Experiment 34: Lab Report About The Equilibrium Constant PDF

Preview

Summary

Lab reports...

Description

Experiment 34

An Equilibrium Constant

CHEM-1112-04

The University of Texas Rio Grande Valley

Spring 2019

Swati Mohan

Emilio Molina

Objective: The experiment performed was meant to determine the equilibrium constant of a chemical system by using a spectrophotometer. Once data was obtained graphing techniques are used to analyze the data and figure out the equilibrium constant.

Introduction: The experiment was completed to determine the equilibrium constant of a chemical reaction using Fe3+ (aq) and SCN-(aq). This chemical reaction created a state of chemical equilibrium. The equilibrium state can be categorized by stating its equilibrium constant as large, greater than one, or small, less than one , but not negative or zero. A large Kc (equilibrium constant) means that at equilibrium, the concentrations of the products will generally be greater than the concentrations of the reactants, meaning it is products favored. A small Kc means that at equilibrium, the concentrations of the reactants will generally be greater than the concentrations of the products, meaning it is reactants favored. The relevant chemical equation for this lab is: Fe3+(aq) + SCN- (aq) ⇌ FeSCN2+(aq) Once the equilibrium concentration of FeSCN+2 (aq) was determined, the equilibrium concentrations of the reactants (Fe+3 (aq) and SCN– (aq) can be calculated.

Procedure: The spectrophotometer was turned prior to the preparing of the solutions so that it had time to become warm. A set of standard of solutions was created according to Table 34.1 below. Table 34.1 Composition of the Set of Standard FeNCS2_ Solutions for Preparing the Calibration Curve

Six 25 mL volumetric flasks were first rinsed and dried to ensure that there was no residue, from previous experiments. The volumetric flasks were then labeled blank and 1-5 so that the order would not be confused. Following Table M Na was added to each volumetric flask. A plastic graduated cylinder was used to measure this solution. It was not cleaned after every used of the same liquid. The solution was poured carefully and slowly into the volumetric flask to ensure that none of it was spilled. Then 10 mL of 0.2 M of Fe(NO3)3 was added to each volumetric flask. Using the same plastic graduated cylinder that was now cleaned and dried, the next solution was added to the volumetric flasks. All of the procedure was done as followed in the lab manual. The solutions needed to be measured by the spectrophotometer. Six cuvettes, a special piece of glassware that holds solutions to be measured in a spectrophotometer, were gathered so that some of the solution could be put into each cuvette. After carefully mixing the solution, each solution was poured into a cuvette. Each cuvette was filled about three-fourths of the way so that there would be enough of the solution to measure. Lids were securely placed on top of

each cuvette to ensure that they were not spilled. All six of the cuvettes were brought to the spectrophotometer so that their absorbance could be measured. After the data were gathered and record, they were plotted in an absorbance versus concentration line slope equation. A calibration curve was created to determine the accuracy and to determine if part one needed to be repeated. Once everything was established, then the next part of the experiment could be accomplished.

Calculation:

Results and Discussion: The significance of knowing an equilibrium constant was to be able to know if the reaction is products favored or reactants favored. From a practical standpoint, producing a given chemical product, it would be essential to know the Kc of a reaction so that the yield of the product could be optimized. If Kc was very large, the concentration of the products was much greater than the concentration of the reactants. The reaction essentially "goes to completion." All or most of the reactants were used up to form the products. If Kc was very small, the concentration of the reactants was much greater than the concentration of the products. The reaction does not occur to any great extent. Most of the reactants remain unchanged and there were few products produced. When Kc was not very large or not very small, then approximately equal amounts of reactants and products were present at equilibrium. The errors that affected the results were due to several items: human factors and equipment or instruments used. There were several different people who measure the liquids. Each one with a different perspective on when volumes have been reached. Each measurement was checked by all, but there was still the chance for errors.

Conclusion: This experiment was to determine the equilibrium constant of a chemical reaction using Fe3+(aq) and SCN-(aq). The concentration of [FeNCS2+] was taken from several known standard concentrations and then graphed to form a line slope equation to determine the Beer’s law equation and the R2 value to be applied to several unknown concentrations and experimentally calculate the average Kc value which was derived from several test solutions. The results of these measurements determined the equilibrium constant for the formation of [FeNCS2+]. This was accomplished by using a spectrophotometer to measure the light absorbency and the data from this was used to graph a calibration curve to determine the molar absorptivity that was proportional to the thickness of the sample, concentrations of the absorbing solution, and the absorptivity of the samples.

