Experiment 4
Determination of the Equilibrium Constant for a Chemical
Reaction
A. Purpose
1.
2.
3.
4.
To review the concepts and principles of chemical equilibrium.
To learn the principles and use of a spectrophotometer.
To learn the principles of chemical analysis by spectrophotometric means.
To determine, spectrophotometrically, the equilibrium constant for the
reaction between iron (III) ion and thiocyanate ion (SCN) at a given
temperature.
B. Theory
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
the products have concentrations which do not change with time: Chemical
Equilibrium is reached.
When a reaction mixture reaches equilibrium at a particular temperature, it obeys
the Law of Chemical Equilibrium which imposes a condition on the
concentrations of reactants and products. This condition is expressed in the
Equilibrium Constant Kc for the reaction at a given temperature.
In this experiment, we study the equilibrium reaction:
Fe3+ (aq) + SCN (aq)  FeSCN2+ (aq)
The equilibrium constant is defined as:
Kc 
[FeSCN2 ]
[Fe3 ][SCN  ]
Kc is constant at constant temperature; this means that mixtures containing Fe3+
and SCN will react until this equilibrium is satisfied, so that the same value of Kc
will be obtained no matter what initial amounts of Fe3+ and SCN are used.
This reaction is particularly suitable for study because Kc has a convenient
magnitude and the color of the FeSCN2+ ion allows for easy spectrophotometric
analysis of the equilibrium mixture. FeSCN2+ is a deep, blood-red complex with
an absorption maximum at 447 nm. Kc will be determined for several mixtures
made up in different ways, and it can thus be shown that Kc has indeed the same
value in each mixture.
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The Beer-Lambert’s Law will be applied, in this experiment, to determine the
FeSCN2+ concentration at equilibrium. A = bc, where A is the absorbance of the
substance,  is the molar absorptivity coefficient (it is constant at a given
wavelength for a particular absorbing substance), b is the width of the light path
that passes through the absorbing substance in centimeters, and c is the molar
concentration of the absorbing species.
In the first part of the experiment, a set of standard solutions of the FeSCN2+
complex is prepared using a concentration of Fe3+ that far exceeds the SCN
concentration. This huge excess of Fe3+ pushes the equilibrium far to the right,
nearly consuming all the original SCN concentration. The reaction is thus
assumed to go to completion. The absorbance of each solution is measured, then
plotted versus the molar concentration of FeSCN2+; this establishes a calibration
curve from which the concentrations of FeSCN2+ are determined for the chemical
systems in the second part.
In the second part of the experiment, you will prepare a set of various solutions
using nearly the same concentration of Fe3+ and SCN ions, and thus have a
number of equilibrium systems. By knowing the initial concentrations of Fe3+ and
SCN, and by determining the equilibrium concentration of FeSCN2+
spectrophotometrically, the equilibrium concentrations of Fe3+ and SCN can be
calculated. Using these equilibrium concentrations, the Kc for the system is
calculated.
Step-by-step Calculation:
Fe3+ (aq)
Initially
+
n0 Fe3+
SCN (aq)
n0 SCN-
 FeSCN2+ (aq)
0
Change
x
x
x
But
___________________________________________________
At equil. (n0 Fe3+  n FeSCN2+) (n0 SCN-  n FeSCN2+)
n FeSCN2+
x = n FeSCN2+
Where: n0 Fe3+ = M0 Fe3+ V0 Fe3+
n0 SCN- = M0 SCN- V0 SCNn FeSCN2+ = x = [FeSCN2+] Vtot = (A/b) Vtot ; A is the measured absorbance for
that particular solution.
Note: In preparing the mixtures for this experiment, the H+ ion concentration will be
maintained at 0.5 M; H+ does not participate directly in the reaction, but its presence
is necessary to avoid the formation of brown color species such as Fe(OH)2+, which
interferes with the analysis of FeSCN2+.
