A coordination compound is formed when groups of atoms, ions, or molecules chemically bond with each other by donating and accepting pairs of electrons. Groups donating electron pairs are called ligands. They are usually Lewis bases. Groups accepting electron pairs are often transition metal cations. These are usually Lewis acids. Chemical bonds formed in this way are called coordinate-covalent, or dative bonds. As in any covalent bond, two electrons are shared between transition metal and ligand. But in a coordination compound, both electrons come from a pair found on the ligand (Figure 1).
|Table 1. Typical Transition Metal Cations, Mt N+. (Thomson Gale.)|
|Typical transition metal cations, MTN+|
|In this context Iron(II) is read or spoken as 1, iron two. Cobalt(III) is referred to as “cobalt three,” and so on.|
|Table 2 . (Thomson Gale.)|
|Typical Ligands, L|
|Name as Ligand||Ordinary name|
|:C=O carbonyl||carbon monoxide|
|Table 3 . (Thomson Gale.)|
|Examples of highly colored coordination compounds|
|Yellow-Orange||Deep Blue||Deep Blue||Blood-Red|
|Coordination||Used to||Used to|
|unit found in||identify||identify|
|blue print ink||copper as||iron as|
|Cu2+ ions||Fe3+ ions|
The metal cation simply acts as the electron pair acceptor, itself donating no electrons to the bond. Because of the complicated nature of these arrangements, coordination compounds are often called coordination complexes or simply complexes.
It is most common to find six ligands coordinated to a single metal cation. The coordination number is then six. Think of ligands as bees swarming about and stinging a victim. The ligand-bee’s stinger is its lone pair or nonbonding pair of electrons. These special ligand-bees attack their victim in groups of six, although ligand groups of four and other numbers
do occur. Six coordination produces the shape of an eight-sided figure or octahedron. Four coordination produces one of two shapes, a flat square planar or a four-sided tetrahedron.
While nearly all cations can form coordination compounds with ligands, those listed in Table 1 are especially common.
Ligands come in all shapes and sizes, though they are usually nonmetals from the right side of the periodic table. Those listed in Table 2 are typical.
The Swiss chemist Alfred Werner (1866-1919) is called the father of coordination chemistry for his work in clarifying how coordination compounds are assembled. Figure 2 names and explains most of the parts.
Several theories help explain the nature of coordination compounds.
Effective atomic number theory matches the total number of electrons of the transition metal cation plus
donated pairs of electrons from the ligands with the stable electron count of a noble gas atom. In [CoIII (NH3) 6]3+ above the central ion Co3+ contains 24 electrons. (A neutral cobalt atom has 27 electrons.) Each ammine ligand donates 2 electrons for a total of 12 electrons coming from the six ligands.
If the total number of electrons associated with a coordination compound is the same as a noble gas, the compound will frequently have increased stability.
Valence bond theory locates and creates empty orbitals on the transition metal cation. The empty orbitals will accept electron pairs from the ligands. A neutral cobalt atom contains 27 electrons: 1s2 2s2 2p63s2 3p6 4s2 3d7. The 4s, 3d, and empty 4p orbitals are the valence orbitals containing the metal’s valence electrons (Figure 3).
When forming cations, the 4s electrons are lost first, ahead of the 3d electrons.
When Co3+ forms a coordination compound with six electron pair donating ammine ligands, it must have six empty receiving orbitals. But Co3+ only shows four empty orbitals, the one 4s and the three 4p. However, if two electrons in the 3d orbitals of Co3+ move over and pair up, two additional empty 3d orbitals are created. Co3+ now has the required six empty orbitals, two 3ds, one 4s, and three 4ps (Figure 4).
This process is called hybridization, specifically d2s1p3 or simply d2sp3 here.
An important magnetic phenomenon is associated with the valence bond theory. Whenever at least one unpaired electron is present, that substance will be very weakly drawn towards a powerful magnetic. This is known as paramagnetism. Co3+ by itself has four unpaired electrons and is paramagnetic. But Co3+ in [Co(NH 3) 6]3+ has no unpaired electrons. Most substances in nature do not have unpaired electrons and are said to be diamagnetic. These substances are very weakly repelled by a powerful magnetic force. It is possible to measure paramagnetic and diamagnetic effects in a chemical laboratory. Theories such as the valence bond theory can thus be tested experimentally.
Crystal field theory is yet another approach to coordination compounds. It treats ligands as negatively charged anions, attracted to the positively charged transition metal cation.
This is what happens when a chloride anion, Cl–, is attracted to the sodium ion, Na+, in forming the ionic, crystalline compound, table salt, Na+Cl–. A crystal “field” is thus the electronic “field,” produced in any ionic compound. Positive charges are attracted to negative charges and vice versa.
According to crystal field theory, when six negative ligands surround a transition metal cation, the energies of the metal’s 3d electron orbitals are split up. Some 3d orbitals are stabilized and some are energized. The amount of stabilization energy predicted by the theory correlates well with that observed in laboratory experiments.
Many coordination compounds have vivid colors.
Crystal field theory predicts that these colors are due to metal 3d electrons jumping from stabilized to
energized orbitals when visible light or daylight shines on them. The bright colors of coordination compounds make them good candidates for dyes, paint pigments, and coloring agents of all sorts.
A very special application of coordination compounds occurs in what is called the chelate effect. Chelation takes place when ligands “bite” or “bee sting” the metal in more than one place at a time. Using dental “biting” terminology, if a ligand has two “teeth” to bite a transition metal, it is called a bidentate (two-teeth) ligand. One example of this is called ethylenediamine, NH 2 CH 2 CH 2 NH 2. Figure 5 shows simultaneous donation of two electrons each from the two nitrogen atoms of ethylenediamine. A five-membered ring is produced involving the metal, two nitrogen atoms, and two carbon atoms.
A hexadentate ligand such as ethylenediamine tetraacetic acid, EDTA, bites a metal cation in six places simultaneously by wrapping itself around the cation (Figure 6). This produces an extremely strong chelated compound. Using EDTA, such ions as calcium, Ca2+, can be extracted from and tested in drinking water.
Chelated transition metal ions are also found in a wide variety of biochemical situations. A basic structural unit called metalloporphyrin is shown in Figure 7. It can be thought of as several ethylenediamine-like units fused and blended together into a single tetraden-tate (four-teeth) chelating ligand. Changing the central metal, M, changes the biochemical activity of the chelated coordination compound.
Research in this area is opening up a new field of chemistry called bio-inorganic chemistry.
See also Chemical bond.