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Spectrophotometric Determination of an Equilibrium Constant & Investigating Le Chatelier’s Principle
INTRODUCTION
Most chemical reactions (e.g., the “generic” A + B→ 2C) are reversible, meaning they have a forward reaction (A + B
forming 2C) and a backward reaction (2C forming A+ B). Initially, when the concentrations of A and B are much higher
than the concentration of C, the forward reaction rate is much faster than the reverse reaction rate. However, when enough
C is formed and A and B consumed, the rate at which A and B are disappearing slows down and becomes the same as the
rate at which A and B are being formed by the reverse reaction. When this happens, the concentrations of reactants (A
and B) become constant, and the same is true of the product, (C). The reaction system is then said to be at
equilibrium. The ratio of the product of the product concentrations to the product of the reactant concentrations at
chemical equilibrium (each concentration raised to a power equal to the coefficients in the balanced reaction) is called the
equilibrium constant, Keq, for the reaction. Thus for the generic example,
Keq for the Formation of FeSCN2+: In some reversible reactions, the forward and reverse reaction rates are fast, so that
equilibrium is rapidly reached. One such reaction is that of iron(III) ion, Fe3+, with the thiocyanate ion, SCN-, that forms a
complex ion, iron thiocyanate, or thiocyanatoiron(III), FeSCN2+:
The double-headed arrow shows that the reaction is reversible. The equilibrium constant in terms of molar concentrations
can be expressed as
In order to determine Keq, it is necessary to determine the equilibrium concentrations of the reactants and products in
solutions, which are prepared by mixing carefully measured volumes of Fe3+ and SCN– solutions of known initial
concentrations. Instead of water, a dilute solution of nitric acid will be used to dilute the solutions. Iron (III) ions can
react with water to form a precipitate, iron (III) hydroxide. The presence of the acid reverses this process and therefore
effectively prevents the precipitate from forming. The equilibrium concentration of the FeSCN2+ ion, [FeSCN2+]eq,
formed in such a solution can be determined from the measured absorbance of the solution using a spectrophotometer.
This is possible because FeSCN2+ absorbs blue and green light, producing a solution that is reddish orange in color while
the Fe3+ and SCN– ions do not absorb visible light. The amount of light absorbed is thus proportional to [FeSCN2+] in
accordance with Beer’s Law. It is therefore necessary to first prepare solutions of known [FeSCN2+] (called standard
solutions) and obtain a calibration plot of absorbance vs. [FeSCN2+].
Preparation of Standard Solutions: To get solutions with known [FeSCN2+], the following process will be used. You
will prepare standard solutions by mixing carefully measured volumes of solutions of Fe3+ (using Fe(NO3)3 stock solution)
and SCN– (using KSCN stock solution) of known concentrations. These volumes for standard solutions A-F are listed in
the Standard Solutions Table. The key to getting a known concentration of FeSCN2+ in each of these solutions is that
the initial concentration of Fe3+ is much greater than the initial concentration of SCN– ion. When the Fe3+
concentration is in large excess, the equilibrium will shift (according to Le Chatelier’s Principle) to the product side until
virtually all the SCN– is converted to FeSCN2+. Thus the equilibrium concentration of FeSCN2+ in a standard solution
will be virtually the same as the initial concentration of SCN– in the solution. This initial value, [SCN–] initial, which
equals [FeSCN2+] in the mixed standard solution, can be calculated from the volume and molarity of the SCN– stock
solution and the final total volume (in L) of the mixed standard solution. With these calculated values for [FeSCN2+] and
the measured absorbances of these standard solutions, a Beer’s Law plot can be obtained.
Calculation of Keq from [FeSCN2+]eq, [Fe3+]eq and [SCN–]eq : A series of sample solutions can then be made up from
varying combinations of the reactant solutions as shown in the Sample Solutions Table. Once the equilibrium
concentration of FeSCN2+ is determined in these solutions using the measured absorbances and the Beers Law plot, it is an
easy matter to determine the equilibrium concentrations of the other reactant species. This is because as equation (1)
shows, for every one FeSCN2+ produced, one Fe3+ ion is used up. Thus the concentration of Fe3+ that remains at
equilibrium is the initial concentration of Fe3+ minus the concentration of Fe3+ that formed product, which is the same as
[FeSCN2+]eq. The equilibrium concentrations of Fe3+ and for the same reasons, SCN– can thus be expressed as
These concentrations are then used to determine Keq from equation (2).