C. Solutions and Setup
Solution A: 2.00  103 M Fe(NO3)3 in 1.0 M HNO3
Solution B: 2.00  103 M KSCN
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Solution C: 0.200 M Fe(NO3)3 in 1.0 M HNO3
Distilled water
D. Procedure (Work in pairs throughout this experiment)
Place 30 mL of solution A (2.00  103 M Fe(NO3)3 in 1.0 M HNO3) in a clean 100
mL beaker. Then in another clean 100 mL beaker, place 20 mL of 2.00  103 M
KSCN (solution B). Finally, in a third beaker, put 40 mL of 0.200 M Fe(NO3)3 in 1.0
M HNO3 (solution C).
Part I:
1. Obtain from the storeroom 12 clean test tubes. Label them 0 to11.
2. Pipet 5.00 mL of solution C into each of the tubes 0 to 5.
3. To the tubes 0 to 5, add 0.00*, 0.50, 1.00, 1.50, 2.00 and 3.00 mL of solution B
respectively. Then add distilled water to each of the tubes so that the final volume
is 20.00 mL in each of them. Deliver the necessary volume of water by means of a
buret.
4. Cover the tubes with parafilm paper and mix the solutions well to make them
homogeneous.
5. Measure the absorbance of each of the prepared solutions (using a
spectrophotometer) at a wavelength  = 447 nm. Record your data in Table 1.
Part II:
6. Pipet 5.00 mL of solution A into each of test tubes 6 to 11.
7. To the test tubes 6 to 11, add 0.00, 1.00, 2.00, 3.00, 4.00 and 5.00 mL of solution
B respectively. Then add distilled water to each of the tubes so that the final
volume is 10.00 mL in each of them. Deliver the necessary volume of water by
means of a buret.
8. Cover the tubes with parafilm paper and mix the solutions well to make them
homogeneous.
9. Determine the absorbance of each of the five prepared solutions, at a wavelength
447 nm. Record your data in Table 2.
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E. CHEM. 203 Lab Report
NAME : _______________________
EXPERIMENT No. : ______________
TITLE : __________________________________
DATE : ________________
Partner : ________________________
Table 1: Standard Solutions and Absorbance Measurements @ =447 nm for the
Calibration Curve
Mixture
V (Fe(NO3)3)
V (KSCN)
V (H2O)
0
5.00
0.00
15.00
1
5.00
0.50
14.50
2
5.00
1.00
14.00
3
5.00
1.50
13.50
4
5.00
2.00
13.00
5
5.00
3.00
12.00
%Transmittance
Absorbance
[FeSCN2+]
Slope of Calibration Curve: b (M1) = __________________
Table 2: Composition of Mixtures and Absorbance Measurements @ =447 nm
Mixture
V (Fe(NO3)3)
V (KSCN)
V (H2O)
6
5.00
0.00
5.00
7
5.00
1.00
4.00
8
5.00
2.00
3.00
9
5.00
3.00
2.00
10
5.00
4.00
1.00
11
5.00
5.00
0.00
%Transmittance
Absorbance
[FeSCN2+]
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Table 3: Equilibrium Concentrations and Equilibrium Constant
Mixture n0 Fe3+
n0 SCN
n Fe3+
n SCN
n FeSCN2+
[Fe3+]
[SCN] [FeSCN2+]
Kc
Show one sample calculation for each calculated quantity (for one mixture only).
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Questions
1. Write the Lewis structure for the SCN ion. Indicate its geometry. What is the
hybridization on the carbon atom?
2. Write an expression for the equilibrium constant of the chemical reaction:
Fe3+ (aq) + H2O  Fe(OH)2+ (aq) + H+ (aq)
3. Consider the chemical reaction:

HClO2 (aq)
H+ (aq) + ClO2 (aq)
What are the equilibrium concentrations of H+, ClO2 and HClO2 in an aqueous
solution with initial concentration [HClO2]0 = 0.260 M, given that the reaction has
an equilibrium constant Kc =1.1  102?
4. Back to the chemical reaction studied in this experiment:
Fe3+ (aq)
+
SCN (aq)
 FeSCN2+ (aq)
We mix 2.00 mL of a 1.00  103 M Fe(NO3)3 with 3.00 mL of a 1.75  103 M
KSCN solution. What are the equilibrium concentrations of Fe3+, SCN and FeSCN2+?
(Use the value Kc = 1.15  102)
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