TOPIC: Equilibrium


Iron(III) - Thiocyanate Ion Equilibria

An equilibrium is set up which has a colored product. Adding either reactant may be observed to displace the equilibrium toward the product side (darken the color), and subtracting either reactant may be observed to displace the equilibrium away from the product side (lighten the color).


Transparency made from template on last page
3% ammonium thiocyanate solution in a dropper bottle
1% iron(III) chloride solution (freshly prepared) in a dropper bottle
1% mercury(II) chloride solution in a dropper bottle
2% sodium dihydrogen phosphate solution in a dropper bottle
5 50-mL beakers or a Corning 6-well Cell Well™
1 100-mL beaker
1 125-mL squirt bottle
stirring rod
overhead projector


Place the Cell Well™ or beakers on the transparency on the overhead projector. Fill the first container with water to a depth of about 1 cm (ca. 10 mL). Allow a moment for the water to become still. Carefully add five drops of FeCl3 solution to one side of the container and five drops of NH4SCN solution to the other side in order to show that individually neither is colored. Stir the solution to form the Fe(NCS)2+ complex. This will serve as a color standard.

Squirt the prepared Fe(NCS)2+ solution into containers 2-5 to a depth of about 1 cm, using just enough solution so that the colors seen on the screen match that of the color standard. Add five drops of FeCl3 solution to container 2, stir, and observe the brown color darken. Add five drops of NH4SCN solution to container 3, stir, and again observe the color darken. Add one drop of NaH2PO4 solution to container 4, and add ten drops of HgCl2 solution to container 5. Stirring will produce a lightening in the color in these two containers.


The principle of LeChâtelier states that if a chemical system at equilibrium is subjected to a stress, the system will react in a way that tends to relieve the stress. "Stress" means a change in concentration, pressure, temperature, etc. imposed on the system. In this demonstration, concentration stresses are being examined. A more detailed statement of the principle based only on concentration can be given. If a chemical reaction is at equilibrium and the concentration of a reactant is increased, the reaction is no longer at equilibrium. In order to achieve equilibrium once more, the extra reactant must be partly used up by turning it (along with the other reactants) into products. On the other hand, if a product concentration is increased, in order to reach equilibrium some product must be turned back into reactants, so the reaction goes backwards. If a reactant is removed, more must be formed, and a part of the products must turn back into reactants. Finally product removal requires more product formation. To sum up, a stress that involves an increase in a reactant concentration or a decrease in a product concentration will cause a reaction to "move to the right", and a decrease in a reactant concentration or an increase in a product concentration will cause it to "move to the left".

The first part of the demonstration established that a reaction did occur between FeCl3 and NH4SCN. Since FeCl3, NH4SCN, and NH4Cl are colorless, the color that appears must result from an interaction between Fe3+ and SCN-. The interaction is in fact the equilibrium

Fe3+(aq) + SCN-(aq) Fe(NCS)2+(aq),

the product being brown.

Adding either Fe3+ or SCN- to the equilibrium caused the solution color to darken because, in order to relieve the stress of added reactant concentration, more Fe(NCS)2+ was formed to use up reactants and re-establish the equilibrium. This darkening establishes another important point. Both free Fe3+ and free SCN- exist in quantity in the original equilibrium mixture. Thus it can be argued that adding H2PO4- or Hg2+ removed free Fe3+ or SCN- from solution rather than destroying the complex directly.

The addition of H2PO4_ removed some free Fe3+ from the equilibrium mixture by the reaction

Fe3+ + H2PO4- Fe(HPO4)2+ + H+

In order to restore Fe3+, some of the Fe(NCS)2+ dissociated and the color of the solution faded.

The addition of Hg2+ removed some free SCN- as the complex Hg(SCN)+, and once again the equilibrium was restored by dissociation of some of the Fe(NCS)2+, resulting disappearance of some of the solution color.

Addition or removal of the reaction product, Fe(NCS)2+, is not useful in this demonstration. Because we are monitoring the position of equilibrium by observing its color, it would not be possible to observe how the reaction responded to a change in its concentration.


Soluble mercuric salts and their solutions are toxic (oral LDLO HgCl2: 25 mg/kg). Toxicity of thiocyanates have not been well studied, but they appear the have a moderate toxicity (oral LDLO: ca. 80 mg).


Doris Kolb, "Introduction to Overhead Projector Demonstrations", Journal of Chemical Education, 1987, 64, 348.

Lee R. Summerlin and James L. Ealy, Jr., "Chemical Demonstrations: A Sourcebook for Teachers", The American Chemical Society, Washington, D. C., 1955.