Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → Inorganic chemistry ↓


Coordination chemistry


Coordination chemistry is an important aspect of inorganic chemistry that deals with the study of coordination compounds or complexes. These compounds are formed by the combination of metal ions with ligands, which are molecules or ions that are able to donate pairs of electrons to the metal center. Understanding coordination chemistry is essential to understand how metals interact in biological, medicinal, and industrial processes.

Basic concepts

At the core of coordination chemistry are coordination compounds, which typically consist of a central metal atom or ion surrounded by an array of bonded molecules or ions, known as ligands. The central metal ion and its bound ligands form a coordination complex.

Central metal atom or ion

The central metal is typically a transition metal, such as iron (Fe), copper (Cu), or nickel (Ni). Transition metals are particularly well suited for forming coordination compounds because they have partially filled d-orbitals, which can easily accept electron pairs from ligands.

Example: Consider the iron(II) ion in [Fe(CN)6]4- complex. Here, iron is the central metal ion.

Ligands

Ligands are ions or molecules that can donate at least one pair of electrons to the central metal. They can be:

  • Monodentate: ligands that donate a pair of electrons, such as the chloride ion (Cl-) or ammonia (NH3).
  • Bidentate: ligands that can form two bonds with the central metal, such as ethylenediamine (en).
  • Polydentate: ligands that can form multiple bonds, such as ethylenediaminetetraacetic acid (EDTA).

The number of attachment points of a ligand is known as the "denticity" of the ligand.

Coordination number

The coordination number of a complex refers to the number of ligand donor atoms bonded to the central metal. Coordination numbers often range from 2 to 12, but are typically found between 4 and 6.

Example: In [Co(NH3)6]3+ complex, the coordination number of cobalt is 6 because it has six ammonia molecules attached to it.

Geometry of coordination compounds

The geometry of coordination compounds varies depending on the coordination number and the nature of the ligand.

Linear geometry

Coordination occurs in number 2. Here's a simple visual representation:

l M l

Example: [Ag(NH3)2]+

Tetrahedral geometry

Found in complexes with coordination number 4. The tetrahedral shape is like a pyramid with a triangular base.

Example: [ZnCl4]2-

Square planar geometry

It is also found with coordination number 4, which is typical for d8 metal ions.

l l l l

Example: [PtCl4]2-

Octahedral geometry

Typical for coordination number 6, where the ligands are evenly distributed around the metal ion.

Example: [Co(NH3)6]3+

Isomerism in coordination compounds

Isomerism in coordination compounds arises from the different ways in which the ligands can be arranged around the metal center. There are two main types: structural isomerism and stereoisomerism.

Structural isomerism

Types include:

  • Linkage isomerism: arises when a ligand can bind to the metal in more than one way. Example: [Co(NH3)5(NO2)]2+ vs. [Co(NH3)5(ONO)]2+.
  • Coordination isomerism: occurs when there is an exchange of ligands between the cationic and anionic parts of the compound. Example: [Co(NH3)6][Cr(CN)6] vs. [Cr(NH3)6][Co(CN)6].

Stereoisomerism

Types include:

  • Geometrical isomerism: Occurs in complexes with certain spatial arrangements. For example, in square planar and octahedral complexes, [Pt(NH3)2Cl2] can have cis and trans isomerism.
  • Optical isomerism: Found in chiral coordination compounds which do not superimpose on their mirror images. Example: some [Co(en)3]3+ complexes.

Stability of coordination compounds

The stability of a coordination compound is an indication of how strongly the metal ion binds to the ligand. Factors affecting stability include:

  • Nature of the metal ion: The charge on the metal and its size affect the stability of the complex. Higher positive charge and smaller size usually form more stable complexes.
  • Nature of the ligand: Some ligands form strong bonds with metals. For example, ligands such as cyanide (CN-) often form very stable complexes.
  • Chelate effect: Polydentate ligands like EDTA form more stable complexes due to the chelate effect, where multiple bonds are formed in the cyclic structure, greatly increasing the stability.

Applications of coordination chemistry

Coordination compounds play important roles in a variety of areas:

Biological systems

  • Hemoglobin: This iron-containing coordination complex is essential for oxygen transport in the blood.
  • Vitamin B12: Contains a cobalt center and is important for DNA synthesis and energy production.

Catalysis

Coordination compounds are used as catalysts in many industrial processes, accelerating reactions without being consumed in the process.

  • Hydrogenation of alkenes: Catalyzed by nickel, platinum or palladium complexes.
  • Wacker process: A palladium complex is used to convert ethylene to acetaldehyde.

Medicine

Some coordination compounds are used directly in medical treatments or diagnostic procedures. For example, cisplatin ([PtCl2(NH3)2]) is a well-known anti-cancer drug.

Environmental chemistry

  • Water softening: EDTA is used in water treatment processes to trap metal ions, thereby reducing water hardness.
  • Heavy metal detoxification: Chelating agents remove toxic metals from living organisms.

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Coordination chemistry

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