_____University of Puget Sound

Department of Chemistry

Chem 230

EXP. 2DETERMINATION OF AN EQUILIBRIUM CONSTANT

INTRODUCTION

This experiment will give you an opportunity to determine the equilibrium

constant for a reaction and use Le Chateliers principle.
FORMATION OF Fe(SCN)2+, IRON (III) THIOCYANATE ION
When chemical substances react, the reaction typically does not go to completion.
Rather, the system goes to some intermediate state in which both the reactants and
products have concentrations, which do not change with time. Such a system is said to
be in dynamic chemical equilibrium. At equilibrium a reaction mixture obeys the Law
of Mass Action, which imposes a condition on the concentrations of reactants and
products. This condition is expressed in the equilibrium constant Kc.
When the reaction between Fe3+ and SCN- (thiocyanate) ions in an aqueous
solution comes to equilibrium, the solution contains reactants and the product,
Fe(SCN)2++, iron (III) thiocyanate ion. The chemical equation for this reaction is
Fe3+(aq)

SCN-(aq)

Fe(SCN)2+(aq)

(1)

The product is a complex ion that has a coordinate covalent bond between the iron
cation and an atom (probably the S atom) from the thiocyanate anion. The color of this
complex ion is so intense that thiocyanate ions can be used to detect very small
quantities of Fe3+. Interestingly, Fe(SCN)2+ appears to exist solely in solution; solid
compounds containing this cation and anion have never been isolated.
The object of this experiment is to determine the value for the equilibrium constant
for reaction (1). The equilibrium constant is given by the expression
Kc =

[Fe(SCN)2+]
[Fe3+] [SCN-]

(2)

where the concentrations of the substances are those at equilibrium. If these

concentrations can be determined, Kc can be calculated easily.
SPECTROPHOTOMETRY
Both reactants, Fe3+ and SCN-, are essentially colorless, whereas the complex ion
product is highly colored. Therefore, you can use a spectrophotometer to monitor the
concentration of Fe(SCN)2+ by measuring the absorbance due to the complex ion.
According to Beers law, the absorbance, A, is proportional to the concentration, c, of
the species that absorbs the light, in this case Fe(SCN)2+.
A = bc

(3)

is called the molar absorptivity . Sometimes this term is expressed as an a and is

called an extinction coefficient. For a particular species, is a constant at a specific
c230 Exp. 2 - Determination of an Equilibrium Constant

wavelength, . b is the path length through the sample. The path length is either the
internal path length of a square cuvette or the inside diameter of a cylindrical cuvette.
The path length can be considered a constant if you use the same cuvette for every
reading or you use a matched set of cuvettes. Because is a constant and b can be
treated as a constant, their product is also a constant which we will call z . Equation
(3) then becomes Equation (4) with z being the slope of the line produced when you
graph absorbance versus concentration.
A = zc

(4)

THE STANDARD OR CALIBRATION CURVE

To determine the relationship between Fe(SCN)2+ and absorbance, the absorbance
of a series of solutions of known concentrations of complex ion must be measured.
But there is a problem. How can solutions of known concentrations of Fe(SCN)2+ be
obtained? After all, this substance is a participant in the equilibrium with Fe3+ and
SCN- ions. Stoichiometric quantities of reactants yield a specific concentration of
product which can be calculated if the equilibrium constant, Kc, is known. But Kc is
the term that we are trying to determine!
The difficulty of making solutions with known concentrations of Fe(SCN)2+ can be
overcome by using Le Chateliers principle. This principle suggests that an
equilibrium position can be shifted to the right, that is, in the forward direction, when
more of a given reactant is added. As more reactant is added, more product will form.
It is possible to add so much of one reactant that essentially all of the other reactant is
converted to product. In this experiment you will use overwhelming amounts of Fe3+
and limited quantities of SCN- to achieve this result. The concentration of Fe(SCN)2+
that is formed will be identical essentially to the starting concentration of the limiting
reactant, SCN-.
In Part 1 of this experiment you will prepare a series of solutions having known,
limited concentrations of SCN-. Therefore, you will know the concentrations of
Fe(SCN)2+. Then you will measure the absorbances of these solutions at max, the
wavelength of maximum absorbance of the complex. Finally, you will plot the
absorbances versus concentrations of Fe(SCN)2+ which should give a straight line.
This plot is called the Calibration Curve (even if it is a straight line.)
DETERMINATION OF Kc, THE EQUILIBRIUM CONSTANT
As noted above, when solutions containing Fe3+ ion and thiocyanate ion are
mixed, reaction (1) occurs forming the Fe(SCN)2+ complex ion which has a deep red
color. As a result of the reaction, the equilibrium amounts of Fe3+ and SCN- are less
than they would have been if no reaction occurred; for every mole of Fe(SCN)2+ that
is formed, one mole of Fe3+ and one mole of SCN- have reacted.
The value of Kc in equation (2) is constant at a given temperature. This means that
solutions containing Fe3+ and SCN- will react until the concentrations of remaining
reactants and formed products satisfy the value of Kc. The same value of the Kc will
be obtained no matter what initial amounts of Fe3+ and SCN- are used. Our purpose
in this experiment will be to determine the value of Kc for this reaction using several
solutions made up in different ways, and to show that Kc indeed has the same value
c230 Exp. 2 - Determination of an Equilibrium Constant

