Hot and Cold Equilibrium—A Demonstration of Le Chatelier’s Principle

Introduction

Three test tubes filled with the same violet liquid are displayed to the class. One tube is placed in cold water—the color of the liquid turns pink. Another tube is placed in hot water—the color of the liquid turns blue. The color of the solution in the third control tube at room temperature remains unchanged. When the tubes are exchanged between the hot and cold water, the colors dramatically reverse!

Concepts

• Equilibrium/Le Chatelier’s principle
• Density
• Absorption/Transmission

Materials

Safety Precautions

This activity requires the use of hazardous components and/or has the potential for hazardous reactions. The isopropyl alcohol in the cobalt chloride/alcohol solution is a flammable liquid and a fire hazard. Avoid open flames while working with this solution. The solution is slightly toxic by ingestion and inhalation and is a body tissue irritant. Avoid contact with all body tissues. 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 cobalt chloride/alcohol solution may be stored in a flammables cabinet in its original bottle and used again from year to year. If necessary, contact a licensed removal company according to Flinn Suggested Disposal Method #27f.

Procedure

1. Fill each of three large test tubes with approximately 50 mL of the violet cobalt chloride solution (about ¾ full). Place the rubber stoppers loosely on the tubes.
2. Fill one large beaker with a mixture of ice and water. Fill a second large beaker with tap water and heat it to 70–80 °C on a hot plate. Note: Do not heat the water higher than 80 °C since the boiling point of isopropyl alcohol is 82.4 °C.
3. Display the three test tubes containing identically colored solutions to the class. Write the equation for the reaction on the board. Emphasize that each solution has a violet color at room temperature.

Ask the students to predict the color of the solution if one tube is cooled and one is heated.

4. Place one of the test tubes into the ice water and one into the hot water. Set the third tube nearby on the demonstration table as a color reference (control).
5. Have students make observations of the solutions in the tubes. Look carefully to determine the location in each tube where the color changes are occurring. The solution in the ice water tube will turn pink, starting at the bottom of the tube. The solution in the hot water tube will turn a deep blue, starting at the top of the solution. Be certain to point out this phenomenon to the students. The colder pink solution has a greater density and sinks (notice that the solution is more pink at the bottom) while the hotter solution has a lower density and rises (notice that the solution is more blue at the top).
6. Students at this point may ask what will happen if the tubes are reversed in the beakers. Allow them to hypothesize and discuss what will occur.
7. Reverse the tubes in the two beakers, placing the tube containing the cold solution into the hot water and the tube containing the hot solution into the ice water. Watch the spectacular color change occur as the solutions reverse colors. Make observations and discuss.
8. When finished with the demonstration, return the solutions from the three test tubes to the empty labeled container provided in the kit. Seal the container tightly and store in a dedicated flammables cabinet for future use.

Student Worksheet PDF

11903_Student1.pdf

Teacher Tips

• The cobalt chloride solution, test tubes, and stoppers provided in this kit may be reused from class to class. Store the cobalt chloride solution in the original container rather than in the test tubes. Seal the bottle tightly and store it in a dedicated flammables cabinet (Flinn compatible chemical family code, Organic #2).
• The initial color of the cobalt chloride solution depends on the concentration of cobalt chloride and on the amount of water in the alcohol solvent. The solution should be violet (purple) to start, with roughly equal amounts of the pink and blue cobalt complex ions.
• If the initial color of the cobalt chloride solution is blue, add a few drops of water to restore the original purple color.
• The demonstration only calls for 150 mL of solution (50 mL per tube) while 300 mL of solution is provided in the kit. The extra solution is provided to allow for any evaporation or spills that may occur.
• If a hot plate is not available, use hot tap water from the laboratory sink. The demonstration will work with 60 °C (warm) water, although the color changes will not be as immediate and dramatic. Do not use a Bunsen burner to heat the water—the solution is flammable.

Answers to Questions

1. When the violet cobalt chloride is placed in the beaker containing cold water, what color does it turn? Where in the tube does the color change begin and why?

The solution in the ice water bath will turn pink starting at the bottom of the test tube because cold water is less dense and therefore sinks.

2. When the violet cobalt chloride is placed in the beaker containing hot water (70–80 °C), what color does it turn? Where in the tube does the color change begin and why?

The solution placed in a hot water bath will turn deep blue starting at the top of the test tube because warm water is less dense and therefore rises to the top of the tube.

3. Le Chatelier postulated that if a stress such as a change in concentration, pressure or temperature is applied to a system at equilibrium, the equilibrium is shifted in a way that compensates for the effects of that stress.

1. Given this information, what effect does an increase in temperature have on this system at equilibrium? Write the chemical equation and use arrows above each reactant and product to indicate whether the concentration is increasing or decreasing.
   {11903_Answers_Equation_2}

   The equilibrium shifts to the right.

2. What effect does a decrease in temperature have on this system at equilibrium? Write chemical equation and use arrows above each reactant and product to indicate whether the concentration is increasing or decreasing.
   {11903_Answers_Equation_3}

   The equilibrium shifts to the left.

4. Use the color wheel and graph below to determine why cobalt(II) chloride appears violet at room temperature.

In the presence of white light the color that a person sees is the color being transmitted. That means its complementary color is absorbed. Looking at the graph one can see at room temperature cobalt(II) chloride absorbs both blue green and orange–red, therefore the color it transmits will lie somewhere in that region. 580 nm is between 525 nm and 675 nm where the graph displays cobalt(II) chloride absorbs light and at 580 nm green/yellow light is absorbed. Therefore the color it transmits at room temperature is violet.

