Fluorescent Oscillating Reaction

Introduction

Three liquids, two colorless and one green, are mixed together in a beaker. The color of the resulting solution begins to oscillate between orange and green, with a period of about ten seconds. When the room is darkened and a black light is suspended behind the solution, the solution fluoresces and a bright orange glow is observed! The oscillations continue with the colors now alternating between dark green and bright orange. Eerie and incredibly cool!

Concepts

• Oscillating reactions
• Reaction mechanisms
• Oxidation–reduction reactions
• Kinetics/catalysts

Materials

Safety Precautions

A small amount of elemental bromine is produced during the reaction—perform the demonstration in a fume hood or wellventilated lab. 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 a strong irritant, slightly toxic, and corrosive to eyes, skin and respiratory tract. Cerium(IV) ammonium nitrate is also a strong oxidizer and a skin irritant. All the substances used in this demonstration contain 1.5 M sulfuric acid, which is corrosive to eyes and skin. 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 may be neutralized with sodium carbonate and flushed down the drain with excess water according to Flinn Suggested Disposal Method #24a.

Prelab Preparation

Solution A (0.14 M KBrO₃ in 1.5 M H₂SO₄ solution): In a 250-mL Erlenmeyer flask, add 70 mL of distilled or deionized water. Set the flask in an ice bath while slowly adding 70 mL of the 3 M H₂SO₄ solution. Once the solution cools, add 3.3 g of potassium bromate, KBrO₃, to the Erlenmeyer flask and stir to dissolve. Adjust the quantities as needed for the number of demonstrations.

Solution B (0.50 M malonic acid, CH(CO₂H)₂, in 1.5 M H₂SO₄ solution): In a 250-mL Erlenmeyer flask, add 70 mL of distilled or deionized water. Set the flask in an ice bath while slowly adding 70 mL of the 3 M H₂SO₄ solution. Once the solution cools, add 7.5 g of malonic acid, CH(CO₂H)₂, to the Erlenmeyer flask and stir to dissolve. Adjust the quantities as needed for the number of demonstrations.

Procedure

1. Add a stir bar to the 400-mL tall form beaker and place the beaker on a magnetic stirrer.
2. Add 140 mL of Solution A and 140 mL of Solution B to the beaker and adjust the stirrer to produce a small vortex in the solution.
3. Place the UV light source to the side of the beaker and turn the light source on.
4. Add 70 mL of the mixed cerium(IV)–ruthenium(II) solution to the beaker.
5. Dim the room lights.
6. The color of the solution will oscillate between dark green and brilliant, fluorescent, red-orange!

Student Worksheet PDF

12675_Student1.pdf

Teacher Tips

• This kit contains enough chemicals to perform the demonstration seven times: 500 mL of the cerium(IV) and ruthenium(II) solution, 1000 mL of 3 M H₂SO₄ solution, 30 g of KBrO₃ and 53 g of malonic acid.
• The explanation of this oscillating reaction can be quite complicated. However, it is not necessary to fully understand the reaction mechanism in order to appreciate the spectacular chemistry that occurs in the demonstration. This demonstration can be performed at any level of science with the explanation suited to the level of the class. A complete discussion is included in this handout; further information can also be found by reviewing the original literature (see the References section).
• The oscillating chemical reaction is very sensitive to chloride ions. Do not use tap water or rinse glassware with tap water.
• A hand-held UV light source like the Flinn Scientific model AP1901 is recommended for this demonstration.

Answers to Questions

1. Describe the observations in this oscillating chemical reaction.

   Two solutions, malonic acid, and sodium bromate are added to a tall form beaker. The solution is clear and colorless. When a yellow ruthenium(II) complex ion solution is added, the color of the solution oscillates between yellow and dark green. When the lights are dimmed and a UV light is shined through the beaker, the yellow appears as a bright fluorescent orange.

