Determination of an Equilibrium Constant
Determination of an Equilibrium Constant
Abstract: In this experiment, two reactions were run to determine the molar absorptivity and the equilibrium constant of FeSCN2+. The main principles used in this lab are equilibrium, LeChatlier’s Principle, Beer’s Law and Spectrocopy. The first reaction was run to completion using LeChatier’s Principle and the second reaction was run to equilibrium. A spectrophotometer was used to measure absorbances. Using a graph of absorbance versus concentration of FeSCN2+ was used to determine that the molar absorptivity constant was 3670. Beer’s Law was used to determine that the average equilibrium constant was 33.1793.
Introduction:
The purpose of this experiment is to determine the value of the equilibrium constant for the reaction:
Fe3+(aq) + HSCN(aq) H+(aq) + FeSCN2+(aq)
In this reaction, iron(III) nitrate, Fe(NO3)3, is mixed with thiocyanic acid, HSCN, to produce the H+ ion and the complex ion thiocyanate iron(III) [FeSCN]2+. This reaction is done twice. The first time it is run to completion and the second time it is run to equilibrium.
The equation for the equilibrium constant, Keq, is given by:
Keq =
The initial concentrations of Fe3+ and HSCN and the equilibrium concentrations of FeSCN2+ will by measured. With all of these concentrations determined, the equilibrium concentrations of all species can be calculated. With the equilibrium concentrations of all the species found, the equilibrium constant can be determined. The major lab techniques used in this lab are equilibrium, LeChatlier’s Principle, spectroscopy, and Beer’s Law. When species react, the concentrations of the products and reactants continuously change until equilibrium is reached. No change of concentration occurs once equilibrium is reached. Equilibrium happens when a reaction is reversible. For the reaction studied in this lab, the double arrows mean that Fe3+ and HSCN can react to form H+ and FeSCN2+ and H+ and FeSCN2+ can react to form
Abstract: In this experiment, two reactions were run to determine the molar absorptivity and the equilibrium constant of FeSCN2+. The main principles used in this lab are equilibrium, LeChatlier’s Principle, Beer’s Law and Spectrocopy. The first reaction was run to completion using LeChatier’s Principle and the second reaction was run to equilibrium. A spectrophotometer was used to measure absorbances. Using a graph of absorbance versus concentration of FeSCN2+ was used to determine that the molar absorptivity constant was 3670. Beer’s Law was used to determine that the average equilibrium constant was 33.1793.
Introduction:
The purpose of this experiment is to determine the value of the equilibrium constant for the reaction:
Fe3+(aq) + HSCN(aq) H+(aq) + FeSCN2+(aq)
In this reaction, iron(III) nitrate, Fe(NO3)3, is mixed with thiocyanic acid, HSCN, to produce the H+ ion and the complex ion thiocyanate iron(III) [FeSCN]2+. This reaction is done twice. The first time it is run to completion and the second time it is run to equilibrium.
The equation for the equilibrium constant, Keq, is given by:
Keq =
The initial concentrations of Fe3+ and HSCN and the equilibrium concentrations of FeSCN2+ will by measured. With all of these concentrations determined, the equilibrium concentrations of all species can be calculated. With the equilibrium concentrations of all the species found, the equilibrium constant can be determined. The major lab techniques used in this lab are equilibrium, LeChatlier’s Principle, spectroscopy, and Beer’s Law. When species react, the concentrations of the products and reactants continuously change until equilibrium is reached. No change of concentration occurs once equilibrium is reached. Equilibrium happens when a reaction is reversible. For the reaction studied in this lab, the double arrows mean that Fe3+ and HSCN can react to form H+ and FeSCN2+ and H+ and FeSCN2+ can react to form
References: 1. Chemistry 203/205: An Introduction to Chemical Systems in the Laboratory. Hayden-McNeil; Plymouth, 2011. 2. Atkins, P.; Jones, L.; Chemical Principles. 5th ed., Freeman: New York, 2008.