Bill and Anthony

Cobalt II chloride is a salt that is well known to change color, as it forms various complexes with water. With this property, it is used in products such as invisible ink, where a dilute, light pink solution of it in water is applied to the paper and appears invisible to the naked eye, but is visible as a sky blue when heat is applied (1). It is also used in humidity test strips, where the strip is blue in a dry environment but turns pink when there is sufficient moisture in the air (2). In these two cases, there is an equilibrium between the anhydrous version of Cobalt II chloride and the hexahydrate version. The blue/anhydrous version is favored by heat, while the pink/hexahydrate version is favored by moisture. It is also known that, at a static temperature, the pink form is less favored by the presence of less polar solvents, such as isopropanol or butanol. The goal of this experiment was to determine the relationship between temperature and the equilibrium concentrations of both tetrahedral and octahedral forms of cobalt II chloride. If it is true that the formation of the blue species of cobalt II chloride requires an activation energy provided by heating, then we hypothesize that the color of the entire solution will change from pink to purple almost instantaneously, once a certain temperature threshold is reached.

Cobalt II chloride is an inorganic salt that has a pale blue color in its anhydrous form. As a hydrate, which it becomes when exposed to moisture, it changes to a pink color (a hydrate is a molecule that contains water as one of its elements). This process is reversible: the pink color can be changed back to blue by heating the salt (4). When the hexahydrate form is dissolved in water, the six water molecules for a coordination complex with the central cobalt atom. The surrounding water molecules are called ligands, and the structure that this complex forms is known as octahedral (5). However, when this octahedral structure is disturbed, for example by adding concentrated hydrochloric acid, some water ligands are forced out by incoming chloride atoms, the complex changes to a tetrahedral structure. Along with this new arrangement of ligands, the color of the solution changes as well, from pink to purple (6). This transition from octahedral to tetrahedral can also be aided with the presence of another solvent or the addition of heat.

Figure 1: Temperature of cobalt II chloride solution versus average red value of sample area. The vertical axis range begins at 100, with the RGB scale spanning from 0 to 255. 5 test tubes were heated from 0 to 78 degrees, and color was sampled from 7 equally spaced time periods from each tube; therefore, each dot represents an average over 5 trials. The uncertainty bars represent average absolute deviation.

Data Interpretation:

Contrary to what was predicted, the color of the solution changed gradually from pink to purple as the temperature increased from 0 to 78 degrees, instead of instantaneously at a specific temperature. The average red value for pixels representing the top surface of the solution in the test tube started at 171, and at first decreased very slightly until about 15 degrees, when the average red value began linearly decreasing until a minimum of 119. The red value best represents the color change of the cobalt II chloride solution because it is the complex’s ability to absorb red light that increases when it changes from an octahedral to tetrahedral structure and replaces the water ligands with chloride ions.

Therefore, the linearly decreasing red value between 15 and 78 degrees represents a linearly increasing amount of cobalt complexes that have a tetrahedral structure. This shift in structure is a result of some of the water molecules leaving the coordination sphere of the cobalt complex. Temperature forced an effect on the other two components of Gibbs free energy because the equilibrium stays in effect, although we have pushed it to favor both the reactant and product sides. The increase in temperature causes the solution to give off heat, increasing the enthalpy. The increase in temperature also causes more movement in the molecules, causing more disordered molecules and rising the entropy. This ambiguous change in Gibbs free energy leads us to hypothesize that rather than the spontaneity of the reaction changing when heat is added, the extra heat is a reactant in the reaction to form the tetrahedral cobalt complex. Our data supports this idea because if the reaction had truly become spontaneous through a positive Gibbs free energy, the shift in color would have happened much more quickly, given the constant stirring of the solution. The slow shift in color suggests that, as more heat becomes available to the molecules in solution, each one individually will use the heat to incorporate chlorine atoms into the coordination sphere while removing water ligands. As fast as heat is introduced from the hot water bath, the molecules will consume it in this endothermic reaction. Furthermore, we speculate that the delayed start to the color change until 15 degrees represents a minimum temperature needed for this reaction to occur. The small change in color before that point may correspond to the molecules that randomly meet the heat threshold, a result of temperature being only an average representation of kinetic energy, rather than an exact representation of every molecule.

