Submission Date: March 1, 2019Section (please circle): M T W Th FRoom (please circle): 204 206Spectrophotometric Determination of The Equilibrium Constant for the Reaction Between Iron (III) and ThiocyanateToni-Ann Brown (100149242), Zhongqi Xiang AbstractTo find the equilibrium constant for the reaction between iron (III) and thiocyanate, absorbances of different equilibrium systems of FeSCN2+(aq) were measured by a spectrophotometer set at wavelength 470 nm. These absorbances were used to calculate the concentration of FeSCN2+(aq) at equilibrium, using Beer’s Law, after which an ICE table was used to calculate the equilibrium concentration of reactants and products and to obtain the equilibrium constant.

The equilibrium constant was found to be 124 at 22.7″. Inaccuracy of results may occur from the inability to maintain the solutions at a constant temperature and leftover deionized water droplets from rinsing causing dilution errors.PurposeThe purpose of this experiment is to determine the equilibrium constant, Kc , for the reaction between iron (III) and thiocyanate ions (Fe3+(aq) + SCN-(aq) ) by finding the absorbances of different concentrations of FeSCN2+ using a spectrophotometer.

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IntroductionThe equilibrium constant is an expression that originates from the combination of concentration of substances in solution that react and are allowed to reach their equilibrium point. The magnitude of the equilibrium constant acts as an estimate of the extent to which a reaction will proceed and the concentration of products and reactants when the reaction reaches equilibrium. This means that if K is a large number, the equilibrium concentration of products will be larger than reactants and if K is a small number, then the equilibrium concentration of the reactants is larger than the products. In the experiment, Beer’s Law and spectrophotometry were used to determine the equilibrium constant. The Beer’s Law states that the absorbance of a solution is proportional to its concentration as indicated by the equation: Absorbance = e L c. This suggests that the ratio of the absorbances is proportional to the ratio of concentration of solutions at the same wavelength. So therefore, the equation for Beers Law can be used to derive an alternative equation: Concentration1 / Concentration2 = Absorbance1 / Absorbance2. This equation will be used to determine the Concentration of FeSCN2+ at equilibrium in format: [FeSCN2+]eq=Aeq/Astd * [FeSCN2+]std. These absorbance values are obtained by using a spectrophotometer. The spectrophotometer indirectly measures the concentration of the orange-red colored iron (III) thiocyanate complex by measuring its absorbance in each solution. Concentrations of reactants at equilibrium can be calculated by subtracting the equilibrium concentration of the product from the initial concentration of the reactants using an Initial-Change-Equilibrium (ICE) table. Materials and MethodMaterials: burette, distilled water, test tubes, graduated cylinder, 0.002 M Fe(NO3)3, 0.002 M KSCN, 0.200 M Fe (NO3)3, spectrophotometer, plastic cuvette, retort stand and clamp, thermometer, beakersProcedureFour burettes were set up with 50 mL of solutions of 0.002 M Fe(NO3)3, 0.002 M KSCN, 0.200 M Fe(NO3)3 and distilled water. Each burette was labelled respectively. Afterwards, four clean, dry test tubes were obtained, labelled 1-4 and used to collect amounts of solution and distilled water as suggested by Table 1.Test Tube Number Fe(NO3)3 (mL)(0.002 M) KSCN (mL) (0.002 M) H2O (mL)1 5.00 2.00 3.002 5.00 3.00 2.003 5.00 4.00 1.004 5.00 5.00 0.00 Table 1. Volumes of solutions and distilled water delivered from each burette into each test tube.Each test tube with solution was tapped repetitively for solutions to be mixed thoroughly. The temperature of the solution in Test tube 1 was measured and recorded to be used for the equilibrium constant. Another solution of FeSCN2+was prepared to be used as the standard solution. 