Reactions Involving the Metal Oxidation State

Because many complexed metals ions have a range of oxidation states accessible in usual solvents it is not surprising to find that syntheses may involve oxidation–reduction reactions. The classical example is cobalt, which is usually supplied commercially as Co(II) salts, but whose complexes are best known as Co(III) compounds. This is because Co (III) compounds are inert and usually robust readily isolable compounds, whereas Co (II) compounds are labile and prone to rapid reactions. This lability is put to good use in synthesis, since it is convenient to use the Co(II) form to rapidly coordinate ligands initially and then oxidize the mixture to the Co(III) form. This oxidation can often be achieved by oxygen in the air alone, depending on the ligand environment and the redox potential (E) of the complex. For example in an aqueous ammonia/ammonium chloride buffered solution, Co(II) reacts in a sequential manner essentially as in (6.26).

Substitution reactions without redox chemistry being involved are available for Co(III) but not commonly met. One example is the use of [Co (CO3)3] as a synthon since the chelated carbonate ion is readily displaced by other better chelating ligands such as polyamines. An example of this type of reaction, where the entering polyamine (cyclam) is a saturated and flexible macrocyclic tetraamine is (6.27).

Where stronger oxidizing agents than air are required to take the Co(II) form to Co(III), which usually applies where fewer N-donor and more O-donor groups are bound, hydrogen peroxide is a particularly useful oxidizing agent, because it leaves no problematical products to separate from the desired complex product. One of the problems with oxidation of Co (II) to Co(III) compounds is that in some cases a bridged peroxo complex, featuring a CoIII-O₂₂-CoIII linkage forms as a stable intermediate. One can envisage its formation through a redox reaction whereby an oxygen molecule is reduced to peroxide ion, and two Co(II) ions are oxidized to Co (III). The O₂ ion in the bridge can be readily displaced by reaction with strong acid and heating, with the use of hydrochloric acid leading to monomers with Coll-Cl- components replacing the bridging group (6.28).

Many other metal complexes can be chemically oxidized to higher oxidation states with an appropriate oxidizing agent. This can occur with complete preservation of the coordination sphere (6.29) which means the oxidation-reduction reaction is reversible if a suitable reducing agent is then employed for reduction of the oxidized form.

Alternatively, the coordination number and/or the ligand set can change substantially in an irreversible oxidation reaction. Any reduction reaction of the product will then not regenerate the original complex.

Not only are oxidation reactions fairly common, but also one may employ reduction reactions in simple synthetic paths provided the reduced form is also in a stable oxidation state. Two examples of reduction reactions with the metal oxidation states included, are given in (6.30) and (6.31).

Reduction reactions occurring along with substitution chemistry are also well known and the examples above are such cases. Another simple example involves the [IrCl] ion, which undergoes both reduction (with hypophosphorous acid) and substitution in (6.32).

Because of the slow substitution chemistry of iridium, sequential reaction steps are well separated in terms of reaction time so that the initial product of the redox-substitution reaction undergoes further simple substitution only with prolonged heating.

Where a reactive lower oxidation state results, a key concern is the necessary protection of the reduced complex from air or other potential oxidants, as they are often readily re- oxidized. Usually, this requires their handling in special apparatus such as inert-atmosphere boxes or sealed glassware in the absence of oxygen. Where active metal reducing agents (such as potassium) are employed, special care with choice of solvent is also necessary. The nickel reduction reaction (6.33) can be performed in liquid ammonia as solvent, since the strongly-bound cyanide ions are not substituted by this potential ligand.

While the Ni(II) is reduced to Ni(0) the potassium metal is oxidized to K(I); it is a standard redox reaction. The Ni (0) complex formed exemplifies the necessity for careful handling of many low-valent complexes; it is readily oxidized by air, and also reacts with water in a redox reaction that liberates hydrogen gas.

Lower-valent metal complexes may be prepared in reduction reactions with full substitution of the coordination sphere. One example results from reduction of the vanadium (III) complex [VCl3(THF)₃] with a suitable reactive metal in the presence of carbon monoxide under pressure (6.34). This is an example of the synthesis of an organometallic compound; more examples occur in Section 6.6.

On some occasions a reducing agent may not be able to reduce the metal centre in a complex, but may be sufficient to initiate chemical change in a ligand. This is, in effect, a reaction of a coordinated ligand (see Section 6.5.2), rather than a metal-centred redox reaction. A simple example involves the reduction of coordinated N₂O to N₂ and H2O using Cr(II) ion as reducing agent (6.35); the metal ion retains its original oxidation state throughout. The reaction is promoted through coordination lowering the N-O bond strength.

Many complexes can be prepared by electrochemical reduction or oxidation under an inert atmosphere. Exhaustive reduction (coulometry) using a large working electrode can lead to a clean formation of the reduced form, if it is sufficiently stable, since electrons alone are used in the reduction process. As an example, the octahedral vanadium(III) complex [V(phen)3]2+ undergoes a series of sequential and reversible reduction steps (6.36) with retention of the three chelate ligands down to oxidation state of formally vanadium (-I).

With unsaturated ligands particularly, the location of introduced electrons is not easily assigned, and they may reside on the ligand (leading to a ligand radical) rather than the metal ion (leading to a lower oxidation state). This does not invalidate the chemistry, but does bring into question the nature of the product. Overall, electrochemistry provides a useful way of performing not only reduction reactions but also oxidation reactions. The sole concern is that exhaustive electrolysis to convert all of one form to another oxidation state takes some minutes to perform, so the kinetic stability of the product does influence the validity of the method.

When undertaking electrochemical reactions, it is important to note that the E for a complex is dependent on the set of ligands coordinated, and can change substantially with donor set. This is represented in (6.37) for the simple substitution of a Maq complex with the potential required to reduce the precursor complex invariably differing from that for the product complex (E°1 ≠ E°2).

The shift in redox potential with ligand substitution is particularly obvious for cobalt as the E for Coq is near +1.8 V making the aqua Co(III) complex inaccessible from Co2+aq with most oxidants whereas introduction of other donor groups lowers the potential substantially and in some cases sufficiently to make even air (oxygen) an effective oxidant of the Co(II) complex.

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مواضيع ذات صلة

Catalysed Reactions

Chiral Complexes

Substitution Reactions without using a Solvent

Substitution Reactions in Nonaqueous Solvents

Reactions Involving the Coordination Shell

Block f : Inner Transition Metals (Lanthanoids and Actinoids)

Block d: Transition Metals

Block p: Main Group Metals

= Block s : Alkali and Alkaline Earth Metals

A New Face - Oxidation-Reduction

A New Body - Stereochemical Change

Lability and Inertness in Octahedral Complexes