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Determining an Equilibrium Constant Using Spectrophotometry

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❶Values of K for acid dissociation reactions K a values are useful for predicting the outcomes of acid-base reactions, which are thermodynamically controlled. Spectrophotometric Determination of an Equilibrium Constant.

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The solution was then mixed with a glass stirring rod. A second clean, dry cuvette was filled three-fourths full with this solution using a disposable Pasteur pipet. The spectrophotometer was blanked with the cuvette filled with 0. The cuvette used for measuring absorbance was filled with this solution using the disposable Pasteur pipet and was rinsed twice.

It was finally filled three-fourth full with the solution in the mL beaker. The spectrophotometer was blanked again with the 0. This process was repeated until For example using the first mixture, 1.

For example using the mixture 1, Using the first mixture as an example,. To find K, I took the concentration and order of the products and divided by the concentration and order of the reactants.

Finally, to find the average K, I added the 10 K values found and divided by My results do not seem very accurate. Even during the experiment, my partner and I agreed that our absorbance readings did not seem right. The readings jumped a significant amount between the first and second readings and the fourth and fifth readings.

After the seventh reading, the absorbance readings seemed to be evening out, as the difference between readings was getting smaller. When reading some of the absorbances, we took a second reading because a few times the absorbance reading was smaller than the previous.

This did not seem right; the absorbance readings should have been going up each reading. Error may be attributed to the fact that a few drops of solution escaped the mL beaker during transfers between the cuvette. Error could have also resulted from the beakers not being totally clean, or the pipetment of solutions could have been inaccurate. This would have caused the volumes and concentrations to be different than they actually were. Abstract This study was on the Mendelian phenotype ratio of corn and sorghum.

According to the Beer-Lambert law, the concentration of a colored species is proportional to its absorbance, which is the amount of energy it absorbs at a specific wavelength of light. This can be expressed mathematically, where A is absorbance, epsilon is the molar attenuation coefficient, which is compound-specific, l is the path length through the sample, and c is concentration.

A calibration curve is created by testing multiple solutions of known concentration, and plotting the resulting absorbance values. With this calibration curve, solutions of unknown concentration can be studied. Absorbance measurements are used to determine the concentration of the colored species. Then, the concentrations of the remaining reactants and products can be calculated.

The following procedure will study the reaction of iron three with thiocyanate to form an iron thiocyanate complex. Once the concentrations have been determined, the value for K can be calculated with an Initial-Change-Equilibrium, or ICE, table which will be explained further in the results. Now that you understand how spectrophotometric methods can be used to determine the equilibrium constant, you are ready to begin the procedure.

To begin, zero a UV-vis spectrophotometer using distilled water as a blank to represent no absorbance. When inserting a cuvette into the spectrophotometer, ensure that it is oriented so light passes through the transparent sides, and that the liquid level is above the path of the beam. Then, prepare 5 test tubes containing the indicated volumes of each reactant solution as shown in the text protocol, which will yield varying concentrations of the product. Cover each tube with a gloved finger, and gently shake to mix.

Allow the tubes to rest for 10 min. Use a Pasteur pipette to transfer a small quantity of the mid-concentration sample, solution 3, to a cuvette, and place it in the spectrophotometer. Use the same cuvette for all measurements, making sure to rinse 3 times in between each sample.

Repeat this process for solutions 2 — 5. Plot the measured absorbance versus concentration of iron thiocyanate for each solution. Determine the line of best fit for the data. The slope of this line is the molar attenuation coefficient. Now that the data for the standard solutions has been acquired, prepare four medium test tubes containing the indicated volumes of solutions as shown in the text protocol.

Cover each tube with a finger and gently shake to mix. Allow them to stand for at least 10 min. This resting period allows the solutions to reach chemical equilibrium.

Use a Pasteur pipette to transfer a small quantity of solution 6 to the cuvette, and place it in the spectrophotometer. Repeat this process for solutions 7 through 9. Once all of the samples have been measured, the molarity and absorbance data for solutions 1 — 5 can be analyzed.

A large excess of thiocyanate was used to ensure that all of the iron reacted, which simplifies the analysis. The data is plotted to create a calibration curve. The path length of light, l, is typically 1 cm, and can be factored out of the calculations.