Discussion

This demonstration illustrates the effect of a change in the temperature on a system at equilibrium, reinforcing Le Chatelier’s Principle [Henry Louis Le Chatelier (1850–1936)]. Le Chatelier postulated that if a stress, such as a change in concentration, pressure, or temperature, is applied to a system at equilibrium, the equilibrium is shifted in a way that compensates for the effects of that stress.

In this experiment, cobalt(II) chloride, CoCl₂, is dissolved in water to form a pink cobalt(II) chloride–water complex ion, CoCl(H₂O)₅⁺. When alcohol is added to the aqueous cobalt chloride solution, three of the water molecules are removed, forming a blue CoCl₂(H₂O)₂ species, according to the following reaction. This process is endothermic so heat is considered a reactant.

What effect does an increase in temperature have on this system at equilibrium?

Here, the increase in temperature is the stress on the system. The system counteracts the effect of the temperature change by absorbing energy and shifting the equilibrium to the right. This change causes a decrease in the concentration of the pink CoCl(H₂O)₅⁺ ion, and an increase in the concentration of the blue CoCl₂(H₂O)₂ species, as shown.
What effect does a decrease in temperature have on this system at equilibrium?

Likewise, when the temperature decreases, the system compensates by releasing energy, causing a shift in the equilibrium position to the left. The result of this change is an increase in the concentration of the pink CoCl(H₂O)₅⁺ ion, and a decrease in the concentration of the blue CoCl₂(H₂O)₂ species, as shown.

Absorption and Transmission of Visible Light

The cobalt(II) chloride/alcohol/water solution appears pink in cold water, violet at room temperature, and blue in hot water. Why? When white light is shined through a colored solution, the molecules in the solution absorb some of the wavelengths of light and transmit others. The color(s) of light that are absorbed are the complementary colors of the colors(s) that are transmitted. Those that are transmitted are the colors that the human eye receives.

The color wheel in Figure 1 displays complementary colors—those colors that are across from one another on the color wheel. Red and green are complementary colors; therefore, a green solution absorbs red light (the complement of green) and transmits green light to the eye. Hence, the solution appears green to the eye.

Absorption Spectra of a Solution of CoCl₂ in Isopropyl Alcohol

The absorption spectra of the cobalt(II) chloride/alcohol/water solution at three temperatures are displayed in Figures 2 and 3. The absorption spectra show the wavelengths of light that are absorbed by the solution at various temperatures.
Figure 2 (cold) shows the absorption spectrum of the cobalt(II) chloride/alcohol/water solution at a low temperature (7 °C). Notice the absorption peak in the blue–green wavelength range. Looking at the color wheel, the complementary color of blue–green is red. Thus, pink (red) wavelengths of light are transmitted to the eye and the solution appears pink in cold water.

Figure 3 (hot) shows the absorption spectrum of the cobalt(II) chloride/alcohol/water solution at a high temperature (40 °C). Notice the absorption peak in the orange–red wavelength range. Looking at the color wheel, the complementary color of orange–red is blue. Thus, blue wavelengths of light are transmitted to the eye and the solution appears blue in hot water.

Figure 4 (room temperature) shows the absorption spectrum of the cobalt(II) chloride/alcohol/water solution at room temperature (24 °C). Notice that there are two absorption peaks, one in the blue–green range and one in the orange–red range. Looking at the color wheel, both red and blue wavelengths of light are transmitted to the eye. Thus, the solution appears violet at room temperature.

Absorption versus Transmission Spectra

For analytical chemists, absorption spectra are important due to the linear Beer’s Law relationship between absorbance and concentration. However, the color of an object or solution is determined by the color of the light it transmits or is reflected from it. For this reason, the colors of the CoCl₂ solution will also be explained based on the transmission spectra at three temperatures superimposed on each other (see Figure 5).

There are two areas of interest in the adjacent spectrum of cobalt(II) chloride dissolved in a 2-propanol-water system. The first is the small blue peak at 500 nm (Point A), which is characteristic of most cobalt compounds and apparently does not change significantly in this experiment, as evidenced by all three lines clustered together at 500 nm. The second area of interest is the pink (orange–red) transmission between 600 and 670 nm, which changes as a function of the temperature, as evidenced by Points B, C and D in Figure 5. This change in the pink coloration can be used to illustrate the shift in equilibrium as predicted by Le Chatelier’s Principle.

Look at the following equation:

It can be seen from the spectrum in Figure 5 that the color of the solution is determined by the concentration of the pink CoCl(H₂O)₅⁺ species. At high temperature (40 °C), the equilibrium has shifted to the right lowering the concentration of the pink CoCl(H₂O)₅⁺ species, leaving only the blue CoCl₂(H₂O)₂ species (Point D). When the reaction is placed in ice water (7 °C), the equilibrium shifts to the left, increasing the concentration of the pink CoCl(H₂O)₅⁺ species, resulting in a pink solution (Point B). The blue species is still present, but the solution color is dictated by the much higher concentration of the pink CoCl(H₂O)₅⁺ species. At room temperature (24 °C), the concentration of the pink CoCl(H₂O)₅⁺ species is low enough to give the solution the violet intermediate color (Point C).

References

Special thanks to Walter Rohr, retired chemistry teacher, Eastchester High School, Eastchester, NY, for providing Flinn with this activity and for measuring and recording the absorption and transmission spectra at various temperatures.

Alyea, Herbert A., “Tested Demonstrations in Chemistry,” Journal of Chemical Education: Easton, PA, 1965; p 222.

Shakhashiri, Bassam Z., Chemical Demonstrations: A Handbook for Teachers of Chemistry, Volume 1, The University of Wisconsin Press: Madison, WI, 1983; pp 280–285.