Le Chatelier’s Principle of Chemical System

Introduction:

According to Le Chatelier’s Principle, when a chemical system at equilibrium undergoes a change in concentration, temperature, volume, or partial pressure, the equilibrium will shift to counteract the change and establish a new equilibrium (Atkins, 1993). Chemical reactions in equilibrium are reversible and any alterations in these factors will cause the equilibrium to adjust and form a new balance.

The purpose of this experiment was to demonstrate the principle of equilibrium and how changing the concentration of specific ions or temperature can shift it. It was hypothesized that observable changes in equilibria would be noticeable if this principle holds true.

To perform this experiment, an electric hot plate, ice cubes, polyethylene transfer pipets, test tubes, a 100 mL beaker, de-ionized water (H2O), and various chemicals were needed. These chemicals included:

• 1 M K2CrO4
• 3 M H2SO4
• 6 M NaOH
• 0.5% methyl orange indicator
• 0.1% phenolphthalein indicator
• 6 M HCl
• 0.1 M CH3COOH (acetic acid)
• 1 M NaCH3COO (sodium acetate)
• <0.1M NH3/li>
• <1M NH4Cl/li>
• <0.1M Fe(NO3)3/li>
• <0.1M KSCN (potassium thiocyanate)/i>
• <015M anhydrous CoCl2 (in methanol)/i>
• <12M HCl/i>
• <54M NaCl/i><>01m BaCl2/LII>> ><>01m CaCl2/LII>> ><>05m H2C2O4 (oxalic acid)/LII>> ><>05m K2C2O4 (potassium oxalate)/LII>> ><>6m NH3/LII>

It is important to note that concentrated chemicals such as 12 M HCl or 6 M HCl are very corrosive, and goggles and gloves must be worn at all times. The mixtures of chemicals must also be disposed of accordingly.

Heating methanol can produce toxic and flammable vapors, making it necessary to perform this procedure under a hood.

Methods:

• The first step in manipulating equilibrium involved using 1 M K2CrO4, 3 M H2SO4, and 6 M NaOH.
• To a test tube, add 3 mL of 1 M K2CrO4 followed by several drops of 3 M H2SO4. Swirl the test tube to mix the two chemicals and record any observed changes.
• Add several drops of 6 M NaOH to the same test tube and stir until a change occurs. Record this change as well.

Several drops of 3 M H₂SO₄ were added to the mixture and stirred. The change was recorded. In the second equilibrium manipulation, a drop of methyl orange indicator was added to 3 mL of water in a test tube. Two drops of 6 M HCl were then added, and any changes were recorded. Following the addition of two drops of HCl, four drops of 6 M NaOH were added to the mixture, and changes in color were recorded. After this procedure, the experiment with these chemicals was repeated using phenolphthalein indicator instead of methyl orange.

In the third equilibrium manipulation, two test tubes were prepared with 3 mL samples of 0.1 M acetic acid and a drop of methyl orange indicator. To one of the samples, 1 M sodium acetate was added gradually while mixing. The changes in color of both test tubes were compared and recorded.

For the fourth equilibrium manipulation, two 3 mL samples of 0.1 M NH3 were prepared with a drop of phenolphthalein indicator. The initial color of these samples was noted and recorded. Then, to one sample, 1 M NH4Cl was added gradually while mixing.

For one sample, 3 mL of 6 M HCl was added drop by drop with mixing. The same was done for the other sample. Any changes in color and odor were recorded.

In the fifth equilibrium manipulation, a 100 mL beaker contained a mixture of 3 mL of 0.1 M Fe(NO₃)₃ and 3 mL of 0.1 M KSCN. Changes in color were recorded after mixing the two chemicals together. The deep red color was reduced by adding water (50-60mL) until it reached an acceptable intensity level, then a test tube containing only 5mL of the solution had an additional twenty-five drops of .01M Fe(NO₃)₃. Any changes in color were noted.

A second sample consisting of diluted solution (5mL) had one milliliter added to it from .01M KSCN and any resulting changes in color were recorded.

A third sample consisting also of diluted solution (5mL) had five to six drops added from a pipet containing methanol-based .15M CoCl2 which turned the blue liquid pinkish-red; this change was also noted.

The fourth and final control sample consisted only of diluted solution (5mL).

In the sixth equilibrium manipulation, three milliliters (.15M CoCl2) dissolved in methanol was placed into a test tube followed by water using a pipet until it changed to pink complex; this mixture was then divided equally into two portions each placed into separate test tubes.

