Four Coordination (ML4)

Coordination number four (ML4) is common and has two major forms, tetrahedral and square planar. The former is the shape predicted by the electron pair repulsion model; the latter is a different shape observed experimentally with many examples known. These are ideal or limiting structures in the sense that they represent the perfect shapes which lie at

Figure 4.7 Examples of complexes with trigonal planar or the rarer T-shaped geometry.

the structural limits for this coordination number; as mentioned already, ideal structures are relatively rare in coordination chemistry, and distorted or intermediate geometries are more likely met, so named since they can be achieved by distorting one or other shape partially towards the other class. The two limiting geometries can be converted one into the other by displacement of groups without any bond breaking being involved, which is energetically less demanding since bond angle deformations with their lower energy demand than bond breaking are then the dominant energy ‘cost’.

The shapes of the two limiting and an intermediate geometry are shown in Figure 4.8, both based on a cubic box frame. The tetrahedral shape is defined by placing two donors on opposite corners at the top of the box, and the other two donors on opposite corners at the bottom of the box, but bottom corners that do not lie directly beneath the corners occupied at the top of the box. This geometry can be converted to the square planar shape by simply ‘sliding’ the top two donors down two edges of the cubic box and ‘sliding’ the other two bottom donors up the other two edges of the box until they are all half-way along the edges. When this occurs, they are all co-planar, and also co-planar with the metal ion placed at the centre of the box; this is, in effect, the square planar shape. If the ‘sliding’ is stopped part way, then an intermediate distorted shape is achieved. In reality all types are well known;

Figure 4.8: The two limiting shapes for four-coordination, tetrahedral and square planar along with an intermediate geometry formed in transition from one limiting shape to the other.

for example, for simple [MCl₄]₂− ions, d7 Co (II) is tetrahedral, d8 Ni (II) is an intermediate geometry, and d9 Cu (II) is square planar in shape. Many complexes described as square planar display small tetrahedral distortions that place the two pairs of donors slightly above or below the average plane including the metal ion; likewise, many tetrahedral complexes display small distortions towards square planarity. Where these distortions are minor it is convenient to ignore them in defining the basic shape.

It is also often convenient to represent shapes in terms of their actual symmetry expressed as the appropriate mathematical ‘label’– Td for tetrahedral, D4h for square planar and   D2d for intermediate geometries in this case– as this defines the shape succinctly and is appropriate for application in spectroscopy. There are simple rules for deciding the symmetry of a complex, and these are described and exemplified in Appendix B.

Tetrahedral or distorted tetrahedral geometries are, from experimental observations

Dominantly found in complexes that are overall neutral or anionic. Simple examples include [CuX₄]₂-[FeX₄]₂- and [CoX₄]₂- (X = halogen anion). Otherwise, it is d0 compounds (such as [TiCl₄]) or d10 compounds (such as [Ni (PF3)₄] and [Ni (CO)₄]) that lean towards tetrahedral geometry. Other dn configurations (except d3) exhibit limited examples and then only usually for the first row transition metals. Ligand arrangement in the tetrahedral geometry minimizes inter-ligand repulsions, so negatively charged ligands prefer this shape because of the more favourable charge-based repulsions when adopting this shape. However this shape does lead to a weak ligand field when compared to what occurs in square planar systems, which can make this geometry less effective. With a balance between favourable repulsions and unfavourable ligand field effects, it is not surprising to find that steric effects or ligand size are an important consideration in this geometry.

Square planar complexes mainly used as examples are those having a d8 metal ion such as Rh(I). Ir(I). Pd(II). Pt (II) and Au (III). Several examples are shown below (Figure 4.9).

One consequence of square planarity is clear in these examples–there are structural isomers possible. The neutral [PtCl₂(NH₃)₂] exists in two geometric isomers, trans (where each pair of groups is as far apart as possible on opposite sides of the molecule) and cis (where the two pairs occupy adjacent sites). Like all geometric isomers, they display distinct chemical and physical properties, including biological properties; the cis isomer is otherwise known as the anti-cancer drug cisplatin but the trans isomer is not effective as a drug.

That square-planarity is common for d8 complexes does not mean that this shape is not  seen for other metal ions, and indeed it is reasonably common amongst metals with from d6 to d9 configurations, It is also important to note that not all d8 systems are square planar. Some metal ions are ambivalent, with the ligand field playing a clear role in the outcome;

Figure 4.9

Examples of four-coordinate square planar complexes.

Figure 4.10

Examples of d8 nickel (II) complexes adopting square planar or tetrahedral geometry, depending on the type of ligand. The square planar complex has no unpaired electrons (diamagnetic) whereas the tetrahedral complex has two unpaired electrons (paramagnetic), allowing easy identification (energy splitting diagrams not to scale).

the d8 Ni(II) ion can formoctahedral, tetrahedral and square-planar complexes, with strong f ield ligands tending to favour square-planarity. They are distinguished readily from their colours and spectroscopic properties; for example, octahedral NiN6 complexes are purple high-spin species and square-planar NiN4 complexes yellow low-spin species (here, N represents an N-donor ligand, not an isolated atom). This ambivalence is a reminder that the square planar geometry can be reached by removal of two trans-disposed ligands from around an octahedron, as well as by bond angle distortion of a tetrahedron. The structural ambivalence of d8 Ni(II) is also exemplified in Figure 4.10, where different ligands lead to different geometry; apart from spectroscopic differences, magnetic properties of square planar and tetrahedral complexes differ, the latter alone having unpaired electrons. Since square-planar complexes are essentially two-dimensional compounds, it is not a surprise to find other molecules coming into reasonably close but still nonbonding po sitions (distances 300pm) above and below the plane in the solid state and in solution, including metal–solvent and even some cases of metal–metal interaction. The square planar cation–anion pair [Pt(NH₃)₄]₂+[PtCl₄]₂− (called in earlier times, Magnus’ green salt) for example, stack cationic and anionic complexes in alternating positions in their crystalline lattice, with a M ...M separation of ∼325pm (Figure 4.11). The separation is close enough to have an effect on physical properties, with the green colour indicating some weak metal–metal interaction, since another form of the salt with a very long separation of over 500pm and no semblance of interaction is pink, colour arising  

Figure 4.11 Stacking of a d8 platinum (II) complex in the solid-state leads to weak Pt···Pt interaction, influencing physical properties.

Figure 4.12

An NO-chelate ligand whose Ni (II) complex displays interconversion between different geometries depending on conditions.

in that case simply from combination of the colours of the two independent ions. When metal ions come very close together, this interaction will also affect the magnetic properties of their assembly significantly. In describing the relationship between tetrahedral and square planar complexes above, we used a model where interconversion could occur without bond breaking. If this is a fair representation of reality, then it should be possible to find some systems that exist either as a mixture of the two forms in equilibrium or can convert between the two forms when a change in conditions is applied. Fortunately, there are indeed some compounds that undergo conversion between tetrahedral and square planar forms in solution, a situation which implies that the stabilities of the two forms are very similar. One now classical example involved the Ni(II) complex of a chelated N,O-donor ligand shown in Figure 4.12, where conversion between dominantly tetrahedral and dominantly square planar forms depends on the temperature, solvent and the type of R-group attached to the coordinated imine nitrogen. Change between the two forms can be readily monitored since their colours and absorption band positions and intensities are different.

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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)

Three Coordination (ML3)

Two Coordination (ML2)

Complex Lifetimes– Together Forever

?Preferences– Do You Like What I Like