2.4 Coordinate covalent bonding

Figure 2.4.1 – Formation of a Lewis acid-base complex.

An example of a coordination complex is hexaaquo cobalt dichloride, Co(H₂O)₆Cl₂. This compound contains a Co²⁺ ion. This electrophilic metal ion is coordinated by six nucleophilic water ligands. Because the water molecules are neutral, the complex still has a 2+ charge overall. There are two chloride counterions to balance the charge, but the chlorides are not attached to the complex ion.

Notice that the usual Lewis conventions are usually abandoned in drawing coordination complexes of transition metals. With so many different nucleophiles sticking to the Lewis acid, the number of formal charges that must be drawn becomes very cumbersome. Usually the oxidation state of the metal cation is denoted. The oxidation state essentially means the charge on the metal cation and it is written in Roman numerals beside the metal atom. In addition, the overall charge on the Lewis acid-base complex is given, with square brackets indicating that the charge belongs to the entire complex within the brackets. Exactly where that charge resides is up to the reader to consider.

Coordination complexes are frequently useful in mining and metallurgy. For example, nickel can be extracted from nickel ore by converting the nickel into Ni(CO)₄ via addition of carbon monoxide. Ni(CO)₄, or tetracarbonyl nickel, is a gas that can be easily separated from the solid ore. When removed from the presence of carbon monoxide, the coordination complex decomposes back into Ni and CO.

Another example is seen in the liquid extraction of metals from ores. Often, a metal is coordinated by some kind of ligand that changes the solubility of the metal atom, so that it can be extracted from the ore with a solvent. The solvent is usually water, but other liquids could be used. One of the easiest cases to picture on paper is the leaching of gold from gold ores with cyanide.

Often, coordination of a ligand to a metal changes many properties of the metal, in addition to changing its solubility. Transition metals have different oxidation states, meaning they can give up different numbers of electrons and become cations with different charges. Sometimes, this event happens more easily upon coordination. In the case of gold, the gold atom can react with air and become a gold cation in the presence of cyanide.

Cyanide is an anion, so it would be added as a salt, such as sodium cyanide, NaCN, or potassium cyanide, KCN. The resulting complex, Au(CN)₂, is actually a complex anion, and it would be associated with one of the counterions added, such as KAu(CN)₂. It is not completely clear why the gold binds two cyanides, but binding two ligands is common in the chemistry of gold and silver, and is based on molecular orbital considerations that are beyond the level of this course.

This complex salt is water-soluble, so it can be removed from the ore with water. Later, it must be converted back into pure gold via electrochemical reactions. These reactions are not related to acid-base chemistry and will not be covered in this course.

Perhaps some of the most interesting coordination complexes involve transition metals in biology. The most familiar of these compounds is hemoglobin. Hemoglobin is a complex protein that contains a crucial iron atom. The iron is present as a 2+ cation, but the heme ring that binds to the iron has a 2- charge, so overall there is no charge on the complex. The iron is still electrophilic, however, and it can bind an oxygen molecule. Much more oxygen can be carried in the bloodstream this way than if the oxygen simply dissolved in the blood.

There is another interesting event that happens when oxygen binds to the iron in hemoglobin. After binding O₂, the iron actually transfers a single electron to the oxygen, becoming a Fe³⁺ cation. Although deoxyhemoglobin, the Fe²⁺ or Fe(II) species, is purple, many Fe³⁺ or Fe(III) compounds are red. Thus, oxyhemoglobin is red.

It is also interesting to note that we are used to seeing iron-containing materials turn a sort of red-brown color when they have been exposed to oxygen for a long time, especially in the presence of water and salts. Our cars, trains, and bridges eventually rust like the iron in their steel turns to reddish iron oxide. William Tolman, a researcher in bioinorganic chemistry at the University of Minnesota, likes to raise the following question: we rely on iron complexes to carry oxygen through our salt-water bloodstream, so why don’t we get rusty?

The answer is that, in a sense, we do. It has been estimated that after several passes through our bloodstream, a hemoglobin molecule meets its end when the oxygen fails to detach and the iron complex decomposes to iron oxide. That’s one reason why you need to eat a diet that contains iron, to replenish your iron stores in order to make new hemoglobin on a regular basis.

To see some additional problems that will help you get to know coordination compounds, go to the main coordination compound chapter. In particular, it is really useful to be able to count electrons in coordination chemistry. Counting electrons is just like learning Lewis structures, except that instead of going to a maximum count of eight electrons, we can go to a maximum of eighteen electrons. That’s because the noble gas configuration for a transition metal would have eighteen electrons, not eight. To learn how to count electrons in coordination complexes, go here.

Much of reactivity can be understood in Lewis acidity terms. Perhaps the simplest examples of reactions are the formation of Lewis acid-base complexes. In a Lewis acid-base complex, a Lewis base has simply shared a pair of electrons with a Lewis acid, forming a new bond.

Frequently, metal atoms or ions act as Lewis acids. They can often accept electrons from a number of different Lewis bases at once, forming “complexes” or “complex ions” (“complex” meaning they are formed from individual parts that connect together). These Lewis bases (also called “ligands”) are said to be “coordinated” to the metal, meaning they are stuck to the metal via the electron pair that they share with it.

A number of examples of ligands (compounds with lone pairs that can bind to metals) are shown in the following table.

                                                                                      Table 2.4.1 – Some Common Ligands.

“Coordination complexes” play important roles in biology as well as economically important processes. Probably the most familiar coordination complex in biology is hemoglobin. It can coordinate with an additional dioxygen molecule and carry the oxygen through the bloodstream, delivering oxygen to tissues much more efficiently.

A very common coordination complex in industrial use is Wilkinson’s catalyst, (PPh₃)₃RhCl. Wilkinson’s catalyst is used to make a number of transformations more efficient; most notably, it is used in hydrogenation reactions. Chemical transformations of this sort are commonly used in making pharmaceuticals and other high-demand materials.

Because coordination compounds can sometimes be anions or cations, there is a convention used to tell the reader which part of the formula is connected together, and which part is the counterion(s). The part listed in square brackets consists of ligands bonding to a central metal; the part outside the brackets is the counterion(s). For example, [(H₂O)₆Co]Cl₂ consists of a Co²⁺ ion bound to six water molecules. Two separate chloride anions are found nearby.