LeChatelier’s Principle – A Simple Equilibrium Demonstration Using CoCl2
LaChatelier’s Principle, the idea of applying a stress to a system at equilibrium and the system
responding to that stress by shifting to reestablish the equilibrium, is a central point of any
introductory chemistry course. It is easy to describe and, in principle, easy to understand.
However, students frequently have difficulty grasping the real meaning of the Principle and its
application as a result of only ‘seeing’ it on a blackboard, rather than experimentally. The
following simple experiment (or demonstration depending on individual laboratory capabilities)
provide a direct way for the students to see the effects of changing concentration and temperature
on an equilibrium by watching the color changes in an aqueous solution of cobalt(II) chloride.
The reaction involved is:
[Co(H2O)6]2+ + nCl - ⇌ [CoCl4]-2 + n H2O
[Co(H2O)6]
+2
+ n Cl-
[CoCl4]
-2
+
n H2O
Cobalt(II) chloride when dissolved in water generates the hexaaquacobalt(II) cation (pink).
Addition of a large enough amount of chloride ion (most easily accomplished by addition of
concentrated HCl) drives the equilibrium to the right, forming the blue tetrachlorocobaltate(II)
ion. The equilibrium is then shifted back to the left two different ways: addition of water (a
product, therefore shifting the equilibrium back to the starting materials), or removal of chloride
(lowering the concentration of a starting material, therefore again shifting the equilibrium to the
left). Finally, the effects of temperature are seen by carefully adjusting the position of the
equilibrium to the “middle” (purple) and then heating and cooling the solution. NOTE: It is very
important that the students recognize that water is not just a solvent in this reaction – it is a
product as well!) It is also important that all beakers are dry before materials are added.
Materials: 0.1 M CoCl2(aq), concentrated aqueous HCl, 0.1 M AgNO3(aq), beakers (2 per …),
graduated cylinders (2 per …), stirring rod, ice and a hot plate.
Procedure: 1) Each student (group…) should measure out 5 ml of the 0.1 M CoCl2 solution and
put it into a small beaker (I like 30 ml beakers for this). Note the pink color. Now, add about 8
ml of concentrated HCl – the color should change to blue, and if the students are paying
attention, they should see the solution pass through an intermediate purple color (which is a nasty
mixture of the aqua, chloro and mixed aqua/chloro species). (An increase in concentration of a
starting material shifts the equilibrium to the products)
2) Take ~ 4 ml of this mixture and place it in a second beaker. Add water dropwise until the
pink color is reestablished ( at about 0.5 ml, the solution is the intermediate purple, and at 1-1.5
ml it should be clearly pink). (The addition of a product forces the equilibrium back toward the
starting materials)
3) Take another ~4 ml of the blue mixture, transfer it to a second beaker. Add 0.1 M aqueous
silver nitrate dropwise to the solution. The first observation is that a precipitate of AgCl forms.
Make sure the students recognize that the precipitate is white and that the color is from what is
still in the solution. Initially, the mixture should still be blue (but opaque). As more silver
nitrate solution is added, the color should change to purple and finally pink (about 1 ml of the
silver nitrate solution should do it). (The removal of a starting material (chloride ion) from the
reaction causes the reaction to shift to the starting materials)
4) Temperature effects. Take the remaining solution and carefully add water dropwise until the
solution is the intermediate purple/lavender color. Divide that solution into two beakers (one to
serve as the reference). Take one of the beakers and place it on the hot plate, observing the color
change; it should become distinctly bluer. The color will also intensify as a result of the different
extinction coefficients for the aqua and chloro species (which you may or may not want to get
into). The color and intensity should be compared to the reference beaker. Then take the hot
solution and cool it in ice. The color should again change, first back to the lavender of the
reference beaker and then to a pinker color. (Equilibria and equilibrium constants are
temperature dependent).
Please note that in this last part, we are cheating a little bit. It is easy to raise the temperature of
a room temperature solution significantly (~25 to as much as 100 C), but harder to lower it by as
much (only ~25 to 0 C) so seeing the effect of cooling the solution is not as obvious. By having
the students heat the solution first, they are also driving off a little of the HCl which helps when
the solution is cooled. I don’t bring this up, but if someone asks I do explain it to them including
the temperature differential problem.
If you are doing this as a demo rather than an experiment (group or individual), you’ll need to
increase volumes so the students can see it. Also, don’t try to ‘project’ it using an overhead
projector. The heat from the projector is enough to cause a bit of chaos for you.
Waste disposal: The small amounts generate very little waste, but it can be dealt with more
easily by adding base. If all the Co and Ag containing solutions are combined, addition of
NaOH until pH ~9 will precipitate the metals. Filtration then provides a very small amount of
solid waste and the filtrate contains only NaCl, NaNO3 and H2O.