One sample was tested by adding 12 M HCl drop by drop until changes were observed and recorded. The other sample of the pink complex was heated to 65-70 degrees Celsius using a beaker of hot water on an electric hot plate, and any color changes were noted. After heating, the complex was cooled in an ice bath and the restoration of its original pink color was observed and recorded. In the seventh equilibrium manipulation, a test tube containing 4 mL of 5.4 M NaCl had 2 mL of 12 M HCl added to it.

Any changes in color resulting from this mixture were noted and recorded. Next, 5 to 6 drops of 1 M K2CrO4 were added to 3 mL of 0.1 M BaCl2 in another test tube, and any resulting changes were recorded. Then, 10 to 12 drops of 6 M HCl were added to this mixture, and any changes were again recorded.

For the eighth and final equilibrium manipulation, a solution was created by mixing together 4 mL of 0.1 M CaCl2 with 4 mL of water. This solution was then divided equally into three separate test tubes.

To the first test tube, we added six or seven drops of a solution containing .5M H2C2O4; in the second test tube we added six or seven drops of a solution containing .5M K2C2O4. The results from both test tubes were then compared and recorded.

First, 10 drops of 6 M HCl were added to test tube 1 and the results were recorded. Next, a slight excess of 6 M NH3 was added to test tube 1 and the results were also recorded. To confirm whether the precipitate from test tube 1 was indeed Ca(OH)2, a few drops of 6 M NH3 were added to test tube 3.

The results showed that when H2SO4 was added to K2CrO4 in the first equilibrium manipulation, the solution changed from light yellow to orange. Then, when NaOH was added, the solution returned back to its original light yellow color. When H2SO4 was once again introduced into the mixture, it turned orange again.

The overall reaction of this equilibrium is:

2 CrO₄^2- + 2 H₃O⁺ -> Cr₂O₇^2-+ 3 H₂O

In the second equilibrium manipulation, when 6 M HCl was added to water with an indicator, the color changed from yellowish to pink. Then, when four drops of NaOH were added, the color changed back to yellow. The reaction for this equilibrium was:

H(indicator) +H₂O – H⁺+ indicator –

In the third equilibrium manipulation, when a drop of methyl orange was added to acetic acid, the color changed from clear to pinkish-purple. When sodium acetate was then added to one of the samples, the color changed from reddish-pink to orange.

The overall reaction of this equilibrium is: CH3COOH + H2O → H3O+ + CH3OO-

In the fourth equilibrium manipulation, phenolphthalein indicator was added to the NH3 samples resulting in a purplish-pink color and pungent odor. When NH4Cl was added to one of the samples, the color changed to a clear solution with no smell. The other sample turned clear with no odor when 6 M HCl was added.

In the fifth equilibrium manipulation, adding Fe(NO3)3 to KSCN caused the solution to change from a clear color to a deep brownish-red.

When more Fe(NO₃)₃ was added to the previously diluted solution, there was no change in color and the solution remained red. Adding KSCN to another diluted sample resulted in a darker blackish-red color. NaOH caused a third sample of diluted solution to turn orange. The fourth sample served as a control, with no changes observed in the red color. The overall reaction for this equilibrium is: Fe³⁺ + SCN^– – Fe(SCN)₂⁺. In the sixth equilibrium manipulation, adding 12 M HCl to the pink CoCl₂ complex caused it to change from pink to purple.

When the second sample of the pink complex was heated, its color changed from pink to purple. The color changed back to purple when the complex was cooled in an ice bath. The reaction for this equilibrium is as follows: CoCl₄^– + 6 H₂O – 4 Cl^– + Co(H₂O)₆²⁺ + energy.

In the seventh equilibrium manipulation, adding 12 M HCl to sodium chloride resulted in a milky white precipitate at the bottom and clear fluid on top. Adding K₂CrO₄ to BaCl₂, resulted in a yellow precipitate. Adding 12 M HCl to this solution dissolved the precipitate into an orange solution.

The reactions for this equilibrium were:

• 4NaCl + H2O → 4Na+ + 4Cl- + H2O
• BaCl2 + H2O + K2CrO4 → BaCrO4 + H2O + 2K- + 2Cl-

In the eighth and final equilibrium manipulation, when CaCl₂ was diluted with water, the solution became clear. When H₂C₂O₄ was added to one of the three samples of diluted solution, a white precipitate formed. The same result occurred when K₂C₂O₄ was added. When 6 M HCl was added to tube 1, the precipitate dissolved and the solution became clear again. Finally, when an excess of NH3 was added to this tube, a white precipitate formed.

