Sarah Kovach
May 19, 2019
Dr. Harvey
Chemistry 102 lab
Title: Equilibria: Measuring on Equilibrium Constant using Spectrophotometry
Purpose:
The purpose of this experiment is to observe the concentration of Fe3+ and SCN- and
measure the equilibrium constant. Equilibrium constant can help us calculate how much of the
substance is needed to find equilibrium. A large k value will have the products favored in the
reaction. A small k will have reactants favored. A saturated solution will not shift to the right
regardless of how much reactant you apply to the solution. The equilibrium can be calculated if
we know the concentration amounts of products and reactants. Kp is used for partial pressure of
gases to find the ratio of the constant. The general reaction is aA+bB=cC+dD with
Kc=[C]c[D]d/[A]a[B]b or substitute partial pressure is such as p[A]a for [A]a. The brackets
indicate molarity and the p stands for partial pressure. Iron(III) ion reacts with Thiocyanate ion to
produce Thiocyanatoiron (III). The reaction is represented by the following equation: Fe3+ +
SCN– <—-> Fe(SCN)2+. The reading from the spectrophotometer, will give a value that needs
the conversion factor from the equations ratio to calculate the absorbance of Fe(SCN)2+ at
different concentrations. Part 1 determined the limiting reactant in the solution to calculate
Fe(SCN)2+. From these values absorbance and Fe(SCN)2+ was used to create a graph that was
increasing with absorbance. The absorbed value and linear treadling calculation helped us
calculate the molarity for each Fe(SCN)2+ solution. The concentrations for Fe3+ and SCN- were
calculated from taking the molarity divided by the total volume. The initial concentration was
subtracted from the equilibrium of Fe(SCN)2+ to create new values for the concentration at
equilibrium for [Fe3+] and [SCN-]. Lastly, Kc was calculated from the equilibrium constants that
were inputted into the equilibrium expression.
Conclusion:
Water and solids are not included in the equilibrium constant because they do not affect
the reactant value at equilibrium. Pure solids and liquids are excluded from this expression
because they concentrations remain constant in the reaction. The Kc in this particular experiment
favored products because the number is much larger than 1. B-1 through B-5 are different values
because they are calculated from different solution concentrations. The results varied a lot from
each other. These values explain how the reaction is moving forward and shows how fast the
reaction will occur to reach equilibrium. Since A-1 through A-7 had absorbance values and
concentration a Kc could be calculated from the reactions stoichiometry. A higher concentration
in the solution will have a higher absorption value with less photons making it through the
solution tube in the spectrophotometer. According to Le Chatelier's principle, if the reaction was
exothermic the Kc would decrease at higher temperature because the reactants are favored. The
Kc can tell us if the reaction is moving forward or backward and this value can be compared by
the other calculated values to tell how fast the solution is reacting. Calculated values became
greater than their previous value which also explains why the graph has a positive slope. As the
solution is more concentrated, the light absorbed becomes a greater value. My calculated values
were similar to a straight line, and the graph’s R^2 value was 9.97E-01 accuracy. My average Kc
value was 15.50. Error could occur from beakers that were not clean, diluting a solution too
much or too little, and not reading the burette and beakers accurately.