Le Chatelier's Principle Learning Goals Source Math ✓ Solved

Le Chatelier's Principle W2021 » Learning Goals » Source Materials

To describe shifts in equilibrium when conditions, such as concentration and temperature, are changed in the General Chemistry II “Le Châtelier’s Principle” laboratory experiment. To explain the observed changes in terms of Le Châtelier’s Principle.

You should refer to the observations recorded in your lab notebook for the CHEM 162 "Le Châtelier's Principle" experiment to complete this assignment. You may also refer to the appropriate sections in your General Chemistry II lecture notes and in the OpenStax textbook.

Your essay should be cohesive and organized. Do NOT simply list all of your explanations and all of your answers to the questions. Be sure to include a topic sentence that effectively introduces the essay's subject. There is a 500 word limit for the essay, so write clearly and concisely. The essay should be impersonal and written in the third person with no first person references—i.e., do not use "I", "we", "my", "our", etc., in the essay.

First, state Le Châtelier's Principle. For each system, describe the observations for each change imposed, and explain your observations. The explanations should include the following information: 1) the change imposed on the system (i.e., the chemical added and the resulting reactant or product concentration that was affected and how, or if heat was added or removed); 2) indicate if the equilibrium shifts right (in the forward direction) or shifts left (in the reverse direction); and 3) indicate the concentration changes for reactants or products that lead to the observed color change or the change in the indicator color based on pH.

Paper For Above Instructions

Le Chatelier's Principle states that if a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the system will adjust to counteract that change, shifting the equilibrium position. This principle is foundational in understanding chemical reactions and their dynamics, especially in various laboratory settings.

In this experiment, we examined several systems to observe how they responded to different imposed changes, such as alterations in temperature and concentration. The first equilibrium system involved the cobalt complex, where the equilibrium shifted due to temperature changes and the addition of chloride ions.

Part I: Cobalt Complex Equilibrium

The reaction can be expressed as follows:

[Co(H2O)6]²⁺(aq) + 4Cl⁻(aq) ⇌ [CoCl4]²⁻(aq) + 6H2O(l)

Initially, the solution appeared pink, indicating the presence of the [Co(H2O)6]²⁺ complex. When the solution was heated, it turned blue. This color change indicates that the equilibrium shifted to the right when heat was added, favoring the formation of the [CoCl4]²⁻ complex, which is blue and confirms that the forward reaction is endothermic—absorbing heat. When cold water was added, the equilibrium shifted left, returning to the pink color, as the system worked to replace the heat that was removed.

Part II: Iron Thiocyanate Equilibrium

The second system investigated was the iron thiocyanate equilibrium:

Fe³⁺(aq) + SCN⁻(aq) ⇌ [FeSCN]²⁺(aq)

Initially, the solution was light yellow. The addition of Fe(NO3)3 created a red solution, indicating that the equilibrium shifted right, favoring the formation of the [FeSCN]²⁺ complex. The further addition of KSCN maintained this red color, still shifting the equilibrium right. However, when Na2SO3 was introduced, the color lightened to yellow, as it reduced Fe³⁺ to Fe²⁺, resulting in an equilibrium shift left to replenish Fe³⁺ that was removed.

Part III: Ammonia-Ammonium Ion Equilibrium

This system can be described by the equation:

NH3(aq) + H2O(l) ⇌ NH4⁺(aq) + OH⁻(aq)

Initially, the addition of phenolphthalein caused the solution to turn dark purple, signaling basic conditions due to the presence of OH⁻ ions. Once HCl was added, the solution became colorless as acidity increased, triggering a shift to the right where NH4⁺ ions were produced to counteract the added protons. Conversely, the introduction of NH4Cl again increased acidity, resulting in a shift left to offset the increase in NH4⁺ concentration.

Part IV: Acetic Acid-Acetate Ion Equilibrium

The last equilibrium under consideration was the acetic acid system:

HC2H3O2(aq) + H2O(l) ⇌ H3O⁺(aq) + C2H3O2⁻(aq)

The initial solution was colorless. The addition of NaOH caused the solution to turn colorless again, indicating a shift to the left to consume H3O⁺. When sodium acetate (C2H3NaO2) was added, the color remained clear but was more basic, reinforcing the shift left due to the increase of acetate ions.

In summary, throughout the experiments, changes in temperature and concentrations facilitated observable shifts in equilibrium, demonstrating the practical applications of Le Chatelier's Principle in a laboratory context.

References

  • OpenStax University. (2021). General Chemistry. OpenStax.
  • Atkins, P. W., & de Paula, J. (2010). Physical Chemistry. Oxford University Press.
  • Brown, T. L., LeMay, H. E., Bursten, B. E., & Murphy, C. J. (2013). Chemistry: The Central Science. Pearson.
  • Zumdahl, S. S., & Zumdahl, S. A. (2017). Chemistry. Cengage Learning.
  • Petrucci, R. H., Harwood, W. S., & Herring, F. G. (2016). General Chemistry. Pearson.
  • Rayner-Canham, M. & Rayner-Canham, G. (2017). Chemistry 2ed. A Scientific Approach. Nelson.
  • University of Wisconsin-Madison. (2012). Twelve Common Errors: A Student Self-Editing Guide. Retrieved from URL: https://content/uploads/sites/535/2018/07/twelve_common_errors_uwmadison_writingcenter_rev_sept2012.pdf
  • Riegert, M. A., & McKinney, E. (2012). Green Chemistry: Chemistry for Sustainable Development. Wiley.
  • Vaida, V. & Ramakrishnan, V. (2013). Principles of Chemical Ecology. Oxford University Press.
  • Garnett, J. & Rogers, P. (2013). Science in Context: Equilibria. Journal of Chemical Education.

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