CHM 152
Last updated Jan. 2012
Lab 6: Experimentally Determining an Equilibrium Constant using
Spectrophotometry
Introduction
In this lab you will experimentally determine the equilibrium constant with respect to concentration, K c, of
the reaction between iron (III) and thiocyanate, SCN -.
Fe3+(aq) + SCN-(aq)  Fe(SCN)2+(aq)
We’ll do this using a spectrophotometer, which is an instrument that can measure the amount of light
absorbed by a sample. The Fe(SCN)2+ ion is a dark red color in aqueous solution. When light is passed
through the sample, some will be absorbed by this colored solution. The more concentrated the solution,
the darker the color and the more light will be absorbed. In a spectrophotometer, light of a certain
wavelength is passed through to a detector that measures the percentage that transmitted (%T) through the
sample. The percent transmittance can be converted to the amount absorbed by the solution (A) using the
following equation.
A = 2.000 – log (%T)
The instrument you’ll be using for this lab can automatically convert percent transmittance to absorption.
In the mid-nineteenth century, German physicist August Beer (no, really, that was his name) discovered
that the absorption of light by a solution is directly proportional to its concentration, a relationship often
referred to as Beer’s Law (or the Beer-Lambert Law, to give additional credit to Swiss physicist Johann
Lambert). The mathematical equation for Beer’s Law is fairly simple and given below.
A = bc or A = abc
In this equation, (or a) is the molar absorptivity coefficient, which relates to how well a solution absorbs
light at particular wavelength, b is the distance (in cm) traveled by the light through the solution, and c is
the molar concentration of the solution. Using this equation, a plot of absorption vs. concentration should
be linear, with a slope of m = b and a y-intercept of b=0.
In Part I of this lab, you will prepare six solutions with a known concentration of Fe(SCN) 2+ and use these
known solutions to prepare a Beer’s Law plot. There is one major point that needs to be considered,
however. This reaction is reversible, so when the equilibrium is established both reactant and product are
typically present in significant amounts. In lecture, we usually use a known equilibrium constant to
determine equilibrium concentrations (using the ICE method), but that’s what we’re trying to determine in
this lab. So how can we know the concentration without first having a K c?
To accomplish this, we’ll “rig” our system by reacting thiocyanate with a large excess of iron (III). This
excess amount of reactant will be so great that, according to Le Châtelier’s principle, the system will shift
right to such a large degree that nearly all of our thiocyanate will be converted to product. There will be
some thiocyanate present, since the equilibrium is still present, but that amount will be too insignificant to
adversely affect our determination of Kc.
In Part II, you will make a new set of solutions, only this time using more equal concentrations of each
reactant. Without a large excess present, both reactants and the product will be present in significant
amounts at equilibrium. After measuring the absorbance of these solutions, we can use our Beer’s Law plot
from Part I to determine the concentration of Fe(SCN) 2+ in each solution then use an ICE table to determine
the equilibrium concentrations of the two reactants. Once the equilibrium concentrations of all three
species are known, the equilibrium constant can be determined.
Concepts to Review
Equilibrium constants
LeChâtelier’s principle
“ICE” tables
Diluting solutions
Plotting with Excel
Procedure
To save time, you and your lab partner should split the work by having one prepare the solutions in Part I
while the other prepares the solutions in Part II. This will allow you to measure the absorptions of both sets
of solutions at the same time.
Part I: Preparation of a Beer’s Law Plot
1. Measure out 70 mL of 0.200 M iron (III) nitrate and 20 mL of 2.00 x 10-3 M sodium thiocyanate and
take them back to your work area. These will be used to prepare your solutions in the next step.
2. Prepare the following set of solutions using a 50.0 0 mL volumetric flask. Use volumetric pipettes to
measure each reactant and use 0.10 M nitric acid (not water) to dilute the solution to 50.00 mL. Mix the
solution by capping the flask and inverting 2-3 times. Once a solution is prepared, transfer it to a beaker or
Erlenmeyer flask so you can use the volumetric flask to make the next solution on the list (don’t forget to
rinse the flask between preps).
Solution
S1
S2
S3
S4
S5
S6
0.200M Fe(NO3)3, mL
10.00
10.00
10.00
10.00
10.00
10.00
2.00 x 10-3M NaSCN, mL
0.00
1.00
2.00
3.00
4.00
5.00
3. Obtain six cuvettes and fill each (approximately ¾ full) with one of the solutions from step 1.
4. Measure the absorption of each solution using one of the lab spectrophotometer.
Operating the Genesys 20 spectrophotometer:
1. Use the nm buttons to set the wavelength to 447 nm (if it isn’t already).
2. Make sure the instrument is set to measure absorption. The screen should display the
wavelength, a number and the letter A (e.g., 447nm x.xxx A) If not, use the “A/T/C” button
to switch to absorption mode.
3. Take sample S1 and wipe the cuvette clean with a Kim-wipe to remove prints or smudges.
Since this sample only contains Fe3+(aq), it will serve as our blank (i.e., zero absorption).
