Determination of an Equilibrium Constant
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.
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 Fe3+ and HSCN. Eventually, all the forward and reverse rates become equal and the concentrations stop changing even though the forward and reverse reactions are still proceeding. Every equilibrium reaction for a certain species has the same equilibrium constant. Once a constant is determined, it stays the same no matter what concentrations of the species are used. The equilibrium constant does however depend on temperature. LeChatlier’s principle states that a change in concentration, temperature, volume, or partial pressure, shifts the equilibrium to counteract the change and a new equilibrium is established2. This is used in the standard solutions. A large amount of Fe(NO3)3 is used to drive the reaction towards the products and to completion. A Spec 20 spectrophotometer is used to measure the absorbance of each solution. This is done because a spectrophotometer works by introducing light of a specific wavelength to a sample. As the light passes through the sample, some of it is absorbed by the sample and some is transmitted through the sample. The spectrophotometer intercepts this transmitted light. By comparing this light to the incident beam, the spectrophotometer finds the percent of light transmitted. The percent transmittance is then converted by the spectrophotometer to absorbance with the equation: -log(%T/100) = Abs
Beer’s Law is also used in this experiment to determine the molar absorptivity constant. Beer's law states that the absorbance is directly proportional to the concentration of a solution. If you plot absorbance versus concentration, the resulting graph yields a straight line. The equation for the straight line can be used to determine molar absorptivity constant. In Beer’s Law, Abs = εl[X] where ε is the molar absorptivity constant, l is the path length and [X] is the concentration of the species. A plot of absorbance vs. concentration of FeSCN2+ in the standard solutions allow for the determination of ε. With the molar absorptivity found,...
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.
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