Coordination Chemistry Complexes

Coordination Complexes

Coordination complexes have at least one metal complex, which contains a metal center and is surrounded by electron-donating ligands. This is known as a complex ion. Counterions balance the charge of the complex ion to form the molecular coordination complex. Coordination complexes are soluble in water, where the counter-ion and metal-ion complex dissociate. The metal ion and ligands behave like a polyatomic ion and do not dissociate.

The geometry of a complex ion takes on standard Valence-Shell Electron-Pair Repulsion Theory (VSEPR) geometries including linear, square planar, tetrahedral, and octahedral. Octahedral complex ions are the most common geometry. Crystal field theory explains energy splitting among d-orbitals in transition metal ions and the VSEPR geometries. Energy splitting is influenced by the shape and orientation of the d-orbital lobes.

Ligands and the Spectrochemical Series

Ligands are classified by the number of bonds, or attachments, they can make with a metal center. A single attachment is known as monodentate (one-toothed). A ligand that makes two attachments is called bidentate (two-toothed), and three attachments is known as tridentate. Ligands donate electron density to the metal center to form the coordinate-covalent bond. Ligands may be charged or neutral. Ligands are classified as being strong or weak according to the spectrochemical series:

(weak) I^- < Br^- < Cl^- < SCN^- < F^- < OH^- < ox^2-< ONO^- < H₂O < NCS^- < EDTA^4- < NH₃ < en < NO₂^- < CN^- (strong)

Orbital Splitting

When six ligands approach a metal center to form an octahedral complex, the five degenerate d-orbitals split into three lower-energy degenerate t_2g orbitals and two higher-energy degenerate e_g orbitals. The distance of the splitting between the t_2g and e_g orbitals is dictated by the strength of the ligand according to the spectrochemical series.

Hund's rule still applies and electrons fill orbitals one at a time, but they fill in accordance to size of the splitting of the t_2g and e_g orbitals. If the split is small, electrons will fill all the orbitals singly before pairing. This maximizes the number of unpaired electrons and is called high-spin. Likewise, a strong-field causes a large t_2g-e_g split: electrons pair in the t_2g set before filling the higher-energy e_g orbitals. This minimizes the number of unpaired electrons and is called low-spin. The drive for electrons to pair is governed by the energy (or size) of the orbital splitting compared to the energy of electron pairing. If the energy of pairing is high compared to the energy of moving into the e_g orbitals, then electrons are high-spin. If the energy of pairing is low compared to the energy of moving into the higher e_g orbitals, then electrons are low-spin.

The distance that electrons have to move from the lower t_2g state to the higher e_g state in the metal center dictates the energy of electromagnetic radiation that the complex absorbs. If that energy is in the visible region (400-700 nm, 1.77 eV - 3.1 eV), the complex generally has a color. Weak-field ligands (I^- → OH^-) cause small splittings and complexes absorb low-energy light (i.e. red) which appear green in color. Strong-field ligands (EDTA → CN^-) absorb high-energy light (i.e. blue-violet) and appear red-yellow in color. Complexes with ligands that are between strong and weak on the spectrochemical series, like ammonia, can adopt either a weak- or strong-field geometry.

The color-ligand relationship is the rationale for the name "spectrochemical series". The number of paired and unpaired electrons also gives rise to paramagnetic diamagnetic properties in metal complexes.

When four ligands coordinate around a metal center, either a square planar or tetrahedral complex can result. The orbital energies in tetrahedral complexes are flipped compared to octahedral complexes, with e_g being lower in energy than t_2g. This is due to the orientations of d-orbitals with respect to the coordinating ligands. In square planar complexes, there are several differences in orbital energies, with d_yz and d_xz being degenerate and lowest in energy (lower than d_z2,), then d_xy, and finally the highest energy d_x2-y2 orbital.

Structure and Color

Because the distance in orbital splitting varies with ligand strength, a coordination complex with the same metal center can have a variety of colors based upon the coordinating ligand. For example, an aqueous solution of Ni(H₂O)₆²⁺ has a light green color, but Ni(NH₃)₆²⁺ is deep blue. The color arises from the change in energy between the t_2g-e_g orbitals. NH₃ is a stronger field ligand, which pushes the orbitals further apart from one another as well as displaces the H₂O ligands from the metal center. We will further explore the effect of ligands on color and coordination complexes in this experiment.

From pigments to people, transitional metals are found throughout fields of chemistry, biology, geology, and engineering. Understanding the behavior of transition metals under different chemical states can be as simple as monitoring color or magnetic behavior. Nearly every 3d (4^th row) transition metal is vital to physiological function and, in all cases, these metals are bound by ligands to form coordination complexes. For example, iron is vital to oxygen transport in all vertebrates. Hemoglobin, a complex protein, contains four heme subunits with Fe²⁺ in the center of each. In hemoglobin, the Fe²⁺ is chelated by a tetradentate ring and a histidine residue, making it square pyramidal (five-sided). When oxygen is present, the subunits become octahedral. O₂ is considered a strong-field ligand, which causes large d-orbital t_2g-e_g splitting, making it low-spin. Relatively high-energy light is required to promote an electron to the e_g state, so blue light is absorbed making oxygenated (arterial) blood appear bright red. In contrast, deoxygenated (venous) blood has a smaller d-orbital splitting and lower energy light red light is absorbed, making deoxygenated blood appear dark, purplish-red. In the same respect, carbon monoxide, CO, is a strong-field ligand and will displace oxygen. It gives blood an even brighter red appearance due to strong-field splitting. The preferential binding for CO over O₂ in blood is often fatal.

Another application of coordination chemistry is in paints and pigments. While many pigments are simple metal oxides, others like Prussian Blue and Phthalocyanine Blue are coordination complexes whose color arises from the splitting in d-orbitals (Figure 2). In Prussian Blue, iron is surrounded by six cyanide ligands, creating the high-spin iron (III) hexacyanoferrate complex, Fe(CN)₆^3-. Another compound, Phthalocyanine Blue, is a square planar complex with a copper (II) ion in the center surrounded by a tetradentate phthalocyanine molecule.

Figure 2. Prussian Blue, an iron-centered coordination complex and Phthalocyanine Blue, a copper-centered coordination complex.

Coordination compounds have a metal ion center with surrounding ligands and a counterion to balance charge. The ligands can be monodentate or chelating with two-four attachment sites. Ligands are also categorized by the spectrochemical series, which classifies the relative strength of the ligands to split a metal's d-orbitals. Both color and magnetic properties are influenced by the metal and the ligands. Large d-orbital splitting requires large energies to promote electrons into the higher energy orbitals and absorbs high-energy light (short wavelength). These are low spin-complexes and have the maximum number of paired electrons. In contrast, a small d-orbital splitting is known as weak-field and absorbs low energy light as well as has the maximum number of unpaired electrons. The charge and identity of the metal ion as well as the bound ligands define both the observed color and magnetic properties in coordination compounds.