4. Place the sample in the instrument and close the cover.
5. Press the “0 ABS/100% T” button to zero the instrument (similar to what you do with a
balance). After a couple of seconds the instrument will read “0.00 A”
6. Remove the sample from the instrument and replace it with sample S2. Record the
absorbance shown.
7. Repeat step 6 for samples S3-S6.
5. Using Excel, prepare a linear plot of Absorption (y) over [Fe(SCN)2+] (x). Force the y-intercept to zero
by right-clicking on the best-fit line, selecting “Format trendline,” and checking box labeled “Set intercept
= 0”
Part II: Determination of the Equilibrium Constant
1. Measure out 40 mL of 2.00 x 10-3 M iron (III) nitrate and 20 mL of 2.00 x 10-3 M sodium thiocyanate
and take the solutions back to your work area.
2. Prepare the following set of solutions in medium-sized test tubes. For each solution add enough 0.10 M
nitric acid (not water) to bring the total volume to 10.0 mL. Mix with a glass stir rod.
Solution
E1
E2
E3
E4
E5
E6
2.00 x 10-3M Fe(NO3)3, mL
5.00
5.00
5.00
5.00
5.00
5.00
2.00 x 10-3M NaSCN, mL
0.00
1.00
2.00
3.00
4.00
5.00
3. Measure the absorption of each solution as outlined in steps 3 and 4 of Part I.
4. Using the linear equation from your Beer’s Law plot, determine the concentration of Fe(SCN) 2+ for
samples E2-E6.
5. Use an ICE table to determine the equilibrium concentrations of the reactants samples E2-E6.
6. Determine the equilibrium constant for samples E2-E6
Waste Disposal
During lab collect all of the solutions to be discarded in a 600mL beaker. When the lab is finished, pour
half of the beaker’s contents down the drain with running water. Dilute the remaining solution by refilling
the beaker with water and pour half of this down the drain. Repeat this until the solution in the beaker is
colorless.
Name: _____________________________
Section: ________
Data
Part 1: Preparation of a Beer’s Law Plot
S1
(blank)
S2
S3
S4
S5
S6
1) Volume of 0.002M NaSCN
added, mL
0
________
________
________
________
________
2) Initial [SCN-] of solution, M
0
________
________
________
________
________
3) Equilibrium [Fe(SCN)2+], M
0
________
________
________
________
________
4) Absorbance
0
________
________
________
________
________
Include a copy of your Beer’s Law plot with your report.
Part II: Determination of the Equilibrium Constant
E2
E3
E4
E5
E6
1) Volume of 0.002M Fe(NO3)3
added, mL
________
________
________
________
________
2) Volume of 0.002M NaSCN
added, mL
________
________
________
________
________
3) Initial [SCN-] of solution, M
________
________
________
________
________
4) Initial [Fe3+] of solution, M
________
________
________
________
________
5) Absorbance
________
________
________
________
________
6) Equilibrium [Fe(SCN)2+], M
________
________
________
________
________
7) Equilibrium [Fe3+], M
________
________
________
________
________
8) Equilibrium [SCN-], M
________
________
________
________
________
9) Equilibrium constant, Kc
________
________
________
________
________
10) Average Kc
________
Sample calculations
Show your work for the following calculations using data from sample E3
a) Initial solution concentrations of Fe3+ and SCN-
b) Equilibrium concentration of Fe(SCN)2+
c) Equilibrium concentrations of Fe3+ and SCN- (include an ICE table)
d) Equilibrium constant
Name: _____________________________
Section: ________
Post-Lab Questions
1. a) Does your mean value of Kc suggest a reactant or product-favored equilibrium?
b) Does this agree with what you observed while prepping the solutions? Explain.
2. A student used a beaker instead of a volumetric flask in Part I, making his final volumes larger than the
50mL he assumed them to be. How would this affect each of the following (too high, too low, or no
effect)? For each case, explain your answer.
a) The initial SCN- concentration
b) The slope of the Beer’s Law plot
c) The calculated product concentration in Part II
d) The calculated reactant concentrations in Part II
e) The calculated value of Kc
Name: _____________________________
Section: ________
Pre-Lab Questions
1. Write the expression of Kc for the reaction being studied in this lab.
2. What assumption is being made in Part I of this lab? Why can’t the same assumption be made in Part
II?
3. When plotting the data in Part I, why does it make sense to set the y-intercept at y=0?
4. A Beer’s Law plot was prepared for the reaction A(aq) + B(aq)  AB(aq), plotting absorption over
AB(aq) concentration. The linear equation for this plot was y = 78.3x.
A solution was prepared by mixing 10.0 mL of 0.100M A with 5.00 mL of 0.100 M B and adding enough
water to bring the total volume to 50.0 mL. The absorption of this solution was measured as 0.646.
Given this information, calculate the following:
a) The initial concentrations of each reactant.
b) The equilibrium concentration of AB
c) The equilibrium concentrations of each reactant
d) The equilibrium constant with respect to concentration