CHEMICAL EQUILIBRIUM: FINDING A CONSTANT, KC
PURPOSE:
To experimentally determine the equilibrium constant for the reaction of Iron (III) ion with thiocyanate
ion to form iron (III) thiocyanate complex using a spectrophotometric method
BACKGROUND INFORMATION:
The principle underlying a spectrophotometric method of analysis involves the interaction of
electromagnetic (EM) radiation with matter. The ultraviolet, visible, and infrared regions of the EM
spectrum are the most common used in analyses; in this experiment the visible region is used. The
wavelength range for the visible spectrum is about 400 nm to 700 nm; the 400 nm radiation
approximates a violet color while the 700 nm region has a red color.
Every chemical species (atoms, molecules, or ion) possesses a characteristic set of electronic,
vibrational, and rotational energy states. Because they are characteristic of a chemical species, energy
transitions between these states are often used to identify their presence and/or concentration in a
mixture. This unique set of energy states for chemical specie is therefore analogous to the unique set
of fingerprints possessed by each person—both can be used for characteristic identifications.
The absorption of EM radiation from the visible spectrum is a result of the excitation of an electron
from a lower to a higher state in the chemical specie. The energy of the radiation that is absorbed is
equal to the difference between these two energy states. The species (atom, molecule, or ion) that
absorbs the radiation is in an excited state. The EM radiation that is not absorbed, and therefore
passes through the sample, is detected by an EM detector (either our own eye or an instrument). The
c
absorbed energy, E, is related to the wavelength, , of the EM radiation by the equation, E = h ,

where H is Planck’s constant and c is the speed of light.
When our eye is the detector, the color that we see is the EM radiation, which the sample does NOT
absorb. The appearance of the sample is that of the complementary color of the absorbed radiation.
For example, if our sample absorbs EM radiation from the yellow region of the visible spectrum, then
the remaining EM radiation is transmitted to our eye and the sample appears violet. The greater the
concentration of the yellow absorbing specie, the darker is its violet appearance. Concord grapes
having a violet appearance absorb yellow light from the EM spectrum. Table 8-1 lists the wavelengths
of the visible spectrum.
Color
red
orange
yellow
green
blue
violet
Table 8-1 Color and Wavelengths in the Visible Region of the Electromagnetic Spectrum
Wavelength (nm)
Color Transmitted
750-610
blue-green
610-595
blue
595-580
violet
580-500
purple
500-435
orange
435-380
yellow
In this experiment the absorption of the visible radiation is used to determine the concentration of the
iron(III) thiocyanate, or Fe(SCN)+2, ion in an aqueous solution. The wavelength at which the
maximum absorption of visible radiation occurs is set on the spectrovis, which is the
spectrophotometer device used in today’s lab. This optimal wavelength must be determined
experimentally first in order to perform the experiment. Today, we know that the optimal
wavelength of absorption for Fe(SCN)+2 is 470 nm. All substances of course have a unique
wavelength at which they absorb.
Several factors affect the amount of visible radiation the Fe(SCN)+2 absorbs:
-
-
The concentration of the Fe(SCN)+2 in solution
The thickness of the solution through which the visible radiation passes (this is determined
by the diameter of the cuvette.
The extent to which the Fe(SCN)+2 absorbs the radiation at a particular wavelength (this is
called its molar absorptivity). This factor is constant for chemical specie at a set
wavelength.
The ratio of the intensities of the transmitted visible radiation, It, through the cuvette to the
incident visible radiation, IO, is the sample’s transmittance, T, or expressed as a percent,
I
%T, %T  t x100 . Frequently a chemist is more interested in knowing the amount of
IO
radiation that the Fe(SCN)+2absorbs rather than the amount that it transmits, the absorption
being directly proportional to the Fe(SCN)+2concentration. The absorbance, A, of the
radiation
is
related
to
the
percent
transmittance
by
the
equation
I
100
. The spectrovis, conveniently, gives you readings of both, so
A   log t x100  log
IO
%T
no calculation is necessary.
The absorbance of radiation is directly proportional to the molar concentration of the Fe(SCN)+2 by the
equation:
A = a · b · Fe(SCN)+2
Where a is the molar absorptivity for Fe(SCN)+2 and b is the thickness of the solution - both of which
are constants at a given wavelength in a given cuvette. If the absorbances of a series of standards are
measured, and their concentrations are known, then a linear relationship can be graphed in order to
obtain an equation for the relationship to predict future concentrations of Fe(SCN)+2, given their
absorbance. This equation and subsequent linear graph are commonly referred to as Beer’s law, the
important relationship being that Absorbance α concentration of Fe(SCN)+2.
