Equations and Calculations

It is not necessary to go through all the calculations given in the lab manual. Instead, we will use a technique that will save you the intermediate steps. Instead of going around from moles and molarity, we will stick to the concentration of the reactants and products. Then we will use the famous, all-powerful table to determine the equilibrium constant. Run one will be used as an example so you can see what numbers you should be close to getting.

We will first find the initial concentrations of the reactants and products. This will be done with the following equation:

Initial Concentration = concentration x reactant volume used / total volume

Run 1: Initial conc. of SCN^- = 2.00x10^-3 M x 1.0 mL / 10.0 mL = 2.00x10^-4 M

Initial conc. of Fe³⁺ = 2.00x10^-3 M x 5.0 mL / 10.0 mL = 1.00x10^-3 M

Initial conc. of FeSCN²⁺ = 0

Next, we run the experiment. Then, using the Beer-Lambert law, you can find the final (equilibrium) concentration of FeSCN²⁺. Record that in the table.

Then, using stoichiometetry find the "change" row of the table.

Finally, use addition to find the final (equilibrium) concentration of the reactants.

With the equilibrium concentrations known, you can use the following equation to find K_eq:

K_eq = (conc. of product 1) / (conc. of reactant 1)(conc. of reactant 2)

K_eq = [FeSCN²⁺] / [Fe³⁺][SCN^-]

K_eq for Run 1 = (2.0x10^-5) / ((1.8x10^-4)(9.8x10^-4)) = 113.4

Your results will vary of course, but they should be somewhere around these numbers.

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