Lab Report (Exp 19) PDF

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Lab Report (Exp 19)...

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Spring 2018 CH 229- Afternoon Class

Experiment #19: The Cobalt Chloride Equilibrium System

Prepared for: Eric Beyerle Prepared by: David McNeely

April 17, 2018

Introduction The purpose of this laboratory experiment was to perform a test in order to physically view Le Châtelier’s principle (explained in section titled “Unique Discussion Topics”). In order to view this principle, the concentration of both the reactants and the products of the reaction were altered, as well as the temperature of the reactant solution. These variables allow one to see the physical manifestations of Le Châtelier’s principle. The reaction at question is the forward reaction of the formula listed under the section titled “Calculations, Graphs, and Chemical Equations” below. This laboratory experiment involves the concept of equilibrium reactions. An equilibrium reaction is once such that both the forward and backward reactions are capable of being performed and have reached a state of equilibrium (the forward and backward rates of reaction are at the same magnitude, and there is no net reaction occurring). This concept could be of interest because it allows for the possibility of a substance or solution to be stored in a state of stable equilibrium and then a certain reactant or product could be added to the solution in order to set either the forward or backward reaction into motion once again. The questions that were decided to be answered after the conduction of this laboratory experiment were: “what effect does heating have on the solutions?”, “what is the difference in composition between aqueous and alcoholic solutions that are the same color?”, and “if the concentration of water to have been decreased, would the system have reacted to form more reactants or products?”. This laboratory experiment is relatively safe but does involve the use of concentrated hydrochloric acid; therefore, care and proper safety equipment should be used during the entirety of this experiment process. Similarly, cobalt(II) chloride is toxic to many different forms of life, so gloves should be worn at all times. This laboratory experiment is to be conducted by preparing one set of solutions of varying concentrations of cobalt(II) chloride, water, and hydrochloric acid and another set of solutions using isopropyl alcohol in place of water from the first set. The solutions are then to be heated and allowed to air cool or cool in an ice bath (depending on the solution). Various characteristics of the solutions can then be determined through qualitative analyses. Experimental Data and Observations Wetted Filter Paper Observations: light pink color observed immediately after cobalt(II) chloride sol’n was added; a blue/purple tinted color was observed after the filter paper was allowed to dry Part B Observations: Sol’n B1-light pink color observed Sol’n B2-murky pink color observed Sol’n B3-blue color w/ pink tones observed Sol’n B4-light blue/purple color observed

Sol’n B5-dark blue/purple color observed All Sol’ns did not initially mix but combined after thorough stirring Part C Observations: Sol’n C1- light pink color observed Sol’n C2- light blue/purple color observed Sol’n C3- link pink color on bottom, blue/purple color on top observed Sol’n C4- blue/purple color observed Sol’n C5- blue/purple color observed All Sol’ns did not initially mix but combined after thorough stirring Part D Observations: -Water/HCl sol’n released small amount of heat after addition of HCl -B4 sol’n turns from light pink to blue/purple color when heated -The C2 alcohol sol’n is the most similar to sol’n B4 in terms of color -C2 sol’n turns from light pink to blue/purple color when heated -Heated B4 sol’n returns to original color when allowed to air cool -Portion of heated C2 sol’n submerged in ice bath quickly returns to original color Calculations, Graphs, and Chemical Equations Cobalt(II) and Chloride Ion Equilibrium Reaction: 2−¿ ( aq ) +6 H 2 O (aq) −¿ ( aq ) CoC l¿4 ¿ 2+ ¿ ( aq )+ 4 C l ¿ H 2 O ¿6 Co¿ Claims and Evidence Both solutions that were heated on the hot plate turned from a light pink color to a dark blue/purple color. The aqueous and alcohol solution that most closely resembled each other were solutions B4 and C2. Solution B4 had 1 mL of 1 M cobalt(II) chloride, 4 mL of water, and 3 mL of hydrochloric acid. Solution C2 had 1 mL of 1 M cobalt(II) chloride, 6 mL of isopropyl alcohol, and 1 mL of hydrochloric acid. If the concentration of water in the system is decreased, the system will react to form more products. This statement can be affirmed through the observational data found in “Part B Observations.” As the chloride ion concentration increases and the water concentration decreases, the solutions turn progressively darker and bluer in color.

Since the results of this experiment are solely qualitative, quantitative and traditional, graphical representations are incapable of providing evidence for the claims made from the

results of this experiment. If pictures were permitted to be taken during the conduction of this experiment, photographic evidence that supported the list of observations would prove beneficial in support of the claims stated above. For example, a picture would provide visual evidence that the C2 alcohol solution was, in fact, the alcohol solution that most closely resembles solution B4 in terms of their blue/purple color compared to the light pink color of the other solutions. Analysis and Discussion Since the results of this experiment are entirely observational, an analysis of the accuracy of the results is not really possible. Observations on color (which comprise much of the results of this experiment) are subjective and should not be the sole type of data collected and reported for any given experiment if at all possible. Mixing different concentrations of the solutions improperly or an improper identification of the color of a solution are the main types of error that I can identify. Both of these types of error are indeterminate because they cannot be accounted for or predicted.

