of a drop of Solution.
Set up a table like the one below in your laboratory notebook in which you can record the
necessary data and calculated values. Show a sample of each type of calculation.
(a)Mass of empty beaker ___ Average mass of 1 drop of water ___
(b)Mass of beaker plus 5 drops water ___ (d)Mass of beaker = 15 drops water ___
(b)-(a)Mass of first 5 drops ___ (d)-(c)Mass of 3rd 5 drops water ___
Average mass of 1 drop of water ___ Average mass of one drop of water ___
(c)Mass of beaker plus 10 drops water ___ Average of masses of 1 drop of water___
(c)-(b)Mass of 2nd 5 drops water ___ Volume of 1 drop: ___
2. Determine the Reaction Rate
In your lab notebook, set up the table similar to the one below.
Trial 1 Trial 2 Trial 3 Trial 4 Avg time Temp Rxn Rate [I-]o [BrO3-]o [H+]o
Exp 1 ________________________________________________________________________
Exp 2 ________________________________________________________________________
Exp 3 ________________________________________________________________________
Exp 4 ________________________________________________________________________
Exp 5 ________________________________________________________________________
Exp 6 ________________________________________________________________________
Exp 7 ________________________________________________________________________
Exp 8 ________________________________________________________________________
2. A) Determine the moles of S2O32- ion that were used in each reaction
In each reaction there is 1 drop of 0.0010 M Na2S2O3 solution. Since you know the
molarity of the solution, and the volume (liters) of one drop, you can calculate the moles
of S2O32- used. (Molarity = moles / liter) It is the same for all eight reactions.
2. B) Determine the moles of BrO31- used in each reaction
Since the rate will be expressed as -∆[BrO3-]/∆t, we will find the moles of BrO31- that
were used before we stopped timing. This represent the change in moles of BrO31- over
time. The blue color begins to appear when all of the thiosulfate ion is consumed.
Examine equations (1) and (2) on the front page and determine the ratio of S2O32- ions to
BrO31- ions. Use this ratio, and the moles of S2O32- you just calculated to determine the
moles of BrO31- ions that reacted. Again, it is the same for each reaction.
2. C) Determine the change in molarity of the BrO31- ion
Since you've just determine the moles of BrO31- used, you can easily calculate the
molarity using the equation molarity = moles / liter. Consider all of the solutions that had
that many moles of BrO31- react, had a total of 12 drops! This is the value of -∆[BrO31-] ,
and it is the same for all reactions.
3. Calculate the Rate for each of the 8 reactions
The rate of each reaction can be found by dividing -∆[BrO31-] that was just found, by the
number of seconds required for each reaction to take place. Write these values in the data
4. Calculate the Initial Concentrations
Use the formula for diluting a solution: (M1) (V1) = (M2) (V2) Where the volume is
measured in drops. This should be calculated before the lab. Because I said so. Write
these in the data table.
5. Calculate the Order of Each Reactant
Just like the problems we've been doing. Baby! Look for two experiments where the
concentration of only one of the reactants is different and see how that effects the rate.
6. Find the Rate Constant
Using the newly found rate law (differential), then plug in the values for each experiment
and find the average k.
7. Check your Data
Experiment 8 is a check on your data. Substitute the concentrations of the reactants for
this experiment into the rate law that you determined and calculate the rate of reaction.
Then determine the percent error between the calculated rate with the actual measured
rate for experiment 8?
1) Why does reaction rate change as concentration changes?
2) Explain the procedure used to find the rate law.
3) Comment on the agreement between measured and calculated rates for Experiment 8.
4) Why does reaction rate change as the temperature changes?
5) Differentiate between reaction rate and specific rate constant.
6) Comment on the effect of a catalyst. How does this effect the activation energy?
7) How could you improve on the results of this lab?