When the conditions of a system at equilibrium are altered, the system responds in such a way as to maintain the
equilibrium. In 1888, Henri-Lewis Le Chatelier described this phenomenon in a principle that states, “when a change in
temperature, pressure or concentration disturbs a system in chemical equilibrium, the change will be counteracted by an
alteration in the equilibrium composition.” You will observe this principle at work by observing the formation of complex
ions between Co2+and water molecules or Co2+ and chloride ions:
A solution of the cobalt (II) ion in water is pink, the color of the complex ion formed between cobalt (II) ions and water
molecules. When chloride ions in the form of HCl are added to the solution, the color changes to blue, corresponding to
the formation of a changed coordination complex between cobalt (II) ions and chloride ions. This reaction is reversible
and quickly reaches a position of chemical equilibrium, which is immediately evident by the color of the solution. In this
lab, you will create a photo lab report of the equilibrium shifts caused by changing concentrations and temperatures.
SAFETY
Do not handle any chemicals with your hands due to possible skin irritation. Nitric acid cannot be poured down the
drain. You must neutralize your solutions with sodium hydroxide until the pH is neutral or in the basic range
before disposal.
MATERIALS & EQUIPMENT
0.200M Fe(NO3)3
0.00200M Fe(NO3)3
0.5M HNO3
0.00200M KSCN
0.1M CoCl2
12M HCl
0.1M AgNO3
dH2O
Spectrophotometer
Cuvettes
Beakers
Volumetric Flasks
Hot Plate
Disposable Pipets
Test tubes
Burets
PROCEDURE
Part A: Standard Solutions and Calibration Curve
1. Turn on the spectrophotometer and, allow it to warm up (at least 20 minutes) while the standard solutions are
prepared.
2. Set the wavelength on the spectrophotometer to 470 nm.
3. Prepare your group’s assigned solution following the volumes described on the table below. Use an
appropriately accurate measurement device for the volumes. Smaller graduated cylinders, burets, and
serological pipets will be provided. Add the KSCN and iron (III) nitrate first. The nitric acid should be
added last without specific measurement—it is poured until you reach the mark on the volumetric flask.
Take your time to ensure accurate measurements. Mix regularly as you add the solutions.
Standard
Total Volume
Volume of
Volume of
Volume of
Solution
0.00200M KSCN
0.200M Fe(NO3)3
0.5M HNO3*
A
100.00mL
1.00 mL
25.0 mL
~74.0 mL
B
100.00mL
2.00 mL
25.0 mL
~73.0 mL
C
100.00mL
4.00 mL
25.0 mL
~71.0 mL
D
100.00mL
6.00 mL
25.0 mL
~69.0 mL
E
100.00mL
8.00 mL
25.0 mL
~67.0 mL
F
100.00mL
10.00 mL
25.0 mL
~65.0 mL
* The volume of nitric acid does not actually need to be measured, as you will simply fill to the volumetric flask line
using nitric acid. Volumes are relative.
4. Calibrate the spectrophotometer with de-ionized water. Be sure to fill the cuvette and properly insert it into
the spectrophotometer by aligning the arrows.
5. Remove and empty the cuvette with the de-ionized water. Transfer about 20 mL of the solution you prepared
into a clean 50 mL beaker, and from the beaker, add the solution to rinse the cuvette. Then fill the cuvette
three-fourths full, and wipe the outside clear sides of the cuvette. Place the cuvette inside the
spectrophotometer and close the lid immediately. Allow 15 seconds for the absorbance reading to stabilize,
and record it in your Part A data table.
6. Measure and record the absorbance of the rest of the standard solutions using the same procedure, making
sure to rinse the cuvette with each solution prior to the measurement.
7. Calculate the concentrations of FeSCN2+ in each of the known solutions.
8. Create a Beer’s Law plot using excel.
Part B: Equilibrium Mixtures and Equilibrium Constant
1. Label five ~150 mL beakers, 1–5, for the equilibrium sample mixtures.
2. To prepare the sample solutions, use the same measurement tool for transferring KSCN and a separate labeled 10.0 mL
graduated cylinder or pipet for transferring the HNO3.