MATERIALS:
0.0020 M KSCN
0.0020 M Fe(NO3)3 (in 1.0 M HNO3)
Three 10.0 mL graduated pipetters
Spectrovis, Logger Pro and computer
0.200 M Fe(NO3)3 (in 1.0 M HNO3)
0.0020 M KSCN
9 small sigma bottles
Kimwipes tissues (preferably lint-free)
PROCEDURE:
1. Obtain three 10.0 mL pipetters, and label them .0020M Fe(NO3)3, .200M Fe(NO3)3, and
.0020 M KSCN. Label three of your sigma bottles in the same manner. Fill each sigma bottle
with approximately 50 mL of each solution.
2. Label each of the remaining 6 sigma bottles Trials 1 through 5, and standard. Each of these
bottles will contain each of the reactions that you will perform. Using the appropriate pipetter,
add the following amounts of each chemical to each sigma bottle:
Sigma Bottle (Trial)
1
2
3
4
5
.0020 M Fe(NO3)3 (mL)
1.0
2.0
4.0
5.0
6.0
.0020 M KSCN (mL)
9.0
8.0
6.0
5.0
4.0
Standard: add 9.0 mL of .200 M Fe(NO3)3 to 1.0 mL of .0020M KSCN into your sigma bottle
labeled “Standard”.
3. Plug your spectrovis into the computer and calibrate it at 472.2 nm, or blue light, with distilled
water.
4. Rinse a cuvette with the solution reaction in the sigma bottle that contains the standard, and
then fill it ¾ of the way full with the solution reaction. Place this into the spectrovis, and
measure the absorbance of the standard. Record this absorbance for later data analysis.
5. Repeat step 4 for each of your 5 trials, making sure you record the absorbance each time. Rinse
the cuvette each time with the solution that contains the reaction.
6. Make sure all of this data is clearly shown in your data table.
7. Using the absorbance of your standard, and each of your six trials, do the following
calculations and data manipulations below.
CALCULATIONS:
ALL CALCULATIONS LISTED BELOW MUST BE CLEARLY SHOWN IN THE
LABBOOK.
1. Write the general equilibrium Kc expression for the reaction.
2. For your standard, calculate the initial concentration of Fe3+ and SCN– in solution, using your
dilution scheme, and then calculate the equilibrium concentration of [Fe3+]eq, [SCN-]eq, and
[FeSCN2+]. Use the equilibrium concentration value of [FeSCN2+] for your standard to
construct an absorbance equation/ratio as outlined in class.
3. Calculate the initial concentration of Fe3+ and SCN– in solution for Trials 1-5. Enter your
results in a clearly labeled table with proper units.
4. Construct an ICE reaction table to find the equilibrium concentrations of [Fe3+]eq and [SCN-]eq
for Trials 1-5. Use your results from Steps 2 and 3. Calculate the equilibrium concentration of
[FeSCN2+]eq for Trials 1-5 using the absorbance equation outlined in class.
5. Using the equilibrium Kc expression for the reaction and your equilibrium concentrations,
calculate Kc for Trials 1-5. Be sure to show the Kc expression and the values substituted in for
each of these calculations.
6. Summarize your equilibrium results by constructing a table of equilibrium concentrations for
each species and the calculated Kc. Clearly label all columns with units. Using your five
calculated Kc values, determine an average value for Kc. Then, calculate a percent deviation for
each of your five trials.
7. Determine a percent error for your average using Kc = 145/1 as the true value.
Percent Deviation =
| Average Value– Measured Value x 100% |
Average Value
DATA TABLES:
Enter the following data from your experimental results and analysis:
Absorbance of your standard solution:
_____________________________
Complete the following table for your lab report.
Solution
Absorbance
[Fe3+] at
equil.
[SCN-1] at
equil.
[FeSCN+2] at
equil.
Kc
1
2
3
4
5
Standard
Average equilibrium constant, Kc:
Temperature:
_____________________________
_____________________________
CONCLUSIONS AND RESULTS:
1. Describe the concept of “dynamic equilibrium”. Was dynamic behavior observed in the
reaction mixture monitored over time?
2. Characterize the iron thiocyanate reaction and subsequent equilibrium as either endothermic or
exothermic. Justify your characterization using Le Chatelier’s Principle.
3. If the equilibrium concentrations of Fe3+ and SCN– are 1.0x10–2 M and 3.0x10–3 M
respectively, what must be the equilibrium concentration of FeSCN2+? Use the above accepted
Kc as a guide.
4. If 5.0-mL of the 0.0020 M iron solution and 10.0-mL of the 0.0020 M KSCN solution are
mixed, what will be the equilibrium concentration of FeSCN2+ in the mixture? Use the above
accepted Kc as a guide.
5. If 0.050 M Fe3+, 0.040 M SCN–, and 0.500 M FeSCN2+ are present in a test tube then:
a. Calculate Qc. Is the system at equilibrium?
b. In what direction (if any) will a net reaction take place (forward or reverse)?