Premed · Premed · General Chemistry 2
Lecture 19: Coordination Chemistry and Transition Metals
General Chemistry II
Learning Objectives
By the end of this lecture, students will be able to:
- Define coordination compounds and identify their components
- Name coordination compounds using IUPAC nomenclature
- Determine the oxidation state of the central metal ion
- Describe common coordination geometries and isomerism
- Explain crystal field theory and predict d-orbital splitting patterns
- Relate d-orbital splitting to the color and magnetic properties of coordination compounds
Lecture Content
I. Transition Metals: Overview
Transition metals occupy the d-block of the periodic table (Groups 3 through 12) and possess several distinctive properties that set them apart from main-group elements. They exhibit multiple oxidation states because both the ns and (n-1)d electrons are available for bonding. They form colored compounds due to electronic transitions between d orbitals. They are often paramagnetic because of unpaired d electrons. They readily form coordination compounds (complex ions). And many transition metals serve as effective catalysts because their variable oxidation states and ability to bind substrates facilitate reaction pathways.
The electron configurations of transition metals follow the general pattern of filling (n-1)d orbitals after the ns orbital, with notable exceptions driven by the special stability of half-filled and fully filled d subshells. Chromium adopts the configuration [Ar]3d^5 4s^1 rather than the expected [Ar]3d^4 4s^2, and copper is [Ar]3d^10 4s^1 instead of [Ar]3d^9 4s^2. When transition metals form cations, electrons are removed from the ns orbital before the (n-1)d orbital. Thus iron, [Ar]3d^6 4s^2, becomes Fe^2+ as [Ar]3d^6 and Fe^3+ as [Ar]3d^5.
II. Coordination Compounds: Components
A coordination compound contains either a complex ion or a neutral complex. The complex ion itself consists of a central metal ion, typically a transition metal cation acting as a Lewis acid, surrounded by ligands, which are molecules or ions that donate electron pairs to the metal (Lewis bases). Monodentate ligands such as NH3, H2O, Cl-, CN-, CO, NO2-, and SCN- donate one electron pair. Bidentate ligands like ethylenediamine (en) and oxalate (C2O4^2-) donate two. Polydentate ligands such as EDTA can donate as many as six electron pairs.
The coordination number is the number of donor atoms directly bonded to the metal. The most common coordination numbers are 2 (linear geometry), 4 (tetrahedral or square planar geometry), and 6 (octahedral geometry). Counter ions are ions that sit outside the coordination sphere and serve to balance the overall charge of the compound.
III. Nomenclature of Coordination Compounds
Naming coordination compounds follows a systematic set of rules. The cation is named before the anion, as in any ionic compound. Within the complex ion, ligands are named before the metal. Anionic ligands receive an "-o" ending (chloro, bromo, cyano, hydroxo, oxalato), while most neutral ligands use their molecule name, with important exceptions: H2O is aqua, NH3 is ammine, CO is carbonyl, and NO is nitrosyl. Greek prefixes (di-, tri-, tetra-, penta-, hexa-) indicate multiple identical ligands, while bis-, tris-, and tetrakis- are used for polydentate ligands. Ligands are listed alphabetically, ignoring prefixes. If the complex is a cation, the metal keeps its element name; if the complex is an anion, the metal name takes the "-ate" suffix, often using the Latin root. The oxidation state of the metal is indicated in Roman numerals within parentheses.
Representative examples include [Cu(NH3)4]SO4, named tetraamminecopper(II) sulfate; K3[Fe(CN)6], named potassium hexacyanidoferrate(III); and [CoCl2(en)2]Cl, named dichloridobis(ethylenediamine)cobalt(III) chloride.
IV. Isomerism in Coordination Compounds
Coordination compounds exhibit both structural and stereoisomerism. Structural isomers have different connectivity of atoms. Linkage isomers arise when an ambidentate ligand can bond through different atoms, as in M-NO2 versus M-ONO or M-SCN versus M-NCS. Ionization isomers interchange ions between the coordination sphere and the outer sphere, as in [Co(NH3)5Br]SO4 versus [Co(NH3)5SO4]Br. Coordination isomers exchange ligands between two complex ions in the same compound.
Stereoisomers share the same connectivity but differ in the spatial arrangement of atoms. Geometric (cis-trans) isomers occur in square planar and octahedral complexes. In cis isomers, identical ligands are adjacent to each other; in trans isomers, they are opposite. This distinction has real biological consequences: cis-[Pt(NH3)2Cl2] (cisplatin) is a potent anticancer drug, while the trans isomer is biologically inactive. Optical isomers (enantiomers) are non-superimposable mirror images that occur in octahedral complexes with bidentate ligands, such as [Co(en)3]^3+. These enantiomers rotate plane-polarized light in opposite directions.
