19.2 Coordination Chemistry of Transition Metals

Coordination Compounds

Coordination compounds consist of a central metal atom or ion bonded to surrounding molecules or ions called ligands. They show up everywhere, from hemoglobin carrying oxygen in your blood to industrial catalysts and anticancer drugs. Understanding how they're built, named, and classified is the foundation for explaining their wide range of properties.

Characteristics of Coordination Compounds

The bonds between the metal center and its ligands are coordinate covalent bonds, meaning both electrons in each bond come from the ligand (not shared one from each atom, as in a typical covalent bond).

• The central metal is usually a transition metal (iron, copper, cobalt) or an inner transition metal (lanthanides, actinides).
• Ligands can be neutral molecules ($\text{H}_2\text{O}$, $\text{NH}_3$) or ions ($\text{Cl}^-$, $\text{CN}^-$).
• The central atom plus all its attached ligands form the coordination sphere, which is enclosed in square brackets in chemical formulas. The metal is listed first, then the ligands: $[\text{Co}(\text{NH}_3)_6]^{3+}$.

Coordination compounds can carry different overall charges depending on the metal's oxidation state and the ligands' charges:

• Cationic: $[\text{Co}(\text{NH}_3)_6]^{3+}$
• Anionic: $[\text{Fe}(\text{CN})_6]^{4-}$
• Neutral: $[\text{Pt}(\text{NH}_3)_2\text{Cl}_2]$

Monodentate vs. Polydentate Ligands

Monodentate ligands bond to the metal through a single atom, forming one coordinate covalent bond per ligand. Common examples include $\text{H}_2\text{O}$, $\text{NH}_3$, $\text{Cl}^-$, and $\text{CN}^-$.

Polydentate ligands bond through two or more atoms, wrapping around the metal like a claw. The word "chelate" actually comes from the Greek word for claw.

• Bidentate ligands bond through two atoms. Ethylenediamine ($\text{H}_2\text{NCH}_2\text{CH}_2\text{NH}_2$, often abbreviated "en") is a classic example, binding through both nitrogen atoms.
• Hexadentate ligands bond through six atoms. EDTA (ethylenediaminetetraacetic acid) is the most well-known hexadentate ligand.

Polydentate ligands form more stable complexes than monodentate ones. This is called the chelate effect: when one polydentate ligand replaces several monodentate ligands, the total number of free particles in solution increases, raising entropy and making the reaction thermodynamically favorable.

Nomenclature for Coordination Compounds

Naming coordination compounds follows a specific set of rules. Here's the process:

1. Name the ligands first, in alphabetical order. Alphabetical order is based on the ligand name itself, ignoring any numerical prefixes.

   • Anionic ligands get an "-o" ending: $\text{Cl}^-$ becomes chlorido (or chloro), $\text{CN}^-$ becomes cyanido (or cyano).
   • Neutral ligands keep their molecule name, with two important exceptions: $\text{H}_2\text{O}$ is called aqua, and $\text{NH}_3$ is called ammine (note the double "m").

2. Use Greek prefixes to indicate how many of each ligand are present: di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6).

3. Name the metal after all the ligands. If the complex is anionic, add the suffix "-ate" to the metal name (e.g., ferrate, cobaltate).

4. Indicate the metal's oxidation state with a Roman numeral in parentheses right after the metal name.

For example, $[\text{Co}(\text{NH}_3)_6]^{3+}$ is named hexaamminecobalt(III) ion. You can figure out the oxidation state by accounting for the charges of all ligands and the overall charge of the complex.

The coordination number is the total number of ligand donor atoms bonded to the central metal. In $[\text{Co}(\text{NH}_3)_6]^{3+}$, the coordination number is 6 because six nitrogen atoms (one from each $\text{NH}_3$) are bonded to cobalt.

Isomerism in Coordination Complexes

Coordination compounds with the same formula can have different arrangements of their ligands, giving rise to isomers with distinct properties.

Geometric isomers differ in the spatial arrangement of ligands around the metal.

• In square planar complexes with the formula $[\text{MA}_2\text{B}_2]$, the two A ligands can be adjacent to each other (cis) or on opposite sides (trans). Cisplatin and transplatin are a real-world example: same formula, but only the cis form works as an anticancer drug.
• Octahedral complexes can also show cis-trans isomerism when they have two or more different types of ligands.

Optical isomers (also called enantiomers) are non-superimposable mirror images of each other, much like your left and right hands.

• Octahedral complexes with bidentate or polydentate ligands commonly exhibit optical isomerism. For example, $[\text{Co}(\text{en})_3]^{3+}$ (where "en" is ethylenediamine) exists as two mirror-image forms.
• These are designated Δ (delta) and Λ (lambda) based on the direction of twist of the ligands around the metal center.
• Optical isomers can have very different biological activities because enzymes and receptors in living systems are themselves chiral and interact differently with each mirror-image form.

Electronic Structure and Properties

When ligands bond to a transition metal, they don't interact equally with all five d orbitals. Some d orbitals end up at higher energy than others. This energy gap is called d-orbital splitting, and its size depends on which ligands are present.

The spectrochemical series ranks ligands by how much splitting they cause, from weak-field to strong-field:

$\text{I}^- < \text{Br}^- < \text{Cl}^- < \text{F}^- < \text{OH}^- < \text{H}_2\text{O} < \text{NH}_3 < \text{en} < \text{NO}_2^- < \text{CN}^- < \text{CO}$

The size of the splitting determines how electrons fill the d orbitals:

• High-spin complexes form with weak-field ligands (small splitting). Electrons spread out across all d orbitals before any pairing occurs, maximizing unpaired electrons.
• Low-spin complexes form with strong-field ligands (large splitting). Electrons pair up in the lower-energy d orbitals first because the energy cost of pairing is less than the cost of jumping to the higher orbitals.

This distinction matters because the number of unpaired electrons directly affects the compound's magnetic properties and color.

Applications of Coordination Compounds

Coordination compounds are found throughout biology, medicine, and industry.

• Hemoglobin contains an iron(II) ion coordinated to a porphyrin ring ligand. Oxygen molecules bind reversibly to the iron center, which is how your red blood cells pick up oxygen in the lungs and release it to tissues.
• Chlorophyll, the green pigment in plants, has a magnesium ion at the center of a porphyrin ring. This complex absorbs light energy to drive photosynthesis.
• Vitamin B12 is a coordination compound with a cobalt ion at its center. It's essential for red blood cell formation and nervous system function.
• Cisplatin ($[\text{Pt}(\text{NH}_3)_2\text{Cl}_2]$) is a square planar platinum(II) complex used as an anticancer drug. It binds to DNA strands, interfering with cell division and triggering apoptosis (programmed cell death) in cancer cells.
• In industry, coordination compounds serve as catalysts. The Monsanto process uses a rhodium complex, $\text{cis-}[\text{Rh}(\text{CO})_2\text{I}_2]^-$, to catalyze the reaction of methanol with carbon monoxide to produce acetic acid.