Revisiting an equilibrium demo from a classic resource

As part of the grade 12 equilibrium unit, I have my students perform a very traditional lab activity demonstrating Le Châtelier's principle. They rotate through a series of stations where they use a semi-microlab approach to investigate a different equilibrium system at each one. For years I have included the chromate/dichromate system1 and the cobalt chloride system2 as regular components of this activity. These reactions are popular among teachers as they produce obvious colour changes, which clearly demonstrate the effects of temperature change, pH and precipitate formation on equilibrium. The problem is that cobalt(II) ions and chromate ions pose fairly severe health hazards and are both classified as possible carcinogens.3 A recent safety review at my school board prompted me to research possible alternatives.

I found an ideal replacement in the classic resource by Summerlin and Ealy (reference 2, pages 83-84). It involves the equilibrium between blue copper(II) ions and the green copper tetrabromide complex:4

heat + Cu2+(aq) + 4Br-(aq) ⇄ CuBr42-(aq)

The student activity outlined below is a modification of the original lecture demonstration.

Student directions

  1. Place 2 dropperfuls (about 10 mL) of 0.2 M CuSO4 solution into each of three 50 mL beakers. These beakers are numbered 1, 2 and 3.
  2. Dissolve small, scoopula-tip amounts of solid KBr in beakers #2 and #3. Repeat until a colour change is observed. Note: This should not take more than two or three additions.
  3. Add 1 M NaOH drop-wise to beaker #3 until a colour change is observed.
  4. Add 1 M H2SO4 drop-wise to beaker #3 until a colour change is observed.
  5. Carefully (don't spill!) place beaker #3 in an ice bath and observe.
  6. Gently heat beaker #3 on a hot plate for a short time and observe.
  7. Add one dropperful (5 mL) of distilled water to beaker #1 and #2 and observe carefully. Repeat to a maximum volume of 30 mL if necessary.

Teaching notes

  1. I have kept quantities and concentrations low to minimize waste while still ensuring that the colour changes are easily observable against a white background.
  2. The formation of the copper tetrabromide complex turns the solution green. Potassium bromide is hygroscopic, so it is important to remind students to put the cap back on the container when not in use. Beaker #1 is used for comparison.
  3. The addition of base forms a copper(II) hydroxide precipitate shifting the equilibrium left resulting in a colour change from green to blue. This step is a simplification of the original demonstration, which calls for the addition of sodium sulfate to increase the pH and form the precipitate.
  4. Addition of acid dissolves the precipitate resulting in a clear green solution again.
  5. The equilibrium is endothermic as written. Cooling results in a colour change from green to blue.
  6. Only slight heating is required for the solution to shift back to green.
  7. Because the equilibrium expression is fifth-order in reactants compared to just first-order in products, dilution of the two solutions results in an easily observable shift in beaker #2 from green to blue.5

Although copper(II) ions are still hazardous and should be handled and disposed of carefully, the solutions in this activity are more appropriate for use at the high school level than solutions of  cobalt or chromate ions. The activity accomplishes everything the other systems do together, but more safely; and with greater clarity and simplicity. Based on a quick internet search, it is apparent that this equilibrium is not commonly used by teachers, which is surprising. It has certainly become a regular component of my toolkit.


  1. This activity is ubiquitous and is readily found using a simple internet search.
  2. Lee R. Summerlin and James L. Ealy, Jr., Chemical Demonstrations: A Sourcebook for Teachers: Volume 1, 2nd Edition, American Chemical Society, 1988, pages 79-82.
  3. This information is readily available through an internet MSDS search.
  4. Summerlin and Ealy state that this complex is green; however my observations suggest that the copper tetrabromide complex is actually yellow. The green colour seems to be a result of the two complexes being present in comparable concentrations.
  5. I credit serendipity for this one.