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Spectrophotometric determination of equilibrium constant for a reaction

  • The calibration curve is shown as Figure 1;
  • Follow procedure for other test tubes;
  • Follow procedure for other test tubes results at Appendix A;
  • The equilibrium constant expressions corresponding to the three possible stoichiometries being considered are given on Page 139 of the lab manual;
  • 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.

Appropriate volumes of 0. Determining the equilibrium constant of a chemical reaction can provide important information about the extent to which it will form products over time. Every chemical reaction is associated with an equilibrium constant, K, which reflects the ratio of the concentrations of the products and reactants when the reaction has stopped progressing. To measure K, these concentrations must be determined. If a reaction contains a single colored component, its interaction with light can be measured to discern its concentration.

The concentrations of the uncolored components can then be calculated indirectly using the balanced chemical equation. This video will illustrate the use of a spectrophotometer to empirically determine the equilibrium constant for an iron thiocyanante reaction.

Most chemical reactions proceed in both forward and reverse directions. As the reaction progresses, it reaches a point where the forward and reverse reactions occur at the same rate. This is known as chemical equilibrium. At this steady state, the ratio of product concentrations to reactant concentrations, each raised to the power of their stoichiometric coefficients, corresponds to the equilibrium constant, K.

To measure K for a system of interest, the coefficients should be known, and the concentrations must be determined, either directly or indirectly. 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.

  1. Most chemical reactions proceed in both forward and reverse directions. This reaction has a very low K value, so it does not proceed forward to a significant degree.
  2. Initial concentration of reactants for Upon setting the zero for the spectrophotometer, calibration.
  3. Now that you understand how spectrophotometric methods can be used to determine the equilibrium constant, you are ready to begin the procedure.
  4. This reaction has a very large K value, indicating that the products form nearly completely over time.
  5. A large excess of SCN- was used in tubes 1 — 5 to ensure that this assumption holds true. K values were determined using the ICE table method.

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. Before measuring the sample, a calibration curve must be generated. 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.

  • To measure K, these concentrations must be determined;
  • Table 5 lists measured absorbances and calculated K values for tubes 6 — 9;
  • The slope of the line, which was calculated to be 7600, 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.

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.

Spectrophotometric Determination of an Equilibrium Constant

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 7600, 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.

  1. The equilibrium constant was calculated by substituting the respective concentrations to the equilibrium constant expression of the reaction.
  2. For test tube 1, substitute values to the expression.
  3. 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.
  4. Results are listed in Table 3. In its more advanced form, the given the measured absorbance of the test tube Beer-Lambert law states that.

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.

  • A [4], [6], [7] and [8];
  • 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;
  • A [4], [6], [7] and [8].

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. Results 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.