Unique Discussion Topics (Questions from page 154 of lab manual) Increasing the chloride ion concentration of the solutions causes them to become darker and bluer in color. This seems to suggest that more product if being formed, therefore increasing the intensity of the blue color of the solutions. This conclusion can be supported by the observational data found in “Part B Observations.” Heating the solutions causes the equilibrium reaction to shift to the right, forming more products and becoming darker in color. This effect is the same for both the aqueous and alcohol solutions (as seen in “Part B Observations” and “Part C Observations”). My observations can be explained in terms of Le Châtelier’s principle. Le Châtelier’s principle states that when an equilibrium reaction becomes unbalanced (either through the addition or removal of a reactant or product) the reaction will immediately begin to reach equilibrium once again. For this experiment, more products are formed when the solution is heated up (causing the darkening and change of hue). Similarly, more reactants are formed when the solution is cooled down (causing the lightening and change of hue). The observational data of solutions B1 and C1 demonstrate a light pink color caused by a 2+¿ −¿ diluted solution of cobalt(II) ( ions) and a dark blue color when the Cl ¿ ions from the C o¿ 2−¿ 2+¿ hydrochloric acid react with the ions from the cobalt(II) solution to form CoC l¿ . The ¿ Co 4 2−¿ 2+¿ ions are pink, and the CoC l¿ ions are blue. C o¿ 4 Reaction 1 is endothermic. This can be determined by the fact that certain solutions did not change from a light pink color to dark blue until they were heated; this addition of heat

serves to prove that the reaction is endothermic, since it required an addition of heat to be performed. Similarly, the removal of heat via an ice bath returns the solution back to its original color; this shows that the removal of heat (an endothermic process) reverses the reaction, again proving its endothermic characteristic. Reaction 2 is reversible. This can be concluded based on the observational data showing the color change of solution B4 from light pink to blue/purple and back to light pink again during its heating and cooling processes. Reflection Topics Extend If this laboratory experiment were to be extended, it would be interesting to see how the results would change if different concentrations of cobalt(II) chloride were used and to measure the rate at which the solutions returned to their original colors during cooling. Different ratios of water/isopropyl alcohol, cobalt(II) chloride, and hydrochloric acid could also be used to see how these different ratios affect the time it takes for the solutions to change color during heating and how fast they return to their original colors during cooling. CH223 Lecture Topics One lecture topic that relates to this experiment is that of equilibrium reactions. Equilibrium reactions are interesting because it appears as though they are stagnant and unchanging. However, the state of equilibrium that exists in this sort of reaction is achieved through the balancing of the forward and backward reactions. Both directions of the reaction are occurring at the same time and rate, so it appears as though no reaction is taking place. When in reality, two different reactions are occurring simultaneously. Applications This sort of reaction can be applied in the real world in the field of medicine. Studies have shown that there exists an equilibrium reaction between hemoglobin and oxygen. As long as there is sufficient oxygen in the bloodstream, hemoglobin molecules will reaction with four oxygen molecules each in order to form two molecules of oxygenated hemoglobin (Chemical Equilibrium). Knowing this attribute of this biochemical reaction allows scientists, doctors, and researchers to be able to determine how certain environmental and chemical factors might affect this process and the effects it might have on the human body. Related Reading According to the Royal Society of Chemistry, two different species of cobalt(II) chloride (one pink and one purple) can “exist together in equilibrium” and change color depending on temperature (The Equilibrium…). This confirms the results of this experiment regarding the color change of the solutions when heated in the hot water bath and returning to their original color when cooled. Green Chemistry The main principles of Green Chemistry that apply to this experiment are that of energy efficiency and real time analysis. Just enough energy (i.e. thermal energy) is used in order to heat the solutions up to the required temperature and then cool them down using the air or ice.

Real time analysis is used in order to prevent excess waste materials in the form of byproducts. Real time analysis of the solutions allows for reactions to be halted when observations are complete and any materials used during the experiment to be cleaned or disposed of.

Works Cited “Chemical Equilibrium - Real-Life Applications.” Science Clarified, www.scienceclarified.com/everyday/Real-Life-Chemistry-Vol-2/Chemical-EquilibriumReal-life-applications.html. “The Equilibrium between Two Coloured Cobalt Species.” The Equilibrium between Two Coloured Cobalt Species- Learn Chemistry, www.rsc.org/learnchemistry/resource/res00000001/the-equilibrium-between-two-coloured-cobalt-species? cmpid=CMP00005957....