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Investigate the Kinetics of the Amazing Iodine Clock Reaction

Difficulty
Time Required Average (6-10 days)
Prerequisites You should be attending or already have taken an introductory chemistry class. This science fair project requires careful attention to detail and will involve some creative problem solving and independent research on your part.
Material Availability You will need to order a chemical kinetics kit online. See the Materials and Equipment list for details.
Cost Average ($50 - $100)
Safety Wear safety goggles when working with all chemicals. Hydrochloric acid is corrosive to hands and eyes. Wear splash goggles and gloves when handling acids.

Abstract

The iodine clock reaction is a favorite demonstration reaction in chemistry classes. Two clear liquids are mixed, resulting in another clear liquid. After a few seconds, the solution suddenly turns dark blue. The reaction is called a clock reaction because the amount of time that elapses before the solution turns blue depends on the concentrations of the starting chemicals. In this chemistry science fair project, you will explore factors that affect the rate of the iodine clock reaction.

Objective

Determine how the concentration of hydrogen peroxide affects the rate of the iodine clock reaction and calculate the reaction order.

Credits

David B. Whyte, PhD, Science Buddies

Cite This Page

MLA Style

Science Buddies Staff. "Investigate the Kinetics of the Amazing Iodine Clock Reaction" Science Buddies. Science Buddies, 16 Nov. 2013. Web. 21 Aug. 2014 <http://www.sciencebuddies.org/science-fair-projects/project_ideas/Chem_p091.shtml>

APA Style

Science Buddies Staff. (2013, November 16). Investigate the Kinetics of the Amazing Iodine Clock Reaction. Retrieved August 21, 2014 from http://www.sciencebuddies.org/science-fair-projects/project_ideas/Chem_p091.shtml

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Last edit date: 2013-11-16

Introduction

Chemical kinetics is the branch of chemistry that is concerned with the mechanisms and rates of chemical reactions. The mechanism of a chemical reaction is a description of what happens to each molecule at a very detailed level—which bonds are broken, which new bonds are formed, and how the three-dimensional shapes of the chemicals changes during the course of the reaction. The rate of the reaction is a measure of its speed. The rate of a chemical reaction can be measured by how quickly the reactants disappear, or by how quickly the products are generated. The iodine clock reaction is a favorite demonstration in chemistry classes because it has an element of drama. Two clear solutions are mixed, producing a new clear solution. Then, after a period of several seconds, the solution turns dark blue. As mentioned, chemical kinetics measures how fast a reaction is occurring. For most chemical reactions, the rate is so fast that special equipment is needed to measure it. For the iodine clock reaction, on the other hand, the rate can be easily measured with a stopwatch.

To perform the iodine clock reaction in this science fair project, you will mix potassium iodide, hydrochloric acid, starch, thiosulfate and hydrogen peroxide. The time it takes for the reaction mix to turn blue will be measured with a stopwatch. For the procedure, you will vary the amount of hydrogen peroxide to see how this affects the time the mixed chemicals stay clear before turning blue.

The reactions that form the basis for the iodine clock reaction are shown below.

Equation 1:

H2O2 + 3 I- + 2 H+ → I3- + 2 H2O

  • H2O2 = Hydrogen peroxide
  • I- = Iodide ion (from potassium iodide)
  • H+ = A proton, from hydrochloric acid (HCL)
  • I3- = Triiodide
  • H2O = Water

This equation states that hydrogen peroxide reacts with iodide ions in acid solution to form triiodide and water.

Triiodide has the very interesting property of reacting with starch to form a dark blue complex. There is starch in the mix of chemicals, so why doesn't the triiodide react with it? The reason the triiodide doesn't react with the starch is that it is immediately consumed in a reaction with the thiosulfate.

Equation 2:

I3S- + 2 S2O32- → 3 I- + S4O62-

  • I3S- = Triiodide
  • S2O32- = Thiosulfate ion
  • 3 I- = Iodide ion
  • S4O62- = Tetrathionate ion

Equation 2 says that triiodide reacts with thiosulfate to form iodide ions and tetrathionate.

The reaction in Equation 2 happens so fast that none of the triiodide has time to form a complex with starch, even though the starch is in the reaction mix. The reactions in Equations 1 and 2 are moving along during the lag time between mixing the chemicals and the dramatic appearance of the blue color. Note that iodide ions are regenerated in Equation 2, so they are available to react with the hydrogen peroxide in Equation 1. The thiosulfate, on the other hand, is consumed as it is turned into tetrathionate. The lag period ends when the thiosulfate is all used up. At this time, the triiodide is able to react with the starch.

Equation 3:

I3- + starch → (I3- starch complex)

  • I3- = Triiodide
  • I3- starch complex, which is blue

This equation says that starch reacts with triiodide to form a blue complex.