Sources of Uncertainty/Suggestions for Improvement:

Our sources of uncertainties are the concentration of the cobalt (II) chloride solution, time in which the solutions were used, the representation of temperature, and the position of the camera. The concentration of cobalt chloride was an uncertainty because with just a small increase in cobalt (II) chloride, the solution would have had to reach greater temperatures to be able to reach the same results. It is suggested that one large solution of the cobalt chloride and isopropyl alcohol mixture and use that one mixture throughout the experiment. This would control the cobalt chloride concentration by making them all the same. Secondly, the age of the solutions was a source of uncertainty because the isopropyl alcohol could have evaporated, changing the concentration of the solution. It is suggested that one large solution of the cobalt chloride and isopropyl alcohol mixture and use that one mixture throughout the experiment. This would control the cobalt chloride concentration by making them all the same. Next, the temperature for the graph was not taken during each trial but after all the trials were taken to see the increase caused by the hot plate. This does not lead to a precise representation of the temperature of the beaker versus the color. It is suggested to film the temperature increase for each test; this will help create an accurate representation of the Temperature versus the RGB. Finally, the current position of the camera makes the stirring magnet visible, adding a brighter, whiter spot to the video that could affect the software into not portraying the proper RGB for the solution. It is suggested that the camera should position at an angle in which the stirring magnet is not visible, so the RGB can be seen more correctly.

Further Study:

As a further study, the effect of repeating the experiment by cycling the complex between its octahedral and tetrahedral forms by placing it back into the ice-salt bath after boiling can be explored. Progressively, the procedure will favor the tetrahedral side of the equilibrium. Because each time that the test tube is heated, it loses a bit of isopropyl alcohol, increasing the concentration of cobalt chloride. In preliminary testing, it was found that higher concentrations of cobalt chloride, the more difficult it is to push the equilibrium toward the tetrahedral structure. To summarize, the repetition of this experiment with the same solution leads to the equilibrium to favor the reactants, keeping the solution looking pink-ish. Another further study that could be explored is how endothermic the reaction is. A possible procedure would have two beakers with hot water of the same temperature in them, and to one beaker a test tube one with the cobalt chloride and isopropyl alcohol solution could be added, and to the other a test tube with only water and isopropyl alcohol. After a minute or so, check the temperature of the water in the beakers, to measure a difference in heat loss from the endothermic reaction.

1. American Chemistry Council. Cobalt Chloride: Colorful Moisture Detector, Aug. 2006, chlorine.americanchemistry.com/Science-Center/Chlorine-Compound-of-the-Month-Library/Cobalt-Chloride-Colorful-Moisture-Detector/.

2. Cotton, Simon. “COBALT CHLORIDE.” COBALT CHLORIDE - Molecule of the Month June 2016, University of Birmingham, June 2016, www.chm.bris.ac.uk/motm/cobalt-chloride/cobalt-chlorideh.htm.

3. Flowers, Paul, et al. Chemistry 2e. OpenStax, 2019, OpenStax, openstax.org/books/chemistry-2e/pages/16-4-free-energy.

4. Kauffman, George B, and Jack Halpern. “Structure and Bonding of Coordination Compounds.” Encyclopædia Britannica, Encyclopædia Britannica, Inc., 3 Aug. 2018, www.britannica.com/science/coordination-compound/Structure-and-bonding-of-coordination-compounds.

5. Macrakis, Kristie. “The Hidden Past of Invisible Ink.” American Scientist, vol. 102, no. 3, 2014, pp. 198–205., doi:10.1511/2014.108.198.

6. “Metal Ion Precipitates and Complexes.” Chemical Demonstrations: a Handbook for Teachers in Chemistry, by Bassam Z. Shakhashiri, vol. 1, University of Wisconsin Press, 1986, pp. 280–285.

1. https://pharmadesiccants.wordpress.com/2014/12/16/what-all-you-need-to-know-about-humidity-indicating-cards/

2. http://www.chem.uiuc.edu/webFunChem/cobalt/CoReview4.htm

3. http://www.chm.bris.ac.uk/motm/cobalt-chloride/cobalt-chlorideh.htm