Geology Reference
In-Depth Information
Richardson and Jeffes (1948) improved the Ellingham diagram by adding a pres-
sure scale. This facilitates the finding of the partial pressure of oxygen on an oxide
line at a given temperature by drawing a line between this point and the graph's
origin at T = 0 and G = 0. Then, if one extends this line to the scale labelled
“p O 2 ” as shown on the right hand side, one is able to obtain oxygen's partial pres-
sure at equilibrium. This in turn leads to the control of reduction processes since
this pressure can be decreased either in vacuum conditions, in an inert atmosphere
such as argon, or with a chemical reactant that traps oxygen.
Likewise, it is possible to add another nomographic scale in order to obtain the
minimum ratio, CO=CO 2 , able to reduce a given oxide. This procedure is similar to
the oxygen scale but now, the line must be drawn between point “C” of the diagram
and that which defines the temperature where a given metal oxide is to be reduced.
The scale is labelled “CO=CO 2 ratio”. This ratio increases as the metal becomes
harder to reduce.
Water decomposition can also be represented on the diagram providing the equi-
librium pressure H 2 =H 2 O ratio for a given metal oxide. Now the line needs to start
at point “H” and connect with that which defines the temperature where a given
metal oxide is to be reduced.
The Ellingham diagram can also represent the variation of the Gibbs function
with temperature for other conversion processes such as the calcination of metal
sulphides. The idea behind the Ellingham diagram may also be extended to other
processes including chlorination or for atmospheres with varied gas mixtures. There-
fore, it constitutes a key tool in the understanding of the chemistry and thermody-
namics of all smelting processes, particularly in pyrometallurgy. For further reading,
see for instance Greenwood and Earnshaw (1984).
9.4.3.2 A thermodynamic overview of hydro- and electrometallurgy
As described in Chap. 8, hydrometallurgy separates metal cations from leached
solutions. This can be done via direct oxidation (for instance CuS with Fe 3+ ) or
reduction (MnO 2 with SO 2 ) of the metal in aqueous solution or with the addition
of acids (such as for ZnO) or bases (for Al 2 O 3 ). Meanwhile, electrometallurgy,
recovers metals via electrolysis, either in an aqueous solution or in a molten salt
acting as an electrolyte. In both cases, a redox process takes place in which the
oxidised (Ox) species gains electrons at the expense of the reduced (Red) one. The
redox reaction, either in an aqueous solution or in an electrolytic cell can be written
as:
Ox + ne ! Red + G
(9.17)
The energy released in this electron exchange is due to the potential disparity
between the species (i.e. their chemical potentials) and the difference between
the chemical potential of products and reactants is G. In a cell this maximum
difference manifests itself as an electromotive force E, measured in Volts, which
 
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