Complex Lifetimes– Together Forever

Metal complexes are not usually unchanging, everlasting entities. It is true that many once formed, do not react readily. However, the general observation is that coordinate bonds are able to be broken with more ease than cleaving a covalent carbon–carbon bond. The strength of metal–donor bonds are typically less than most bonds in organic molecules, but much greater than the next most stable type, hydrogen bonds. We recognize carbon–carbon bonds as strong and usually unchanging and hydrogen bonds as weak and easily broken;the coordinate covalent bond lies in the middle ground, albeit nearer the carbon–carbon end of the park. Breaking a metal–ligand bond is almost invariably tied to a follow-up process of making another metal–ligand bond in its place, so as to preserve the coordination number and shape of the complex. When the same type of ligand is involved in each sequential process, ligand exchange is said to have occurred. Where the ligands and solvent are one and the same there is no opportunity for any alternative reaction. Even at equilibrium metal complexes in solution display a continuous process of ligand exchange where the rate of the exchange process is driven by the type of metal ion, ligands and/or solvent involved.

Figure 3.27

Metal–ligand preferences for key ligands. In any column a Class A metal ion prefers ligands from the top whereas a Class B metal ion prefers ligands from the bottom.

From the perspective of the central metal, ligand exchange can vary with metal ion from extremely fast (what we refer to as ‘labile’ complexes) to extremely slow (termed ‘inert’ complexes). Whereas direct exchange of one ligand by another of exactly the same type is the simplest process inherently the fact that such exchanges can occur suggests that if other potential ligands of a different type are present, they may intercept the process and be inserted in the place of the original type– ligand exchange has become ligand substitution.

Chemistry is full of ‘opposites’– resulting from often setting two extremes as the limits for defining behaviour (the ‘black or white’ approach). Of course, these limits are not isolated options, but are usually the extremes of a continuum of behaviour (the chemical equivalent of saying that something is rarely ever black or white, but more a shade of grey). How rapidly metal ions take up or lose ligands is a good example of this characteristic. There are two extreme positions; very fast reactions of labile compounds or very slow reactions of inert compounds. Labile systems are the party animals of metal complexes– making and breaking relationships rapidly. We can measure the rate at which ligand exchange with the same type of ligand occurs, even though it appears that nothing changes, by using radioactive isotopes to allow the rate of the process to be monitored, provided the process is not too rapid. It’s a bit like painting a white house with fluorescent white paint– nothing appears to have changed until you see it at night. At the molecular level we simply measure the uptake of radioactive ligand into the coordination sphere of the metal over time as it replaces nonradioactive ligand which allows us to define the rate of ligand exchange.

An inert system is like modern marriage– maybe not joined forever but willing to give it a good try. These are complexes which, once formed, undergo any subsequent reaction very slowly. Inertness can be so great that it overcomes thermodynamic instability. This means that a complex may be pre-disposed towards decomposition, but this will happen so very, very slowly that to all intents and purposes the complex is unreactive, or inert. Cobalt (III) amine complexes are the classic example of this; thermodynamically they are unstable in aqueous solution, but they are so inert to ligand substitution reactions that they can exist in solution with negligible decomposition for years.

Even for a metal ion regarded as inert, however, the rate of substitution or replace ment of particular ligands will differ, and may differ significantly. For example, the half-life for replacement of the coordinated perchlorate ion from the cobalt(III) complex [Co(NH3)5(OClO3)]2+ by a water molecule in acidic aqueous solution is about seven seconds whereas the half-life for replacement of an ammonia from the related [Co(NH3)6]3+ is about 3 800 years! Chemistry is astounding in its diversity, if nothing else. Note one important aspect of the above discussion– lability and inertness are kinetic terms and all about the rate at which something reacts. A species that is inert is kinetically stable. This does not require that it be thermodynamically stable however (although it may be). Thermodynamic stability is about being in a form which has no other readily accessible species lower in energy. The reason we can have kinetic stability in a system that is thermodynamically unstable is that the two differ. Kinetics is about transition to equilibrium; thermodynamics is about the situation at equilibrium. For a molecule to convert from one form to another by any process, it is considered necessary for it to overcome an activation energy barrier whereby it must proceed through a higher energy transition

Figure 3.28

A reaction coordinate defining the barrier to reaction (activation energy) as well as the energy difference between reactants (initial state) and products (final state) or reaction energy.

state (represented in Figure 3.28). Reactants convert to products through this transition or activated state, which is a short-lived ‘half-way house’ between the two forms. A simple macroscopic example may illustrate this. Consider a thin flat piece of board painted white on one side and red on the other. Now arrange to turn this over on a tabletop with the board touching the table throughout. You can do this, of course, by turning the board on to its edge then continuing the motion until it lies flat on its other side. In doing this transformation, the position where it is precariously standing on its edge can be considered a transition (or activated) state, and if you let go of it in that position it may fall back so the white side is up (equivalent to no reaction) or fall forward so the red side is up (equivalent to a completed reaction). To get it from lying flat to on its side costs you effort or energy what we would call the activation energy at the molecular level. Molecules acquire this activation energy mainly through collision as a result of their motion; if the collision energy suffices to take the reactants to an arrangement where they can proceed to products without further energy, they have achieved the status of an activated or transition state species. In our macroscopic example, your physical effort takes the board to the upright activated state. For a very heavy piece of board, the amount of effort is significant and you may not get it into an upright position every time you try. This amounts to a high activation barrier or large activation energy which equates with a slower rate of reaction. If the board is small and light, the effort required is small and the task easy and rapidly completed; this amounts to a low activation barrier or small activation energy, consistent with a fast rate of reaction. If you next consider your board starting on a different level to where it finishes (table to floor, for example), it is obvious that one is the lower level; the board could fall from table to floor, but not easily the reverse way. On the molecular level, being on the ‘floor’ would be a thermodynamically more stable arrangement for the molecule than being on the ‘table’. The difference in energy between the reactants and products is the reaction energy and the stability of a complex depends on the size of this energy difference. This additional consideration (thermodynamics) doesn’t change the process by which you turn the board over (kinetics), although there is a relationship between them at the molecular level, which we won’t explore here. At the molecular level, however, raising temperature increases molecular velocities and increases the probability of a collision achieving the transition state, so the rate of a reaction increases with increasing temperature. Suffice to say that kinetics and thermodynamics, in combination, govern all of our chemical reactions, but obviously tempered by what exactly is available chemically to react.

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مواضيع ذات صلة

Constitutional (Structural) Isomerism

Isomerism– Real 3D Effects

Chameleon Complexes

Moulding a Relationship - Ligand Influences

Metallic Genetics– Metal Ion Influences

Higher Coordination Numbers (ML7 to ML9)

Six Coordination (ML6)

Five Coordination (ML5)

Four Coordination (ML4)

Three Coordination (ML3)

Two Coordination (ML2)

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