in each circumstance. The reaction is a good one to study because Kc is a convenient

magnitude and the color of the Fe(SCN)2+ ion makes the analysis of the equilibrium
mixture amenable to spectrophotometric analysis.
In Part 2 of this experiment you will prepare solutions of known initial
concentrations of iron(III) nitrate, Fe(NO3)3, and potassium (or sodium) thiocyanate,
KSCN (or NaSCN.) Then you will determine the absorbance at max of each solution.
The absorbance can be used to determine the concentration of Fe(SCN)2+. Knowing
the initial concentrations of the reactants and the equilibrium concentration of
Fe(SCN)2+, you can calculate the equilibrium concentrations of the remaining
reactants. You can then determine the value of Kc.
CALCULATIONS
Since the calculations required in this experiment may not be apparent, we will go
through a specific example. Let us assume that we prepare a solution by mixing 10.0
mL of 2.00 x 10-3 M Fe(NO3)3 with 10.0 mL of 2.00 x 10-3 M KSCN. As a result of
reaction (1) some red Fe(SCN)2+ ion forms. The concentration of the Fe(SCN)2+ in this
solution can be determined by finding the absorbance of the solution at 447 nm and
using the standard curve to determine its concentration. Let's say the concentration at
equilibrium is found to be 1.50 x 10-4 M. Our problem is to determine Kc for the
reaction from this information. To do this we can carry out the following steps and
summarize the information on a reaction table.
1.
2.
3.
4.
5.

Write a balanced equation for the reaction.

Calculate the initial concentrations of the reactants and/or products.
Using stoichiometry, find the change during the reaction.
Calculate the equilibrium (final) concentrations.
Write the equilibrium constant and calculate Kc for the reaction.

Here are the steps for solving Kc for reaction (1) from the information given above.
Fe3+(aq)

SCN-(aq)

Fe(SCN)2+(aq)

Calculate the initial [Fe3+], [SCN-] and [FeSCN2+]:

(0.00200 x 0.0100)
= 0.00100 M
0.020
(0.00200
x
0.0100)
[SCN-]init =
= 0.00100 M
0.020
[FeSCN2+] = 0

[Fe3+]init =
and of course,

init

Set up the reaction table, with x = [FeSCN2+], the concentration of Fe(SCN)2+ at

equilibrium.
Concentration (M)
Initial
Change
Equilibrium

Fe3+
+
SCN0.00100
0.00100
-x
-x
0.00100 - x
0.00100 - x

c230 Exp. 2 - Determination of an Equilibrium Constant

FeSCN2+
0
+x
x
3

To calculate the equilibrium concentration of all the components we must solve for
x. From the reaction table, x is the equilibrium concentration of FeSCN2+ which you
will determine from the absorbance versus [Fe(SCN)2+] standard curve. In the above
problem it is given as the following:
x = [FeSCN] = .000150 M
Therefore, the equilibrium concentrations are:
[Fe3+] = 8.5 x 10-4 M

[SCN-] = 8.5 x 10-4 M

and

[FeSCN2+] = 1.50 x 10-4 M

Substituting into the mass-action expression:

Kc =

[Fe(SCN)2+]
[Fe3+] [SCN-]

