A solution of cobalt(II) chloride in ethanol appears blue due to formation of [CoCl4]2–. When water is added, the equilibrium shifts toward the hydrated [Co(H2O)6]2+ complex, producing a pink solution.

1. Dissolve 5 g of anhydrous cobalt(II) chloride in 250 ml of ethanol in a 500 ml glass flask or bottle.
2. Prepare at least 250 ml of water in a separate bottle (a plastic bottle can be used for dramatic effect).
3. Show the audience the flask containing the blue solution.
4. Add water to the blue solution and observe the color change to deep pink as the equilibrium shifts.

Cobalt chloride equilibrium lab (Le Chatelier's Principle) - BAMChem:

Cobalt equilibrium and thermodynamics - Tommy Technetium:

📄 “Cobalt colors” experiment - MEL Science: https://melscience.com/AU-en/articles/cobalt-colors-experiment/?srsltid=AfmBOopDSnZSBpzZQGb52NU-j5oCeoZGGFxEhD5v4pO5kdXVM_GJGhl6

None

• Wear safety goggles, gloves, and a lab coat.
• Cobalt(II) chloride is toxic, irritant, and a suspected carcinogen; avoid skin contact and inhalation.
• Ethanol is flammable; keep away from flames or sparks.
• Dispose of cobalt-containing solutions as hazardous chemical waste; do not pour down the sink.

• Why is the ethanol solution blue before adding water? (Ethanol lowers solvent polarity, favoring [CoCl4]2– formation.)
• Why does adding water cause the solution to turn pink? (Water increases polarity, favoring the hydrated [Co(H2O)6]2+ complex.)
• How does this demonstration illustrate Le Chatelier’s principle? (The equilibrium shifts to counter changes in solvent composition.)
• How would increasing chloride ion concentration affect the color? (It would favor [CoCl4]2–, making the solution bluer.)