The Overhead Oscillating Clock

Introduction

Oscillating clock reactions are mesmerizing due to their colorful and unpredictable changes. This variation of a classic oscillating chemical reaction is easily performed on the overhead projector to give travelling waves of blue through a red solution.

Concepts

• Oscillating reactions
• Oxidation–reduction
• Reaction mechanisms

Materials

Safety Precautions

A small amount of bromine gas is released from the reactions in this demonstration; adequate ventilation is necessary. Potassium bromate is a strong oxidizing agent and poses a fire risk in contact with organic material; it is a strong irritant and moderately toxic. Malonic acid is slightly toxic and a strong irritant; it is corrosive to eyes, skin and respiratory tract. Sodium bromide is slightly toxic by ingestion and a severe body tissue irritant. Ferroin solution contains ferrous sulfate and 1,10-phenanthroline. Ferrous sulfate is slightly toxic by ingestion and 1,10-phenanthroline is highly toxic by ingestion. Sulfuric acid solution is corrosive to eyes, skin, mucous membrane and other body tissue. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

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 reaction mixture should be neutralized with sodium carbonate and flushed down the drain with excess water according to Flinn Suggested Disposal Method #24a.

Procedure

1. Add 5 mL of solution A to a clean Petri dish.
2. Add 1 mL of solution B and 0.5 mL of solution C to the Petri dish.
3. Carefully swirl the Petri dish until the solution becomes clear and colorless.
4. Add 2 mL of ferroin solution. Swirl the dish. The solution should initially oscillate between pale red and pale blue.
5. Place the Petri dish on the overhead.
6. After the color oscillations stop (approximately 5 minutes), small blue dots will begin to appear.
7. The dots will expand, creating concentric rings of blue.
8. When the rings start overlapping, swirl the Petri dish to restart the process. Oscillations will continue for about 30 minutes.

Student Worksheet PDF

12548_Student1.pdf

Teacher Tips

• The kit contains enough chemicals to perform the demonstration seven times. Additional amounts are provided to allow for demonstration in a beaker using scaled-up amounts of chemicals.
• The oscillating reaction can be performed in a beaker by scaling up the reagent amounts. For example, by multiplying the amounts by 10, you get 50 mL of Solution A, 10 mL of Solution B, 5 mL of Solution C and 20 mL of Solution D. The oscillating colors are blue, violet, and red. They should oscillate every two to three minutes and last over 30 minutes.
• The reaction is extremely sensitive to chloride ions. Be sure to rinse all glassware with distilled or deionized water.
• Bubbles of carbon dioxide gas will be seen to form throughout the solution and will appear as black dots on the overhead screen.
• As the blue dots grow, a red dot will form inside the expanding blue dot, followed by a blue dot inside the red one, and so on.
• In the shallow Petri dish, the colors will apear faint blue and orange.
• To darken the color of solution, add a few additional drops of ferroin solution.

Answers to Questions

1. Describe what happened in this demonstration.

Three solution, sodium bromate, malonic acid and sodium bromate are all added to a Petri dish. The solution is clear and colorless. Then ferroin solution is added, and the solution oscillates between red and blue. After the color changes stop, blue rings begin to appear, and when the Petri dish is swirled the red-blue oscillations begin again.

2. This oscillating reaction involves two competing processes in which bromate ions are reduced. The first, Process A, occurs when bromide ion concentration is above a certain level, and the second, Process B, occurs when the same concentration is below a certain level. Write the chemical equation for the following steps in both processes.

a. Process A, Part 1. Bromate ions are reduced by bromide ions in the presence of hydrogen molecules.

BrO₃^– + 5Br^– + 6H⁺ → 3Br₂ + 3H₂O

b. Process A, Part 2. Bromine reacts with malonic acid [CH₂(CO₂H)₂].

Br₂ + CH₂(CO₂H)₂ → BrCH(CO₂H)₂ + Br^– + H⁺ 

c. Process B, Part 1. Bromate ions are reduced to bromine by iron(II) ions. 

BrO₃^– + 12H⁺ + 10Fe²⁺ → Br₂ + 10Fe³⁺ + 6H₂O 

3. Write the equation for the overall chemical reaction, in which bromate ions are reduced to bromide ions and malonic acid is oxidized to carbon dioxide and water.

3CH₂(CO₂H)₂ + 4BrO₃^– → 4Br^– + 9CO₂ + 6H₂O

4. The Fe(II) complex is red, while the Fe(III) complex is blue. During Process A the solution is red, but during Process B the solution is blue. Knowing this, explain how red, violet and blue are all produced in the solution.

Iron remains in its reduced Fe(II) state during the first process, giving the solution its red color. Once Process B begins, Fe(II) is oxidized to Fe(III). Right now both complexes exist, and the solution is a mixture of the red and blue colors combining to form violet. When most of the Fe(II) has been oxidized to Fe(II), the solution is blue.

