Bonding in Coordination Compounds

Bonding In Coordination Compounds

Werner’s theory was the first practical attempt to define the nature of bonding in coordination compounds, and it remains the most widely accepted today. However, the Valence Bond Hypothesis, introduced by Linus Pauling in 1931 and is the most straightforward theory that explains the nature of bonding inside the coordination sphere, is the most straightforward. 

According to the theory:

• The core metal must lose a sufficient number of electrons (equivalent to its valency).
• For covalent (coordinate) bonds to form with the ligands, the cation must have equivalent unoccupied orbitals as the number of coordination units in its ionisation energy.
• It is known as hybridised orbitals when a new set of equivalent bonding orbitals with specified directional features is generated by combining the cation orbitals of a cation, such as s, p, or D (depending on the total number of bonds that will be made with the ligands) and a ligand.
• Those metal electrons present in the inner d-orbitals that are not capable of bonding do not participate in hybridization (i.e., in chemical bonding).
• The vacant hybrid orbitals of the cation can overlap with those of the ligand orbitals, which can donate an electron pair for bonding, resulting in the formation of coordinate bonds between the metal and ligand. In this case, the number of bonds created equals the empty orbitals provided by the centre metal ion, the cation.
• Depending on the situation, the d-orbitals engaged in hybridization can be inner (n-1)d orbitals or outer ‘and’ orbitals. Inner orbital or low spin complexes are those generated by utilising inner d-electrons. Those formed using outer d-electrons are known as Outer Orbital or high spin complexes. Inner orbital or low spin complexes are those formed using inner d-electrons (because they contain more unpaired electrons).
• The complexes that contain unpaired electrons are paramagnetic, whereas those that do not contain unpaired electrons are diamagnetic.
• Because of the action of a potent ligand, it is possible to drive electrons to couple up in opposition to Hund’s law of maximum multiplicity.

Bonding In Coordination Compounds: Theories

Valence bond theory in coordination compounds

Coordination compound bonding is presently the subject of a slew of hypotheses. American physicists Linus Pauling and John C. Slater devised an important component of the valence bond theory for coordination compounds, which explains bonding by examining how many empty electrons the metal ion’s hybridised orbitals have. A coordinate-covalent bond is formed when an unfilled metal ion orbital and a filled ligand orbital overlap, and each ligand provides an electron pair to do so.

As an example, sp (linear), sp3 (tetrahedron), dsp2 (planar), and D2sp33 (octahedron) are all examples of hybridizations in which orbitals of a given type and number participate in the hybridization. In many cases, the number of unpaired electrons measured is in agreement with the theoretical forecast. By differentiating between inner orbital complexes (d2sp3) and outer orbital complexes (sp3d3), Henry Taube, a Canadian-born American Chemistry Nobel winner, further developed the theory in 1952 to account for discrepancies between octahedral complexes.

The absence of antibonding molecular orbitals formed during complex formation is a major flaw in the basic valence bond theory for coordination compounds. As a consequence, the vivid hues of many complexes, which come from their selective absorption of just certain wavelengths of light, cannot be explained in this way. After World War II, VB theory was used to explain almost all coordination occurrences. Chemists at the time were worried about geometric and magnetic susceptibility issues, and this response was obvious.

The primary shortcoming of the simple valence bond theory for coordination compounds is the omission of antibonding molecular orbitals generated during complex formation. Thus, it cannot account for the vibrant colors of many complexes resulting from their selective absorption of only specific wavelengths of light. From the early 1930s through the early 1950s, VB theory was utilized to explain practically all coordination phenomena. It provided straightforward answers to the geometric and magnetic susceptibility concerns that concerned chemists at the time.

Limitations of valence bond theory for coordination compounds:

1. It was unable to explain the precise magnetic characteristics of the compounds in detail.
2. It could not provide an adequate explanation for creating low spin and high spin complexes.
3. It does not provide a precise explanation for the stability of coordination molecules in this context.
4. It makes no distinction between ligands with a weak field and those with a strong field.
5. In terms of the color of the complexes, there is no explanation given in the text.

Crystal Field Theory: bonding in coordination compounds

The Crystal Field Theory (CFT) is a model for transition metal and ligand binding interactions. It depicts the impact of the ligand’s non-bonding electrons being attracted by the positive charge of the metal cation and the negative charge of the metal cation. Due to the static electric field produced by a surrounding charge distribution, the degeneracy of electronic orbital states, commonly d or f orbitals, is broken as the ligands approach the core metal ion. CFT can account for various magnetic characteristics, hues, and hydration energies of transition metal complexes, but it fails to account for bonding.

Due to repulsion between like charges, the electrons in the d orbitals of the core metal ion and those in the ligand repel each other. As a result, the d electrons closest to the ligands will have enormous energy than those further away, dividing the energy of the d orbitals. This division is influenced by:

1. The metal ion’s composition.
2. The oxidation state of the metal (a higher oxidation state leads to a more extensive splitting).
3. The ligands’ arrangement around the metal ion.
4. The ligands that surround the metal ion’s nature.

Excluding the dz2 orbital, which has two opposed lobes and a doughnut of electron density in the center, all d orbitals have four lobes of electron density. The d orbitals can be subdivided into two groups. The dx2–y2 and dz2 points are parallel to the x, y, and z axes. They make up an, e.g., set. The lobes of the dxy, dxz, and dyz, on the other hand, all line up in the quadrants, with no electron density on the axes. The t2g set is made up of these three orbitals. The d orbitals are usually degenerate, but they can sometimes split, with the, e.g., and t2g subsets having differing energies. The CFT explains this.

The stability that occurs from placing a transition metal ion in the crystal field created by a collection of ligands is the crystal field stabilization energy (CFSE). It occurs because when the d orbitals in a ligand field are split, some of them become lower in energy than previously. For example, in the instance of an octahedron, the t2g set has lower energy. As a result, the metal ion is more stable in the ligand field by the amount known as the CFSE if any electrons occupy these orbitals. The energy of the, e.g., orbitals, on the other hand, is higher. As a result, the amount of CFSE is reduced by injecting electrons into them.

Transition-metal complexes of all geometries can benefit from crystal field stabilization. The high amount of crystal field stabilization that this shape produces with this number of electrons explains why many d8 complexes are square-planar.

Conclusion: 

Werner’s concept was the first to define the nature of bonding in coordination compounds, and it is still widely accepted today. Valence bond theory is a synthesis of early covalent bond theories. A crucial concept in valence bond theory is resonance. Several types of orbitals can be bonded together to form hybrid orbitals. Valence bond structures are utilized when Lewis structures cannot adequately reflect the many resonance-induced coordination molecules. When ligands approach the core metal ion, the static electric field breaks d- or f-subshell degeneracy. Determining the energy of d electrons near the ligands causes the d orbitals to divide. The CFSE is the stability that occurs from ligand binding.