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2.6 Pi Donor and Acceptor Ligands

The identity and thus unique properties of a metal complex is the result of the interaction, as modeled using molecular orbitals, between coordinated ligands and the metal center. This interaction and thus shape of orbitals is governed by the nature of the metal center's electron density and the ability of the ligand to donate or accept electrons to the metal center that will determine the molecular orbitals. This system can be simplified by looking at the metal center's oxidation state and electron count and by categorizing ligands based on their donating/accepting capacities.

The spectralchemical series shows the trend of compounds as weak field to strong field ligands based on how they affect the difference in energy between e_g and t_2g energy levels of the molecular orbitals, causing a large split or a small split. Furthermore, ligands can be characterized by their π-bonding interactions.

Examples of Weak Field Ligands:

X^-, OH^-, H₂O

Examples of Strong Field Ligands:

H^-, NH₃, CO, PR₃

The affect of both π donor and acceptor ligands on the d-orbitals of the metal center is illustrated in this MO diagram.

In a π-donor ligand, the SALCs are occupied, hence it donates the electrons to the molecular σ σ* and π π* orbitals. The orbitals associated to e_g orbitals are not involved in πinteractions therefore it stays in the same energy level (figure 1). On the other hand, the occupied ligand SALCs t_2g orbitals thatwould form molecular orbitals with the metal t_2g orbitals (ie. d_xy, d_xz, d_xy) are lower in energy than its metal counterparts. The resulting MO has π* orbitals that are energetically lower than the σ* orbitals that are formed from the non-bonding orbitals (e_g). The difference between the t_2g π* and e_g σ orbitals is donated as Δ, split. In the π-donor case, the Δ is small due to the low π* level.

Conversely, the t_2g SALCs of a π accepting ligand are higher in energy than the metal t_2g orbitals because they are unoccupied. The resulting t_2g π* orbitals are higher than the σ* orbitals. This creates a larger Δ between the e_g and t_2gπorbitals, making these π-accepting orbitals high split ligands.

The metal d electron configurations for low and high spin.

Finally, the magnitude of Δ as influenced by the identity of the ligand can dictate how electrons are distributed in the metal d orbitals (figure 2) if the number of valence electrons of the metal center causes an odd number of electrons to be present in the metals T2g and Eg orbitals, causing unevenly filled, degenerate orbitals.^[1] Weak field ligands produce a small Δ hence a high spin configuration. Strong field ligands produce a large Δ hence a low spin configuration on the d electrons. This behavior of the d-electron configurations is referred to as the Yahn-Teller Distortion and is discussed further in future chapters.^[2]

2.7 Metal-Metal Bonds

Atomic orbitals of metals form bonding and antibonding molecular orbitals in a complex. These combinations of MOs include σ, π, δ orbitals. These terms describe the manner in which orbitals of the same symmetry overlap and creating bonding and antibonding interactions. Based on the symmetries of metals, the MO diagrams of two bonded metals will produce different combinations of bonding and antibonding orbitals. A prime example is the Re₂Cl₈^2- complex.

The Re₂Cl₈^2- complex is the first compound where quadruple bonds were observed. In the case of this compound, we can learn about all three kinds of bonding. The equivalent Re metals in this particular complex have four respective symmetries for all five d orbitals. Due to the high symmetry of the compound (an eclipsed D₄h point group), the five d orbitals produce four bonding and four antibonding orbitals. One is a σ bond, where the d_z² orbitals on each metal have a head on overlap along the z-axis. There are two π bonds where d_xz and d_yz overlap by two lobes out of the four lobes in each d orbitals along the z-axis. The last interaction is the δ bond, where all four lobes from the d_x²_-y² are overlapping.

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^ Pfenning, Brian (2015). Principles of Inorganic Chemistry (1 ed.). Hoboken, New Jersey: Wiley & Sons Inc. p. 562. ISBN 978-1-118-85910-0.
^ Pfenning, Brian (2015). Principles of Inorganic Chemistry (1 ed.). Hoboken, New Jersey: Wiley & Sons Inc. p. 562. ISBN 978-1-118-85910-0.

[1] Pfenning, Brian (2015). Principles of Inorganic Chemistry (1 ed.). Hoboken, New Jersey: Wiley & Sons Inc. p. 562. ISBN 978-1-118-85910-0.

[2] Pfenning, Brian (2015). Principles of Inorganic Chemistry (1 ed.). Hoboken, New Jersey: Wiley & Sons Inc. p. 562. ISBN 978-1-118-85910-0.

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