R07 Table showing several standard reduction potentials

Using illustration R5 it has been shown that the reducing power of redox couples i.e. the potentiality to donate electrons increases in the order:

and using illustration R6 it has been shown that the oxidising power and halogen redox couples i.e. the potentiality to accept electrons increases in the order:

this qualitative information on the relative reducing and oxidising power of redox couples is quantified by measuring their standard reduction potentials. Standard reduction potentials of half-reactions are referred to the standard hydrogen electrode.

Standard hydrogen electrode

The standard hydrogen electrode has, by convention, a standard reduction potential of zero volts at unit effective concentration, a temperature of 298 K and a hydrogen pressure of one atmosphere. The hydrogen electrode consists of a platinised platinum electrode immersed in a 1 mol/L solution hydrogen ions. Hydrogen gas at a pressure of 1 atm is bubbled over the platinum electrode. On the surface of the platinum an equilibrium is established between hydrogen gas and hydrogen ions.

A potential develops on the surface of the platinum. It is assigned a value of zero volts.

Determination of standard reduction potentials

Standard reduction potentials are determined by connecting the system represented in the half-reaction equation concerned to a standard hydrogen electrode via a salt bridge and reading off the potential difference on either a previously calibrated potentiometer or a high resistance voltmeter. The standard hydrogen electrode is by convention always placed on the left.

The configuration of the cells whose standard reduction potential is being measured depends upon the species involved. In the case of a metal ion/metal couple the metal is present as an electrode immersed in a 1mol/L solution of its ions at 25° C. In the case of an ion/ion couple a platinum electrode is immersed in a solution containing 1 mol/L of each of the ions at 25° C. Finally, in the case of a gas/ion couple the gas at a pressure of 1 atm is bubbled over a platinum electrode immersed in a solution of 1 mol/L of ions at 25° C.

Table of standard reduction potentials

When the values for are tabulated it is the convention to write the most strongly reducing redox couple at the top of the table and the most strongly oxidising couple at the bottom.
he table of standard reduction potentials in illustration R7 shows the values for redox half-reactions which have been measured with respect to the standard hydrogen electrode as reference.
It will be noticed that the value of the redox half-reaction equations at the top of the table are large and negative, while those for the half-reaction equations at the bottom of the table are large and positive. Strong reducing agents and hence the least oxidising redox couples, such as Na⁺/Na and Zn²⁺/Zn have large negative values, whereas the least reducing and hence most oxidising redox couples such as Cl₂/2Cl^- and F₂/2F^- have large positive values and are to be found at the bottom of the table.

Use of standard reduction potentials

This table can be used to predict the outcome of classroom experiments and also to explain various industrial processes which are based on redox half-reaction equations by using their values.

Bromine is produced industrially by bubbling chlorine gas through a solution of bromide ions. The same principle cannot be used to produce chlorine industrially, because the installation necessary for the production of fluorine is too expensive. Other oxidising agents such as the , which can be used in the laboratory scale production of chlorine, are also too expensive for large scale use.

Chlorine is produced industrially by the electrolysis of an aqueous solution of sodium chloride. The process is not self-sustaining and electrical energy must be constantly supplied. Electrolysis is the only method of producing fluorine because there are no oxidising agents which are more powerful than fluorine.
The table also clarifies why the alkali metals such as sodium (Na) react so violently with the hydrogen ions in neutral water. The half-reaction equations of these metals are above that of hydrogen ions and so they reduce them. Zinc only reduces water in the presence of additional H⁺ ions, i.e. in an acid medium ( see illustration R9 for the influence of ion concentration on reduction potentials).

The table also quantifies the qualitative difference in oxidising power of the Zn²⁺/Zn, Pb²⁺/Pb and Cu²⁺/Cu

and quantifies the redox reactions depicted in illustration R5. It is the convention for measuring the cellpotential (E_cell) or electromotive force (e.m.f.) of electrochemical cells always to place the more negative electrode system (the anode) on the left :

This leads to :

because we substractthe reduction potential of the left electrode from the reduction potential of the right electrode.

