Finding the Equilibrium Constant - Lab Report Example

Finding the equilibrium constant, Kc, of the formation of thiocyanoiron The aim of the experiment was to determine the equilibrium coefficient (Kc) of the formation of thiocyanoiron. Using absorbance measurements of FeSCN2+ and initial concentrations of the reagents, the Kc value was found to be almost constant (134-137) for reactant concentrations not in a ratio of 1:1. A Fe3+: SCN- ratio of 1:1 resulted in a Kc value of 1.0019. The standard deviation was found to be ±58.322. Introduction The equilibrium constant of any reaction will always have the same value and at any value of molar coefficient, as long as the temperature is held constant. Using this idea, this experiment will find the equilibrium constant in the reaction that forms thiocyanoiron. This formation occurs as follows: The Kc is given by: After equilibrium if established, a colorimeter is used to determine the absorbance (and by extension the concentration) of the thiocyanoiron complex formed. This is then used to determine the concentration of the Fe3+ and SCN-. The concentrations obtained are then used in the Kc equation above to find the equilibrium constant. The colorimeter makes use of Beer’s law (Abs = εlc) to determine the concentration1. The molar absorptivity of FeSCN2+ will be obtained by making use of a standard solution containing Fe3+ ions in excess, and by employing Le Châteliers principle to describe the shift of the reaction to the right (thus promoting the formation of the product), and in the process utilizing the all of the available SCN- ions2. As a result, the [FeSCN2+] in the equilibrium is equivalent to the amount of SCN- that was initially available3. Because the mole ratio is 1:1, and also because the initial concentrations are known, the concentration of the unreacted reactant can be found be finding the difference between initial concentration of the reactant and the equilibrium concentration of the product4. Methods 5.00 mL of 0.00200 M Fe(NO3)3 was pipetted into four different labeled test tubes. The volumes were recorded. 25 mL of 0.00200 M KSCN was poured into a dry 100 mL beaker and then 2.00, 3.00, 4.00 and 5.00 mL pipetted into the test tubes labeled initially. The volume used was recorded. 25 mL of distilled water was placed in a 100 mL beaker and then 3.00, 2.00, 1.00 and 0.00 mL of this water pipetted into the four labeled test tubes. The actual volumes of the test tubes were recorded. Each solution was mixed thoroughly with a plastic dropper. The temperature of one of the four solutions was measured and recorded; this was to be used in measuring of the equilibrium constant, Kc­. A standard solution of FeSCN2+ was prepared by pipetting 18 mL of 0.200 M Fe(NO3)3 into a 25 x 150 mm test tube labeled “5”. 2.00 mL of 0.00200 M KSCN was pipetted into the same test tube. The mixture was then thoroughly stirred with a plastic dropper. Colorimeter #10 Vernier Software was connected to the computer interface and then set at 470 nm and then let to warm up for 5 minutes. A blank was prepared by filling a cuvette ¾ full with water. The colorimeter was then calibrated. The cuvette was then emptied of the water and then rinsed twice with 2 mL portions of the test tube 5 standard solution. The cuvette was then filled ¾ full with the solution from test tube 5. The outside of the cuvette was wiped with a tissue and it was then placed in the colorimeter. The absorbance reading of the solution was then measured and recorded. The cuvette’s contents were discarded and the cuvette rinsed twice with the test tube 1 solution and then filled ¾ full. The procedure above was followed so as to measure and record the absorbance of the solution. This was repeated using the solutions in the test tubes 2, 3 and 4. The absorbance values obtained were recorded. Results Test tube # Vol. of Fe(NO3)3 (mL) Vol. of KSCN (mL) Vol. of distilled water (mL) Total volume (mL) Measured absorbance 1 5.00 2.00 3.00 10.00 0.157 2 5.00 3.00 2.00 10.00 0.234 3 5.00 4.00 1.00 10.00 0.304 4 5.00 5.00 0.00 10.00 0.370 5 18.00 2.00 0.00 20.00 0.690 Table 1: Table of the volumes of the reagents used, distilled water added and the colorimeter absorbance