RQ, Background Information

research question

'The Determination of an Equilibrium Constant by varying the concentration of KSCN in a coloured complex reaction'

Background information

The Reaction we are studying

When chemical substances react, the reaction typically does not go to completion. Rather, the system goes to some intermediate state in which the rates of the forward and reverse reactions are equal. In this state, the reactants and the products have concentrations which do not change with time. Such a system is said to be in chemical equilibrium. When in equilibrium at a particular temperature, a reaction mixture obeys the Law of Chemical Equilibrium, which imposes a condition on the concentrations of reactants and products. This condition is expressed in the equilibrium constant K_c for the reaction.

In this experiment we will study the equilibrium properties of the reaction between iron(III) ion and thiocyanate ion:

Fe³⁺ (aq) + SCN^– (aq) ⇆ FeSCN²⁺ (aq)

When solutions containing Fe³⁺ ion and thiocyanate ion are mixed, they react to some extent, forming the FeSCN²⁺ complex ion, which has a deep red color. As a result of the reaction, the equilibrium amounts of Fe³⁺ and SCN^– will be less than they would have been if no reaction had occurred; for every mole of FeSCN²⁺ formed, one mole of Fe³⁺ and one mole of SCN^– will react. The equilibrium constant expression K_c for Reaction 1 is:

[FeSCN²⁺]

[Fe³⁺][SCN-] ^{= K}c

The value of K_c is constant at a given temperature. This means that mixtures containing Fe³⁺ and SCN^– will come to equilibrium with the same value of K_c, no matter what initial amounts of Fe³⁺ and SCN^– were used. Our purpose in this experiment will be to find K_c for this reaction for several mixtures that have been made up in different ways, and to show that K_c indeed has the same value in each of the mixtures. The reaction is a particularly good one to study because Kc is of a convenient magnitude and the color of FeSCN²⁺ ion makes for an easy analysis of the equilibrium mixture.

The mixtures will be prepared by mixing solutions containing known concentrations of iron (III) nitrate, Fe(NO₃)₃, and potassium thiocyanate, KSCN. The color of the FeSCN²⁺ ion formed will allow us to determine its equilibrium concentration. Knowing the initial composition of a mixture and the equilibrium concentration of FeSCN²⁺, we can calculate the equilibrium concentrations of the rest of the pertinent species and then determine K_c.

Colorimetry and Beers Law - including Determining [FeSCN²⁺]

In this experiment, you will use a spectrophotometer to determine [FeSCN²⁺] in the equilibrium mixtures. Instructions for use of the spectrophotometer are attached. The spectrophotometer measures absorbance, the amount of light absorbed by the complex at a given wavelength. Beer’s law expresses the relationship between absorbance, A, and concentration of a colored species, c.

A = εbc

Here, ε is a constant that depends on the wavelength of light and on the substance that is absorbing the light; b is the distance that the light travels through the sample of the absorbing substance.

The FeSCN²⁺ complex absorbs blue light. That is why it has a reddish orange color. It absorbs the most light at a wavelength of 447 nm. Therefore, at this wavelength absorbance measurements will have the highest sensitivity to [FeSCN²⁺]. In this experiment, you will measure the absorbance of all solutions at 447 nm.

In the first part of the experiment, you will determine the relationship between the absorbance and [FeSCN²⁺] at 447 nm. You will do this by measuring the absorbance of three standard solutions, in which [FeSCN²⁺] is known. There is a problem here: how can known concentrations of FeSCN²⁺ be obtained? FeSCN²⁺ participates in the equilibrium with Fe³⁺ and SCN^– ions. Known amounts of the reactants will not necessarily yield a known amount of the product.

This difficulty can be avoided. According to Le Châtelier’s principle, a net reaction from left to right (that is, in the forward direction) occurs when more of a reactant is added. As more and more of the same reactant is added, more and more of the product will be formed. It is possible to add so much of this reactant that essentially all of the other reactant will be converted to product. You will use limiting quantities of SCN^– and overwhelming amounts of Fe³⁺ to produce known amounts of FeSCN²⁺ in your standard solutions. The amount of FeSCN²⁺ that is formed will then be essentially equal to the starting amount of the limiting reactant.

Once you have measured the absorbances of the standard solutions, you will plot the absorbances against the concentrations of FeSCN²⁺ on a graph, or “calibration curve”. The points on the calibration curve will fall on a straight line, which has a slope of εb, the constant of proportionality between absorbance and concentration. You can use this calibration curve to find [FeSCN²⁺] in other solutions.

In the second part of the experiment, you will measure the absorbance of a different set of solutions, in which substantial amounts of both reactants and the product are present. You will use your calibration curve to convert the measured absorbance to the equilibrium concentration of FeSCN²⁺ in each solution. From the initial concentrations of the reactants and the equilibrium concentration of the product, you will calculate the equilibrium constant for the reaction.

In preparing each of the mixtures in this experiment you will maintain the concentration of H⁺ ion at 0.5 M. The hydrogen ion does not participate directly in the reaction, but its presence is necessary to avoid the formation of brown-colored species such as FeOH²⁺, which would interfere with the analysis of [FeSCN²⁺].

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