3. Rinse a buret with distilled water and then fill it with 0.002 M Fe(NO3)3. Use the buret to transfer 10.00 mL of the
0.002 M Fe(NO3)3 into each beaker for the equilibrium sample solutions.
4. Carefully add the appropriate amounts of KSCN and HNO3 into each beaker, according to the table below.
Equilibrium
Total Volume
Volume of
Volume of
Volume of
Solutions
0.002 M Fe(NO3)3
KSCN Solution
0.5 M HNO3
1
20.00 mL
10.00 mL
2.00 mL
8.00 mL
2
20.00 mL
10.00 mL
4.00 mL
6.00 mL
3
20.00 mL
10.00 mL
6.00 mL
4.00 mL
4
20.00 mL
10.00 mL
8.00 mL
2.00 mL
5
20.00 mL
10.00 mL
10.00 mL
0.00 mL
5. Swirl each beaker to mix the solutions.
6. Recalibrate the spectrophotometer using DI water.
7. Measure and record the absorbance of the equilibrium solutions, 1 – 5. Start with the most dilute solution
first, rinsing the cuvette or using a new cuvette with each solution prior to the measurement.
8. Neutralize all solutions containing nitric acid until the pH is neutral or in the basic range. Chemicals can be disposed
of down the front sinks.
9. Thoroughly clean all cuvettes and return them to the back bench.
Part C: Demonstration of Le Chatelier’s Principle
1. Obtain four medium-sized test tubes and label them: P1, P2, B1, and B2.
2. Measure approximately 5 mL of 0.1 M CoCl2 solution into each of the four test tubes.
3. Slowly and carefully add 5 mL of concentrated HCl to test tubes B1 and B2. Note: The solution in test tubes B1 and B2
should turn blue.
4. Set aside test tubes P1 and B1 as control solutions.
5. Obtain a 400-mL beaker and fill it with tap water about half full. Use a hot plate to heat the water to just under boiling.
6. Obtain a second 400-mL beaker and prepare an ice water bath filled about half full.
7. Add 5 mL of concentrated HCl to test tube P2 in approximately 0.5 mL increments until the solution turns blue in color.
8. Once the solution has turned blue in test tube P2 add 5 mL of distilled water in approximately 0.5 mL increments until
the solution reverts to its original pink color.
9. Add 2 mL of 0.1 M AgNO3 solution to test tube B2.
10. Obtain test tube P2, which contains roughly 15 mL of pink solution, and place it in the hot water bath. Keep it in the
hot water bath until the solution gradually changes in color.
11. Remove test tube P2 from the hot water bath and immediately immerse it in the ice water bath. The solution should
change colors.
12. Consider what each moment demonstrates in terms of equilibrium shifts. Repeat the procedure and take
photographs as you proceed—you will use them in your lab report!
13. Neutralize any solutions that are strongly acidic. Dispose of solutions.
DATA
Part A
[FeSCN2+]
Solution
A
B
C
D
E
F
Part B
Beaker
1
2
3
4
5
Part C
Step
3
7
8
9
10
11
[Fe3+] initial
Observations
[SCN-] initial
Absorbance
Absorbance
[FeSCN2+] equilibrium
Equilibrium Position
CALCULATIONS & POST LAB QUESTIONS
Part A: Calibration Curve:
1. Show a sample calculation of [FeSCN2+] for standard solution A.
2. Attach an Excel spreadsheet giving the table of data and the clearly labeled chart of the Beer’s Law plot showing the
equation of the line and a trend line.
Part B: Equilibrium Solutions
1. Calculate the [FeSCN2+]eq using the equation of the line from the Beer’s Law plot.
2. Calculate [Fe3+]eq and [SCN–]eq
3. Calculate Keq.
4. Attach a table showing the calculation of the average of the Keq values for sample solutions 1-5 and state,
“The mean value of Keq is __________.”
5. Does your data indicate that Keq is independent of the initial concentrations of the reactants?
6. List important sources of error in the determination of Keq and the effect that each would have on your results.
Part C: Le Chatelier’s Principle
Create a picture flow that demonstrates the color shifts in equilibrium. Explain what is going on in each picture and how
the equilibrium is being affected.