The slope of the line, which was calculated to be , is therefore the attenuation coefficient. For the test solutions 6 — 9, this value and the absorbance are used to calculate the iron thiocyanate concentrations at equilibrium.

With this data, the ICE table could then be utilized. The initial reactant concentrations are based on the known molarities of iron and thiocyanate added to the solution, and the total volume of the reaction. Because the product is formed from the 1: The equilibrium concentration of each species is now known. These values are used to calculate the equilibrium constant for each solution. The values are roughly constant over the range of concentrations studied. The concept of the equilibrium constant is important to a wide range of scientific fields.

The equilibrium constant can be used to provide useful information about the extent to which a reaction will form products over time. In this example, two reactions containing crystal violet were observed. The first solution was composed of crystal violet and sodium hydroxide. The color was observed to rapidly change from purple to colorless. This reaction has a very large K value, indicating that the products form nearly completely over time. Crystal violet was then reacted with sodium acetate.

This solution remained purple indefinitely. This reaction has a very low K value, so it does not proceed forward to a significant degree. Finally, the dissociation constant — a specific type of equilibrium constant — can be used to describe protein behavior.

In this example, changes in the structure of RNA were monitored in magnesium reaction buffers. Purified RNA was mixed into solution with known concentrations of magnesium, and allowed to reach equilibrium.

Then, the resulting RNA structure was plotted. In this case, higher concentrations of magnesium caused reactive sites on RNA to be less protected, producing a Kd that was half the value. You've just watched JoVE's introduction to spectrophotometric determination of the equilibrium constant. You should now understand the relationship defined by the Beer-Lambert law, how to determine concentration from absorbance using a spectrophotometer, and how to calculate an equilibrium constant using equilibrium concentrations.

Table 4 lists the absorbance and concentration data for solutions 1 — 5. A large excess of SCN - was used in tubes 1 — 5 to ensure that this assumption holds true. The measured absorbances agree well with Beer's law. Table 5 lists measured absorbances and calculated K values for tubes 6 — 9. K values were determined using the ICE table method. Because all of the product was formed from the 1: Table 6 shows the process for test tube 6.

The equilibrium constant is calculated from the concentrations in the equilibrium row. For test tube 6,. Measured absorbance values and calculated K for the reaction of iron III with thiocyanate. The ICE table that illustrates the process used for test tube 6. The equilibrium constant provides useful information about the extent to which a reaction will proceed to form products over time. Reactions with a large value of K , much larger than 1, will form products nearly complete given enough time Figure 3.

Reactions with a value of K less than 1 will not proceed forward to a significant degree. The equilibrium constant thus serves as a measure of the feasibility of a chemical reaction.

The equilibrium constant of this reaction is greater than 1. A significant amount of colored product forms in each case, even though the initial concentrations of reactants differ. The equilibrium constant also provides useful thermodynamic information about the changes in free energy, enthalpy, and entropy in the course of a chemical reaction.

The equilibrium constant is related to the free energy change of reaction:. The free energy change of reaction is in turn related to the enthalpy and entropy changes of reaction:.

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Spectrophotometric Determination of an Equilibrium Constant. Spectrophotometric methods for the determination of K involve direct measurement of the concentration of one colored component and indirect measurement of the others. Visible spectroscopy of the colored component at known concentrations reveals the .

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At equilibrium, the reactants turn into product and the products decompose into reactants at the same rate. This ratio of the products to reactants at equilibrium is represented by the equilibrium constant, or K. K is found by taking the concentration and order of the products and dividing by the concentration and order of the reactants.

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The method used to determine which stoichiometry is correct involves using the three potential equilibrium constant expressions for the three possible reactions. (Updated ) 1 Determining An Equilibrium Constant Using Spectrophotometry and Beer’s Law Objectives: 1.) To determine the equilibrium constant for the reaction of iron (III) and thiocyanate to.

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Spectrophotometric Determination of an Equilibrium Constant - Free download as Word Doc .doc), PDF File .pdf), Text File .txt) or read online for free. Results and Discussion Report for Experiment 5. The equilibrium constant, Kc We can now substitute the equilibrium concentrations of all the species into the equilibrium expression (Eq. 2) and calculate Kc. Le Chatelier’s Principle (Tro, pp ) When the conditions of a system at equilibrium are altered, the system responds in such a way as to maintain the equilibrium.