2. This oscillating reaction involves two competing processes in which bromate ions are reduced. The first, Process A, occurs when the bromide ion concentration is low and the second, Process B, occurs when the bromide concentration is high. Write the chemical equation for the following steps in both processes.
   1. Process A, Part 1. Bromate ions are reduced to bromine and water by the ruthenium(II) complex ions in the presence of hydrogen ions.

      BrO₃^– + 12H⁺ + 10[Ru(II)(bipy)₃]²⁺ → Br₂ + 10[Ru(III)(bipy)₃]³⁺ + 6H₂O

   2. Process B, Part 1. Bromate ions are reduced to bromine and water by bromide ions in the presence of hydrogen ions.

      BrO₃^– + 5Br^– + 6H⁺ → 3Br₂ + 3H₂O

   3. Write the equation for the overall chemical reaction of malonic acid and bromate ions to produce bromide ions, carbon dioxide and water.

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

Discussion

The fluorescent oscillating reaction is a modified version of the classic Belousov–Zhabotinsky (BZ) oscillating reaction. The overall reaction involves the ruthenium-catalyzed oxidation of malonic acid by bromate ions in dilute sulfuric acid. The bromated ions are reduced to bromide ions, while the malonic acid is oxidized to carbon dioxide and water.

The overall reaction is 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, that is, the pathway by which reactants are transformed into products.

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

Process A

Bromate ions are reduced by [Ru(II)(bipy)₃] complex ions to produce bromine and [Ru(III)(bipy)₃]ions through a simple redox reaction as shown in Equation 2a. The bromine reacts with malonic acid (Equation 3b) to form bromomalonic acid and bromide ions. The bromomalonic acid in turn reacts with the [Ru(III)(bipy)₃] complex ions (Equation 2b) to form carbon dioxide, bromide ions and formic acid. As the concentration of bromide ions increases, the rate of Equation 3a increases until eventually Process B dominates.


Process B

Bromate ions are reduced by bromide ions as shown in Equation 3a. The amber tint which may develop is caused by the production of elemental bromine. This tint soon disappears as the bromine reacts with malonic acid as shown in Equation 3b.

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

As the reaction oscillates between Process A and Process B, triggered by changes in the bromide ion concentration, concentrations of the ruthenium complex ions oscillate as well—these concentration changes explain the observed color changes. During Process A, the ruthenium complex ions are oxidized from the orange Ru(II) complex ion to the dark green Ru(III) complex ion.

The solution stays orange until the most of the Ru(II) complex ions have been oxidized, at which time the solution color switches to dark green. When process B takes over, The Ru(III) complex ions are now reduced, and the solution flashes back to orange. These oscillations can last up to an hour, until all of the malonic acid has reacted. The Ru(II) complex ion is fluorescent.

The reaction can also be visualized as shown in Figure 1: (Winfree, 1974)

Color Changes and Steps for Figure 1

1. Start at the middle of the diagram. Initially, the solution is orange. Bromate ions are reduced to bromine by Ru(II) complex ions (Equation 2a), which are oxidized to Ru(III). When the ratio of Ru(III) to Ru(II) complex ions is large enough, the solution turns dark green. Ru(II) complex ions are responsible for the fluorescence.
2. Bromine (Equation 2a) reacts with malonic acid (Equation 3b) to form bromomalonic acid. During this process, Ru is in the oxidized Ru(III) state, and the solution is dark green.
3. Bromomalonic acid now reduces the Ru(III) complex ions to Ru(II) complex ions, (Equation 2b), releasing carbon dioxide and bromide ions. Once formed, the Ru(II) complex ions continue to reduce bromate ions (Equation 2a) and the solution remains dark green.
4. After a time the high bromide ion concentration shuts off reaction 2a and reaction 3a starts. The concentration of Ru(II) complex ions builds until the solution switches to bright fluorescent orange again.
5. When most of the bromide ions have been consumed, reaction 3a shuts off and 2a begins again.
6. The process repeats itself until either bromate or malonic acid is exhausted.

References

Flinn Scientific would like to thank Rhonda Reist, chemistry teacher at Olathe North High School in Olathe, KS, for sharing the original idea and procedure for this demonstration with us.

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.