On the basis of the standard half-cell redox potentials of Zn²⁺/Zn and Cu²⁺/Cu one would expect Zn metal to play the role of the anode (Zn oxidised to Zn²⁺) and Cu²⁺ the role of the cathode (oxidising agent, being reduced).
Then :

This positive value for E_cell ensures a self-sustaining redox reaction between Zn and Cu²⁺, even when energy is not (continously) supplied to the system.

Important reducing and oxidising agents.

The table given below shows many more redox halfreaction equations than have been considered in the preceding series of experiments. They have been written in accordance with the IUPAC convention, all values are standard reduction potentials and the letters (g) (l) (s) and (aq) have been added to show the physical state: gas, liquid, solid and aqueous respectively. These letters may be written on the same level as the symbol and immediately next to it; e.g. Li+(aq) or

Compounds with both oxidising and reducing properties

There are situations in which the ions indicated by a halfreaction equation can function as both an oxidising and reducing agent depending on the other ions in the reaction and the type of medium. Hydrogen peroxide can, for example, function as both an oxidising and as a reducing agent. In the table (left) the substance H₂O₂ is found top right and bottom left :

Simultaneous oxidation and reduction of a compound is known as disproportionation. The fact that it can behave in this way does not affect neutral aqueous hydrogen peroxide when being stored . Aqueous solutions only containing H₂O₂ remain stable. Decomposition only occurs when H₂O₂ comes into contact with a catalyst such as Fe²⁺, Fe³⁺, I-, MnO₂, blood, etc. It is commonly used for cleaning contact lenses, bleaching hair and as a disinfectant.
Ammonium nitrate (NH₄NO₃) is another compound that can disproportionate, but this is not linked to the type of medium being used and a real catalyst does not need to be present. This can be explained by using the above table from which it can be seen that the ion is at the top right-hand side of the table and is therefore a strong reducing agent. The nitrate ion () is at the bottom left-hand side of the table and is therefore a possible strong oxidising agent. In an aqueous solution or in the solid, the and ions can remain stable.
A jolt however can be sufficient to start a disproportionation or autoxidoreduction reactions.
Ammonium nitrate was used for many years as a fertilizer but it is also a powerful explosive, decomposing on heating:

Ammonium dichromate (NH₄)₂Cr₂O₇(s) also reacts in a similar way on heating decomposing spectacularly from orange crystals to a large heap of green flakes.

Thermodynamic capability versus kinetic reality

Above we were clearly confronted with examples of reductors and oxidants which, even strongly mixed, do not proceed to an electron transfer or a redox reaction.
Consider H₂O₂, that does not disintegrate spontaneously, or NH₄NO₃, that does not explode spontaneously.

The fact that one reagent is an element to be found at the top right and the other reagent is an element to be found at the bottom left of the standard potential table does not always guarantee a self-sustaining redox reaction.
Some cases require something more: an impulse, a push or a poke, in the form of temperature or in the form of a catalyst. In this way H₂O₂, which can safely be stored in (dark) bottles, will yet start to spontaneously disintegrate as soon as a pinch of a catalysing substance is added to it. Interesting to see here is the possible impact of iron (II) salts and iron (III) salts: their half-reaction is somewhere between the two half-reactions for hydrogen peroxide.

Useful as it may be, the table of standard potentials does not allow prediction with certainty of which selfsustaining redox reactions are actually possible.
It is said that the table of standard potentials is normative for the thermodynamic tendency of reductors and oxidants to react in a self-sustaining way. No matter how big the difference between their E°-values is, this does not say anything about the kinetics of the redox conversion, i.e. the ease and the speed with which this can happen in practice.

A lack of visible reaction upon combination of redox couples from the tables, which thermodynamically should produce a self-sustaining reaction, may indicate a kinetic problem or could also indicate the formation of an insoluble salt on the surface of the metal, especially if we think of lead in an aqueous solution of zinc ions. The zinc salt should be choses with care and zinc nitrate is probably the most suitable for use here.