Catalytic Spectrum Show

Demonstration Kit


Catalysts are vital components in the world economy. Approximately 90% of all manufactured goods employ a catalyst in some step of the production process. Show students the dramatic effect of a catalyst on the rate of a reaction as two similar solutions undergo the same color changing reaction at vastly different speeds.


  • Catalyst
  • Rate of reaction
  • Collision theory
  • Redox reaction
  • Acid–base indicators


(for each demonstration)
Ammonium molybdate, (NH)6Mo7O24•4H2O, 0.08 g*
Hydrogen peroxide solution, H2O2, 10%, 40 mL*
Sodium acetate, CH3CO2Na•3H2O, 3.8 g*
Sodium hydroxide solution, 0.2 M, 50 mL*
Sodium thiosulfate, Na2S2O3•5H2O, 8.7 g*
Universal indicator solution, 3 mL*
Balance, centigram (0.01-g precision)
Distilled or deionized water
Erlenmeyer flask, 1000-mL
Erlenmeyer flasks, 500-mL, 3
Graduated cylinder, 10-mL
Graduated cylinders, 50-mL, 2
Graduated cylinder, 250-mL
Magnetic stirrer and stir bar
Marking pen
Timer, seconds
Universal indicator color charts, 30*
Weighing dishes, 3
*Materials included in kit.

Safety Precautions

Ammonium molybdate is a skin and eye irritant and is moderately toxic by ingestion. Hydrogen peroxide solution is an oxidizer and a skin and eye irritant. Sodium hydroxide solution is slightly corrosive. Sodium thiosulfate is slightly toxic by ingestion and is a body tissue irritant. Universal indicator solution is alcohol-based and therefore flammable. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please consult current Safety Data Sheets for additional safety, handling and disposal information.


Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The final reaction solution in each flask may be combined, neutralized, if needed, and disposed of according to Flinn Suggested Disposal Method #26b. Universal indicator solution may be disposed of according to Flinn Suggested Disposal Method #26b. 

Prelab Preparation

  1. Add 400 mL of distilled or deionized water to a 1000-mL Erlenmeyer flask.
  2. Measure the following amounts of chemicals into separate weighing dishes.
    1. 8.7 g of sodium thiosulfate.
    2. 3.8 g of sodium acetate.
  3. Place the Erlenmeyer flask on a magnetic stirrer and add a stir bar.
  4. Add the sodium thiosulfate and the sodium acetate to the water in the Erlenmeyer flask and stir to mix the solution.
  5. Measure 50 mL of 0.2 M sodium hydroxide solution in a graduated cylinder and add the solution to the Erlenmeyer flask.
  6. Add distilled or deionized water to bring the solution volume to 1000 mL. Stir to mix thoroughly.
  7. Add 3–5 mL of universal indicator solution to obtain a bright purple solution that is easily visible to all students in the class.
  8. Measure 0.08 g of ammonium molybdate into a clean weighing dish and set aside.
  9. Distribute the universal indicator color charts to the class.


  1. Label three 500-mL Erlenmeyer flasks, “Catalyst,” “No Catalyst” and “Control.”
  2. Place the three Erlenmeyer flasks in a row on the demonstration table.
  3. Add 225 mL of the prepared solution (from step 7 in the Prelab Preparation section) to each of the Erlenmeyer flasks.
  4. Add the 0.08 g of ammonium molybdate the flask marked “Catalyst” and swirl the flask to dissolve.
  5. Add 20 mL of 10% hydrogen peroxide solution to each of two separate 50-mL graduated cylinders.
  6. Simultaneously add 20 mL of the hydrogen peroxide solution to each of the flasks marked “Catalyst” and No Catalyst,” respectively. Note: The third flask is a control for color comparison.
  7. Start the timer. Observe and compare the color changes in the two reaction flasks. The solution in the “Catalyst” flask will change from purple to blue to green to yellow to orange and, finally, to red-orange, over the course of one to two minutes. The solution in the “No Catalyst” flask will undergo the same color changes, but at a much slower rate, about 10 to 15 minutes.

Student Worksheet PDF


Teacher Tips

  • This kit contains enough chemicals to perform the demonstration as written seven times: 5 g of ammonium molybdate, 65 g of sodium thiosulfate, 30 g of sodium acetate, 350 mL of 0.2 M sodium hydroxide solution, 280 mL of 10% hydrogen peroxide solution and 50 mL of universal indicator solution.
  • A student worksheet is included as an optional assessment tool for use by the instructor.
  • To speed up the reaction time in the “No Catalyst” flask, add less sodium acetate to the original solution. Adding 0.8 g of sodium acetate, instead of 3.8 g to the original solution, reduces the time needed for the uncatalyzed solution to change from purple to red-orange to nine minutes.
  • The plot of pH versus time does not represent a change in product concentration, Δ[H+], over time, but rather the change in the negative logarithm of the product concentration, Δ(–log[H+]), over time. The graph will, however, show the relative speeds of the catalyzed and uncatalyzed reactions.
  • While the colors change in the uncatalyzed solution beaker, you may have students collect the data and plot the graph in the worksheet.

