rChemistry 201 Experiment 1a Measurement of an Equilibrium Constant1 - Week Two Objective: To calculate the concentration of thiocyanate in human saliva using a previously calculated Kc. Background: Some chemical species can be detected based on their color. Since they absorb light in the visible spectrum we can track their concentration as a function of their absorbance or transmittance. This methodology will be applied to the iron – thiocyanate system. Fe3+(aq) + SCN-(aq)  Fe(SCN)2+(aq) The iron(III)thiocyanate complex is deep red unlike either of the other species in solution, so its appearance in a chemical reaction can be easily tracked. In the previous solution, the equilibrium constant of iron thiocyanate was calculated using a calibration curve several solutions with known concentrations. In this experiment, the same equilibrium constant will be used. The concentration of iron thiocyanate will be determined spectroscopically, and the amount of thiocyanate can be calculated.  Fe( SCN ) 2   Kc   3   Fe   SCN    

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Adapted from the Journal of Chemical Education Vol. 76 No. 9 September 1999 (Lahti, Vilpo, Hovinen).

Procedure: 1. Perform a qualitative determination of the amount of thiocyanate in saliva. Collect some saliva in a beaker, and transfer a few drops of stock iron nitrate (0.2M) solution in nitric acid. Observe any color changes. Set aside or discard. 2. Obtain enough 0.200M Fe(NO3)3 and 2.00x10-3M KSCN to make each of the solutions in the table below. (Note: Both species are dissolved in 0.1M HNO3) Prepare each of the solutions as shown in a test tube, noting the final volume of the solution. Table 1 Calibration Solution # C1 – Blank C2 C3 C4 C5 0.200M Fe(NO3)3 in (0.1 M HNO3) (mL) 5.00 5.00 5.00 5.00 5.00 2.00x10-3M KSCN in (0.1 M HNO3) (mL) 0.00 1.00 2.00 3.00 4.00 0.1M HNO3

5.00 4.00 3.00 2.00 1.00

3. Make all solutions before getting any absorbance readings. Determine the absorbance (near 460nm) of each solution