A Color Shifting Equilibrium of Cobalt
Samantha Kistler and Ken Overway
Department of Chemistry, Bridgewater College, 402 East College Street, Bridgewater, VA 22812-1599, USA
Introduction:
Transition metal complexes are typically very colorful because
the differences between the d-orbital energy levels match the energy
of visible radiation. Cobalt(II) is a d7 transition metal that changes
color when involved in the following reaction
Results
By measuring the Keq as a function of temperature, eqn. (3) yields
the H and S from the slope and y-intercept.
Others who have studied this reaction have only focused on it as
an example of Le Chatelier’s Principle.1-3 The purpose of this
experiment is to quantitatively determine the entropy and enthalpy
of the cobalt reaction by studying the equilibrium as a function of
temperature.
I- <
Br- <
Smaller
Splitting
Cl- <
F- <
OH- <
H2O < NH3 < en < NO2 <
CO
Larger
Splitting
Figure 1. Pattern of the Spectrochemical Series and absorbed
color.
-Grxn = RT ln Keq
(1)
Since G is very temperature dependent, it can be replaced by
enthalpy (H) and entropy (S)
G = H - TS
(2)
Substituting (2) into (1) yields the van’t Hoff relationship.
 ΔH 1 ΔS
ln K eq 
 
R T R
y  m x b
(3)
Figure 2. Spectral Scans of Aqua and Chloro Complexes of
Cobalt.
Experimental:
Calibration Curve
A calibration curve of standards were prepared from a stock
solution of CoCl2·6 H2O in 12 M HCl (to produce a solution
of only CoCl42-) via a concentration shift.
Temperature Trials
A 138 mM solution of Co(NO3)2·6 H2O was prepared in 6 M
HCl. For hot trials, the cobalt solution, stir bar, and
thermocouple were placed in a cuvette and placed inside a
boiling water bath. When the solution reached 80 ºC or more it
was placed into a spectrophotometer (Vernier LabPro
Colorimeter or Spectronic 20D+), where absorbance was
measured using computer software as the solution cooled to
room temperature. Cold trials involved placing the cuvette
into ice water and repeating the general process.
y = -4465x + 2.5835
R2 = 0.9969
-10.8
-11.2
-11.6
-12
0.0029
0.003
0.0031
0.0032
0.0033
1/Tem perature (1/K)
CN- <
Co(H2O)62+ + 4 Cl- + heat  CoCl42- + 6 H2O
where the octahedral aqua complex is pink and the tetrahedral
chloride complex is blue. Crystal Field Theory explains the color
change using what called the spectrochemical series (Figure 1).
Since H2O is a moderately strong ligand, the [Co(H2O)6]2+ complex
will have a larger splitting energy that the [CoCl4]2- complex. The
color observed is the complementary color of the absorbed color, so
if [Co(H2O)6]2+ has a larger splitting energy, its absorbed
wavelength will be bluer and its observed wavelength will be redder
compared [CoCl4]2-.
The reaction under investigation is also temperature dependent.
Adding heat shifts the reaction to the blue complex and removing
heat shifts it to the pink complex (Figure 2). Since temperature can
shift the reaction equilibrium, is it possible to measure some
thermodynamic parameters of this reaction, namely entropy and
enthalpy. The equilibrium constant (Keq) is related to the Gibb’s
Free Energy of the reaction via
-10.4
ln Keq
Abstract
Transition metal complexes are typically very colorful because
the differences between the d-orbital energy levels match the energy
of visible radiation. This makes studying their chemical properties
easy using spectroscopic methods.
In this experiment the
equilibrium between two cobalt complexes was studied in order to
determine thermodynamic properties of the reaction.
The
quantitative results confirmed that the reaction was endothermic and
increases entropy, as expected by inspecting the reaction. The
results seem to also vary in some ways and are consistent in others,
indicating that an experimental parameter is not consistent, however
the method was stable enough to be successfully employed in the
General Chemistry II lab.
Figure 3. Plotting ln(Keq) vs. 1/T to determine H and S, where
the slope = -H/R and S/R (where R = 8.3145 J/K mol).
Table 1. Pooled results from researchers and class data.
condition§
VC, from cold to RT
VC, from cold to RT
VC, from hot to RT
VC, from hot to RT
Spec 20, from cold to RT
Spec 20, from cold to RT
Spec 20, from hot to RT
Spec 20, from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Class Data (VC), from hot to RT
Enthalpy (J/mol)
H
error
22,834  0.28%
23,953  0.47%
37,124  0.16%
32,220  0.10%
70,322  0.60%
89,392  0.79%
35,880  0.37%
35,346  0.33%
72,131  0.10%
47,187  0.03%
41,668  0.05%
51,853  0.15%
43,146  0.13%
41,756  0.11%
46,153  0.05%
68,959  0.08%
46,825  0.04%
40,390  0.04%
46,825  0.32%
40,096  0.18%
29,681  0.52%
Entropy (J/K mol)
S
error
-28.19  -0.09%
-22.91  -0.20%
21.48  0.10%
5.86  0.20%
128.08  0.13%
193.71  0.15%
15.63  0.33%
14.44  0.31%
142.68  0.16%
73.37  0.06%
46.74  0.12%
87.27  0.05%
62.35  0.06%
59.32  0.11%
72.52  0.10%
134.69  0.12%
71.10  0.08%
42.68  0.11%
71.10  0.08%
53.60  0.05%
23.33  0.23%
§ VC=Vernier Colorimeter, RT=room temperature
Discussion
The results showed consistency for hot trials (red text) using either
spectrophotometer. Cold trials (blue test) show inconsistency from hot
trials and between spectrophotometers. Class data ranged widely. Antifogging measures were employed for cold trials to no avail. Future
work will require the tightening of experimental parameters (6.0 M
HCl standardization, stirring) and a continuous cold to hot trial from
10 °C to 80 °C. Unfortunately the results cannot contribute to
thermodynamic tables for CoCl42- because the concentration of
chloride is so large (6 M) that H and S values cannot be considered
to be standard state, which requires 1 M concentrations for all
chemicals.
References
1. Grant, A.W. J. Chem. Educ. 1984, 61, 446.
2. Ophardt, C.E. J. Chem. Educ. 1980, 57, 453.
3. Barrera, N.M.; McCarty,J.L.; Dragojlovic, V. Chem. Educator 2002,
7, 142-145.