Discussion

This oscillating reaction demonstrates the modified Belousov-Zhabotinsky (BZ) reaction, which is an iron-catalyzed bromate–malonic acid reaction.

The overall reaction occurring in this demonstration is the iron-catalyzed oxidation of malonic acid by bromate ions in dilute sulfuric acid. The bromate ions are reduced to bromide ions, while the malonic acid is oxidized to carbon dioxide and water. The overall reaction can be represented by Equation 1:

In order to gain some understanding and appreciation for how this overall reaction can produce the amazing, repetitive color changes observed in the demonstration, it is necessary to look at the reaction mechanism or, in other words, how the reactants are transformed into products.

The mechanism involves two different competing processes—Process A involves ions and two-electron transfers; Process B involves radicals and one-electron transfers. The dominant process at any particular time depends on the bromide ion concentration. Process A (see Equation 2a) occurs when the bromide ion concentration rises above a certain critical level, while Process B (see Equation 3a) is dominant when the bromide ion concentration falls below a critical level. Oscillations occur because Process A consumes bromide ions, leading to conditions which favor Process B. Process B (indirectly) produces bromide ions, leading to conditions that favor Process A.

Process A

Bromate ions are reduced by bromide ions through a series of oxygen transfers (two-electron reductions) as shown in Equation 2a. This reaction occurs when Solutions A and B are mixed. The amber color which may develop is caused by the production of elemental bromine. This color soon disappears as the bromine reacts with malonic acid as shown in Equation 2b.

Process A results in an overall decline in the bromide ion concentration. Once the necessary intermediates have been generated and most of the bromide ions have been consumed, the rate becomes negligible and Process B takes over.

Process B
Bromate ions are reduced by iron(II) ions to produce bromine through the overall redox reaction shown in Equation 3a. Process B produces Fe(III) ions and Br2. Both of these species react at least in part to oxidize the malonic acid (see Equation 2b) and the bromomalonic acid (see Equation 3b) to form additional bromide ions. As the concentration of bromide ions increases, the rate of Equation 2a increases until eventually Process A once again dominates.

As the reaction oscillates between Process A and Process B, triggered by changes in the bromide ion concentration, the concentrations of the two different iron ions in solution oscillate as well—these concentration changes will explain the color changes observed. While Process A occurs, the iron ions are in their reduced state, Fe(II). During Process B, some iron ions are oxidized to Fe(III) and thus the ratio of Fe(II) to Fe(III) oscillates as well.

The indicator used in this demonstration is ferroin, which is tris(1,10-phenanthroline) ferrous sulfate. The Fe(II) complex is red while the Fe(III) complex is blue; thus the color of the solution changes as the iron is oxidized or reduced.

The color changes are outlined as follows:
Red = All of the blue Fe(III) has been reduced to the red Fe(II) complex; the solution appears red.
Violet = Equation 3a is oxidizing the red Fe(II) complex to the blue Fe(III); the mixture of blue and red appears violet.
Blue = Fe(II) has been oxidized to the blue Fe(III) complex; the solution appears blue.

The reaction can also be visualized using Figure 1: (Winfree, 1974)

Steps for Figure 1
(Start at about 7 o’clock on Figure 1.)

1. Initially, bromide reduces bromate (Equation 2a) to form bromine, which in turn reacts with malonic acid (Equation 2b) to form bromomalonic acid. During this process, Fe is in its reduced Fe(II) state, and the solution is red.
2. When the bromide ion concentration drops below a threshold level, the second set of reactions starts to dominate. Bromate is now reduced to bromine by the ferrous ion [Fe(II)], instead of bromide (Equation 3a). Fe(II) is oxidized to Fe(III). This combination of ferrous complex and ferric complex produces the violet color.
3. The solution turns blue when most of the Fe(II) complex is oxidized to Fe(III).
4. Bromomalonic acid now reduces the ferric phenanthroline to ferrous phenanthroline, (Equation 3b), releasing carbon dioxide (CO₂), and bromide ions. The ferrous phenanthroline continues to reduce bromate (Equation 3a) to produce ferric phenalthroline and bromine. After time, the high concentration of bromide shuts off reaction 3a and starts 2a. The ferrous phenanthroline builds up and the solution switches to red.
5. The process repeats itself until either bromate or malonic acid is exhausted.

References

Field, R. J.; Schneider, F. W., J. Chem. Ed., 1989, 66, pp. 196–204

Kolb, D. J. Chem. Ed., 1988, 65, 1004.

Pojman, J. A., Craven, R., Leard, D. C., J. Chem. Ed., 1994, 71, pp. 84–90.

Shakhashiri, B. Z., Chemical Demonstrations: A Handbook for Teachers of Chemistry; University of Wisconsin Press: Madison; 1985; Vol. 2, pp. 297–300.

Winfree, A. T., Sci. Am., 1974, 230, 82.