Method of Initial Rates
The reaction is first run with specific concentration of reactants and the induction period is timed. Then the same reaction is run with the concentration of potassium iodide halved and the induction period is timed again. Lastly the hydrogen peroxide concentration is halved but the same concentration of hydrogen peroxide as a the first run is used and the induction period is timed. The results of the kinetics experiment can be used to determine the order of reaction with respect to the iodide ion and hydrogen peroxide, and to determine the rate law for this reaction Rate = k [H2O2]x [KI]y
Summary of Lecture Slide and the Student Worksheet:
Independent Variable _________ Dependent Variable ____________ Control Variable _______________
| Experiment | [H2O2], M | [KI], M | Reaction Time | Rate of Reaction |
| H | 0.045 | 0.100 | ||
| I | 0.045 | 0.050 | ||
| J | 0.0225 | 0.100 |
This page is under construction. Web page author: T. Greenbowe, University of Oregon
Discussion of the Chemistry of this Demonstration
While this demonstration is effective, the overall understanding of the reaction mechanism is complicated for students. The thiosulfate ions quickly consume the triiodide ions. The blue-black color does not appear until all of the thiosulfate ions are consumed. The thiosulfate ions are the limiting reactant. The rate of reaction is first-order in potassium iodine and first order in hydrogen peroxide. For the Method of Initial Rates option, the details of the mechanism are revealed to the students in order to have the students determine the overall chemical reaction.
For the entire explanation of this demo see Bassam Shakhashiri.Chemical Demonstrations: A Handbook for Teachers of Chemistry, Vol. 4, pp 42-43. U. of Wisconsin: (1992). The following is a excerpt from this volume.
The sudden change from colorless to deep blue-black solutions in this demonstration can be explained with the following sequence of equations:
- 3 I-(aq) + H2O2 + 2 H+(aq) ==> I3-(aq) + 2 H2O(l)
- I3-(aq) + 2 S2O3 2-(aq) ==> 3 I-(aq) + S4O6 2-(aq)
- 2 I3-(aq) + starch ==> starch-I5- complex + I-(aq)
The first equation indicates that, in an acidic solution, iodide ions are oxidized by hydrogen peroxide to triiodide ions. These triiodide ions are reduced back to iodide ions by thiosulfate ions, as indicated in equation (2). This reaction is much faster than the reaction of equation (1). It consumes triiodide ions as fast as they are formed. This prevents any readily apparent reaction of equation (3). However, after all the thiosulfate ions have been consumed by the reaction of equation (2), triiodide ions react with starch to form the blue starch-pentaiodide complex .
The "A" beakers contain sodium thiosulfate, potassium iodide, and a little bit of starch. The "B" beakers contain hydrogen peroxide and sulfuric acid.
- Reaction 1: [H2O2] = 0.045 M; [KI] = 0.050 M
- Reaction 2: [H2O2] = 0.045 M; [KI] = 0.025 M
- Reaction 3: [H2O2] = 0.0225 M; [KI] = 0.050 M
The concentrations of the other reactants remain constant. ([H2SO4] = 0.10 M; [Na2S2O3] = 0.0028 M). By comparing the length of the induction periods, conclusions can be drawn regarding the dependence of the rate of reaction on the concentration of reactants.
The rate of reaction is how fast or slow a reaction occurs relative to a standard. The greater the rate of reaction, the less time it takes for the reaction to go to completion, i.e. the less time it takes for reactants to be converted to products. When focusing students' attention on the role of concentration of reactants with respect to the rate of reaction it is important to discuss and or make the connection to collision theory. The collision model assumes that molecules must collide in order for a reaction to occur. The more molecules present, the more chance for frequent "effective" collisions, the more often a reaction occurs. Reaction rate is proportional to the number of effective collisions, which depends on the concentration of the reactants. Inclusion of several picture diagrams or "molecular scenes" as well as showing one of our computer animations comparing two systems with different initial concentration of reactants is recommended. The reaction rate of the system with double the initial concentration is twice that of the other. However, the time it takes for the reaction to go to completion of the system with double the initial concentration is half the time. Students have a difficult time with the concept of a greater reaction rate taking less time.
This demonstration and accompanying discussion of collision theory should help students visualize what is occurring on the molecular level as well as make note of the inverse relationship between reaction rate and time.
Chemicals and Materials
- Three 600 mL beakers containing clear solutions labeled "1A", "2A", and "3A". (See "Discussion" section for contents of the beakers.)
- Three 400 mL beakers containing clear solutions labeled "1B", "2B", and "3B". (See "Discussion" section for contents of the beakers.)
- Three glass stirring rods.
- A clock or watch for timing the reactions.
Procedure for Performing the Demonstration
Upon mixing of the two colorless solutions, the resultant solution is initially colorless. Record the time that it takes for the resultant solution to turn blue-black in color.
- Pour the contents of beaker 1B into the beaker labeled 1A. Begin timing as soon as the solutions are poured. Stir vigorously for a few seconds.
- Pour the contents of beaker 2B into the beaker labeled 2A. Begin timing as soon as the solutions are poured. Stir vigorously for a few seconds.
- Pour the contents of beaker 3B into the beaker labeled 3A. Begin timing as soon as the solutions are poured. Stir vigorously for a few seconds.
Safety Precautions
Avoid getting any of the solutions on you. If you do get chemicals on your skin, wash thoroughly with soap and water. If you get chemicals in your eyes, flush with water for 15 minutes and seek medical attention.
References
1. Bassam Shakhashiri. (1992). Chemical Demonstrations: A Handbook for Teachers of Chemistry, Vol. 4, pp 42-43. U. of Wisconsin.
2. Creary, Xavier and Morris, Karen M. (1999). A New Twist on the Iodine Clock Reaction: Determining the Order of a Reaction. J. Chem. Educ., 1999, 76 (4), p 530. DOI: 10.1021/ed076p530
The analysis of the data collected from a series of experiments of the iodine clock reaction is used to determine the order of a reaction with respect to the reactants and to determine the rate law for the reaction.
Education Resources to Accompany this Demonstration
http://introchem.chem.okstate.edu/DCICLA/iodine_clock.html
An excellent guided-inquiry resource for an alternative approach to Iodine Clock Reaction using potassium iodate and sodium bisulfite available from Flinn Scientific ©2016. This is a guided inquiry experiment.
https://www.flinnsci.com/api/library/Download/1fec5df4c5a64a7f98df4730ab...
“The Order of Reaction” experiment in Kinetics, Volume 14 in the Flinn ChemTopic™ Labs series; Cesa, I., Editor; Flinn Scientific: Batavia, IL (2003).