Falling through the activity series of the metals

In this diagram, electron donors (otherwise known as reducing agents or reductants) are shown on the left, and their conjugate oxidants (acceptors) are on the right.

The vertical location of each redox couple represents the free energy of an electron in the reduced form of the couple, relative to the free energy of the electron when attached to the hydrogen ion (and thus in H2).

An oxidant can be regarded as a substance possessing vacant electron levels; the "stronger" the oxidizing agent, the lower the energy of the vacancy (the sink). If a reductant is added to a solution containing several oxidants, it will supply electrons to the various empty levels below it, filling them from the lowest up. Note however, that electron transfer reactions can be very slow, so kinetic factors may alter the order in which these steps actually take place.

H2 and H2O. Locate the couples involving these two elements within the vertical section labeled "water stability range" (light blue background). The metals above the H2/H+ couple are known as the active metals because they can all donate electrons to H+, reducing it to H2 and leaving the metal cation. In other words, H+ can serve as an electron sink to these metals, which are therefore attacked by acid. But since some H+ is always present in water, all of these metals can react with water. Generally the higher they are, the more readily they react. With zinc and below, reaction with water is so slow at room temperature as to be negligible, but these metals will be attacked by acidic solutions, in which the concentration of H+ ions is much greater. Those metals that are below hydrogen in this table are not attached by H+ and are referred to as the noble metals. (Gold, Au, just below chlorine, is the noblest of all.)

The species on the right side below the H2O/O2 couple can all serve as electron sinks to water and will oxidize it to O2. However, this reaction can be extremely slow; only F2, the strongest of all the oxidizing agents (at the very bottom of the table) reacts quickly.

It turns out, then, that only those redox pairs situated within the water stability region are thermodynamically stable in aqueous solution; all others will tend to decompose the water.

Three scales of free energy are shown. The one on the left corresponds to the standard reduction potential of the couple, which is the free energy per electron-mole (recall the relation
G = –nFE°.)

The rightmost scale gives the corresponding energy in kiloJoules per mole of electrons transferred, so it applies directly only to a half reaction written as a one-electron reduction.

About the pE scale. It can be shown that pE = (log K)/n for the n-electron reduction of the oxidant by H2

Just as pH is a measure of the availability of protons in the solution, so the pE represents the availability of electrons; thus the more negative the pE, the more ``reducing'' is the solution, and the greater will be the fraction of each couple in its reduced form, with the lower ones being most strongly affected.

Although pE is almost never encountered in introductory courses, it is widely used in environmental chemistry and geochemistry.

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