values of FeSCN2+ To find [FeSCN2+]1 = (Abs1/Absstd)[FeSCN2+]std = y = 0.1800 mol Fe3+ Therefore, [FeSCN2+]std = 4.00x10-6 mol/ 0.00200 = 2.00x10-4 mol SCN- The [FeSCN2+]eq is given by: Solution #1: [FeSCN2+]1 = (Abs1/Absstd) x [FeSCN2+]std = (0.157/0.690)x2.00x10-4M = 4.55x10-5 M at equilibrium Solution #2: [FeSCN2+]2 = (Abs2/Absstd) x [FeSCN2+]std = (0.234/0.690)x2.00x10-4M = 6.78x10-5 M at equilibrium Solution #3: [FeSCN2+]3 = (Abs3/Absstd) x [FeSCN2+]std = (0.304/0.690)x2.00x10-4M = 8.81x10-5 M at equilibrium Solution #4: [FeSCN2+]4 = (Abs1/Absstd) x [FeSCN2+]std = (0.690/0.690)x2.00x10-4M = 1 M at equilibrium The initial concentration of solution #1: The initial concentration of solution #2: The initial concentration of solution #3: The initial concentration of solution #4: To calculate Kc values: Solution # Fe3+ + SCN- FeSCN2+ 1 Initial conc. 1.00x10-3 4.00 x10-4 0.00 Change in conc. -x -x +x Equilibrium conc. 1.00x10-3-x 4.00 x10-4-x 4.55 x10-5 2 Initial conc. 1.00x10-3 6.00 x10-4 0.00 Change in conc. -x -x +x Equilibrium conc. 1.00x10-3-x 6.00 x10-4-x 6.78 x10-5 3 Initial conc. 1.00x10-3 8.00 x10-4 0.00 Change in conc. -x -x +x Equilibrium conc. 1.00x10-3-x 8.00 x10-4-x 8.81 x10-5 4 Initial conc. 1.00x10-3 1.00 x10-3 0.00 Change in conc. -x -x +x Equilibrium conc. 1.00x10-3-x 1.00 x10-3-x 1.00 Table 2: Table of the ICE values of the Fe3+,SCN- and FeSCN2+ contained in the labeled test tubes 1, 2, 3, 4. For solution #1: [Fe3+]1 = 1.00x10-3 – 4.55 x10-5 = 9.545 x10-4 [SCN-]1 = 4.00 x10-3 - 4.55 x10-5 = 3.545 x10-4 = 134.46 =134 For solution #3: [Fe3+]2 = 1.00x10-3 – 6.78 x10-5 = 9.322 x10-4 [SCN-]2 = 6.00 x10-3 – 6.78 x10-5 = 5.322 x10-4 = 136.66 =137 For solution #3: [Fe3+]3 = 1.00x10-3 – 8.81 x10-5 = 9.119 x10-4 [SCN-]3 = 8.00 x10-3 – 8.81 x10-5 = 7.119 x10-4 = 135.70 =136 For solution #4: [Fe3+]4 = 1.00x10-3 – 1.00 = -0.999 [SCN-]4 = 1.00 x10-3 – 1.00 = -0.999 = 1.0019 Average = (137+136+134+1)/4 = 102 Variance = (352+342+322+ (-1012)) = 13606/4 = 3401.5 Standard deviation = = ±58.322 Discussion The experiment aimed to calculate the equilibrium constant for the reaction between thiocyanate (SCN-) and Fe3+ ions to form thiocyanoiron (FeSCN2+). Combination of the two ions results in the establishment of an equilibrium. By measuring the absorbances of these three ions, starting with different initial concentrations it is possible to determine the equilibrium constants. If an equilibrium is formed when the SCN- and Fe3+ ions react, then the constants obtained should be almost similar. From the results obtained, it is evident that as the concentration of KSCN added to the Fe(NO3)3 increased there was an increase in the Kc value obtained from 134 to 137. This value then dropped to 136 and then to 1.0019. The rise in the Kc value was from the fact that the increased KSCN volume provided more SCN- ions to react with the Fe3+ ions. This led to the concentration of the product (FeSCN2+) being higher than the reactants, causing the equilibrium of the reaction to lie to the right side of the equation5. The equilibrium constant then reduced to 136 because addition of more KSCN caused the Fe(NO3)3 to become the limiting reagent. The relatively high values of the equilibrium constants, Kc, for the different concentrations of solutions 1-3 mean that the reactions went into completion. Solution 4 resulted in an equilibrium constant of 1.0019. This value is close to 1 and it indicates that equilibrium was attained at an intermediate point, where the concentrations of both the reactants and products were the same6. Conclusion This experiment successfully determined the value of the equilibrium constant established during the formation of thiocyanoiron. When the SCN- concentrations were lower than that of Fe3+ but increasing, the constants were at least within ±2 of each other (134,137,136). At a reactant ratio of 1:1, the constant was 1.0019. This is slightly different from the ideal value of 1.0000. The slight discrepancy could have been caused by experimental errors during the transfer of the reagents to the labeled test tubes. This can be avoided in future experiments by ensuring that all transfers of reagents involving pipettes are carried out with caution. Read More