March 2014

# Sharing chemistry with the community: A colorful catalyst

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Meredith Rahman*, Duke University, Durham NC

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The catalytic oxidation of tartrate ions by hydrogen peroxide has been a staple of the Duke chemistry instructional program for many years. The demonstration clearly and colorfully brings to life the textbook description of what a catalyst is and how it enables a reaction to proceed at a faster rate by becoming involved in the reaction, yet in the end returns to its original state. A pink-colored solution of cobalt(II) chloride (the catalyst) is added to a warm mixture of sodium potassium tartrate and hydrogen peroxide (the reaction mixture). One can visually see the color change from pink to green while, at the same time, the rate of gas generation (carbon dioxide) steadily increases. Later, the effervescence slows to zero, and the pink coloration of the solution returns illustrating that the catalyst, although involved in the reaction, is not consumed. The net reaction for the oxidation of tartrate ions by hydrogen peroxide is:

5H2O2(aq) + C4H4O62-(aq) → 4CO2(g) + 2OH1-(aq) + 6H2O(l)

and is believed to occur in two steps:

3H2O2(aq) + C4H4O62-(aq) → 2CO2(g) + 2HCO21-(aq) + 4H2O(l)

2HCO21-(aq) + 2H2O2(aq) → 2CO2(g) + 2H2O(l) + 2OH1-(aq)

The reaction proceeds slowly, even when warmed to around 50oC.

Cobalt(II) ions serve as a homogeneous catalyst for this reaction. The hydrated cobalt(II) ions (pink) are oxidized to cobalt(III) ions (green) by the hydrogen peroxide. The tartrate ions bind with the cobalt(III) ions and are subsequently oxidized as the cobalt(III) ions are reduced back to cobalt(II). These reactions continue until the hydrogen peroxide (limiting reagent) is consumed and all the cobalt has been reduced by the remaining tartrate. The addition of a small amount of the hydrogen peroxide to the remaining mixture will enable the reactions to proceed anew.

## A colorful catalyst

### Concepts

- Catalysis, homogeneous catalyst

### Materials

- 400 mL of 0.3 M sodium potassium tartrate solution
  (84.6 g solid sodium potassium tartrate tetrahydrate diluted to 1.0 L)
- 130 mL of 6% hydrogen peroxide solution (combine 60 mL of 30% hydrogen peroxide with 240 mL of water)
- 15 mL of 0.15 M cobalt(II) chloride solution (5.0 grams of solid cobalt(II) chloride hexahydrate diluted to 125 mL)
- Two large beakers (I use 600-mL tall-form beakers)
  Thermometer
- Hot plate
- 3 graduated cylinders for measuring solutions

Reaction mixture without the catalyst.

Catalyst added.

Green bubbles.

Peroxide almost used up.

Cobalt (III) converting to cobalt (II).

Reaction at completion.

### Safety

- Wear safety glasses/goggles
- Wear protective gloves
- Hydrogen peroxide is an oxidizer and an irritant; avoid contact.
- Cobalt(II) chloride hexahydrate is moderately toxic and a suspected carcinogen; avoid contact and/or ingestion.
- The reaction is exothermic; use caution when handling the beakers to avoid burns.

### Advance preparation

- Combine 200 mL of 0.3 M sodium potassium tartrate and 65 mL of 6% hydrogen peroxide into each of the two large beakers.
- Warm the mixtures on a hot plate to a temperature between 45 and 50oC and then maintain the temperature within this range.

### Procedure

- Remove one of the two beakers containing the reaction mixture from the hot plate and have the audience note whether a reaction is taking place. There likely will be just a few visible bubbles of carbon dioxide.
- Measure 15 mL of the 0.15 M cobalt(II) chloride solution; have the audience note the pink coloration due to the presence of hydrated cobalt(II) ions in the solution.
- Add the cobalt(II) chloride solution to the large beaker containing the reaction mixture; have the audience observe and note the changes. The mixture changes to a deep green coloration with vigorous effervescence; eventually the effervescence ceases and the color of the mixture returns to being pink. (see photos on previous page)
- If desired, you can add a small sample of the hydrogen peroxide (5-10 mL) to the resulting mixture. The mixture again turns green with vigorous effervescence. When the effervescence stops the color changes back to pink. This shows that the hydrogen peroxide was the limiting reagent and is the reason why the reaction ceased. One can also deduce that the cobalt(II) ions are first interacting with the hydrogen peroxide, not the tartrate ions.
- Add a portion of this pink solution to the second beaker containing the reaction mixture; the process repeats as before but the colors will be less intense due to dilution. However, the changes will take place much faster. This is due to the reaction being exothermic. The temperature of the solution from the first trial is now significantly above the original temperature range.

### Disposal

- Place the contents of the beakers in a designated waste container for disposal.
- Consult local regulations for disposing of the chemical waste.

Meredith Rahman took on the challenge of learning this demonstration, the underlying chemistry and means to convey the chemistry appropriate for a high school audience. She has subsequently performed this demonstration for students at two different high schools. The following are her thoughts on learning and performing the Colorful Catalyst demonstration.

The first time I presented the Colorful Catalyst demonstration was for the Kinston High School physics class. I practiced the mechanics of the demonstration and in rehearsal had felt comfortable with the flow of my talk. However, when I actually performed it, the reaction seemed to go much more slowly. I became uncomfortable with this unanticipated pace and feared that the audience was losing interest. I felt unsure of what to say during this extended period of waiting, leaving me feeling less than successful in conveying the chemistry. Upon reflection, I realized the potential benefit of this “down time”. Having students quietly observe the process without distraction actually enables them to have the time to experience the reaction and ponder catalysts.

A few weeks later, I performed Colorful Catalyst at St. Mary’s School, an all girls’ high school in Raleigh. I incorporated what I had learned previously into this performance. I crafted my script to be better sequenced with the reaction, allowing the girls enough time to carefully focus on the process without interruption. These modifications were beneficial, enhancing the level of engagement of the students. And based on the question responses, the girls demonstrated a greater understanding of the underlying chemical concepts. Students not only made better observations, but had time to consider the significance of these observations. For example, the students correctly connected the change in color to the catalytic involvement and the increased effervescence to the faster rate of reaction.

Unlike lectures or reading texts, a demonstration enables the students to learn from actual experience. This made the definition of a catalyst real and memorable. The demo presentation was satisfying for me because it clarified a chemical concept that I had once found confusing. I was able to share this insight with the students. Learning, practicing and performing this demonstration helped the students — and me — to better understand catalysis. Careful planning and sequencing of one’s talk, and allowing time for students to have uninterrupted observation and thought, enhance its effectiveness of the demonstration. As a result of these experiences, I have more confidence as a presenter and look forward to my next opportunity to share Colorful Catalyst with the community.

*Meredith Rahman is a senior at Duke University majoring in Evolutionary Anthropology with a minor in biology. She is considering a career in medicine upon graduation.

**Dr. Kenneth Lyle is a lecturing-fellow in the Department of Chemistry at Duke University. All inquiries should be addressed to kenneth.lyle@duke.edu.

The Powell Family Trust, the Duke-Durham Neighborhood Partnership, and Biogen Idec – Research Triangle Park, fund the Duke Chemistry Outreach Program.

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