The faster the reaction in Equation 1 goes, the faster the triiodide uses up the thiosulfate and the faster the triiodide is free to react with the starch. What is the rate of the first reaction? The rate of the reaction in Equation 1 is a measure of how the concentration of hydrogen peroxide changes per unit time:

Equation 4:

Rate = [Change in (H2O2)]/sec

  • (H2O2) = Concentration of hydrogen peroxide

Equation 4 indicates that the rate of the reaction is proportional to the reciprocal of the time.

The rate of a reaction depends on the concentration of the reactants. In Equation 1, for example, increasing the amount of hydrogen peroxide will increase the rate at which it reacts with iodide. The concentrations of iodide and acid remain the same, so the rate will depend only on the changes in hydrogen peroxide concentration. (The iodide is recycled between Equations 1 and 2, and the concentration of acid is high enough that the change in its concentration is small. Note the concentrations of the reactants in the Materials and Equipment section). The rate actually depends on the concentration of hydrogen peroxide raised to a power, called the "reaction order."

Equation 5:

Rate = k(H2O2)x

  • k = Rate constant, in 1/seconds (s)
  • (H2O2) = Concentration of hydrogen peroxide, in moles/liter
  • x = Order of the reaction for hydrogen peroxide, unitless

The good news from Equation 5 is that the rate depends on the concentration of hydrogen peroxide, and you will know what the concentration of hydrogen peroxide is when the reaction starts. You will use the number of hydrogen peroxide drops as a measure of its concentration.

There are many ways to explore the chemistry of the iodine clock reaction. The details of the reaction mechanisms can be studied by varying the concentrations of other reactants, in addition to hydrogen peroxide. And it is ideal for investigating the effect of temperature on reaction rates. Some of these areas are touched on in the Variations. Now, let's get started.

Terms and Concepts

  • Chemical kinetics
  • Chemical reaction
  • Mechanism
  • Rate
  • Reactant
  • Product
  • Iodine clock reaction
  • Solution
  • Potassium iodide
  • Hydrochloric acid
  • Starch
  • Thiosulfate
  • Hydrogen peroxide
  • Complex

Questions

  • What would happen to the iodine clock reaction if the reaction shown in Equation 2 were very slow instead of very fast?
  • What is the acid that is used in the reactions described above?
  • Based on your research, what is chemical complex?
  • What is the structure of the starch and triiodide complex? Why is it blue?
  • Why does the rate of Equation 2 depend on the rate of Equation 1?
  • How would adding more thiosulfate affect the lag period of the clock reaction?
  • Some sources state that triiodide is not really the form of iodine in the starch complex. What other forms have been proposed and which way does the evidence point (triiodide or not) in your opinion?
  • How can Equations 4 and 5 be combined? Hint: The result should show that the reciprocal of the time is proportional to the hydrogen peroxide concentration raised to the value of x, the rate constant.

Bibliography

Materials and Equipment

These items can be purchased from Carolina Biological Supply Company, a Science Buddies Approved Supplier:

  • Chemical splash safety goggles
  • Reaction rates kit
    • The kit contains the following solutions and materials: 0.5 M potassium iodide solution, 3% hydrogen peroxide, 0.01 M sodium thiosulfate, 0.1 M hydrochloric acid, 5 g starch, 1 mL syringes.
    • Note: If you are ordering this kit through Carolina Biological Supply Company, the kit must be ordered by a teacher and shipped to a school or business address, so plan accordingly.
  • Beaker, 50 mL
  • Microplates, 24-wells each. Alternatively, you may mix the drops on a piece of plastic wrap.
  • Thermometer
You will also need to gather these items:
  • Disposable gloves. Can be purchased at a local drug store or pharmacy, or through an online supplier like Carolina Biological Supply Company. If you are allergic to latex, use vinyl or polyethylene gloves.
  • Newspaper to protect work area
  • Distilled water (1 gallon); available at grocery stores
  • White sheet of printer paper
  • Toothpicks
  • Stopwatch or other timer
  • Permanent marker
  • Lab notebook

Disclaimer: Science Buddies occasionally provides information (such as part numbers, supplier names, and supplier weblinks) to assist our users in locating specialty items for individual projects. The information is provided solely as a convenience to our users. We do our best to make sure that part numbers and descriptions are accurate when first listed. However, since part numbers do change as items are obsoleted or improved, please send us an email if you run across any parts that are no longer available. We also do our best to make sure that any listed supplier provides prompt, courteous service. Science Buddies does participate in affiliate programs with Amazon.comsciencebuddies, Carolina Biological, and AquaPhoenix Education. Proceeds from the affiliate programs help support Science Buddies, a 501( c ) 3 public charity. If you have any comments (positive or negative) related to purchases you've made for science fair projects from recommendations on our site, please let us know. Write to us at scibuddy@sciencebuddies.org.

