Equilibrium Reaction Lab

The purpose of the experiment is to determine the equilibrium constant of the reaction between Fe3+ ions and SCN2+ ions. In solutions one to five (the reference solutions), a large volume of concentrated Fe3+ ions was used in order to react with and convert all of the SCN2+ ions into FeSCN2+ ions. In solutions six to ten (the test solutions), a constant volume of Fe3+--in lesser concentration--and varying volumes of SCN2+ ions were diluted with DI water and reacted to produce an unknown amount of FeSCN2+ ions in an equilibrium reaction. The absorbance of each solution was found using a spectrophotometer, the first five of which was used to create a calibration curve, whose equation relates the concentration and absorbance of FeSCN2+ ions.
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Although the reaction continued, the concentration of both reactants--Fe3+ ions and SCN- ions--and the concentration of the product--FeSCN2+ ions--remained the same. A spectrophotometer was used to find the absorbances of the solutions by shining a beam of light through the solution and measuring how much light was absorbed at a specific wavelength. As the concentration of SCN- ions increased in each of the reference reactions, the concentration of FeSCN2+ ions increased, turning the solution a blood-red color and increasing the absorbance of the solution. The relationship between absorbance and concentration of FeSCN2+ ions is explained by Beer’s Law, which states that as the concentration increases, the absorbance increases. The law is further demonstrated by the calibration curve of the FeSCN2+ ions. Using the calibration curve and the absorbance values for the reference solutions, the concentrations at equilibrium and as a result, the equilibrium constant can be found. The equilibrium constant for this particular experiment is a value representing the ratio of the concentration of FeSCN2+ ions at equilibrium to the concentration of Fe3+ ions and SCN- ions at