Coordination chemistry studies compounds in which a central metal ion is bonded to surrounding ions or molecules called ligands. This cheat sheet helps students identify complex ions, name ligands, determine oxidation states, and predict common shapes. These skills are important for understanding transition metal reactions, colors, magnetism, and biological metal complexes.
The core ideas are coordination number, ligand denticity, complex charge, and metal oxidation state. Students also need to connect geometry with common coordination numbers such as and . Crystal field splitting explains why many transition metal complexes are colored and why some are paramagnetic.
Formation constants such as describe how strongly ligands bind to metal ions.
Key Facts
- A coordination complex has a central metal ion bonded to ligands, often written in brackets such as .
- The oxidation state of the metal is found by charge balance: .
- The coordination number is the number of donor atoms directly bonded to the metal, so has coordination number .
- Monodentate ligands donate one electron pair, while bidentate ligands donate two electron pairs through two donor atoms.
- Common geometries include linear for coordination number , tetrahedral or square planar for coordination number , and octahedral for coordination number .
- The formation constant is for the equilibrium .
- In octahedral crystal field splitting, the orbitals separate into lower-energy orbitals and higher-energy orbitals with energy gap .
- A complex is paramagnetic if it has one or more unpaired electrons, and it is diamagnetic if all electrons are paired.
Vocabulary
- Coordination complex
- A species made of a central metal ion bonded to surrounding ligands through coordinate covalent bonds.
- Ligand
- An ion or molecule that donates an electron pair to a metal ion to form a coordinate covalent bond.
- Coordination number
- The number of ligand donor atoms directly bonded to the central metal ion.
- Denticity
- The number of donor atoms in one ligand that can bond to the same metal ion.
- Chelate
- A complex formed when a multidentate ligand bonds to a metal ion at more than one site.
- Crystal field splitting
- The separation of metal orbital energies caused by the electric fields of surrounding ligands.
Common Mistakes to Avoid
- Ignoring ligand charge when finding oxidation state is wrong because neutral ligands such as and charged ligands such as affect charge balance differently.
- Confusing coordination number with the number of ligands is wrong because one ligand can attach through more than one donor atom, as in a bidentate ligand.
- Assuming all coordination number complexes are tetrahedral is wrong because some metal ions commonly form square planar complexes.
- Forgetting brackets around the complex ion is wrong because brackets show which ligands are directly coordinated to the metal, as in .
- Treating color as random is wrong because the observed color depends on , ligand strength, and electronic transitions between split orbitals.
Practice Questions
- 1 Find the oxidation state of cobalt in .
- 2 What is the coordination number of the metal in if is bidentate?
- 3 Write the expression for for .
- 4 Explain why a complex with strong-field ligands may have fewer unpaired electrons than a similar complex with weak-field ligands.
Understanding Coordination Chemistry & Ligands
A ligand forms a coordinate bond by supplying both electrons in the shared pair. The metal ion accepts that pair because it has available orbitals and a positive charge that attracts electron density. This does not mean every ligand binds equally well.
Charge, size, donor atom type, and the solvent all affect binding. A negatively charged ligand often binds strongly because electrostatic attraction helps. Water molecules can occupy sites around a dissolved metal ion, then be replaced when a different ligand is added.
This replacement process is called ligand substitution. Some metal complexes exchange ligands quickly, while others remain unchanged for a long time. Stability describes which mixture is favored at equilibrium.
Speed describes how fast the change happens. These are different ideas.
Multidentate ligands can wrap around a metal ion and make ring structures. This is called chelation. Chelated complexes are often more stable than similar complexes made from separate one point ligands.
Part of the reason is entropy. When one multidentate ligand attaches, it may replace several small molecules from the metal's surroundings. More separate particles are then free in solution, which can favor the products.
This chelate effect matters in medicine and environmental chemistry. EDTA is a ligand used to bind metal ions in water treatment and in some treatments for metal poisoning.
In blood, hemoglobin holds an iron ion within a large ring shaped ligand system. The surrounding structure controls how iron binds oxygen without releasing the iron ion.
The shape around a metal affects more than appearance. For complexes with four donor atoms, tetrahedral and square planar arrangements can have very different electron arrangements and chemical behavior. Many platinum two complexes are square planar.
This geometry is important for cisplatin, a platinum containing drug. Its cis arrangement allows it to bind to parts of DNA. Spatial arrangement also creates isomers.
Cis and trans isomers have matching formulas but place ligands next to each other or opposite each other. Some complexes can form mirror image versions called optical isomers. These versions may interact differently with biological molecules because many biological structures are themselves not symmetrical.
Crystal field ideas explain patterns that students can test against observations. When ligands approach a metal, they repel electrons in the metal d orbitals. Orbitals pointing more directly toward ligands are raised in energy more than orbitals pointing between ligands.
If visible light has the right energy, electrons can absorb it and move to a higher d orbital. The remaining transmitted or reflected light gives the complex its observed color. Ligands differ in how large an energy gap they produce.
A small gap can leave more unpaired electrons, while a large gap can cause electrons to pair in lower energy orbitals. When solving problems, first find the metal oxidation state, then determine its d electron count. Next identify the likely geometry and ligand strength before deciding electron pairing, magnetism, or color.
Keep the charge on a ligand separate from its number of donor atoms. Those two properties are commonly confused.