1
Historical Background

Anmol Singh, Himadri Priya Gogoi, and Pranjit Barman

National Institute of Technology Silchar, Department of Chemistry, Silchar, Assam, 788010, India

1.1 Introduction

Over the past few decades, chemistry of coordination compounds has expanded and diversified considerably. Coordination chemistry accomplished a significant interdisciplinary place at the interface of numerous fields, which serves as a connecting link between the different classes of chemical, physical, and biological sciences. The precise date that the first coordination compound was prepared is challenging to ascertain. Johann Jacob Diesbach, a color manufacturer, discovered the blue pigment known as Prussian blue in the seventeenth century, which is the beginning of coordination complex chemistry [1]. Chromium was discovered in 1798 by the German researchers Lowitz and Klaproth separately in a sample of a heavy black rock located near north of the Beres of Mines. The chromium in the same mineral from a small deposit in the Var region of South-eastern France was discovered by another German chemist Tassaert in 1799. He identified this mineral as the Cr-Fe spinel, also known as chromite (FeO·Cr2O3) [2].

Alfred Werner's findings marked the beginning of an insight into coordination compounds and their varied properties. Werner's coordination theory in 1893 was the first effort which focused on the quantity and type of groups connected to the central metal ion to characterize the bonding and interpretation of coordination complexes. In 1913, Werner's idea and meticulous work over the ensuing 20?years earned him the Noble Prize in Chemistry. Due to the development of sophisticated physicochemical techniques of high accuracy and precision, his fundamental theories regarding the stereochemistry of metal complexes, mechanisms of isomerization, and racemization, etc. prevail indisputable even now despite all the advancements that have occurred since his time and through the last 50?years [3, 4]. These have greatly improved our knowledge of the attributes of metal-ligand connection, the configuration and stereochemistry of metal complexes, and their stabilities, liabilities, and other characteristics [5, 6]. Werner's theory became a key component of "The Electronic Theory of Valency" proposed by G.N. Lewis (1916) and extensively related to coordination compounds by N.V. Sidgwick (1927), replacing the majority of the earlier concepts presented by Berzelius (1819), Grohen (1837), Claus (1856), Blomstr (1869), and Jorgenson (1894).

Coordination chemistry has expanded in three dimensions, taking into consideration its breadth, depth, and applications. The five Noble awards that significantly impact the topic reflect the sustained respect for the emerging science (A. Werner, 1913; M. Eigen, 1967; Wilkinson and Fischer, 1973; H. Taube, 1983; Cram, Lehn and Pedersen, 1987). In the list, Werner and Cram, Lehn & Pedersen acknowledged the ancient and new domains of coordination chemistry [7].

A coordination compound, also known as a metal complex, comprises an ensemble of bonded ions or molecules, regarded as ligands, and a central ion or atom called a coordination center, which is typically a metal ion. There are reasonably profound changes in the properties of a metal ion on the complex formation, which reflects in structures, stereochemistry, stability, and many other properties. It is necessary to understand the properties of bond between the metal ion and ligand to interpret the design of complex.

1.2 Theories of Coordinate Bond

1.2.1 Valence Bond Theory

According to Lewis and Sidgwick theory, ligands donate an electron pair to a metal ion to establish a coordinate bond [8, 9]. In 1931, L. Pauling expanded this theory to comprehend and predict various coordination complex traits, including magnetic behavior, stereochemistry, kinetics, and other physical and other chemical characteristics [10]. This hypothesis states that a coordinate bond between the metal and ligand leads to the complex formation of the duo. It describes how s, p, or d orbitals, as well as hybridized orbitals, i.e. sp3, d2sp3 or sp3d2, or dsp [2], etc. generate strong bonds. Tetrahedral, octahedral, and square planar stereochemistry, each dealt by coordinate covalent bonds, formed by metal atoms having sp3, d2sp3/sp3d2, or dsp [2] hybrid orbitals with the ligand orbitals. Through semi-quantitative calculations, Pauling and Gould elucidated the stability of complexes and concluded that the formation of sigma bonds among ligand and metal from electron donation occurs in carbonyl and nitrosyl compounds [10]. Back donation, i.e. dp-pp linkage amid metal and ligand, justifies the stability of carbonyl and nitrosyl compounds.

