At this point in your chemistry career, you should be able to predict the products of chemical reactions, the states of the products, and whether the reaction will occur spontaneously at any given set of conditions.  You should even be able to determine the rate at which the reactants are consumed and predict the amount of time it would take to produce a given amount of product.  While this is extremely useful information, it only applies to a limited set of reactions, namely those that occur in one direction only:

Here the reactants A and B collide with sufficient energy and the proper geometry to form the products C and D.  What about a reaction in which C and D now become reactants in the opposite direction and form the products A and B?

Initially, when A and B were mixed, the reaction proceeds in the forward direction to produce C and D.  However, as time progresses, the concentration of C and D increases causing an increase in the rate of the reverse reaction.  Concurrent with this increased rate of the reverse reaction is a reduction of the forward rate due to the decrease in the concentration of A and B.  At some point, the rate of the forward and reverse reactions will become the same and we will reach a state of dynamic equilibrium:

This state of dynamic equilibrium does not mean that the forward and reverse reactions have stopped.  Molecules of A and B are still reacting to form C and D and molecules of C and D are reacting to form A and B.  However, since the rate of the forward and reverse reactions is the same, it will appear that nothing is happening.  As such, all quantifiable physical and chemical properties such has pH, color, and concentration will remain constant.

In 1888, Henri Louis Le Chatelier stated that if a stress were applied to a system at equilibrium, the system would respond in such a way to reduce that stress and establish a new equilibrium.  Le Chatelier's principle, as it is now known, is a very powerful tool for understanding equilibrium reactions.

Purpose:

The purpose of this experiment is to familiarize you with the concept of an equilibrium reaction and the use of Le Chatelier's principle in predicting what effect a given stress will have on a system at equilibrium.  We will study five chemical equilibrium reactions and observe the effect of applying various stresses to them:

These reactions have been chosen so that the application of stress to the system will result in visually discernable changes.   The particular stresses we will be studying, will include changing the temperature of the system, changing the pH, and changing the concentration of the reactants and products.

Procedure:

The HIn Equilibrium:

1. Add 10 mL of distilled water and 4 drops of 6 M HCl to a beaker and swirl.  Mark this beaker 'A' (acid).
2. Add 10 mL of distilled water and 4 drops of 6 M NaOH to another beaker and swirl.  Mark this beaker 'B' (base).
3. Add approximately 1 mL of distilled water to a clean test tube.  Now add 4 drops of the methyl orange (C₁₄H₁₄N₃NaO₃S) acid/base indicator solution and 2 drops of the dilute acid solution 'A' to the test tube and shake gently.  Record the color.
4. Add drops of the dilute base solution 'B' to the test tube until the color changes.  Shake the solution between additions of the base.  Record the color.
5. Add drops of the dilute acid solution 'A' to the test tube until the color changes.  Shake the solution between additions of the acid.   Record the color.

1. Add 20 mL of distilled water, 20 drops of 0.1 M Fe(NO₃)₃ and 20 drops of 0.1 M KSCN to a 100-mL beaker.  Stir the solution thoroughly; the color of this solution is due to the Fe(SCN)²⁺ ion.
2. Add 3 mL of this solution to three separate, clean and dry test tubes.
3. Add 20 drops of distilled water to one of the test tubes and gently shake.  This is the reference for rest of the tests in this section.   Record the color.  For the best results, stand the test tube up on a piece of white paper and look straight down through the solution.
4. Add 20 drops of 0.1 M Fe(NO₃)₃ to one of the test tubes and gently shake.  Record the color.
5. Add 20 drops of 0.1 M KSCN to another test tube and gently shake.  Record the color.
6. Compare the colors in the test tubes from 4) and 5) with the test tube in 3).  The intensity of the color in each test tube will indicate the relative concentration of the Fe(SCN)²⁺ ion in that test tube.

1. Add 10 drops of 0.1 M Ni(NO₃)₂ to a clean test tube. Record the color.
2. Now add 6 M NH₃ drop wise to this test tube until the color changes and intensifies.  Record the color.
3. Now add 6 M HCL drop wise to this test tube until the color changes again.  Record the color.

1. Add 10 drops of 0.1 M Co(NO₃)₂ to a clean test tube.  Record the color.
2. Add 10 drops of concentrated HCl (NOT 6 M HCl!).  Shake the test tube gently, and record the color.
3. Add 10 drops of distilled water, shake gently, and record the color.
4. Carefully waft the test tube in a low temperature (no inner blue cone) Bunsen burner flame.  Record all of the color changes that occur.
5. Cool the test tube in an ice water bath until the color changes.  Record the color.

1. Add 1 mL of 6 M NaOH to a clean medium (not your largest, not your smallest) test tube.
2. Now add 1 mL of 1 M Ca(NO₃)₂ to this test tube and stir thoroughly.  A white precipitate of Ca(OH)₂ should form.
3. Isolate the precipitate using the centrifuge.  Remember to counter balance the centrifuge with a test tube of equal size and volume of distilled water.
4. Pour the supernatant down the drain and use your scoopula to transfere a small amount of the precipitate into 50 mL beaker. 
   
5. Add 10 mL of distilled water to the beaker and stir to suspend the solid.
6. Add 5 mL of 6 M HCl to the beaker and stir thoroughly.  Record your results.
7. Now add 10 mL of 6 M NaOH to the beaker and stir thoroughly.  Record your results.

Results/Calculations:

• Using the balanced equations for each of the equilibrium reactions you have run, and your observations, explain what 'stresses' were applied to each reaction.  How did the reactions respond to these 'stresses'?  Were they what you expected?  Why?
• Were any of the reactions endothermic or exothermic?   How would you use this information to alter the equilibrium?

(Updated 10/31/12 by C.R. Snelling)