Redox/Electrochemical Properties

Oxygen Reactivity

Molecular oxygen is a remarkable molecule. The oxygen molecule exists in a reduced form and therefore can serve is a resonably strong oxidant. The reactive potenital or more specifically the electrochemical desire to react means oxygen can drive a range of reduction chemical reactions; yet kinetically oxygen prevented from doing so. In summary the molecule is thermodynamically reactive yet kinetically unreactive. The reason for this is unusual behaviour the triplet ground state of the molecule which prevents reactions with singlet state molecules proceeding as they are "spin forbidden" and occur only slowly at room temperature. Many processes in biology utilizing the thermodynamic reactivity of oxygen have evolved and implement metal centred catalysts to accelerate the reaction. The limited reactivity of O₂ due to the triplet/singlet spin problem, and the fact that O₂ has limited solubility in water ([O₂] 1 atm = 1mM, or from air ~ 250 µM at 25°), severely limits the chemical reactions of this molecule. These factors are the main reasons there is some 4 x 10¹⁹ mols (1 x 10¹⁵ tons) of O₂ accumulated in the atmosphere.

Oxygen Intermediates

Oxygen and water serve as opposite ends of a redox scale. The 4 electron oxidation of water to O₂ is reversed by the 4 electron reduction to H₂O. Along this pathway there is a possibility for the generation of the superoxide anion (O₂•–), hydrogen peroxide (HOOH), or hydroxyl radicals (OH•). In biology <link> these intermediates are generally undesirable. In other applications the generation of these intermediates is used for synthetic chemical means. Reactions with O₂ proceed via electron or H-atom transfer reactions.

Redox reactions can be summarised as two half reactions: an oxidation and a reduction half reaction. If the sum of the half potentials is positive then the reaction will proceed spontaneously. The concepts on the left summarise the reduction reaction and the oxidation reaction.

Redox potentials for Oxygen

The standard redox 1/2 potentials (E_M°) of O₂ redox chemistry are summarised below. The O₂ redox potentials are strongly pH dependent (use pH menu below to illustrate this). Additionally, the O₂ redox potentials dependent strongly on the pathway. Thus the lowest energetic pathway is the concerted 4-electron pathway from O₂ to H₂O with only 0.82 V at pH 7. When reduction involves a metal electrode, the chemistry is generally is associated with a metal-hydroxide activated surface that acts as a catalyst.

Sawyer, DT (1991) "Oxygen Chemistry" Oxford University Press

The potentials can be expressed as

Where (Delta) G is the free energy of the reaction, n the number of moles, F the Farraday constant (96.485 kJ/ V mol ) and E the standard potential. Aspects relating to the pH dependence of the reaction are mathematically derived in the Nernst equation. This equation rationalises the pH on the equilibrium of the reaction and translates to -59 mV change per unit increase in pH (thus -59 mV per decade in equilibrium).

Ozone redox potentials

Ozone (O₃) is also a potent oxidant and is generated via electrochemical discharge in O₂ (not via direct electrochemistry on an electrode). A variety of ozone generators exist for commercial production and large amounts are generated on site, at low cost, at literally thousands of companies around the world. It is a very useful product.

The large redox potentials for Ozone reduction are indicative of the use of this molecule as a oxidant. This property translates to a useful chemical industry uses for sterilisation, cleaning, detoxification processes and has no direct waste products itself. Many of the applications that have traditionally used Chlorine over the years have been able to make a Green Chemistry transition using ozone.

Redox Chemistry A fantastic summary of electrochemical terms and concepts:

Nernst Equation defined as

Molecular Oxygen (O₂)