Complexes or coordination compounds are molecules that posess a metal center that is bound to ligands (atoms, ions, or molecules that donate electrons to the metal). These complexes can be neutral or charged. When the complex is charged, it is stabilized by neighboring counter-ions. Introduction Coordination chemistry emerged from the work of Alfred Werner, a Swiss chemist who examined different compounds composed of cobalt (III) chloride and ammonia. Upon the addition of hydrochloric acid, Werner observed that ammonia could not be completely removed. He then proposed that the ammonia must be bound more tightly to the central cobalt ion. However, when aqueous silver nitrate was added, one of the products formed was solid silver chloride. The amount of silver chloride formed was related to the number of ammonia molecules bound to the cobalt (III) chloride. For example, when silver nitrate was added to CoCl 3 ·6NH 3 , all three chlorides were converted to silver chloride. However, when silver nitrate was added to CoCl 3 ·5NH 3 , only 2 of the 3 chlorides formed silver chloride. When CoCl 3 ·4NH 3 was treated with silver nitrate, one of the three chlorides precipitated as silver chloride. The resulting observations suggested the formation of complex or coordination compounds . In the inner coordination sphere , which is also referred to in some texts as the first sphere , ligands are directly bound to the central metal. In the outer coordination sphere , sometimes referred to as the second sphere , other ions are attached to the complex ion. Werner was awarded the Nobel Prize in 1913 for his coordination theory. The following table is a summary of Werner's observations: Initial compound Resulting compounds upon adding AgNO3 CoCl3·6NH3 [Co(NH3)6]3+(Cl-)3 CoCl3·5NH3 [Co(NH3)5Cl]2+(Cl-)2 CoCl3·4NH3 [Co(NH3)4Cl2]+(Cl-) CoCl3·3NH3 [Co(NH3)3Cl3] As the table above shows, the complex ion [Co(NH 3 ) 6 ] 3+ is countered by the three chloride ions. The multi-level binding of coordination complexes play an important role in determining the dissociation of these complexes in aqueous solution. For example, \(\ce{[Co(NH3)5Cl]Cl2}\) dissociates into three ions while \(\ce{[Co(NH3)4Cl2 ]Cl}\) dissociates into two ions. By applying a current through the aqueous solutions of the resulting complex compounds, Werner measured the electrical conductivity and thus the dissociation properties of the complex compounds. The results confirmed his hypothesis of the formation of complex compounds. It is important to note that the above compounds have a coordination number of 6, which is a common coordination number for many inorganic complexes. Coordination numbers for complex compounds typically range from 1 to 16. Properties of Coordination Complexes Some methods of verifying the presence of complex ions include studying its chemical behavior. This can be achieved by observing the compounds' color, solubility, absorption spectrum, magnetic properties, etc. The properties of complex compounds are separate from the properties of the individual atoms. By forming coordination compounds, the properties of both the metal and the ligand are altered. Metal-ligand bonds are typically thought of Lewis acid-base interactions. The metal atom acts as an electron pair acceptor (Lewis acid), while the ligands act as electron pair donors (Lewis base). The nature of the bond between metal and ligand is stronger than intermolecular forces because they form directional bonds between the metal ion and the ligand, but are weaker than covalent bonds and ionic bonds. Common Ligands Monodentate ligands donate one pair of electrons to the central metal atoms. An example of these ligands are the halide ions (F - , Cl - , Br - , I - ). Polydentate ligands, also called chelates or chelating agents, donate more than one pair of electrons to the metal atom forming a stronger bond and a more stable complex. A common chelating agent is ethylenediamine (en), which, as the name suggests, contains two ammines or :NH 2 sites which can bind to two sites on the central metal. An example of a tridentate ligand is bis-diethylenetriamine. An example of such a coordination complex is bis-diethylenetriamine cobalt(III). Complex compound/ion Coordination number Oxidation State of Metal Atom [Fe(CN)6]4- 6 2+ [Co(NH3)4SO4]- 5 1+ [Pt(NH3)4]2+ 4 2+ [Ni(NH2CH2CH2NH2)3]2+ 6 2+ Complex ions can form many compounds by binding with other complex ions in multiple ratios. This leads to many combinations of coordination compounds. The structures of certain coordination compounds can also have isomers, which can change their interactions with other chemical agents. The binding between metal and ligands is studied in metals, tetrahedral, and octahedral structures. There are many pharmaceutical and biological applications of coordination complexes and their isomers. References Kleinberg, J.; Argersinger, W J; Grisworld E. Inorganic Chemistry.; D. C. Heath and Company: Boston, MA, 1960; Chapter 6. Martell, Arthur E. Chemistry of the metal chelate compounds.; Prentice-Hall inc: New York, NY, 1952. Oxtoby, David W.; Gillis H. P.; Campion, Alan. Principles of Modern Chemistry, 6th ed.; Thomson Brooks/Cole: Belmont, CA, 2008; Chapter 8. Myers, Thomas R. "Rules for coordination number of metal ions (CEC)." J. Chem. Educ. 1981 , 58 , 681. Syamal, A. "Some improper terms in coordination chemistry." J. Chem. Educ. 1985 , 62 , 143. Toma, Henrique E.; Araki, Koiti; Dovidauskas, Sergio. "A Cyclic Voltammetry Experiment Illustrating Redox Potentials, Equilibrium Constants, and Substitution Reactions in Coordination Chemistry." J. Chem. Educ. 2000 77 1351. 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