9.00 mL of 0.200 M Fe(NO3)3 alongside 1.00 mL 0.002 M KSCN was delivered into another test tube which was labelled as test tube 5 and stirred thoroughly (the standard solution). Following this, the spectrophotometer was calibrated and set at a wavelength of 470 nm. A blank was prepared by filling a cuvette ѕ of its full holding capacity with distilled water. This blank cuvette was placed in its respective slot in the spectrophotometer which was then closed. Set ref was selected and a reading of 0 on the monitor appeared. The cuvette was then removed and emptied and rinsed twice with the solution from test tube 1 and then filled ѕ with the solution. The cuvette was wiped with soft tissue and placed in the slot in the spectrophotometer. After closing the lid, the absorbance was observed and recorded. The cuvette was once more removed, emptied and rinsed with solution from test tube 2. It was filled with solution 2 and the absorbance for this was measured and recorded as well. Absorbances for test tubes 3,4 and5 were measured in a similar fashion. ResultsObservationsUpon mixing, the solutions in all test tubes changed from colorless to orange-red. DataAbsorbanceTrial 10.189 Trial 20.281 Trial 30.378 Trial 40.461Absorbance of standard (Trial 5)0.911Temperature22.7°C Table 2. Absorbances and temperature for each trial Kc expressionKc= [FeSCN2+]/([Fe3+] [SCN-])[Fe3+]i /M1.00*10-3 1.00*10-31.00*10-31.00*10-3[SCN-]i /M4.00*10-4 6.00*10-4 8.00*10-4 1.00*10-3[FeSCN2+]eq /M4.15*10-5 6.17*10-5 8.30*10-5 1.01*10-4[Fe3+]eq /M9.59*10-4 9.38*10-4 9.17*10-4 8.99*10-4[SCN-]eq /M3.59*10-4 5.38*10-4 7.17*10-4 8.99*10-4Kc value121122 126 125 Average of Kc values Kc= 124 at 22.7″Table 3. concentration of each ion and Kc values at different trialsCalculationsTrial 1: Initial concentration of Fe3+Initial concentration of SCN-Concentration of [FeSCN2+] eq using absorbance values at equilibrium and standardization Concentration of Fe3+ at equilibrium Concentration of SCN- at equilibrium Equilibrium constant, KcAverage Kc value[Fe3+] (M) [SCN-] (M) [FeSCN2+] (M)Initial 1.00*10-3 4.00*10-4 0.00Change -4.15*10-5 -4.15*10-5 +4.15*10-5Equilibrium 9.59*10-4 3.59*10-4 4.15*10-5Table 4. ICE table for equilibrium concentrations for Test Tube 1. DiscussionThe Fe3+ ion that contributes to the red orange colour of FeSCN2+, exists as a hydrated octahedral complex at equilibrium in aqueous solution. Fe (H2O)63+(aq) + SCN-(aq) ” Fe (H2O)5SCN2+(aq)A complex is a substance formed from the combination of ions or molecules that are attached to and surround a central metal ion. In this case, the complexes are Fe(H2O)63+(aq) and Fe(H2O)5SCN2+(aq) with Fe3+ being the metal centre. The molecules or ions that encircle the central metal ion are referred to as ligands. In this case, H2O and SCN- would be considered as the ligands.Figure 1. 3D structure of Fe(H2O)63+(aq) and Fe(H2O)5SCN2+(aq). In the reaction, the Fe3+ ions were orange-red and SCN- ions were colorless while the FeSCN2+ ions were red-orange. The concentration of FeSCN2+ complex ions formed was proportional to the intensity of the red color which was proportional to absorbance. That is, as the concentrations increased from trial to trial, so did the absorbances. Despite an increase of concentration, the equilibrium constant remained in the same limit between 121 and 126. Therefore, increasing the equilibrium concentration at constant temperature does not affect the equilibrium constant. The average Kc was found to be 124 at 22.7″. There were many possible errors that could have affected the accuracy of this experiment. If the temperature at which the reaction occurred could not be kept constant throughout all the trials, then the equilibrium constant may be affected, since equilibrium constant is dependent on temperature. In addition to this, when the cuvettes were rinsed with distilled water after each trial, water droplets may have been left over in the cuvette, causing an error in dilution which affects the absorbance by lowering it and as a result, lowering the Kc too.Conclusion It can be concluded that the equilibrium constant of the chemical reaction between iron (III) and thiocyanate is 124 at 22.7″.ReferencesClark, J. (2015, July). Equilibrium Constants: Kc. Retrieved February 26, 2019, from J. (2014, November). An Introduction to Complex Metal Ions. Retrieved February 26, 2019, from A. (2018). Chemical Equilibrium: Finding a Constant, Kc. In Chemistry 1020 Lab Manual (pp. 45-50).Colorimetric Analysis. (n.d.). Retrieved February 26, 2019, from equilibrium constant K. (n.d.). Retrieved February 26, 2019, from

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