Experiment # 20
Nichole Morgan
Chemistry 106, Section 17
1
The purpose of this experiment was determine the equilibrium constant for the
reaction between the Ferris ion (Fe3+), and the Thiocyanate ion (SCN-), which produces
Ferric Thiocyanate ion [Fe(H2O)5SCN]2+. To do this a series of standard solutions were
created from two stock solutions of Fe(NO3)3 and NaSCN. From the standard solutions, the
absorbency was measured and a Beer’s Law plot was produced. The Beer’s Law plot also
revealed a best-fit line, slope, and y-intercept. For the second part of the experiment the
absorbency of five solutions was measured via a spectrophotometer, only this time
different values of the less concentrated Fe(NO3)3 was used as well as both NaSCN, and
HNO3. The absorbency values allowed us to calculate the equilibrium concentrations of Fe3+
and SCN- from the reaction stoichiometry. Then from the equilibrium concentrations we
were able to determine the equilibrium constant.
The equilibrium value for Ferric Thiocyanate is 113 (+-) 3, this information was
discovered by looking up the scholarly article Colorimetric Determination of the
Iron(III)−Thiocyanate Reaction Equilibrium Constant with Calibration and Equilibrium
Solutions Prepared in a Cuvette by Sequential Additions of One Reagent to the Other. This
information is of good quality because it stems from the American Chemical Society
website, where scholarly articles and journals related to the study of Chemistry are
published. As well, this particular information pertains to the experiment we conducted
and the value seems within range.
The value I calculated our equilibrium to be from the data we collected was 192.01,
while the literature value indicates a somewhat lower value at 110-116. The difference in
the calculated and literature value is about 82-88, not a significant difference, but much
lower than what was calculated and even lower than our average which was determined to
Experiment # 20
Nichole Morgan
Chemistry 106, Section 17
2
be 247.35. Based on the literature value alone, I believe our results to be inaccurate and I
primarily reduce this inaccuracy to not calibrating the spectrophotometer again after our
first round of collections. We did recalibrate it but only took the data for our first sample
which was significantly lower than the previously collected value, thus if we were to redo
our entire series of solutions, our equilibrium value would be much lower, more accurate,
and closer to the provided literature value.
As aforementioned the primary issue we dealt with during this experiment was the
spectrophotometer relaying inaccurate data because it was not calibrated between trials,
reducing the accuracy of our collected data. This issue was from human error and a mistake
we should have taken into consideration, although we calibrated the spectrophotometer at
the beginning of the experiment it has been proven that it is necessary to calibrate it
throughout the experiment so our data is accurate and precise. By not calibrating it, our
standard deviation was much higher than 20% and our absorbency values were about 200
points of their actual values. We noticed this and recalibrated the spectrophotometer and
quickly re-measured all of our solutions and got our standard deviation down to 19%,
however I still believe there were some issues with the machine and it’s ability to collect
accurate data.
The use of the blank for this experiment differed from experiment 19 in that for
experiment 19 DI water was used as the blank while for this experiment the blank solution
consisted of 25 mL of .21 Fe(NO3) and 75 mL of .1 HNO3. The purpose of using this solution
as the blank was to measure how SCN- reacts with Fe(NO3), since NaSCN was not added to
the first solution it was used as the basis or starting point to base the future reactions off of.
For the colorimetric determination experiment the “zero” or starting part was water
Experiment # 20
Nichole Morgan
Chemistry 106, Section 17
3
because it does not absorb any light, thus it had to be zero, while the solutions measured all
had varying amounts of color. In other words, for this experiment all the solutions had
varying amounts of NaSCN added which was measured by the spectrophotometer, while in
the previous experiment all had varying amounts of the color added and the absorbency
measured.
In the first part of the experiment there was an excess of Fe3+ causing the
equilibrium to shift to the right and allowing the reaction to reach completion. As well the
as a result the concentration of the products increased. The shift occurred because Q was
less than the K value. For the second part of the experiment however, all of the reactants
were about the same so the equilibrium remain unchanged, but also because the
stoichiometric ratio between SCN- and [Fe(H2O)5SCN]2+ was 1 to1 which proved to be vital
when calculating the equilibrium constant.
Experiment # 20
Nichole Morgan
Chemistry 106, Section 17
4
Works Cited
Iron(III) Thiocyanate Revisited: A Physical Chemistry Equilibrium Lab Incorporating Ionic
Strength Effects Cathy L. Cobb and G. A. LoveJournal of Chemical Education 1998 75 (1), 90