Laboratory Questions: 1. All spectrophotometers are different. The spectrophotometer is to be set at 447 nm. What experiment could you do, what data would you collect, and how would you analyze the data to ensure that 447 nm is the best setting for measuring the absorbance of FeNCS2_ in this experiment? A scan of wavelengths using a standard solution of the substance you are interested in measuring. Assuming that the maximum absorbance occurs at 447 nm, then continue to use that wavelength. As the concentration increases, the absorbance will also increase. You can then plot these data and you should get a straight line. From this, you can actually calculate the molar extinction coefficient. Another experiment you might be able to do is to watch the absorbance decrease with time assuming that the compound of interest is not stable. As the compound degrades, you should be able to see a decrease in A, and then plot that against time to get a rate of decay. 2. In a hurry to complete the experiment, Joseph failed to calibrate the spectrophotometer. As a result, all absorbance values for the standard solutions that are measured and recorded are too high. How will this affect the following for the Test Solutions in Parts B and C? a. Will the equilibrium concentrations of FeNCS2_ be too high, too low, or unaffected? Explain. Too high absorbance values will give too high concentrations of Fe(NSC)2+ according to a Beer's law. b. Will the equilibrium concentrations of Fe3_ be too high, too low, or unaffected? Explain. The equilibrium concentrations of Fe3+ will be too low, because the calculated reacted number of moles will be to high according to a reaction c. Will the calculated equilibrium constants be too high, too low, or unaffected? Explain. The equilibrium constants are directly proportional to the concentration of a complex Fe(NSC)2+, therefore, they will be too high as well. 3. One of the standard solutions had an abnormally low absorbance reading, causing a less positive slope for the data plot. a. Will the equilibrium concentrations of FeNCS2_ in the Test Solutions (Part B) be too high or too low? Explain. Beer's law, where A is absorbance, e is molar absorptivity, c is concentration and l is path length Plot of A vs c gives straight line with slope = el, If slope is less positive => calculated el value is too low => calculated c will be too high for a fixed value of A from test solution (c x el = fixed A value so low el gives high c) => equilibrium concentration [Fe(SCN)2+] will be too high b. Will the calculated Kc for the equilibrium be too high, too low, or unaffected by the erred data plot? Explain. Fe3+(aq) + SCN-(aq) Fe(SCN)2+(aq), Kc = [Fe(SCN)2+]/[Fe3+][SCN-] If [Fe(SCN)2+] is too high (due to erred data plot), calculated Kc will be too high

4. Fingerprint smudges are present on the cuvette containing the solution placed into the spectrophotometer for analysis. a. How does this technique error affect the absorbance reading for FeNCS2_ in the analysis? Explain. The fingerprint smudges will result in apparent higher absorbance. Therefore, the values of absorbance will be greater. b. Will the equilibrium concentration of FeNCS2_ be recorded as being too high or too low? Explain. Too high absorbance values will give too high concentrations of Fe(NSC)2+ according to a Beer's law: c. Will the equilibrium concentration of SCN_ be too high, too low, or unaffected by the technique error? Explain. The equilibrium concentrations of SCN- will be too low, because the calculated reacted number of moles will be to high according to a reaction d. Will the Kc for the equilibrium be too high, too low, or unaffected by the technique error? Explain. The equilibrium constants are directly proportional to the concentration of a complex Fe(NSC)2+, therefore, they will be too high as well.

5. For the preparation of Test Solution 8 (Table 34.2), the 2.0 mL of 0.1 M HNO3 is omitted. a. Will this technique error cause the absorbance reading for FeNCS2_ to be too high or too low? Explain. This would result into a very low concentration of FeNCS2+. HNO3 ensures that all of Fe is present in the +3 oxidation state necessary for the complex formation. If HNO3 is avoided we would have lower concentration of Fe3+ ion in solution which would yield lower amount of FeNCS2+. b. Will the Kc for the equilibrium be too high, too low, or unaffected by the technique error? Explain. Kc which is the ratio of product complex concentration to the concentration of substrates present in solution at equilibrium would be too low. This is because the complex FeNCS2+ is low and thus give a lower Kc value. 6. The equation, A = a • b • c (see footnote 1), becomes nonlinear at high concentrations of the absorbing substance. Suppose you prepare a solution with a very high absorbance that is suspect in not following the linear relationship. How might you still use the sample for your analysis rather than discarding the sample and the data? Since the absorbance is very high, we have to dilute the solution to a known concentration where the absorbance is within the limits....