1.50 x 10-4 M

= (8.5 x 10-4 M) x (8.5 x 10-4 M)

= 2.1 x 102 M-1

In this experiment you will determine the absorbance at max of different solutions.
That data will be transformed into concentrations similar to those shown in the
example above. The calculations involved in processing the concentration terms are
completely analogous to those illustrated. (Your value for Kc will differ from the one
we obtained, since the data in our example were obtained at a different temperature
thus relate to a different value of Kc.)

c230 Exp. 2 - Determination of an Equilibrium Constant

EXPERIMENTAL PROCEDURE
WORK IN PAIRS
During this experiment you and your partner will be working with a
spectrophotometer which is an expensive, delicate analytical instrument. Operate it
with care. Follow the instructions of your laboratory instructor.
All stock solutions are prepared in 0.10 M HNO3, nitric acid, and dilutions are
done using 0.1 M HNO3, Therefore the H+ concentration is maintained at a constant
concentration. The hydrogen ion does not participate directly in the reaction you are
studying, but its presence is necessary to avoid the formation of brown-colored
species such as Fe(OH)2+, which would interfere with the analysis of [Fe(SCN)2+].
You will be using pipets and a volumetric flask to make up solutions. Practice
pipetting skills using water before starting the following procedure,
All solutions in this experiment can be disposed in the sink with plenty of
water.
Part 1.
1.

The Standard Curve

Prepare solutions of varying Fe(SCN)2+ concentration. Use a 50 mL
volumetric flask to make a series of five solutions. Make the five
solutions by pipetting 0, 1, 2, 3 and 4 mL aliquots of 0.00200 M KSCN
into the 50 mL volumetric flask. For each solution pipet 10.00 mL of
0.200 M Fe(NO3)3 into the flask and fill the flask to the 50 mL mark with
0.10 M HNO3. Mix thoroughly by inverting the flask and swirling to
ensure that equilibrium is established. Place the solution into another
clean, dry, labeled container and thoroughly rinse the volumetric flask
before making the next solution. These solutions are used to establish a
standard curve of absorbance at the max versus Fe(SCN)2+
concentration. (See Table 1.) Note: Prepare all five solutions and then
obtain the spectral data.

Table 1. Composition of Solutions for Standard Curve

Solution

0.200 M Fe(NO3)3
(in 0.1 M HNO3), mL

0.00200 M NKSCN
(in 0.1 M HNO3), mL

0.10 M HNO3 , mL

1
2
3
4
5

10.00
10.00
10.00
10.00
10.00

0.00
1.00
2.00
3.00
4.00

add to total 50.00

add to total 50.00
add to total 50.00
add to total 50.00
add to total 50.00

2.

Solution 1, the blank. Rinse a cuvette with several portions of Solution 1

and then fill it about three-fourths full with Solution 1. Solution 1 is
called the blank and is used to calibrate the spectrophotometer. Carefully
dry the outside of the cuvette with a clean Kimwipe to remove water
and fingerprints. Thereafter, handle the cuvette by its top edge. Any
foreign material on the cuvette window affects the intensity of the
transmitted light.
After the instrument has been connected to the computer and the
application has connected to the spectrophotometer, an absorption
c230 Exp. 2 - Determination of an Equilibrium Constant

spectrum must be obtained to determine the max but first the

spectrophotometer must be calibrated with a blank. We will use
Solution 1 as the blank. Select the Calibrate menu item under the
Experiment. Proceed with the calibration as prompted by the LoggerPro
application.
3.

Determine the maximum wavelength for Fe(SCN)2+.

a. Empty the blank cuvette and rinse it twice with small amounts of
the Fe(SCN)2+ solution from Mixture#3. Fill the cuvette 34 full with the
Fe(SCN)2+ solution and place it in the cuvette holder of the
spectrometer.
b. Click the Collect button. A full spectrum graph of the Fe(SCN)2+
solution will be displayed. Note that one area of the graph contains a
peak absorbance. Click Stop to complete the analysis. Record the max in
your lab notebook.
c.
To save your graph of absorbance vs. wavelength, select Store
Latest Run from the Experiment menu.
d. Click the Configure Spectrometer Data Collection icon on the
toolbar. A dialog box will appear.
e.
Select Absorbance vs. Concentration under Set Collection Mode.
The peak absorbance will be automatically selected. Click the Collect
button to proceed. If you need to start over, click Clear and select the
desired wavelength (or wavelengths) again.
All solutions in this experiment can be disposed in the sink with
plenty of water.

