Advanced Chemistry with Vernier ©Vernier Software & Technology 10 - 1 LabQuest 10 The Determination of an Equilibrium Constant The equilibrium state of a chemical reaction can be characterized by quantitatively defining its equilibrium constant, K eq . In this experiment, you will determine the value of K eq for the reaction between iron (III) ions and thiocyanate ions, SCN – . Fe 3+ (aq) + SCN – (aq) ↔ FeSCN 2+ (aq) When you mix amounts of Fe 3+ and SCN – , a reaction occurs to produce FeSCN 2+ , but not all of the reactants react. Thus, your beaker (or flask or cauldron) will contain some of each of these three species, which is your equilibrium system. To learn more about the system, we need to figure out a way to count the number of different ions in the reaction mixture. That is the major objective of this experiment, and to achieve this objective you will take advantage of something about FeSCN 2+ – in aqueous solution it has a reddish color. The two reactants, Fe 3+ and SCN – , are essentially colorless in solution, thus the red color you will see when you conduct the reaction is produced by the FeSCN 2+ ions. One of the more important numbers that help us understand an equilibrium system is called the equilibrium constant, K eq . For the reaction between Fe 3+ and SCN – , the K eq is defined by the equation shown below. ] SCN ][ Fe [ ] FeSCN [ 3 2 eq K To find the value of K eq at a given temperature, it is necessary to determine the molar concentration of each of the three species in solution at equilibrium. You will determine the concentrations by using a Vernier Colorimeter or Spectrometer to measure the amount of light of a specific wavelength that passes through a sample of the equilibrium mixtures. The amount of light absorbed by a colored solution is proportional to its concentration. The red FeSCN 2+ solution absorbs blue light, thus the Colorimeter users will be instructed to use the 470 nm (blue) LED. Spectrometer users will determine an appropriate wavelength based on the absorbance spectrum of the solution. The wavelength will be close to, but not exactly, 470 nm. In order to successfully evaluate this equilibrium system, it is necessary to conduct two separate tests. In Part I of the experiment, you will prepare a series of standard solutions of FeSCN 2+ from solutions of varying concentrations of SCN – and constant concentrations of H + and Fe 3+ that are in stoichiometric excess. The excess of H + ions will ensure that Fe 3+ engages in no side reactions (to form FeOH 2+ , for example) which could interfere with your measurements. In an excess of Fe 3+ ions, the SCN – ions will be the limiting reagent, thus all of the SCN – will form FeSCN 2+ ions. The FeSCN 2+ complex forms slowly, taking at least one minute for the color to develop. It is best to take absorbance readings after a specific length of time has passed, between two and four minutes after preparing the equilibrium mixture. Do not wait much longer than five minutes to take readings, however, because the mixture is light sensitive and the FeSCN 2+ ions will slowly decompose.
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