Correlation to Next Generation Science Standards (NGSS)

Science & Engineering Practices

Developing and using models
Analyzing and interpreting data
Constructing explanations and designing solutions

Disciplinary Core Ideas

MS-PS1.A: Structure and Properties of Matter
MS-PS1.B: Chemical Reactions
HS-PS1.B: Chemical Reactions

Crosscutting Concepts

Energy and matter

Performance Expectations

MS-PS1-2. Analyze and interpret data on the properties of substances before and after the substances interact to determine if a chemical reaction has occurred.
HS-PS1-4. Develop a model to illustrate that the release or absorption of energy from a chemical reaction system depends upon the changes in total bond energy.
HS-PS1-2. Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.

Answers to Questions

  1. Using a timer, record in the Data Table the time each color change occurs, corresponding to the chart colors.
    Solution ColorpHCatalyzed Time (sec)Uncatalyzed Time (sec)
    Purple 10.0 0 0
    Blue 9.0 16 25
    Green 8.0 26 39
    Light Green 7.0 34 53
    Yellow 6.0 44 188
    Orange 5.0 47 720
    Red-orange 4.0 57 >1000
  2. On the graph below, plot pH versus time for both the catalyzed and the uncatalyzed reaction.
  3. The overall balanced equation for this oxidation–reduction reaction is

    Write the half-cell reactions for oxidation and for reduction.

    1. Oxidation
    2. Balance O atoms by adding H2O to side with less O atoms.
    3. Balance H atoms by adding H+ to deficient side.
    4. Add electrons, e, to balance charges on each side of half-cell reactions.
    5. To balance electrons in overall reaction, multiply reduction half-cell reaction by 4, then add half-cell reactions.


In this demonstration, hydrogen peroxide is added simultaneously to two separate flasks, both of which contain sodium acetate, sodium thiosulfate, sodium hydroxide and universal indicator in solution. One of the flasks also contains a small amount of ammonium molybdate, a catalyst while the other does not.

The overall equation for the thiosulfate ion oxidized to the sulfate ion by hydrogen peroxide is

The hydrogen ion produced (H+) is rapidly neutralized by the hydroxide ions in solution. The acetate ion acts as a buffer by replacing the hydroxide ions consumed in the neutralization reaction.
Collision theory offers a simple explanation for how fast a reaction will occur—in order for a reaction to occur, reactant molecules must first collide. Not all collisions, however, will lead to products. In order for colliding molecules to produce products, the collision energy must exceed a certain critical energy level, called the activation energy, for the reaction (see Figure 1). If the activation energy is low, almost all colliding molecules will have sufficient energy to overcome the energy barrier for the reaction. These reactions will occur very fast. In contrast, if the activation energy is high, only a small fraction of the colliding molecules will have sufficient energy to overcome the energy barrier for the reaction. These reactions will occur much more slowly.
To increase the rate of a reaction, one of two things must occur: (1) more molecules with sufficient kinetic energy to overcome the activation energy barrier must collide, or (2) the height of the activation energy barrier must be reduced.

A catalyst increases the rate of a reaction because it decreased the activation energy that is needed for reactants to be transformed to products. In general, a catalyst provides a modified or new pathway of the reaction to occur. The new reaction pathway has a lower activation energy. When the activation energy for the reaction is reduced, the fraction of colliding molecules that have enough energy to overcome this energy barrier increases (see Figure 2).
The pH value of the solution remains relatively constant until all the hydroxide ions are consumed. The pH then steadily falls until one of the other reactants (thiosulfate ion or hydrogen peroxide) is consumed. The pH changes can be seen as the universal indicator changes color from purple to blue to green to yellow to red-orange. The catalyzed solution color changes from purple to red-orange in 1–2 minutes. The uncatalyzed solution will take about ten to fifteen minutes to change from purple to orange, and about one hour to change from orange to red-orange.

As the reaction proceeds, the concentration of the hydrogen ion increases. By plotting the times at which the color changes occur versus the pH values of the indicator colors, the relative speeds of the catalyzed and uncatalyzed reactions can be compared.

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