Structure in complex chemistry begins with a focus on "coordination number", the number of ligands attached to the metal (more specifically, the number of σ-type bonds between ligand(s) and the central atom). Usually one can count the ligands attached but sometimes even the counting can become ambiguous. Coordination numbers are normally between two and nine, but large numbers of ligands are not uncommon. The number of bonds depends on the size, charge, and electron configuration of the metal ion. Most metal ions may have more than one coordination number.
Typically the chemistry of complexes is dominated by
interactions between s and p molecular orbitals of the ligands and the d (or f) orbitals of the metal ions. The s, p and d orbitals of the metal give the possibility to allocate 18 electrons (see 18-Electron rule; for f-block elements this extends to 32 electrons). The maximum coordination number for a certain metal is thus related to the electronic configuration of the metal ion (more specifically, the number of empty orbitals) and to the ratio of the size of the ligands and the metal ion. Large metals and small ligands lead to high coordination numbers, e.g. [Mo(CN)8]4-. Small metals with large ligands lead to low coordination numbers, e.g. Pt[P(CMe3]2. Due to their large size, lanthanides, actinides, and early transition metals tend to have high coordination numbers.
Different ligand structural arrangements result from the coordination number. Most structures follow the points-on-a-sphere pattern (or, as if the central atom were in the middle of a polyhedron where the corners of that shape are the locations of the ligands), where orbital overlap (between ligand and metal orbitals) and ligand-ligand repulsions tend to lead to certain regular geometries. The most observed geometries are listed below, but there are many cases which deviate from a regular geometry, e.g. due to the use of ligands of different types (which results in irregular bond lengths; the coordination atoms do not follow a points-on-a-spere pattern), due to the size of ligands, or due to electronic effects
Linear for two-coordination:
Trigonal planar for three-coordination
In chemistry, trigonal planar is a molecular geometry with one atom at the center and three atoms at the corners of a triangle all in one plane. In a nonpolar molecule all the bond angles are 120°, although polar molecules - such as H2CO - will deviate from this ideal geometry. In general the atomic orbitals of a trigonal planar molecule are sp2 hybridized. Examples of molecules with a trigonal planar geometry are boron trifluoride and formaldehyde.
Pyramidalization is a distortion of this molecular shape towards a tetrahedral molecular geometry. One way to observe this distortion is in pyramidal alkenes
A generic trigonal planar molecule showing ideal bond angle.
Tetrahedral or square planar for four-coordination
1-Tetrahedral molecular geometry
In a tetrahedral molecular geometry a central atom is located at the center with four substituents located at the corners of a tetrahedron. The bond angles are 109.5°. This molecular geometry is found for saturated compounds of carbon and silicon. Other molecules with this particular geometry are the perchlorate ion ClO4- and the sulfate ion SO42-.
Inverted tetrahedral geometry
Geometrical constraints in a molecule may cause a severe distortion of a tetrahedral geometry towards an inverted one. In inverted carbon for instance all 4 substituents are now on the same side
organic molecules displaying inverted carbon are tetrahedranes and propellanes. The penalty usually is increase in strain energy for the molecule resulting in increased reactivity.
Note that inversion also takes place in so-called Walden inversion and Nitrogen inversion but with different meanings
The square planar molecular geometry in chemistry describes the spatial arrangements of atoms in a chemical compound
In this geometry co-ordination can be imagined to result when two ligands on the z-axis of an octahedron are removed from the complex, leaving only the ligands in the x-y plane. As the z-ligands move away, the ligands in the square plane move a little closer to the metal. So the orbital splitting diagram for square planar coordination can thus be derived from the octahedral diagram. dz2 falls most in energy. dxy dyz also drop in energy but not as much. Conversely dx2-y2 and dxy increase in energy.
Examples of organometallic compounds with this geometry are Vaska's complex and Crabtree's catalyst