The purpose of this experiment was to determine the equilibrium constant of the reaction between iron (III) chloride (FeCl3) and potassium thiocyanate (KSCN) through the use of spectrophotometry.

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Spectrophotometric Determination of the Equilibrium Constant of a Reaction

Angelica Seribo Baluyut

Institute of Chemistry, University of the Philippines, Diliman, Quezon City 1101 Philippines

Department of Chemical Engineering, College of Engineering, University of the Philippines, Diliman, Quezon City 1101 Philippines

ABSTRACT

The purpose of this experiment was to determine the equilibrium constant of the reaction between iron (III) chloride (FeCl3) and potassium thiocyanate (KSCN) through the use of spectrophotometry. Five different standard solutions were prepared and their respective absorbance was measured with a spectrophotometer. From the results, a calibration curve was plotted and the equilibrium concentrations of the blood-red complex,   FeSCN2+ in the five unknown solutions were attained. The relationship between the equilibrium concentration and the absorbance was studied. Through further computations, the equilibrium concentrations of the iron (III) chloride (FeCl3) and potassium thiocyanate (KSCN) were calculated and were plugged in the equilibrium expression. An equilibrium constant of 254 was found. Compared to the literature value of 890[1], there was a 71.46% difference. Although a high percentage error was obtained, the result was enough to verify that spectrophotometry is an effective way of calculating the equilibrium constant of a certain reaction.

Introduction

        Chemical reactions that are reversible reach a state called chemical equilibrium. In this state, the concentrations of the products and reactants are constant or stable. When in equilibrium, there is no net change in the amounts of the products and reactants if not disturbed by any means. For this reason, the proportion of the concentrations of the products over the reactants raised to their respective coefficients is also constant. This proportionality is known as the equilibrium constant, Keq. For the reaction in this experiment,

Fe(aq)3++SCN(aq)-↔FeSCN(aq)2+          [1]

The equilibrium expression is,

     Keq= [FeSCN2+]eq[Fe3+]eq[SCN-]eq                             [2]                   

The product,  FeSCN2+, is a blood red complex that absorbs visible light. The main objective of this experiment was to know the equilibrium constant for the formation of  FeSCN2+. Spectrophotometry, therefore, may be used to determine the system’s equilibrium constant through the use of a machine called a spectrophotometer. The spectrophotometer measures the reflection or transmission of a material as a function of wavelength. In this experiment, a  UV-Vis spectrophotometer is used to measure the amount of light that the colored substance in the solution absorbed. With the data gathered, the concentration of the substance may be obtained through the use of the Beer-Lambert’s law which directly relates the concentration of the substance and the amount of light it absorbed. The relationship is expressed by the equation:

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                                A=εbc                            3

Where,

A = absorbance (from the spectrophotometer)

ε = molar absorptivity coefficient M-1cm-1

b = path length in cm

c = analyte molar concentration

Note: Deviations from Beer's Law can be caused by:[2]

(a)        Stray light: there should be no stray light that is outside the wavelength used.

(b)        Unequal light path lengths across the light beam.

(c)        Unequal absorber concentration across the light beam.

(d)        Changes in refractive index of the solution at high ...

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