Coordination Chemistry

Coordination compounds contain a central metal ion (usually a transition metal) bonded to surrounding molecules or ions called ligands through coordinate covalent bonds. In these bonds, both electrons come from the ligand, which acts as a Lewis base donating an electron pair to the metal (Lewis acid).

The metal ion plus its attached ligands form the coordination sphere, enclosed in square brackets in formulas. Species outside the brackets are counter ions, not directly bonded to the metal. The coordination number is the total number of donor atoms bonded to the metal — common values are 2, 4, and 6, producing linear, tetrahedral or square planar, and octahedral geometries respectively.

Monodentate ligands (e.g., NH₃, Cl⁻, H₂O) bind through a single donor atom. For these ligands the coordination number equals the number of ligands attached.

Polydentate ligands bind to a metal through two or more donor atoms simultaneously. Bidentate ligands such as ethylenediamine (en, H₂NCH₂CH₂NH₂) have two donor atoms; for example, [Co(en)₃]³⁺ has three bidentate ligands giving a coordination number of six.

A complex formed with a polydentate ligand is called a chelate (from the Greek for “claw”). Chelates are generally more stable than analogous complexes with monodentate ligands because dissociation requires breaking multiple bonds simultaneously.

The hexadentate ligand EDTA binds through six donor atoms and is widely used in water softening, food preservation, and medicine. In chelation therapy, drugs such as dimercaptosuccinic acid (DMSA) form water-soluble chelates with toxic heavy metals (arsenic, mercury, lead), allowing the kidneys to excrete them.

Werner nomenclature follows systematic rules for naming coordination compounds:

1. In ionic compounds, name the cation before the anion.
2. Name ligands alphabetically before the metal. Anionic ligands add -o to the root (chlorido, hydroxido, cyano). Neutral ligands use their molecule name, with four exceptions: H₂O = aqua, NH₃ = ammine, CO = carbonyl, NO = nitrosyl.
3. Use Greek prefixes (di-, tri-, tetra-, penta-, hexa-) for multiple identical ligands. If the ligand name itself contains a prefix, use bis-, tris-, tetrakis- with the ligand in parentheses.
4. State the metal name followed by its oxidation state in Roman numerals. When the complex is an anion, the metal name takes the suffix -ate (ferrate for iron, cuprate for copper, plumbate for lead).

Example: [Co(NH₃)₄Cl₂]Cl is tetraamminedichloridocobalt(III) chloride. The oxidation state of cobalt is +3 because the complex ion has a +1 charge: +1 = Co + 0 + 2(−1).

The coordination number of a complex largely determines its geometry. The three most common arrangements are:

• Octahedral (coordination number 6): six ligands arranged at 90° angles around the metal, e.g., [Co(H₂O)₆]²⁺.
• Tetrahedral (coordination number 4): four ligands at 109.5° angles, common for d⁰ or d¹⁰ metals with low oxidation states, e.g., [Zn(CN)₄]²⁻.
• Square planar (coordination number 4): four ligands in a flat arrangement at 90°, typical for d⁸ metals such as Pt²⁺ and Ni²⁺, e.g., [Pt(NH₃)₂Cl₂].

Less common geometries include linear (coordination number 2, e.g., [Ag(NH₃)₂]⁺), trigonal bipyramidal and square pyramidal (coordination number 5). Unlike main-group compounds, nonbonding d electrons do not influence the arrangement of ligands.

Coordination compounds can exist as isomers — species with the same formula but different structures or spatial arrangements.

Geometric isomers (cis/trans) occur when two identical ligands can occupy adjacent positions (cis) or opposite positions (trans). For example, [Co(NH₃)₄Cl₂]⁺ has a violet cis isomer and a green trans isomer, demonstrating that spatial arrangement affects physical properties including color, polarity, and solubility.

Optical isomers (enantiomers) are non-superimposable mirror images, analogous to left and right hands. Complexes such as [M(en)₃]ⁿ⁺ are chiral and rotate plane-polarized light in opposite directions.

Other isomer types include linkage isomers (same ligand bound through different atoms, e.g., –SCN vs. –NCS) and ionization isomers (exchange of an ion between the coordination sphere and the counter ion).

Crystal field theory (CFT) explains the electronic structure, color, and magnetic behavior of coordination complexes. It treats metal–ligand interactions as purely electrostatic: ligand electron pairs repel the metal’s d electrons, causing the five d orbitals to split into groups of different energy.

In an octahedral complex, ligands approach along the axes, repelling the e_g orbitals (d_z² and d_x²−y²) more than the t_2g orbitals (d_xy, d_xz, d_yz). The energy gap between these sets is called Δ_oct.

The magnitude of Δ_oct depends on the ligand. The spectrochemical series ranks ligands by increasing field strength: I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO. When Δ_oct is large (strong-field ligands), electrons pair in the lower t_2g set before occupying e_g, producing a low-spin complex. When Δ_oct is small (weak-field ligands), electrons spread across all five orbitals before pairing, producing a high-spin complex.

Transition metal complexes are often vividly colored because they absorb visible light with energy matching Δ_oct. An absorbed photon promotes an electron from the t_2g to the e_g set. The color you see is complementary to the absorbed wavelength — for example, [Cu(NH₃)₄]²⁺ absorbs orange-red light and appears blue.

Strong-field ligands produce large Δ_oct values, so these complexes absorb higher-energy violet or blue light and appear yellow, orange, or red. Weak-field ligands yield small Δ_oct, absorbing lower-energy red or orange light, making the complex appear blue or green. Complexes with d⁰ or d¹⁰ configurations (no possible d–d transitions) are usually colorless.

CFT also explains magnetic behavior. Complexes with unpaired electrons are paramagnetic (attracted to a magnetic field), while those with all electrons paired are diamagnetic. The measured magnetic moment reveals the number of unpaired electrons, confirming whether a complex is high-spin or low-spin.

Use this workflow to keep naming, geometry, and electronic reasoning consistent:

1. Identify the coordination sphere and any counterions first.
2. Determine oxidation state from ligand charges and overall complex charge.
3. Name ligands alphabetically (ignoring multiplicative prefixes for alphabetization), then metal with oxidation state.
4. Check geometry from coordination number, then evaluate possible isomerism.
5. Use ligand-field strength to reason about splitting magnitude, spin state, color, and magnetism.

Frequent errors: losing charge balance when assigning oxidation state, mixing ligand naming conventions, and treating color predictions as exact wavelength calculations rather than qualitative d-orbital splitting outcomes. A final consistency check is whether nomenclature, oxidation state, and electron count all tell the same chemical story.