4.

Collect absorbance-concentration data of solutions 1 through 5.

a. Remove the cuvette from the spectrometer and dispose of the
solution as directed. Use the solution in Test Tube 1 to rinse the cuvette
twice with ~1 mL amounts, and then fill it 34 full. Wipe the cuvette
with a paper towel and place it in the spectrometer cuvette holder. Click
the Collect button. When the absorbance reading stabilizes, click the
Keep button. Enter the concentration of the solution and click OK.
b. Discard the cuvette contents as directed. Using the solution in Test
Tube 2, rinse and fill the cuvette 34 full. Wipe the cuvette and place it in
the spectrometer. When the absorbance reading stabilizes, click Keep.
Enter the concentration.
c.
Repeat Step 4b for the remaining test tubes of the standard
Fe(SCN)2+ solution. When you have finished testing the standard
solutions, click Stop. Record each solution's absorbance at the max in
your laboratory notebook in Table 1.

5.

Prepare standard curve. Calculate the concentration of Fe(SCN)2+ for

solutions 2 through 5. Check that the absorbance, A, (ordinate) versus
[Fe(SCN)2+] (abscissa) has a reasonable correlation to a Linear Fit line.
Display the Linear fit line through the points. Look at your graph with
your instructor to determine if your results are reasonable.
LEAVE THE Application OPEN with the data and proceed with Part 2.
c230 Exp. 2 - Determination of an Equilibrium Constant

Part 2 . Determination of Equilibrium Constant (Kc)

1.
Prepare 0.00200 M Fe(NO3)3. A solution of 0.00200 M Fe(NO3)3 is not
available. To make this solution, dilute 0.50 mL of 0.200 M Fe(NO3)3,
measured with a 1.00-mL pipet, to 50 mL with 0.1 M HNO3 in a
volumetric flask. This is the stock 0.00200 M Fe3+ solution used in Part
2.
2.

Prepare test solutions. Table 2 indicates the volumes of each stock

solution needed to prepare test solutions. Use pipets for the volumetric
measurements . Dry medium size test tubes are suitable containers for
these solutions. Stir each test solution for approximately one minute
until the color does not change.
Table 2. Test Solutions for Kc Determination
Solution

0.00200 M Fe(NO3)3*
(in 0.1 M HNO 3), mL

0.00200 M KSCN
(in 0.1 M HNO 3), mL

0.1 M HNO3, mL

6
7
8
9
10

5.00
4.00
3.00
2.00
1.00

1.00
2.00
3.00
4.00
5.00

4.00
4.00
4.00
4.00
4.00

3.

Collect absorbance data of solutions 6 through 10.

a. Rinse the cuvette twice with the unknown solution and fill it about
34 full. Wipe the outside of the cuvette and place it into the
spectrometer.
b. Record the absorbance of each of the unknown in the data table
(Table 2) in your laboratory notebook. (Do not click on the Collect
button.)
c.
Repeat Step 3b for the remaining test tubes solution (tubes 6-10) of
the equilibrium solutions of Fe(SCN)2+. Record each solution's
absorbance at the max in your laboratory notebook in Table 2.

4.

Determine Fe(SCN)2+ equilibrium concentration. From the absorbance

value of each solution (6-10) and the standard curve plotted from Part 1,
the Fe(SCN)2+ equilibrium concentration for each solution can be
determined.

5.

Calculate Kc for each solution. Determine the average Kc and its

standard deviation and report with appropriate significant figures.