<image>A diagram showing geometric isomerism in a square planar complex [Pt(NH3)2Cl2]. Panel A: The cis isomer (cisplatin) with the two NH3 ligands adjacent (90 degrees apart) and the two Cl ligands adjacent. The structure is drawn with Pt at the center and ligands at the four corners of a square. Label: "cis-diamminedichloridoplatinum(II) -- active anticancer drug." Panel B: The trans isomer with NH3 ligands opposite each other (180 degrees apart) and Cl ligands opposite. Label: "trans-diamminedichloridoplatinum(II) -- biologically inactive." A note explains that geometric isomers have different physical and biological properties despite the same formula.</image>
V. Crystal Field Theory (CFT)
Crystal field theory explains the electronic structure, color, and magnetic properties of coordination compounds by considering how the electric field of the surrounding ligands affects the energies of the metal's d orbitals. In a free (gaseous) metal ion, all five d orbitals are degenerate, meaning they have the same energy. When ligands approach, the d orbitals split into groups of different energies.
In an octahedral complex, six ligands approach along the x, y, and z axes. The two d orbitals that point directly along these axes (d_z^2 and d_x^2-y^2) experience greater electrostatic repulsion from the ligands and are raised in energy, forming the e_g set. The three d orbitals that point between the axes (d_xy, d_xz, d_yz) experience less repulsion and are lowered in energy, forming the t_2g set. The energy gap between these two sets is the crystal field splitting energy, Delta_o.
In a tetrahedral complex, four ligands approach between the axes, and the splitting pattern is inverted: d_xy, d_xz, and d_yz are higher in energy. The tetrahedral splitting, Delta_t, is smaller than the octahedral splitting, approximately (4/9) Delta_o.
VI. The Spectrochemical Series and High-Spin vs. Low-Spin
The spectrochemical series ranks ligands by the magnitude of the crystal field splitting they produce: I- < Br- < Cl- < F- < OH- < H2O < NH3 < en < NO2- < CN- < CO. Ligands on the left are weak-field ligands that produce a small Delta, while those on the right are strong-field ligands that produce a large Delta.
For octahedral complexes with d^4 through d^7 electron configurations, the relative magnitudes of Delta_o and the electron pairing energy determine whether the complex is high-spin or low-spin. When Delta_o is smaller than the pairing energy (weak-field ligand), electrons occupy the higher-energy e_g orbitals before pairing in the t_2g set, resulting in a high-spin complex with more unpaired electrons. When Delta_o is larger than the pairing energy (strong-field ligand), electrons pair in the t_2g set before occupying e_g, producing a low-spin complex with fewer unpaired electrons.
The magnetic properties follow directly from the electron configuration. Paramagnetic complexes have unpaired electrons and are attracted by a magnetic field. Diamagnetic complexes have all electrons paired and are slightly repelled. The magnetic moment can be estimated from the number of unpaired electrons using mu = sqrt(n(n+2)) BM (Bohr magnetons).
<image>An energy level diagram comparing d-orbital splitting in octahedral and tetrahedral crystal fields. Left: Free ion with five degenerate d orbitals. Center: Octahedral field showing t_2g (lower, three orbitals: d_xy, d_xz, d_yz) and e_g (upper, two orbitals: d_z^2, d_x^2-y^2) with the splitting energy labeled as Delta_o. For a d^6 ion, two filling schemes are shown: high-spin (4 unpaired electrons with weak field ligand like H2O) and low-spin (0 unpaired electrons with strong field ligand like CN-). Right: Tetrahedral field showing the inverted splitting with e (lower) and t_2 (upper), and Delta_t labeled as smaller than Delta_o. The spectrochemical series is listed below from weak to strong field.</image>
VII. Color of Coordination Compounds
Many transition metal complexes display vivid colors because they absorb visible light. When a photon has energy exactly equal to Delta, the crystal field splitting energy, it promotes an electron from the lower-energy d-orbital set to the upper set. The color we observe is the complementary color of the light absorbed. A complex that absorbs red light appears green or blue; one that absorbs violet or blue light appears yellow or orange.
The relationship Delta = E = hv = hc/lambda connects the splitting energy to the wavelength of absorbed light. A larger Delta means the complex absorbs shorter wavelengths (higher energy) of light, while a smaller Delta means it absorbs longer wavelengths (lower energy). For example, [Ti(H2O)6]^3+ absorbs green-yellow light at approximately 500 nm and appears purple.
The color of a complex depends on the nature of the ligand (through the spectrochemical series), the oxidation state of the metal (higher charge generally produces larger Delta), and the coordination geometry (octahedral versus tetrahedral). Complexes with d^0 or d^10 configurations are colorless because no d-d transitions are possible.
