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Areas of Science Chemistry
Time Required Average (6-10 days)
Prerequisites You should be familiar with chemical reactions and their rates.
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 Hydrochloric acid is corrosive to hands and eyes. Wear splash goggles and gloves when when working with all chemicals. Avoid direct contact with chemicals as hydrogen peroxide and the iodine-starch complex can stain skin and clothing.


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 project, you will explore factors that affect the rate of the iodine clock reaction.


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

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David B. Whyte, PhD, Science Buddies
Edited by Svenja Lohner, PhD, Science Buddies

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General citation information is provided here. Be sure to check the formatting, including capitalization, for the method you are using and update your citation, as needed.

MLA Style

Science Buddies Staff. "Investigate the Kinetics of the Amazing Iodine Clock Reaction." Science Buddies, 21 Apr. 2021, Accessed 17 May 2021.

APA Style

Science Buddies Staff. (2021, April 21). Investigate the Kinetics of the Amazing Iodine Clock Reaction. Retrieved from

Last edit date: 2021-04-21


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 change 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. A demonstration of this reaction is shown in the video below.

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 by monitoring the color change of the reaction. 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 using 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 = 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 does the triiodide not react with it? The reason the triiodide does not react with the starch is that it is immediately consumed in a reaction with the thiosulfate.

Equation 2:

  • I3 = 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- = 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:

  • (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:

  • 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
  • Reaction order


  • 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 a 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.


  • Wikipedia Contributors. (2009, September 28). Iodine clock reaction. Wikipedia: The Free Encyclopedia. Retrieved October 30, 2009.
  • Department of Chemistry, The University of North Carolina at Chapel Hill. (2008). Kinetics. Retrieved October 30, 2009.

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Materials and Equipment

  • Iodine clock reaction kit; available at Carolina Biological Supply Company
    • You will need following solutions from the kit: 0.05 M potassium iodide solution, 3% hydrogen peroxide, 0.01 M sodium thiosulfate, 0.1 M hydrochloric acid, 1% starch solution.
    • 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.
  • Chemical splash safety goggles; available at Carolina
  • Mini cup, 2 oz.
  • Disposable gloves. If you are allergic to latex, use vinyl or polyethylene gloves.
  • Newspaper to protect work area
  • Distilled water; available at grocery stores
  • 3 mL graduated pipettes (6); available at Amazon
  • Stopwatch or other timer
  • A white sheet of paper
  • Permanent marker
  • Lab notebook

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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. Start the iodine clock reaction (beginning with reaction 1) by adding the potassium iodide, hydrocholoric acid, starch, sodium thiosulfate, and distilled water to the mini cup using a fresh pipette for each chemical. The volumes of each reactant are given in Table 1 below.
ChemicalReaction 1Reaction 2Reaction 3Reaction 4Reaction 5Reaction 6Reaction 7Reaction 8
Potassium iodide (KI)2 mL2 mL2 mL2 mL2 mL2 mL2 mL2 mL
Hydrochloric acid (HCl) 1 mL1 mL1 mL1 mL1 mL1 mL1 mL1 mL
Starch solution 0.5 mL0.5 mL0.5 mL0.5 mL0.5 mL0.5 mL0.5 mL0.5 mL
Sodium thiosulfate (Na2S2O3) 0.5 mL0.5 mL0.5 mL0.5 mL0.5 mL0.5 mL0.5 mL0.5 mL
Distilled Water (H2O) 1 mL1.5 mL2 mL2.5 mL3 mL3.5 mL4 mL4.5 mL
Hydrogen peroxide (H2O2) 4 mL3.5 mL3 mL2.5 mL2 mL1.5 mL1 mL0.5 mL
Table 1. Reaction volumes of each chemical for all the iodine clock reactions you will perform.
  1. Swirl the cup briefly to mix the chemicals. The solution should be clear as shown in Figure 1.
Bottles of potassium iodide, hydrochloric acid, starch solution and sodium thiosulfate poured into a plastic cup
Figure 1. Mixing all the reactants (except hydrogen peroxide) results in a clear solution.
  1. Place the mini cup on a white sheet of paper to make the color change more visible.
  2. Then use a fresh pipette to suck up the respective volume of hydrogen peroxide for your reaction.
  3. Add all the hydrogen peroxide at once to the mini cup and start your stopwatch as soon as the hydrogen peroxide is added.
  4. Look at the solution against a white background and stop the stopwatch once you can first see some blue color developing.
  5. Record this time (the reaction time) in your lab notebook.
  6. Carefully pour the dark blue reaction solution down the sink with plenty of cold running water to wash it down. Be careful with this solution, as it contains iodine and will stain surfaces with which it comes in contact.
  7. Repeat steps 3–10 two more times with the same hydrogen peroxide concentration to make sure you get reproducible results. Keep the temperature the same for each reaction as temperature also affects the reaction time.
  8. Once you have finished reaction 1, start varying the hydrogen peroxide concentration. Repeat steps 3–11 and change the hydrogen concentration according to Table 1 until you have tried all eight reactions. Remember, each time, you decrease the amount of hydrogen peroxide, you have to increase the amount of water respectively. This will keep the total volume of the reaction constant, while changing the concentration of the hydrogen peroxide.

Analyzing Your Results

  1. Use your data to make a data table that shows the following:
    1. The amount of added hydrogen peroxide, in mL.
    2. The determined reaction time (or time lag) for each test, in seconds (s).
    3. The average of the reaction times calculated from your three repeated trials for the same H2O2 concentration. For example, average the three trials for the reaction with 3 mL of hydrogen peroxide.
    4. The reciprocal of the average reaction time for each test, in 1/s.
  2. Graph the amount of hydrogen peroxide (in mL) 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 Amount 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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  • Graph the amount of hydrogen peroxide (in mL) 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 versus hydrogen peroxide concentration.

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