c230 Exp. 2 - Determination of an Equilibrium Constant

QUESTIONS
Q.1. When Fe3+ and SCN- react and reach an equilibrium with Fe(SCN)2+, what
happens to the conc. of Fe3+? How are the number of moles Fe(SCN)2+
produced and the number of moles Fe3+ used up related to each other?
Q.2. A student mixes 5.0 mL of 2.00 x 10-3 M Fe(NO3)3 with 5.0 mL of
2.00 x 10-3 M NaSCN. She finds that in the equilibrium mixture the
concentration of Fe(SCN)2+ is 1.4 x 10-4 M. Use the data to put together a
reaction table and calculate Kc.
Q.3. What effect does a dirty cuvette (caused by fingerprints, water spots, or lint)
have on the absorbance reading for a Fe(SCN)2+ solution?
How does this error affect the reported equilibrium constant?
Q.4. In your calculations, the path length, b, and the molar absorptivity, , of the
absorbing species are not considered, see Equation (3). Explain why neither of
these terms has to be determined.
Q.5. If the absorbance of a solution reads greater than 2.0 on the spectrophotometer,
how could the preparation of the solution be modified to obtain a more accurate
reading?
Q.6. Why can you calculate [Fe(SCN)2+] from the amounts of reactants mixed in Part
1, but you can not calculate [Fe(SCN)2+] from the amounts of reactants mixed in
Part 2?

c230 Exp. 2 - Determination of an Equilibrium Constant

WHAT TO DO

Before Lab:
Read this experiment. Fill in the Title bar to include the
experiment #, title, your name, partner's name, date, lab section, and TA's name. This
information should be included on each page in your laboratory notebook for every
experiment.
All prelab material should be written in your laboratory notebook; the pre-lab
assignment will be graded by your TA at the start of the lab.
For this weeks experiment your Prelab write-up should include, in addition to the
completed title bars, the following sections.
1.
Purpose
After carefully reading the experiment, briefly state what you
understand to be the purpose of this lab exercise.
2
Table of Reagents for Fe(NO3)3, KSCN, and HNO3. BE SURE TO NOTE
THE DISPOSAL PROCEDURES.
Your lab book should also have the following information and tables ready for
use:
3.
Prepare the following two large tables in your notebook before coming
to lab (please note that these tables are extended versions of Tables 1 and
2 in the experiment). Each of these tables should be one page and should
be oriented sideways in your notebook (landscape). Before lab you
should have completed columns 1-6 for both of the tables shown below.
Table 1. Calibration Data using the Absorbance of Fe(SCN)2+
*Volumes in mL
3+
trial 0.200 M Fe , mL 0.00200 M SCN , mL
1
2
3
4
5

10.0

0.00

Initial Conc, M
3+
[Fe ]
[SCN ]

Total
50.0

0. 0400

Equil Conc, M
[Fe(SCN)2+]

0.00000

Table 2. Test Solutions for Determination of the Equilibrium Constant

*Volumes in mL
trial
6
7
8
9
10

3+

Fe
5.00
4.00
3.00
2.00
1.00

SCN
1.00
2.00
3.00
4.00
5.00

Initial Conc, M
Total
10.00
10.00
10.00
10.00
10.00

3+

[Fe ]

Equilibrium Conc, M
-

[SCN ]

2+

[Fe(SCN) ]

3+

[Fe ]

[SCN ]

During the Lab: Carry out the procedure as written. Modify your procedure if
you find it necessary, but be sure to record the modifications. Record all observations
and measurements as you do the experiment. Be sure to obtain the data for all
solutions in Part 1 and Part 2. Put the data in the appropriate tables that you have
prepared in your laboratory notebook. Plot the standard curve on a sheet of good
quality graph paper (10 mm to the cm) and discuss the graph with your instructor
before leaving lab.
To Be Turned In: The report for this experiment is to be done as a complete
formal laboratory report. It must be typewritten according to the Guidelines (you may
download them from the CHEM 230 Lab Information website). When making your
final calibration curve you will want to use fine grid graph paper. Click on the link
c230 Exp. 2 - Determination of an Equilibrium Constant

found on the Chem 230 Lab Information web page to download a PDF file of fine grid
graph paper that you can print out. (Be sure that any "shrink to fit" print settings are
turned off.) A table containing the information summarized in the Calculations
section that follows the Experimental Procedure should be included in your results
section.
Two appendices should be included with the report. These appendices can be
done in your lab book and you can turn in the copy pages. One appendix should have
the answers (in complete sentences) to the questions listed at the end of the
Experimental Procedure section. The second appendix should include the following
sample calculations: one sample calculation for the concentration of Fe(SCN)2+ in
solution 2, one reaction table related to solution 6, and one sample calculation for Kc
for solution 6.

c230 Exp. 2 - Determination of an Equilibrium Constant

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