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Experimental Procedure

Performing the Experiment

  1. Be sure to wear your safety goggles and latex gloves.
  2. Protect the surface on which you are working with newspaper.
  3. Prepare a 1% starch solution by dissolving the 5 grams (g) of starch from the reaction rates kit in 50 mL of distilled water in the beaker.
  4. Set up the microplate for the first set of experiments.
    1. As an alternative, mix the drops on a flat piece of plastic wrap that is placed over a piece of white printer paper. Separate the drops by 2–3 centimeters (cm).
    2. Look at the mixed drops against a white background; for example, against a white piece of paper.
    3. Experiments will be run in duplicates (two wells per test).
    4. If you do not get reproducible results, run the experiment again.
  5. In two adjacent wells, add the following (use the syringes that came with the kit) to each well:
    1. 4 drops of the potassium iodide (KI) solution
    2. 2 drops of water
    3. 2 drops of hydrochloric acid
    4. 1 drop of starch solution
    5. 1 drop of thiosulfate solution
    6. 8 drops of hydrogen peroxide
  6. Mix well with a toothpick.
  7. Start the timer. Begin timing as soon as the hydrogen peroxide is added.
  8. Record the time that you can first see some blue color developing in your lab notebook.
    1. Call this the time lag.
    2. Also record the room temperature.

Varying the Hydrogen Peroxide Concentration

Note: You will now continue the procedure using one less drop of hydrogen peroxide and one more drop of water. The first two sets are detailed below.

For 7 drops hydrogen peroxide:

  1. In two adjacent wells, add the following to each well:
    1. 4 drops of the potassium iodide (KI) solution
    2. 3 drops of water
    3. 2 drops of hydrochloric acid
    4. 1 drop of starch solution
    5. 1 drop of thiosulfate solution
    6. 7 drops of hydrogen peroxide
  2. Mix well with a toothpick.
  3. Start the timer.
  4. Record the time that the solution turned blue.
  5. Record the room temperature.

For 6 drops hydrogen peroxide:

  1. In two adjacent wells, add the following:
    1. 4 drops of the potassium iodide (KI) solution
    2. 4 drops of water
    3. 2 drops of hydrochloric acid
    4. 1 drop of starch solution
    5. 1 drop of thiosulfate solution
    6. 6 drops of hydrogen peroxide
  2. Mix well with a toothpick.
  3. Start the timer.
  4. Record the time that the solution turned blue.
  5. Record the room temperature.

Completing the Experiment

  1. Repeat steps 1–5 for 5, 4, 3, 2, and 1 drop of hydrogen peroxide. Remember, each time, add one drop of water and subtract one drop of hydrogen peroxide. This will keep the total volume of the reaction constant, while changing the concentration of the hydrogen peroxide.
  2. Perform the entire procedure two more times with clean materials. This will demonstrate that your results can be repeated.

Analyzing Your Results

  1. Make a data table that shows the following:
    1. The number of drops of hydrogen peroxide.
    2. The time lag for each test, in seconds (sec).
    3. Average of the time lags for similar samples, in seconds. For example, average the three trials for the set with seven drops of hydrogen peroxide.
    4. The reciprocal of the average time lag for each test, in 1/sec.
  2. Graph the number of drops of hydrogen peroxide on the x-axis and the reciprocal of the time lag on the y-axis.
    1. The reciprocal of the time lag is proportional to the rate, so the y-axis is a measure of the reaction rate.
    2. Label the y-axis Reaction rate and the x-axis Concentration of hydrogen peroxide.
  3. Note the shape of the curve. From Equation 5 in the Introduction, the curve shape depends on "x," the order of the reaction for hydrogen peroxide.
    1. If x = 1, the curve will be a straight line. (Doubling the hydrogen peroxide will double the reaction rate).
    2. If x = 2, the curve will look like the curve for y = x2, which is a parabola (doubling the hydrogen peroxide will make the reaction go four times faster).
  4. Based on your results, what is the order of the reaction for hydrogen peroxide?

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Variations

  • Graph the number of drops on the x-axis and the time lag on the y-axis. What does this graph show?
  • Calculate the concentration of hydrogen peroxide for each sample and use this number in your analysis.
  • Vary the concentration of other reagents, such as potassium iodide or thiosulfate. Predict how the changes will affect the rate then design a procedure to test your hypothesis.
  • Devise a procedure to test how the reaction rate changes with temperature.
  • Use a graphing program to determine the value of "x" in your graph of the rate vs. hydrogen peroxide concentration.

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