Coordination Complexes	Werner's Thoery of Coordination Complexes	Typical Ligands
Typical Coordination Numbers	Lewis Acid-Lewis base Approach to Bonding in Complexes

Test table for Per - File 1 of 1 - rigdoc.zip (09/19) Binaries.

Coordination compounds, such as the FeCl₄^- ion and CrCl₃ 6 NH₃, are called such because they contain ions or molecules linked, or coordinated, to a transition metal. They are also known as complex ions or coordination complexes because they are Lewis acid-base complexes. The ions or molecules that bind to transition-metal ions to form these complexes are called ligands (from Latin, 'to tie or bind'). The number of ligands bound to the transition metal ion is called the coordination number.

Although coordination complexes are particularly important in the chemistry of the transition metals, some main group elements also form complexes. Aluminum, tin, and lead, for example, form complexes such as the AlF₆^3-, SnCl₄^2- and PbI₄^2- ions.

A ski lodge in the Pyrenees has used a helicopter to shift 50 tonnes of snow downhill in an attempt to keep slopes open after an exceptionally mild winter left their runs bare. A temperature rise. The Co atom is on a center of symmetry and the 1,2-di-4-pyridylethane (bpe) ligand also sits across a crystallographic inversion center. The Co atom is octahedrally coordinated by two aqua ligands.

Alfred Werner developed a model of coordination complexs which explains the following observations.

At least three different cobalt(III) complexes can be isolated when CoCl₂ is dissolved in aqueous ammonia and then oxidized by air to the +3 oxidation state. A fourth complex can be made by slightly different techniques. These complexes have different colors and different empirical formulas.

CoCl₃ 6 NH₃	orange-yellow
CoCl₃ 5 NH₃ H₂O	red
CoCl₃ 5 NH₃	purple
CoCl₃ 4 NH₃	green

The reactivity of the ammonia in these complexes has been drastically reduced. By itself, ammonia reacts rapidly with hydrochloric acid to form ammonium chloride.

NH₃(aq) + HCl(aq) NH₄⁺(aq) + Cl^-(aq)

These complexes don't react with hydrochloric acid, even at 100^oC.

Solutions of the Cl^- ion react with Ag⁺ ion to form a white precipitate of AgCl.

Ag⁺(aq) + Cl^-(aq) AgCl(s)

When excess Ag⁺ ion is added to solutions of the CoCl₃ 6 NH₃ and CoCl₃ 5 NH₃ H₂O complexes, three moles of AgCl are formed for each mole of complex in solution, as might be expected. However, only two of the Cl^- ions in the CoCl₃ 5 NH₃ complex and only one of the Cl^- ions in CoCl₃ 4 NH₃ can be precipitated with Ag⁺ ions.

Measurements of the conductivity of aqueous solutions of these complexes suggest that the CoCl₃ 6 NH₃ and CoCl₃ 5 NH₃ H₂O complexes dissociate in water to give a total of four ions. CoCl₃ 5 NH₃ dissociates to give three ions, and CoCl₃ 4 NH₃ dissociates to give only two ions.

Werner explained these observations by suggesting that transition-metal ions such as the Co³⁺ ion have a primary valence and a secondary valence. The primary valence is the number of negative ions needed to satisfy the charge on the metal ion. In each of the cobalt(III) complexes previously described, three Cl^- ions are needed to satisfy the primary valence of the Co³⁺ ion.

The secondary valence is the number of ions of molecules that are coordinated to the metal ion. Werner assumed that the secondary valence of the transition metal in these cobalt(III) complexes is six. The formulas of these compounds can therefore be written as follows.

[Co(NH₃)₆³⁺][Cl^-]₃	orange-yellow
[Co(NH₃)₅(H₂O)³⁺][Cl^-]₃	red
[Co(NH₃)₅Cl²⁺][Cl^-]₂	purple
[Co(NH₃)₄Cl₂⁺][Cl^-]	green

The cobalt ion is coordinated to a total of six ligands in each complex, which satisfies the secondary valence of this ion. Each complex also has a total of three chloride ions that satisfy the primary valence. Some of the Cl^- ions are free to dissociate when the complex dissolves in water. Others are bound to the Co³⁺ ion and neither dissociate nor react with Ag⁺.

The [Co(NH₃)₆]Cl₃ complex dissociates in water to give a total of four ions, and all three Cl^- ions are free to react with Ag⁺ ion.

One of the chloride ions is bound to the cobalt in the [Co(NH₃)₅Cl]Cl₂ complex. Only three ions are formed when this compound dissolves in water, and only two Cl^- ions are free to precipitate with Ag⁺ ions.

H₂O
[Co(NH₃)₅Cl][Cl]₂(s)	Co(NH₃)₅Cl²⁺(aq) + 2 Cl^-(aq)

Once again, the three Cl^- ions are free to dissociate when [Co(NH₃)₅(H₂O)]Cl₃ dissolves in water, and they precipitate when Ag⁺ ions are added to the solution.

H₂O
[Co(NH₃)₅(H₂O)]Cl₃(s)	Co(NH₃)₅(H₂O)³⁺(aq) + 3 Cl^-(aq)

Two of the chloride ions are bound to the cobalt in [Co(NH₃)₄Cl₂]Cl. Only two ions are formed when this compound dissolves in water, and only one Cl^- ion is free to precipitate with Ag⁺ ions.

Werner assumed that transition-metal complexes had definite shapes. According to his theory, the ligands in six-coordinate cobalt(III) complexes are oriented toward the corners of an octahedron, as shown in the figure below.

Any ion or molecule with a pair of nonbonding electrons can be a ligand. Many ligands are described as monodentate (literally, 'one-toothed') because they 'bite' the metal in only one place. Typical monodentate ligands are given in the figure below.

Other ligands can attach to the metal more than once. Ethylenediamine (en) is a typical bidentate ligand.

Each end of this molecule contains a pair of nonbonding electrons that can form a covalent bond to a metal ion. Ethylenediamine is also an example of a chelating ligand. The term chelate comes from a Greek stem meaning 'claw.' It is used to describe ligands that can grab the metal in two or more places, the way a claw would.

Linking ethylene- diamine fragments gives tridentate ligands and tetradentate ligands, such as diethylenetriamine (dien) and triethylenetetramine (trien). Adding four -CH₂CO₂^- groups to an ethylenediamine framework gives a hexadentate ligand, which can single-handedly satisfy the secondary valence of a transition-metal ion.<

Co Dmg 2 4-cn Pyr Br