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Determination of Kc for a Complex Ion Formation
• • Find the value of the equilibrium constant for formation of FeSCN2+ by using the visible light absorption of the complex ion. Confirm the stoichiometry of the reaction.
In the study of chemical reactions, chemistry students first study reactions that go to completion. Inherent in these familiar problems—such as calculation of theoretical yield, limiting reactant, and percent yield—is the assumption that the reaction can consume all of one or more reactants to produce products. In fact, most reactions do not behave this way. Instead, reactions reach a state where, after mixing the reactants, a stable mixture of reactants and products is produced. This mixture is called the equilibrium state; at this point, chemical reaction occurs in both directions at equal rates. Therefore, once the equilibrium state has been reached, no further change occurs in the concentrations of reactants and products. The equilibrium constant, K, is used to quantify the equilibrium state. The expression for the equilibrium constant for a reaction is determined by examining the balanced chemical equation. For a reaction involving aqueous reactants and products, the equilibrium constant is expressed as a ratio between reactant and product concentrations, where each term is raised to the power of its reaction coefficient: aA (aq) + bB (aq) ! cC (aq) + dD (aq)
[C] [D] [A] a [B] b
When an equilibrium constant is expressed in terms of molar concentrations, the equilibrium constant is referred to as Kc. The value of this constant at equilibrium is always the same at a ! given temperature, regardless of the initial reaction concentrations. Whether the reactants are mixed in their exact stoichiometric ratios or one reactant is initially present in large excess, the ratio described by the equilibrium constant expression will be achieved once the reaction composition stops changing. We will be studying the reaction that forms the reddish-orange iron(III) thiocyanate complex ion, Fe(H2O)5 SCN2+(aq). The chemical equation describing the formation of this complex ion is Fe(H2O)63+(aq) + SCN–(aq) ! Fe(H2O)5SCN2+(aq) + H2O(l) (2)
Notice that the reaction involves the displacement of a water ligand by a thiocyanate ligand, SCN–. However, for simplicity, and because water ligands do not change the net charge of the species, water can be omitted from the formulas of Fe(H2O)63+(aq) and Fe(H2O)5SCN2+(aq) in our chemical equation. Thus, the chemical formula Fe(H2O)63+(aq) may be simplified as Fe3+(aq) and Fe(H2O)5SCN2+(aq) as FeSCN2+(aq). Making these changes, we can rewrite the simplified chemical equation for the reaction and its corresponding equilibrium constants expression as:
Determination of Kc
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Fe (aq) + SCN (aq) ! FeSCN (aq)
3+ – 2+
[FeSCN ] = [Fe ][SCN ]
In this experiment you will create several different aqueous mixtures of Fe3+ and SCN–. Because this reaction reaches equilibrium nearly instantly, these mixtures turn reddish-orange very quickly due to the formation of the product ! FeSCN2+(aq). The intensity of the color of the mixtures is proportional to the concentration of product formed at equilibrium. As long as all mixtures are measured at the same temperature, the ratio described in Equation (3) will be the same because the value of Kc is a constant. Measurement of [FeSCN2+]: Because the complex ion product is the only strongly colored species in the system, its concentration can be determined by measuring the intensity of the orange color in equilibrium systems of these ions. This method relies on Beer's Law, which is given by the equation:
A = "!c
where A is the absorbance of light by the solution; c is the molar concentration of the solution; ! is the path length, or the distance that the light...