| dc.description.abstract | Nature of Mononuclear Complexes in Cobalt–Ethanolamine Systems
The work presented in this thesis indicates that cobalt forms both pure and mixed hydroxy complexes. The formation constants of complexes such as CoM²?, CoM?²?, CoM?²?, CoM?²?, CoM?OH?, CoMOH?, CoD²?, CoD?²?, CoDOH?, CoT²?, CoTOR?, and CoT(OH)? have been obtained by graphical analysis of the n–pA curves in the presence of 0.1 M, 0.5 M, and 1.0 M ethanolammonium ion. The step constants for individual pure complexes are given in Table VIII.1. For comparison, the step constants of ethanolamine (mono-, di-, and tri-) complexes of nickel and copper obtained in this laboratory, as well as ammonia complexes, have been included.
Comparison of the step constant of CoM²? with that of Co(NH?)²? indicates that the constant for CoM²? is slightly greater, suggesting a slight chelate effect. However, the step constants for CoM?²? and CoM?²? are lower than the corresponding ammonia complexes, indicating no significant chelate effect. Being a primary amine, the formation constants are lower than those of the corresponding ammonia complexes. In nickel and copper complexes, the chelate effect is observed up to ML? and ML? respectively. It should be noted that it would have been more appropriate to compare ethylamine complexes instead of ammonia complexes. Unfortunately, the constants for cobalt–ethylamine complexes are not available. However, it is expected that the constants of cobalt–ethylamine complexes would be lower than those of cobalt–ammonia complexes. Therefore, the conclusions drawn are considered valid.
The formation constants of cobalt–diethanolamine and cobalt–triethanolamine are also included in the same table. Diethanolamine and triethanolamine are secondary and tertiary amines. It is not possible to draw conclusions regarding the chelate effect based on step constants since the corresponding constants for cobalt–diethylamine and cobalt–triethylamine are not available in the literature.
It is expected that with an increase in the number of ethanol groups in the amine, the successive formation constants of all complexes should decrease. In the present study, this is true for monoethanolamine and diethanolamine complexes. In the case of the cobalt–triethanolamine system, where only the first complex is produced, the step constant is slightly higher than CoD²?, possibly due to a stronger chelating effect, as observed in the nickel–triethanolamine system.
Crystal field splitting provides extra stability to the system. In a weak octahedral field, the ligand field stabilization effect follows the order:
d? < d? < d? < d?
Here, d? occupies a higher position than d? due to the Jahn–Teller effect. Irving and Williams and Mellor and Maley have shown experimentally that the stability constants of similar complexes follow the order:
Mn²? < Fe²? < Co²? < Ni²? < Cu²? > Zn²?
The data presented in Table VIII.1 indicate that the ethanolamine (mono-, di-, and tri-) complexes of cobalt, nickel, and copper studied in this laboratory follow this trend.
Nature of Hydroxy Complexes
Another important result presented in this thesis is the use of Bjerrum formation curves to obtain evidence for the formation of hydroxy complexes and their stability constants. In Table VIII.2, the mixed hydroxy complexes obtained and their stability constants are given:
Mixed Hydroxy Complexlog Stability ConstantCoMOH?6.80CoM?OH?8.67CoDOH?7.72CoT(OH)?14.83
It is interesting to compare the behavior of cobalt with nickel. Under similar conditions, nickel does not form mixed hydroxy complexes. Since nickel is stabilized by the ligand field to a greater extent than cobalt, one might expect greater ease of hydrolysis of cobalt–amine complexes. This possibly explains the difference between cobalt and nickel.
There are several differences between mono-, di-, and triethanolamines, and there is a gradation in the tendency for the formation of mixed hydroxy complexes. The tendency follows the order:
monoethanolamine < diethanolamine < triethanolamine
In the case of monoethanolamine, in the presence of 1 M monoethanolammonium ion solutions, it is possible to go up to n = 4, and evidence for the formation of MA?OH? is obtained only between i = 3.4 and 3.8. For 1 M diethanolammonium ion solutions, pure complexes are produced only up to i = 1.4. Beyond this, analysis is not possible since the i–pA curves become very steep due to the formation of higher mixed hydroxy complexes. In the case of 1 M triethanolammonium ion solutions, MA(OH)? is produced in the i range 0.85 to 1.2. Beyond this range, analysis is not possible due to a steep rise in the curves, evidently caused by the formation of higher mixed hydroxy complexes. | |
