Iron in aqueous solution has ready access to two stable oxidation states, the ferrous, Fe (II), and the ferric, Fe (III). This property underlies the participa tion of iron in reactions spanning allof biochemistry, including those con trolling the flow of electrons through bioenergetic pathways; the activation of molecular oxygen, nitrogen, and hydrogen; the decomposition of noxious derivatives of oxygen such as peroxide and superoxide; the synthesis of DNA; and the binding of oxygen by hemoglobin, myoglobin, and heme rythrins. Indeed, Neilands (1) has hypothesized that" life, in any form, without iron is in all likelihood impossible." But the same redox properties that make iron so useful in biocatalysis and oxygen transport present hazards to the organism. At the pH, ionic compo sition, and oxygen tension of most physiological fluids the stable state of iron is Fe (III). Fe (II) is readily oxidized to Fe (III) by molecular oxygen, an event easily verifiable each time a rusty precipitate forms in a beaker of ferrous salt standing in air. The hydrolytic propensity of the ferric ion is so great that the equilibrium concentration of free, aquated Fe3+ cannot ex ceed about 1< r17 M in neutral solution (2). Throughout the biological universe, therefore, organisms have beencompelled to evolve specific iron sequestering molecules to maintain the element in soluble form available for transport and for biosynthesis of essential iron proteins and enzymes. By suppressing the catalytic activity of iron in one-electron redox reactions, such iron-binders may also reduce exposure of cells to potentially damaging species such as peroxide, superoxide, and hydroxyl radical (3). In the verte-