Iron

Iron is the most abundant transition element in the Earth's crust and, in general, in all life forms. An outline of the distribution of iron in the Earth's crust 20 ,21 is shown in Table 1.2. As can be seen, approximately one-third of the Earth's mass is estimated to be iron. Of course, only the Earth's crust is relevant for life forms, but even there it is the most abundant transition element. Its concentration is relatively high in most crustal rocks (lowest in limestone, which is more or less pure calcium carbonate). In the oceans, which constitute 70 percent of the Earth's surface, the concentration of iron is low but increases with depth, since this iron exists as suspended particulate matter rather than as a soluble species. Iron is a limiting factor in plankton growth, and the rich

fisheries associated with strong up-welling of ocean depths result at least in part from the biological growth allowed by these iron supplies. Properties that dominate the transport behavior of most transition metal ions are: (l) redox chemistry, (2) hydrolysis, and (3) the solubility of the metal ions in various complexes, particularly the hydroxides.

As an example of the effects of solubility, consider the enormous variation in the concentration of iron in rivers, depending on whether the water is from a clear mountain stream running over rock or a muddy river carrying large amounts of sediment. However, the amount of dissolved iron in the form of free ferric ion or its hydrolysis products, whatever the source of water, is extremely low. As can be seen from the solubility of hydrated Fe(III) (Ks ~ 10 -18 M) (Table 1.2), the concentration of free ferric ion is extraordinarily low at neutral pH; so significant concentrations of soluble iron species can be attained only by strong complex formation.

One example of the versatility of iron as a function of its environment is how the ligand field can strongly alter the structural and ligand exchange properties of the metal ion (Figure 1.3). The ligand field can also alter the redox properties. For high-spin ferric ion, as found in the aquo complex or in many other complexes (including the class of microbial iron-transport agents called siderophores, to be discussed later), the coordination geometry is octahedral or pseudo-octahedral. In the relatively weak ligand field (high-spin ground state), the complex is highly labile. In a strong ligand field, such as an axially ligated porphyrin complex of ferric ion, or the simple example of the ferrocyanide anion, the low-spin complex is exchange-inert. Similarly, the high-spin octahedral ferrous complexes are exchange-labile, but the corresponding axially ligated porphyrin complexes, or the ferrocyanide complexes, are spin-paired (diamagnetic) and ligand exchange-inert. Large, bulky ligands or constrained ligands, such as those provided by metalloprotein and enzyme sites, can cause a tetrahedral environment, in which both ferrous ion and ferric ion form high-spin complexes.

The distribution of specific iron complexes in living organisms depends strongly on function. For example, although there are many different iron complexes in the average human, the relative amounts of each type differ more than 650-fold (Table 1.3). The total amount of iron in humans is quite large, averaging more than three and up to five grams for a healthy adult. Most of the iron is present as hemoglobin, the plasma oxygen-transport protein, where the function

of the iron is to deliver oxygen for respiration. A much smaller amount of iron is present in myoglobin, a muscle oxygen-storage protein. For transport, the most important of these iron-containing proteins is transferrin, the plasma iron-transport protein that transfers iron from storage sites in the body to locations where cells synthesizing iron proteins reside; the major consumers of iron in vertebrates are the red blood cells. However, at any given time relatively little of the iron in the body is present in transferrin, in much the same way that at any given time in a large city only a small fraction of the population will be found in buses or taxis. Other examples of iron-containing proteins and their functions are included in Table 1.3 for comparison.

An example of different iron-coordination environments, which alter the chemical properties of iron, is the difference in the redox potentials of hydrated Fe3+ and the electron-transport protein cytochrome c (Table 1.4). The coordination

 

environment of iron in cytochrome c is illustrated in Figure 1.4. For example, the standard reduction potential for ferric ion in acid solution is 0.77 volts; so here ferric ion is quite a good oxidant. In contrast, cytochrome c has a redox potential of 0.25 volts. A wide range of redox potentials for iron is achieved in biology by subtle differences in protein structure, as listed in Table 1.4. Notice the large difference in the potential of cytochrome c and rubredoxin (Figure 1.5), 0.25 volts vs. -0.06 volts, respectively. In polynuclear ferredoxins, in which each iron is tetrahedrally coordinated by sulfur, reduction potentials are near - 0.4 volts. Thus, the entire range of redox potentials, as illustrated in Table 1.4, is more than one volt.