1.2.2 Crystal Field Theory

The splitting of d or f orbitals is explained by crystal field theory (Figure 1.1). Degeneracies of electron orbital states are typically brought on by a static electric field induced by the distribution of charges nearby. Hans Bethe and John Hasbrouck van Vleck, two scientists, proposed a theory that describes several spectroscopies of transition metal coordination complexes, specifically optical spectra (colors) [11, 12]. Although it makes an effort to characterize and explain bonding, crystal field theory (CFT) is effective in defining colors, magnetic characteristics, and hydration enthalpies of transition metal complexes. The attraction among the positively charged metal cation and ligands leads to the interaction between a transition metal and ligands. Owing to the attraction between like charges, the electrons in the ligand's d-orbital resist one another, splitting the energy of the d-orbital. The following factors have an impact on this splitting:

Figure 1.1 Crystal field splitting in tetragonally elongated, octahedral, and square planar filled.

1. Concerning the spherical field, the distribution of ligands encircling metal, the coordination number of the metal ion, such as tetrahedral or octahedral, and the type of ligands, a greater oxidation state of the metal causes considerable splitting.
2. Tetrahedral symmetry hinders the ligand electrons from approaching d-orbitals directly. The higher energy orbitals are dz2 and dx2-y [2], while lower energy orbitals are dxy, dyz, dxz. The energy splitting will be less than in an octahedral field. CFT can also be used to explain complex geometry.

1.2.3 Molecular Orbital Theory

The electronic structure of molecules is described using quantum mechanisms. The electrons are not allocated to specific atom-to-atom linkage; instead, they are regarded as moving under the influence of the nuclei in the entire molecule. In 1935, molecular orbital theory (MOT) stated that molecular orbitals were formed by linear combinations of atomic orbitals. Atomic orbitals must overlap within space and cannot form molecular orbitals if they are too apart. The electrons in an atom may be either waves or particles; therefore, electrons can be considered as accommodating an atomic orbital or a wave function ?, which is a solution of Schrodinger wave equation. The overlapping atomic orbitals must have nearly the same energy; maximum overlapping must have symmetry concerning the molecular bonding axis [13, 14].

1.2.4 Ligand Field Theory

The orbital configuration, bonding, and other properties of coordinating complexes are explained by ligand field theory (Figure 1.2). To formulate the more extensive and realistic ligand field theory, which describes the process of chemical bonding in transition metal complexes, the MOT and CFT were later integrated. This theory which elucidates the loss of degeneracy of metal d-orbitals in transition metal complexes was developed by incorporating the tenets of CFT and MOT. Ligand field theory originated in the 1930s by John Hasbrouck Van Vleck and Orgel to explain the results using CFT and visible spectra of transition metal complexes [12,15-18]. The vacant d-orbitals of transition metals facilitate bonding, which affects the colors they absorb in solution. Depending on the intensity of the surrounding ligands, the various d-orbitals are influenced differentially and have their energy raised or diminished. The symmetry orbitals of the ligands in octahedral complexes generate bonding and antibonding combinations through the dz2 and dx2-y2 orbitals as a result of approaching along the x-, y-, and z-axes. In contrast, the dxy, dxz, and dyz orbitals remain nonbonded [19, 20].

Figure 1.2 Molecular orbital energy-level diagram of an octahedral complex.

References

1. 1 Kraft, A. (2008). On the discovery and history of Prussian blue. Bull. Hist. Chem. 33 (2): 61-67.
2. 2 Kauffman, G.B. (1977). Il'ya Il'ich Chernyaev (1893-1966) and the trans effect. J. Chem. Educ. 54 (2): 86.
3. 3 Banerjee, D., Ray, U., Jasimuddin, S.K. et al. (2006). Copper(II) complexes of thioether containing an azoimidazolyl system. X-ray structure of Cu (SEtaaiNEt) Cl2 (SEtaaiNEt...