Magma Redox Geochemistry. Группа авторов

which participates with half‐reaction 1.7 in giving Reaction 1.3 and does not involve explicitly the water solvent.

      4 vertical lines, representing no change of oxidation state but only acid–base reactions (a sole exchange of protons for aqueous solutions), such as(1.23) with the boundary plotted at the pH value for which Q23 = K23 with aFe2O3(s) = (aFe3+)2 = 1 (and also aH2O = 1).

We now see that Reactions 1.3 and 1.4 are both related to half‐reaction 1.22, but they refer to different redox conditions and then have different meanings. In Reaction 1.3 water is the oxidizing agent in acidic conditions, whereas it is the reducing agent in Reaction 1.4 (Appelo and Postma, 1996), which represents a regulating mechanism of O₂, probably the oxidative alteration of rocks containing Fe²⁺, in particular at oceanic ridges.

Figure 1.2 The E‐pH (Pourbaix) diagram for the Fe‐S‐H₂O system at 25°C at 1 bar total pressure and for total dissolved sulfur activities of 0.1 (panel a) and 10^–6 (panel b). Superimposed are the stability fields for H₂S, HS^–, HSO₄^2– and SO₄^2– dissolved species (red lines, traced from data in Biernat and Robins, 1969). Note how the field of stability of pyrite, FeS₂, shrinks and that of magnetite, Fe₃O₄, expands with decreasing total sulfur activity.

      Modifed from Vaughan (2005).

      Basically, E‐pH diagrams demonstrate that breaking of a redox reaction into half‐reactions is one of the most powerful ideas in redox chemistry, which allows relating the electron transfer to the charge transfer associated with the speciation state and the acid–base behavior of the solvent. Superimposing E‐pH diagrams allows a fast recognition of the existing chemical mechanism occurring in an electrolyte medium. For example, Figure 1.2 on the Fe‐H‐O‐S system can be seen as the result of the superposition of stability diagrams for H‐O‐S and Fe‐O‐H system. The resulting diagram in Figure 1.2 shows that the pyrite–magnetite boundary has a negative slope due to half‐reaction:

      (1.24)

      but also a positive slope well visible in Figure 1.2b due to sulfur reduction and dissolution in water as HS^–:

(1.25)

Reaction 1.25 implies of course a positive slope, because H⁺ appears on the right side and electrons on the left. We can also appreciate the reduction of sulfur from pyrite to pyrrhotite at pH > 7:

      (1.26)

      which has a negative slope of –0.0295pH because the number of exchanged electrons is double than protons.

      These concepts can then be transferred to other solvents in which ligand–metal exchanges lead to a different speciation state and are governed by a different notion of basicity, i.e. oxobasicity, such that (see Moretti, 2020 and references therein):

(1.27)

      which can be also related to redox exchanges via the normal oxygen electrode (Equation 1.6), in the same way the normal hydrogen electrode (Reaction 1.7) can be put in relation with the Bronsted‐Lowry definition of acid–base behaviour in aqueous solutions (see Moretti, 2020):

      (1.28)

Figure 1.3 Limit of equilibrium potential‐pO^2– graphs in molten alkali carbonates and sulfates, at 600°C

      (modified from Trémillon, 1974).

It is then possible to define pO^2– = ‐logaO^2– and introduce E‐pO^2– diagrams, in which acid species will be located at high pO^2– values. These diagrams were first introduced by Littlewood (1962) to present the electrochemical behaviour of molten salt systems and provide an understanding of the stability fields of the different forms taken by metals in these systems. Reference potential for molten salt is chosen either from anion or from cation, but anion, making up the ligand, is normally selected because there may be several different cations in the system.

For molten solvent diagrams, such as carbonate and sulfate melts, the stability area of the bath depends on the salt itself and can be seen by using as examples oxyanion solvents (Figure 1.3). Limitations on the pO^2– scale of oxoacidity (Reaction 1.27) are given by the values of the Gibbs free energy of the formation reactions of alkali carbonates or sulfates at the liquid state, which depends on temperature as well as on pressure. On the basic side (low pO^2– side) the limit is imposed by the solubility threshold of the generic M^ν+O_ν/2 oxide in the electrolyte medium, i.e., pO^2–min ≈ M^ν+O_ν/2 solubility, whereas on the acidic side the limit is imposed by P_CO2 or P_SO3 = 1 bar. For example, it is 11 units in the case of the ternary eutectic Li₂CO₃+Na₂CO₃+K₂CO₃ at 600°C and 19.7 units in the case of the ternary eutectic Li₂SO₄ + Na₂SO₄ + K₂SO₄ at the same temperature (Trémillon 1974; Figure 1.3).

      The upper stability limit is related to the O^–II/O_2(g) redox system (Reaction 1.6), i.e., to the oxidation of CO₃^2– and SO₄^2– anions:

(1.29)

(1.30)