Valence Bond Theory of Coordination Compounds

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  • 0:04 Overview of Valence…
  • 1:32 Bonding and Properties…
  • 2:25 Discussion of Bonding…
  • 5:04 Bonding in Cobalt Complexes
  • 7:06 Lesson Summary
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Lesson Transcript
Instructor: Saranya Chatterjee

Saranya has a masters degree in Chemistry and in Secondary Education. She has taught high school, AP chemistry for 2 years and is teaching undergraduate college chemistry for 3 years.

This lesson will talk about coordination compounds or transition metal complexes and Valence Bond Theory. It will discuss bonding and magnetic properties of a few coordination compounds.

Overview of Valence Bond Theory

One of the early models of chemical bonding is the valence bond theory, which was introduced by Linus Pauling. In this theory, the formation of a covalent bond between two atoms occurs through the build-up of an electron density between the nucleii of the two atoms. The valence atomic orbital of one atom shares space or overlaps with the valence atomic orbital of another atom. The overlap of orbitals allows two electrons of opposite spins to share the space between the two atoms, leading to the formation of a covalent bond. In this model, the bonding electron pairs are located in the region between atoms, and the non-bonding (the lone pair) electrons are located in directed regions in space which imparts particular geometry to the molecule.

In this lesson, we will focus on coordination compounds, which are transition metal complexes where the transition metal is the central atom surrounded by other atoms. We often assume that atomic orbitals of the central atom in a compound mix to form hybrid orbitals. The process of mixing is called hybridization. Transition metals have s, p, and d orbitals, which undergo hybridization. The molecular geometries of the transition metal complexes can be stabilized through different types of orbital hybridization of the central metal ion as shown here:

Molecular Geomgetries and Hybridization Table

Bonding and Properties of Coordination Compounds

A coordination compound contains a central metal or ion, which is usually a transition metal surrounded by neutral molecules or anions called ligands. The transition metal is a Lewis acid and a d-block element in the periodic table containing valence d electrons, and ligands are usually Lewis bases containing at least one electron pair to donate to the central metal atom. Ligands may be neutral like water (H2 O), ammonia(NH3), and ethylene diammine(en); or anionic like chloride(Cl-) and cyanide(CN-). Also, another way of classifying ligands depends on their strength: they may be strong or weak. The spectrochemical series shown here determines the strength of a ligand:

O22-< I- < Br-< S2- < SCN- (S-bonded) < Cl- < N3 - < F-< NCO- < OH- < C2 O42- < NCS- < CH3 CN < py (pyridine) < NH3 < en (ethylenediamine) < bipy (2,2'-bipyridine) < phen (1,10-phenanthroline) < NO2- < PPh3 < CN-

Discussion of Bonding in Ni Complexes

Looking at the series, we can understand that oxide is the weakest and CO, CN- is the strongest ligand. Usually strong field ligands form inner orbital, diamagnetic complexes, and weak field ligands form outer orbital, paramagnetic complexes. Let us discuss this in detail with examples. Look at the following two coordination compounds, also called transition metal complexes, of nickel, NiCl4 2- and Ni(CN)4 2-. Chloride and cyanide being anionic ligands, Ni in these two complexes is in +2 oxidation state with electron configuration Ar 3d8, which looks like this:

Electron distribution in free Nickel ion

Now, with 4 incoming ligands, 2 geometries are possible, either tetrahedral or square planar, as we saw in the earlier table. Chloride being a weak field ligand according to the spectrochemical series will form an outer orbital complex with sp3 hybridization leading to a tetrahedral NiCl4 2- complex. Four chloride ligands bring in 8 electrons, and the existing transition metal ion Ni2+ has 8 electrons (6 being spin paired and 2 unpaired) leading to a paramagnetic complex. The dipole moment of a paramagnetic complex can be calculated by using the formula:

Dipole moment or magnetic moment µ = square root{n(n+2)} B.M. where n = number of unpaired electrons and B.M. = Bohr Magneton.

Hence, µ of NiCl4 2- = square root{2(2+2)} B.M. = 2.828 B.M.

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