Finally, when NH3 was added to the third test tube, a very light precipitate also formed. The reactions for this equilibrium manipulation were:

• Ca2+ + C2O42- – CaC2O4 (s)
• K2C2O4 – 2 K+ + C2O42-
• H2C2O4 + H2O – H3O+ + HC2O4-
• H3O+ + C2O42- NH3 + H3O+ – NH4+ + H2O

Discussion:

The first manipulation was the manipulation of chromate ion-dichromate ion equilibrium. The higher concentration of H₂SO₄ pushed the color to change to orange. When NaOH was added, the concentration was pushed back to the initial equilibrium. Then when H₂SO₄ was added again, the concentration was pushed back to the orange color.

These changes occurred due to the addition of hydroxide ions, which affected the equilibrium. The second equilibrium that was manipulated was the weak acid-base indicator equilibrium. The color changes of the indicators in these reactions were caused by adding HCl, which made the equilibrium more acidic. When NaOH is added, the color changes indicate that the indicator has returned to its basic form. The third manipulated equilibrium was a weak acid-weak base equilibrium, where adding acetate ions resulted in a change in color of the indicator.

This shows that the indicator has changed from its acidic form to its base form. It also means that the concentration of hydrogen ions decreased when sodium acetate was added. The fourth equilibrium manipulated was also a weak acid-weak base equilibrium. The indicator was added to a weak base of NH3, which changed its color and odor. When NH4Cl was added, the ion concentrations changed and pushed the indicator back to its acidic form, which is colorless and odorless. The addition of HCl also pushed the equilibrium towards a more acidic solution resulting in similar consequences.

The manipulation of the fifth equilibrium involved a complex ion equilibrium, where cations attract negatively charged ions to form complexes. Specifically, the Fe3+ cation attracted the negatively charged SCN- ion, resulting in the formation of Fe(SCN)2+ and a dark red color. The addition of more Fe(NO3)3 to this solution did not cause any change since positive cations do not attract each other. However, adding more KSCN led to a higher concentration of SCN- ions and pushed the equilibrium towards a darker red solution.

When NaOH was added to the third sample, it caused the formation of the complex Fe(OH)3, which is highly insoluble. This resulted in a change to an orange color. The sixth equilibrium was manipulated by changing the temperature-dependent equilibrium of complex ions. The cobalt complex in methanol is blue and tetrahedral, while the aquo complex is octahedral and has a pink color. The equilibrium between these two forms involves a significant energy change and is temperature dependent.

When HCl was added, the pink aquo complex was pushed to change its complex pack to its tetrahedral form. This change also occurred when the solution was heated, as it required substantial energy. However, when the solution was cooled, the equilibrium shifted back to its original octahedral pink form. Another manipulation involved an equilibrium of saturated solutions. Adding HCl to a NaCl solution increased the concentration of Cl- ions, which in turn pushed Na+ ions within the solution to react more and form a NaCl precipitate.

The same explanation could be applied to saturated barium chromate. When K2CrO4 is added to barium chloride, the Ba2+ ions bind with the CrO42- ions to create a barium chromate complex. When HCl is added, the Cl- ion concentration increases, thus pushing the equilibrium back to the formation of barium chloride. The eighth manipulation applied the law of chemical equilibrium to solubility equilibrium. When H2C2O4 and K2C2O4 are added to CaCl2, the increase in C2O42- concentration causes a precipitate of Ca2C2O4.

When HCl was added, the concentration of Cl- ions increased, thus pushing the equilibrium back to CaCl2. When NH3 was added, it was possible that some OH- ions were generated, thus pushing the equilibrium towards a precipitate of Ca(OH)2.

Ultimately, Le Chatelier’s Principle was apparent in the experiments performed. The principle states that if a chemical system at equilibrium experiences a change in concentration, temperature, volume or partial pressure, then the equilibrium shifts to counteract the imposed change and a new equilibrium is established” (Atkins 1993). The hypothesis was supported in the sense that when equilibria were pushed in one direction by a change in concentration or temperature, there was an obvious change in equilibrium. This change in equilibrium was noted and observed by a color change or formation of precipitate.

References:

Atkins, P. W. (1993). The Elements of Physical Chemistry (3rd ed.). Oxford University Press.