CHEMICAL EQUILIBRIUM LABORATORY 3 Essay

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GEORGIA MILITARY COLLEGENATURAL SCIENCE DEPARTMENTONLINE CAMPUSLABORATORY 3 - CHEMICAL EQUILIBRIUMNAMESTUDENT NUMBERCLASSPROFESSOR’S TITLE AND NAMEIntroductionGenerally, equilibrium is a state of balance between opposing forces. In chemical reactions, equilibrium is achieved when the concentrations of the products and reactants are in balance, so no further changes are observed in the system (Smith, 2024). Chemical equilibrium plays a fundamental role in industrial processes as well as in human life. Why is it important to understand the concept of equilibrium in industrial processes? This question is best answered using the Le Chatelier’s principle, which states that a change in one of the elements of a system in dynamic equilibrium will trigger a shift in the equilibrium position in an attempt to counter the change and reestablish equilibrium (Smith, 2024). Factors that can cause changes to a system in equilibrium include reactant concentrations, temperature, and pressure (Smith, 2024). According to Le Chatelier’s principle, an increase or decrease in any of these factors will trigger a shift in the equilibrium point in the opposite direction (Smith, 2024). Using this knowledge, industrial chemists can adequately manipulate chemical reactions to increase or decrease the production of certain products.This laboratory uses the reaction between Cobalt (II) and chloride ions to observe how the equilibrium point changes due to changes in temperature and concentration of reagents. Figure 1 below presents the chemical equation for the above reaction:[Co (H2O)6]+2 + 4Cl- ? CoCl4-2 +6H2OFigure 1: [Co (H2O)6]+2 forms a pink complex, while Co Cl 4-2 is a blue complex.Combined with six water molecules, Cobalt (II) forms a pink complex that turns into a blue complex, CoCl42 upon reacting with chloride ions. This laboratory seeks to realize three objectives:i) To enhance the ability to apply Le Chatelier’s principle.ii) To enhance understanding of the equilibrium constant concept.iii) To analyze the effects of changes in temperature or concentration on the equilibrium constant.From figure 1 above, the equilibrium constant (k) for the reaction is given by:The general hypotheses established at the start of the lab were:i) Addition of chloride ions will increase the reactants above the products in the equilibrium constant, thus shifting the equilibrium in figure 1 to the right (the products side), leading to more production of CoCl42, and causing the solution to turn blue.ii) Addition of distilled water to the reaction in figure 1 will decrease the concentration on the products side of the constant. Thus, there will be more reactants than products, causing the equilibrium to shift to the left (the reactants side), leading to more production of [Co (H2O)6]+2, which causes the solution to turn pinkiii) Addition of silver Nitrate (AgNo3) will reduce chloride ions (Cl-) due to the formation of AgCl. This will imply that there will be more products than reactants and the equilibrium will shift to the right (towards the products) to produce more chloride ions. Consequently, the solution turns blue.Materials and MethodsPreparing the Labi) Select ‘Virtual Lab’ on the home page of the course to load the lab environment.ii) Wait for the lab environment to load, then select ‘File’, and subsequently, ‘Load an Assignment’.iii) Choose the category labeled ‘Chemical Equilibrium’ and then the assignment titled ‘Cobalt Lab.’ At this point, the lab preparations are complete and one is ready to perform the experiment.

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Performing the Experimentiv) In the stockroom, choose ‘Glassware’ and then ‘Empty 1000mL Erlenmeyer flask.’v) Return to the stockroom and in the ‘Solutions’ tab, select ‘Cobalt (II) Chloride Exp’ solutions. Move to the work area the flask containing 1M Cobalt (II) Chloride solution.vi) Return to the ‘Solutions’ tab, select HCL, and move the flask containing 12M HCL to the work area.vii) Return to the ‘Solutions’…

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…to absorb the extra heat caused by the temperature increase (Smith, 2024). Thus, the equilibrium shifts to the right, producing more CoCl42, which leads the solution to turn blue. The accuracy of the procedure could be increased by ensuring all instruments are properly calibrated and conducting parallel experiments to help identify anomalies in the collected data.ConclusionThis laboratory used the reaction between Cobalt (II) and chloride ions to study changes in the equilibrium position resulting from changes in temperature and concentration of reagents. The experiment followed Le Chatelier’s principle, which states that a change in one of the elements of a system in dynamic equilibrium will trigger a shift in the equilibrium position in an attempt to counter the change and reestablish equilibrium. At equilibrium and room temperature, the solution of Cobalt (II) Chloride forms a pink complex. Addition of concentrated HCL increases the concentration of chloride ions in the solution, causing the equilibrium point to shift to the right, leading to the production of more CoCl42, which forms a blue complex at equilibrium. Thus, the solution turns blue. Similarly, addition of distilled water reduces the concentration of CoCl42 , causing the equilibrium to shift to the left, and the solution turns back to pink. Addition of silver nitrate (AgNo3) to the equilibrium blue solution takes away Cl- ions from the solution through formation of silver chloride (AgCl), causing the equilibrium to shift to the left and the solution turns pink. Finally, heating the pink equilibrium mixture causes the equilibrium point to shift to the right as the system seeks to reestablish equilibrium by absorbing the excess heat caused by the temperature change. Consequently, the solution turns blue. While measures were taken to ensure accurate results, one could improve the accuracy of the procedure by conducting parallel experiments to check for anomalies….....

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