Energy and Equilibrium
Lab #2: Spectrophotometric Determination of an Equilibrium Constant
Part 1: Introduction to Spectrophotometry
Part 2: Determining K for a Reaction
Purpose
To construct a calibration graph and calculate the extinction coefficient for known
concentrations of a compound in solution that absorbs visible light and use those data to perform an
experiment to determine the equilibrium constant of a reaction.
Introduction (Part I)
A spectrophotometer is an instrument designed to measure the percentage of light of a given
wavelength that is transmitted through a liquid sample. For a sample that absorbs no light of that
wavelength, the transmittance will be 100%, and vice versa. The Beer-Lambert law related the
concentration of the substance to its absorbance (a function of transmittance):
𝐴 = 𝜀𝑐𝑙
A = absorbance (absorbance = log (100/T), where T = transmittance)
ε = molar absorptivity coefficient (M-1cm-1)
c = concentration (mol/L)
l = path length light travels through the substance (usually 1 cm)
Many factors affect the value of the molar absorptivity coefficient, so a calibration must be
done, preferably under the same or similar conditions as the experiment. A calibration is done by
plotting absorbance values vs concentration values for a reference set of samples of known
concentration. The resulting plot is subject to a linear regression. Because the path length is 1, the BeerLambert law will simplify to:
𝐴 = 𝜀𝑐
y = mx
The resulting slope, therefore, is the molar absorptivity coefficient.
Procedure (Part I)
1. Ensure that your spectrophotometer is turned on, the wavelength is set to 600 nm, and the
correct filter is selected.
2. Rinse and fill a cuvette 2/3 of the way with distilled water. Insert the cuvette into the
spectrophotometer and use the knob on the front to set the absorbance to 0.
3. Gather a 25 or 50 mL graduated cylinder and a 100 mL (or less) beaker from the glassware
cabinet. Rinse both with distilled water.
4. Add 10.0 mL of 0.5 M CuSO4 to your graduated cylinder.
5. Fill a cuvette approximately 2/3 of the way with 0.5 M CuSO4 and find the absorbance. If the
absorbance reads 1 or higher, the absorbance is out of range and unusable.
6. Empty and clean your cuvette.
7. Add 10.0 mL of distilled water to your graduated cylinder, diluting the CuSO4.
8. Fill your cuvette approximately 2/3 of the way with the solution from the cylinder and find
the absorbance.
9. Transfer 10 mL of the solution from your graduated cylinder to the empty cylinder. Discard
the remainder and the solution in the cuvette.
10. Repeat steps 7-9 until you have 7 points between 1.0 and 0, or your solution becomes too
dilute to register an absorbance (an absorbance of 0).
Analysis (Part I)
1. Calculate the concentration of CuSO4 at each dilution.
2. Construct a plot of absorbance vs concentration and perform a linear regression.
3. Determine the value of the molar absorptivity coefficient for your sample. Comment on the
reliability of this value (it is related to the correlation coefficient of the linear regression: R2).
Introduction (Part II)
Today we will investigate the equilibrium between the iron (III) ion and the thiocyanate ion. Iron
(III) forms a pale yellow solution and reacts with colorless aqueous thiocyanate to form an equilibrium
with the blood red FeNCS+2 complex ion as shown below.
Fe+3 (aq) + SCN– (aq) ↔ FeNCS+2 (aq)
Pale yellow Colorless
Blood red
The procedure for Part II has two stages. In the first, we make use of Le Chatelier’s Principle and
flood the system with an excess of iron (III) ions in order to force the equilibrium to favor the products.
In doing so, we may assume that the reaction essentially goes to completion – a single arrow rather than
a double arrow. In the second part, the amount of iron (III) is kept constant while the amount of
thiocyanate is varied. Because the concentrations are relatively close to each other, equilibrium will be
established. We will measure absorbance and use data from the first part to calculate an equilibrium
constant for the reaction.
Procedure (Part II)
1. Obtain ~15 mL of 0.10 M HNO3. Use approximately half to rinse a spectrophotometer cuvette,
then drain and fill the cuvette 2/3 with the solution.
2. Using the cuvette with 0.10 M HNO3, blank the spectrophotometer at 447 nm.
3. In the rear of the room are containers labeled with Roman numerals I through VI. Obtain an ~10
mL sample of the solution from one container in a small beaker, fill the cuvette 2/3 of the way
and measure the absorbance at 447 nm.
4. Repeat for each solution labeled with a Roman numeral. These are solutions of known
concentration for making a calibration curve.
a. Each solution contains 40.0 mL of 0.200 M Fe(NO3)3
b. The solutions from I-VI contain 0, 4.0, 8.0, 12.0, 16.0, and 20.0 mL of 0.00200 M KSCN,
respectively.
c. Each solution contains 0.10 M HNO3 to a total volume of 200. mL
5. Rinse your cuvette with 0.10 M HNO3.
6. Prepare samples as described below using your spectrophotometer cuvette. Wear a glove on
one hand for this part. After adding all substances to the cuvette, cover the top with your
gloved thumb and invert the cuvette several times to ensure mixing occurs. Rinse and dry the
glove that came in contact with the solution. For each sample, measure absorbance at 447 nm
and rinse with 0.1M HNO3
Tube #
1
2
3
4
5
6
mL 2.00x10-3 M Fe(NO3)3 in
0.10M HNO3
5.0
5.0
5.0
5.0
5.0
5.0
mL 2.00x10-3 M KSCN
in 0.10 M HNO3
0
1.0
2.0
3.0
4.0
5.0
mL 0.10 M HNO3
5.0
4.0
3.0
2.0
1.0
0
Analysis (Part II)
1. Calculate the concentration of FeNCS+2 in each solution (I-VI). You may assume that all of the
thiocyanate ions react to form the complex ion.
2. Construct a calibration curve according to the Beer-Lambert Law and determine the molar
absorptivity coefficient for the complex ion.
3. Calculate the initial concentration of iron (III) nitrate and potassium thiocyanate in each
experimental tube (1-6).
4. Using the Beer-Lambert law and your coefficient from Question #2, calculate the
concentration of complex ion in each tube.
5. Construct an ICE table for each experimental tube (1-6) and calculate the value of K for each
tube.
Teacher’s Notes
-
Part I of this lab is an introduction to spectrophotometry. Feel free to omit if your students have
already been trained.
-
If your equipment allows the gathering of an absorbance spectrum, rather than a single
wavelength measurement, strongly consider having students determine the absorbance
maximum of each compound as compared to a blank.
-
This is also an excellent lab to review the role of spectator ions and writing net ionic equations.
Often students fall into the “trap” of thinking that a solution can have only a cation or only an
anion in it without a counter-ion.
-
Equilibrium constant values tend to vary widely, with the tubes 2 and 3 giving K values far lower
than the rest. It is possible that the reaction kinetics at such low concentration require longer
times to come to equilibrium. Regardless, tubes 4-6 give excellent consistent data.