Equilibrium and Le Chatelier's Principle

Le Chatelier's Principle states that: If an equilibrium system is subjected to a stress, I the system will react to remove the stress. To remove a stress, a system can only do one of two things: form more products using up reactants, or reverse the reaction and form more reactants, using up products. In this experiment you will form several equilibrium systems. Then, by putting different stresses on the systems, you will observe how equilibrium systems react to a stress.

Before you carry out each section, predict which way you think the equilibrium will shift. Then, carry out the reaction to verify your prediction.

Chemicals

Sodium chloride, NaCI(s) Hydrochloric acid, HCI, 12 M

Potassium thiocyanate, KSCN, 0.002 M Hydrochloric acid, HCI, 0.1 M

Bromothymol blue indicator solution Sodium hydroxide, NaOH, 0.1 M

Potassium thiocyanate, KSCN(s) Iron(M) nitrate, Fe(N03)3, 0.2 M

Silver nitrate, AgN03. 0.1 M Ethanol, C2H5OH(l)

Cobalt(II) chloride, CoCl2*6H2O(s) Disodium hydrogen phosphate, Na2HP04(S)

Equipment

Test tubes, 13 x 150-mm Test tube rack

Beaker, 100 mL Graduated cylinders

Stirring rod Funnel, filter paper, and holder for funnel

Procedure

1. Equilibrium in a Saturated Solution

You will investigate the equilibrium in saturated sodium chloride solution:

NaCl(s) ó Na+(aq) + Cl-(aq)

Pour some solid NaCl into a 13 x 100-mm test tube and fill the tube 3/4 full of distilled water. Cork and shake to form a saturated solution. If all the NaCl dissolves, pour some additional NaCl in the, tube and shake until a saturated solution with some excess solid is obtained.

Filter the solution into a second test tube. To this saturated solution of NaCl, add some Cl- ions in the form of concentrated HCl Record and explain the results.

2. An Acid-Base Indicator Equilibrium

Acid-base indicators are large organic molecules that can gain and lose hydrogen ions to form substances that have different colors. The reaction of the indicator bromothymol blue can be illustrated as follows:

HIn(aq) ó H+(aq ) + In-(aq)

yellow blue

In this reaction HIn is the neutral indicator molecule, and In- is the indicator ion after the molecule has lost a hydrogen ion. Equilibrium reactions can easily be forced to go in either direction.

Reactions like this are said to be reversible.

Fill a small test tube about half-full of distilled water. Add several drops of bromothymol blue indicator solution. Add 5 drops of 0.1 M HCl and stir This will increase the amount of H+ in solution. Note the color of the indicator.

Next add 0.1 M NaOH drop by drop with stirring until no further color change occurs. Adding OHions causes the H+ ion concentration to decrease as the ions combine to form water molecules. Again, note the color. See if you can add the right amount of acid to this test tube to cause the solution to be green in color after it is stirred (half of the indicator is blue and half is yellow).

3. A Complex Ion Equilibrium

An equilibrium system can be formed in solution with the following ions:

Fe3+(aq) + SCN-(aq) ó FeSCN2+(aq)

colorless colorless red-brown

The iron ion (Fe3l) and the thiocyanate ion (SCN-) are both colorless; however, the ion that forms from their combination, the FeSCN2+ ion, is colored a dark red-brown. It is the color of this ion that will indicate how the equilibrium system is being affected.

Pour about 25 mL of 0.0020 M KSCN solution (a source of SCN- ion) into a beaker. Add 25 mL of distilled water and 5 drops of 0.20 M Fe(N03)3 solution. Swirl the solution and note the following: the color of the KSCN solution, the color of the Fe(N03)3 solution, and the color of the resulting complex ion.

You will stress the equilibrium system that has resulted in several ways. Pour equal amounts of the solution from the beaker into four test tubes. The solution in the first test tube will be the reference solution.

To the second test tube add 2-3 crystals of solid KSCN. Describe the results.

To the third test tube add 6 drops of Fe(N03)3 solution. Stir and describe the results.

I

To the fourth test tube add small crystals of NaHP04, a few at a time. Stir and note the results.

Phosphate ions PO43-, have the ability to form complex ions with Fe3+, which has the same effect as removing Fe3+ from solution.

4. An Equilibrium with Cobalt Complex Ions

In this section we will investigate the equilibrium between two different complex ions of cobalt. The reaction is endothermic:

Co(H2O)62+(aq) + 4 Cl-(aq)ó CoCl42-(aq) + 6 H20(l) D H = +50 kJ/mol

pink blue

Safety Alert

Ethanol is flammable. Turn off all flames.

Silver Nitrate causes stains on skin and clothing

Wash off spills with soap and water immediately.

 

Measure about 10 mL of ethanol into a beaker.

Examine solid cobalt(M chloride, noting both its color and the formula of the compound. Dissolve a small amount of cobalt(II) chloride (about half the size of a pea) in the beaker of ethanol. The solution should be purple. If it is pink, add a little concentrated HCI until it is purple.

Put about 2 mL of the alcoholic cobalt solution into each of three small test tubes. To one of the test tubes, add 3 drops of distilled water, one drop at a time with stirring, noting what happens with each drop. Add 3 drops of distilled water to each of the other two test tubes. Make a note of the effect of this stress on the system.

The first test tube is the control. To the second test tube, add 5 drops of concentrated HCI,

12 M, one drop at a time with stirring. Note the results.

To the third test tube add a few crystals of solid sodium chloride. Stir and note the results.

Put the remainder of alcoholic cobalt solution from the beaker into a fourth test tube. Add 10 drops of 0.1 M silver nitrate solution, one drop at a time. Silver and chloride ions combine to form a precipitate of AgCI. Note the color of the solution as the chloride ions precipitate. You may wish to let the precipitate settle to observe the solution color more easily, or you may centrifuge the test tube.

Obtain a sealed Beral pipet containing some of the alcoholic cobalt chloride-water system. Note its color. Immerse the large end of the pipet in some hot water (about 60'Q and see if there is a color change.

Lastly, chill the Beral pipet in an ice bath to see if the color change in the previous step is reversible. Explain the effect of the temperature change on the equilibrium in terms of the fact that the value of D H for the reaction is +50 kJ/mol.

 

Discussion

In your laboratory report describe the results of each part of the experiment and interpret them using LeChatelier's Principle.