Oxidation-reduction potentials in bacteriology and biochemistry

Oxidation-reduction systems play so intimate and so essential a part in living organisms that life itself might be defined as a continuous oxidation-reduction reaction.It is not surprising, therefore, that theoretical speculations and experimental studies on oxidation and reduction processes in animals and plants have been actively pursued since the isolation of oxygen over 150 years ago.The dependence of animal life on the maintenance of an adequate oxygen supply has, rather naturally, dominated the outlook and led to the assumption that oxygen itself is an essential OXIDATION-REDUCTION POTENTIALS involving the taking up of electrons.An academic definition of this sort is isolated in a mental vacuum until it is accepted and absorbed as an integral part of the ordinary point-of-view.By itself such a definition, although possibly of strict accuracy, may appear paradoxical until it is harmonised with every-day laboratory experience.It is therefore a useful exercise to apply the electronic concept of oxidation-reduction reactions to processes not usually regarded electronically.For example, since conversion of metallic silver to silver ion involves loss of an electron it is an oxidation : oxidation Ag~> Ag®In the light of ordinary experience this does not appear an obvious example of oxidation.The silver ion in a silver nitrate solution is not obviously in a higher state of oxidation than metallic silver ; but if the reverse reaction is considered^-for example, the conversion of an alkaline silver nitrate solution to a metallic mirror of silver by warming with a sugar solution-the process is quite obviously a reduction :reduction Ag© > AgFurther consideration of similar processes serves only to confirm the logical basis of the electronic concept of oxidation-reduction processes.One further point should perhaps be emphasised here.It is already clear that oxidation is the reverse of reduction (and reduction is the reverse of oxidation), but another fact is that every oxidation is accompanied by a reduction and vice versa.In the case mentioned above of the ferrous-ferric chloride oxidation-reduction system, the ferrous ion gives up an electron and is thereby oxidised to the ferric ion :oxidation -pe®® > Fe®©© + e but at the same time a chlorine atom takes up that electron and is thereby reduced to a chlorine ion : reduction CI + e > CI© Oxidation cannot proceed unless there is a corresponding reductant to take up the electrons liberated, and, conversely, a substance cannot be reduced unless there is a corresponding oxidation to liberate the necessary electrons.As we have seen, oxidising agents are substances capable of taking up electrons and reducing agents are those able to part with electrons.The readiness with which substances take up, or part with, electrons determines the intensity level of their oxidising or reducing functions.In order to measure the functions quantitatively it is necessary to find a method of measuring electronic escaping tendency, or " fuga- city " as Clark (1923, 1) has described it.Since oxidation and reduction reactions are, by definition, electronic migrations involving exchanges of electric charges, it becomes clear that the quantitative study of oxidation-reduction processes will be effected by measurements of electric potential differences. ELECTRODE POTENTIAL RELATIONSHIPSIt is found, in fact, that if an " unattackable electrode " (such as platinum metal) be immersed in a reversible oxidation-reduction system, a potential difference is set * Ionic concentrations are referred to throughout although, more accurately, ionic activities should be considered in accordance with modern views.This has been done in order to assist lucidity by the avoidance of more complicated symbols and expressions which confuse the issue without any propor- tionate advantage accruing save that of emphasising the complex relationship between Et and pH.OXIDATION-REDUCTION POTENTIALS in which may be substituted the value of [Ox.]/[R0] derived from the mass action equation ( 23).(25) E, = E|] + ^In ^That is to say, the electrode potential may be regarded as being established by the equilibrium between hydrogen ions and hydrogen atoms.The hydrogen concen- tration ([H]) is a function of the partial pressure of hydrogen gas (P), and, in fact, (26) OXIDATION-REDUCTION POTENTIALS with systems of unknown properties, to avoid the use of the term rH on account of its spurious effect of apparent simplicity.With reversible oxidation-reduction systems of the simple type : oxidation (35) Eed ©^==^O x. + e the E^varies by 0-06 volt per unit pH, that is the Ej, -pH curve is said to have 0-06 slope (at 30°C.).In such a system as (36) Ked.e© ^Ox.+ 2e the variation per unit pH can be 0-03, 0-06 or 0-09 volt.These facts apply only to ideal systems when hydration may be neglected.The variation of E^with pH occurs only in those ranges of pH in which the dissociation constants are effective, and the systematic study of all the possible cases is outside the scope of this introductory treatment.A very complete study is reported by Clark and Cohen (1923).It is shown above that in the simplest case at 30°C. (equation 31) :