The Spectrophotometric Determination of an Equilibrium Constant
The Spectrophotometric Determination of an Equilibrium Constant.
Abstract:
The report presents determination of equilibrium constant for the formation of a complex ion FeSCN2+. This was accomplished using a colorimeter to measure absorbance of some known concentration solutions in order to generate the calibration curve. The equation of the graph was used to compute the equilibrium concentrations of the reactants and products, needed to calculate the equilibrium constant for the reaction. I. Introduction:
The purpose of the experiment is to determine the equilibrium constant for the formation of a complex ion iron (III) thiocyanate (FeSCN2+). This ion is formed in the reaction of iron (III) ion with thiocyanate ion SCN-):
Eq.1 Fe3+ (aq) + SCN- (aq) FeSCN2+ (aq,) The equilibrium constant is the ratio between the reactants and products when the forward and reverse reaction rates are equal. It is calculated from the mass action expression:
Eq. 2 Kc = [FeSCN2+] / [Fe3+][SCN-] In order to calculate the equilibrium constant it is necessary to know the concentrations of all ions at equilibrium. The concentration of FeSCN2+ ions can be determined colorimetrically as the FeSCN2+ ions are red colored (the reactant ions are colorless), and therefore they are the primary absorber in the mixture. According to Beer’s law the higher the concentration of the FeSCN2+ ion in the solution the more intense the red color (The Columbia Encyclopedia, 6th ed., 2008). By pushing the reaction in equation 1 to completion using LeChatelier’s principle with different volumes of the reactants, the Beer’s curve of absorbance versus concentration can be generated and used to determine the concentration of FeSCN2+ in an equilibrium mixture. If the initial concentrations of the reactant ions are known, their equilibrium concentrations can be calculated using the ICE table, and then the equilibrium constant can be calculated (Kotz,
Abstract:
The report presents determination of equilibrium constant for the formation of a complex ion FeSCN2+. This was accomplished using a colorimeter to measure absorbance of some known concentration solutions in order to generate the calibration curve. The equation of the graph was used to compute the equilibrium concentrations of the reactants and products, needed to calculate the equilibrium constant for the reaction. I. Introduction:
The purpose of the experiment is to determine the equilibrium constant for the formation of a complex ion iron (III) thiocyanate (FeSCN2+). This ion is formed in the reaction of iron (III) ion with thiocyanate ion SCN-):
Eq.1 Fe3+ (aq) + SCN- (aq) FeSCN2+ (aq,) The equilibrium constant is the ratio between the reactants and products when the forward and reverse reaction rates are equal. It is calculated from the mass action expression:
Eq. 2 Kc = [FeSCN2+] / [Fe3+][SCN-] In order to calculate the equilibrium constant it is necessary to know the concentrations of all ions at equilibrium. The concentration of FeSCN2+ ions can be determined colorimetrically as the FeSCN2+ ions are red colored (the reactant ions are colorless), and therefore they are the primary absorber in the mixture. According to Beer’s law the higher the concentration of the FeSCN2+ ion in the solution the more intense the red color (The Columbia Encyclopedia, 6th ed., 2008). By pushing the reaction in equation 1 to completion using LeChatelier’s principle with different volumes of the reactants, the Beer’s curve of absorbance versus concentration can be generated and used to determine the concentration of FeSCN2+ in an equilibrium mixture. If the initial concentrations of the reactant ions are known, their equilibrium concentrations can be calculated using the ICE table, and then the equilibrium constant can be calculated (Kotz,
References: Kotz, J.C., Townsend, J.R., Treichel, P. (2008). Chemistry and chemical reactivity (7th ed.). Florence, KY: Brooks Cole. The Columbia Encyclopedia, Sixth Edition. Beer’s Law. (2008). Retrieved October 26, 2010 from Encyclopedia.com: http://www.encyclopedia.com/doc/1E1-Beerslaw.html