Chapter 6: Coordination Chemistry- Bonding
- Page ID
- 440835
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)With the exception of molecular orbital theory, the bonding models discussed in previous courses for main group complexes are not easily applied to transition metal complexes. This chapter describes bonding models that are used to explain the bonding and electronic structures of coordination complexes. As with many theories, the crystal field and ligand field theories discussed in this chapter were developed because existing bonding theories couldn't explain the observed properties of some transition metal complexes.
As an example let us consider the following two transition metal complexes, [Fe(NH3)6]2+ and [Co(NH3)6]3+. The Fe2+ and the Co3+ centers both have a valence electron configuration of 3d6 and have the same ligands coordinated. Lewis, VSEPR, and valence bond theories predict that two isoelectronic molecules should have very similar properties, but in the case of [Fe(NH3)6]2+ and [Co(NH3)6]3+ that is not true. The light absorption properties (and colors) of the two are quite different. The iron complex absorbs very little visible light and therefore appears a very pale yellow color and the cobalt complex strongly absorbs visible light and has a dark red-orange color. Also, the iron complex is paramagnetic (having unpaired electrons) and the cobalt complex is diamagnetic (having no unpaired electrons). Crystal field theory, and later ligand field theory were developed in part to explain the spectroscopic and magnetic properties of transition metal complexes.
| Complex | [Fe(NH3)6]2+ | [Co(NH3)6]3+ |
| Valence electron configuration | [Ar] 3d6 | [Ar] 3d6 |
| Color | pale yellow | Red-orange |
| Magnetic properties | paramagnetic | diamagnetic |
- Draw a crystal field diagram for a given metal complex and use to identify the number of unpaired electrons
- Identify whether a given metal complex will be high or low spin
- Calculate the crystal field stabilization energy for a given metal complex
- Use ligand field theory to explain the spectrochemical series
Thumbnail image shows pi donation from a ligand to a metal
- Section 6.1: Crystal Field Theory
- Crystal field theory describes how orbital degeneracy is broken in the presence of ligands. It is based on the electrostatic interaction between negatively charged orbitals and negative point charges as ligands.
- Section 6.2: Crystal Field Stabilization Energy
- A consequence of crystal field theory is that the distribution of electrons in the d orbitals may lead to net stabilization of some complexes depending on the specific ligand geometry and metal d-electron configuration. It is a simple matter to calculate this stabilization knowing the electron configuration and the crystal field splitting diagram.
- Section 6.3: Factors That Affect Crystal Field Splitting
- The magnitude of the crystal field splitting (Δ) dictates whether a complex with four, five, six, or seven d electrons is high spin or low spin, which affects its magnetic properties, structure, and reactivity. Large values of Δ (i.e., Δ > P) yield a low-spin complex, whereas small values of Δ (i.e., Δ < P) produce a high-spin complex.
- Section 6.4: Ligand Field Theory
- Ligand field theory is an extension of crystal field theory which includes orbital overlap between ligand orbitals and the metal d orbitals. It allows us to explain the differences between strong field and weak field ligands.
- Section 6.6: Jahn-Teller Distortions
- The Jahn-Teller effect is a geometric distortion of a molecular system that reduces its symmetry and energy. This distortion is typically observed among octahedral complexes where the two axial bonds can be shorter or longer than those of the equatorial bonds.