Experiment 8 4m5w38

Lauren Lee Lab Report 8 Lab Partner: Nadia Underwood 25 March 2014

Le Châtelier’s Principle Introduction Le Chatelier’s Principle shows how concentration, pressure, and temperature affect the equilibrium position for a reaction. With a constant and given temperature, the Kc value remains the same however the amount of reactants and products may shift to accommodate the changes made to the reaction to create equilibrium again. When the temperature is changed, the Kc value changes as well and the reactants and products shift to fit the new Kc value. Using colored substances can help the understanding of this principle. Procedure Part 1: Thiocyanatoiron(III) System To create the system, the bottom of a Petri dish was covered with the KSCN solution. Two to three drops of Fe(NO3)(aq) was added to the dish. The Petri dish was tilted and lightly swirled to mix the solution. The Petri dish was set on top of a white piece of paper. Observations were recorded. A very small crystal of solid KSCN was put into the dish carefully. The dish was swirled to dissolve the crystal in the solution. The liquid in the Petri dish was divided into 2 clean and different test tubes. One percent AgNO3 was added drop by drop into one of the test tubes until the solution’s color changed. The precipitate was ignored; the color changes were recorded. Some small crystals of Na2HPO4 were added to the remaining test tube. Observations were recorded. Part 2: Methyl Orange System In one test tube, 3 mL DI H2O and 2-3 drops of In- (methyl orange: an acid-base indicator) were mixed together in a small, clean test tube. Drop by drop, 1 M HCl was added to the test tube. Observations were recorded. After the color change, the solution was swirled to mix. To the same test tube, 1 M NaOH was added dropwise. Observations were recorded.

Part 3: Cobalt(II) Chloride System One millimeter of [Co(H2O)6]Cl2 was added to a small clean test tube. The color was recorded. A hot plate and a 50 mL beaker of water was set up. The water was brought to a simmer. The test tube was set in the beaker of hot water, held by test tube tongs. Observations were recorded. Data Table 1: Part 1 – Recorded Observations and Color Changes Additions/ Changes to System Observations Dish was tilted, solution was mixed  Light yellow to light orange to dark orange ( dark orange increased with swirling) Small crystal of KSCN was added  Darker orange or orange-red 1% AgNO3 drops were added  Milky substance  Precipitate was made  The final color was milky white to a polluted light yellow Na2HPO4 crystals were added  Clear Table 2: Part 2 – Recorded Observations and Color Changes Additions/ Changes to System Observations Initial Color  Light yellow or orange 1 M HCl drops were added  Light pink or red as soon as HCl drops were added 1 M NaOH drops were added  Gradually became yellow Table 3: Part 3 – Recorded Observations and Color Changes Additions/ Changes to System Observations Initial  Pink Adding heat (test tube in simmering water)  Slowly changed to dark pink, to purple, to blue

Discussion Part 1: Thiocyanatoiron(III) System During this experiment, the first stressor was introduced to the equilibrium through the form of the 1% AgNO3 drops. The initial reaction observed is this: 𝐹𝑒 3+ (𝑎𝑞) + 𝑆𝐶𝑁 − (𝑎𝑞) ⇌ 𝐹𝑒𝑆𝐶𝑁 2+ (𝑎𝑞) (light yellow) (colorless) (red) But adding the AgNO3 yielded the following reaction: 𝐴𝑔+ (𝑎𝑞) + 𝑆𝐶𝑁 − (𝑎𝑞) ⇌ 𝐴𝑔𝑆𝐶𝑁(𝑠) The AgNO3 reacted with the reactant SCN- to form an unreactive precipitate, removing the SCN- from the initial reaction. Although SCN- is colorless, its removal caused the solution to become a milky white and light yellow. This observation can be explained by Le Chatelier’s Principle. Because the concentration of one of the reactants decreased (or some of the reactant was removed from the reaction), the equilibrium shifted to the left to accommodate the change made in the system. Since the equilibrium favored the reactants more in this case, more reactants were being made and so the color of the system was more light yellow (color of the reactant Fe3+) than red (color of the products). If 1 drop of 1 M FeCl3 was added to the solution of Fe(NO3)3 and KSCN at equilibrium, the equilibrium would shift to the right, favoring the products. This is because the addition of 1 M FeCl3 increase the concentration of reactants (Fe3+) in the reaction. According to Le Chatelier’s Principle, adding reactants to a system pushes the equilibrium to the right to for the increase. Thus the amount of products would increase and the solution would become a darker red. Part 2: Methyl Orange System From the devised experiment, it can be concluded that red form of methyl orange was the protonated indicator and the yellow form was the un-protonated indicator. When methyl orange was in the presence of 1 M HCL, a highly acidic substance, the solution immediately became red. Water is always at a neutral pH of about 7 so it did not create the high acidity or protons in the reaction. HCl completely or nearly completely dissociated, creating many protons or H+. When the H+ was present, it bound to the N of the indicator and made the solution red. The HCl was the acid because it donated protons while the indicator became the base or the conjugate acid because it accepted protons. Knowing that HCl made protons which bound with the indicators that created protonated indicators, it was seen that the red solution contained the protonated methyl orange.

Similarly, when NaOH was added to the solution, it became yellow, the color of the unprotonated indicator. NaOH dissociated into OH- ions which reacted with the H+ protons to create H2O. This removed the protons from the indicator and created colorless water. Thus the indicator was left and its color showed – yellow. Part 3: Cobalt(II) Chloride System This reaction in the forward direction was endothermic. The reaction from the system looked like the one below: [Co(H2O)6]Cl2(aq, pink) ⇌ Co𝐶𝑙2 (𝐻2 𝑂)2 (𝑎𝑞, 𝑏𝑙𝑢𝑒) + 4 𝐻2 𝑂(𝑙) This reaction was endothermic because endothermic reactions absorb heat to form bonds. When the test tube of [Co(H2O)6]Cl2 sat in the beaker of simmering water, the solution was absorbing heat. When the solution absorbed enough heat, it changed from pink to blue because its new bonds had changed and formed. Heat in this situation was also similar to a reactant that helped to form the product. Therefore, the increase of the heat reactant shifted the equilibrium towards the products and increased the blue color. Based on that information, the increase of temperature also increased the value of Kc. Since the equilibrium shifted to the right to absorb more energy, the equilibrium constant increased as well. The products formed at a faster rate Conclusion These experiments helped apply Le Chatelier’s Principle to different reactions. The experiments also explored different aspects of the Principle such as the effects of changes in concentrations and temperature. Application to different reactions demonstrated the usefulness of this Principle. This lab also allowed students to use critical thinking skills to create experiments to test for certain objectives. Not only did it make students apply knowledge but also allowed a degree of freedom in the lab procedures. One problem in the lab was that the heating of the [Co(H2O)6]Cl2 took a long time to convert it into the products. The color change progressed slowly so perhaps there was a mistake in the concentrations of the reactants or in the heating process. Human error could be found in the precision of the amounts of reactants or products added to the test tubes. Additionally, because the data was qualitative so the data could have been subjective, lending itself to human error. Although qualitative data served its purposes for this lab, the data could not be quantified in an exact universal way.