Activities of phosphorus in copper-nickel liquid alloys saturated with solid copper-nickel solid solutions at 1573 K
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9/12/03
10:26 AM
Page 741
Authors’ Reply H.Y. SOHN and S. PALDEY Five years since its publication, it is gratifying that our article is still attracting the interest of researchers from different parts of the world. With regard to the question about the application of the Gibbs phase rule in choosing our experimental conditions, we fixed the inlet partial pressures of Mg, TiCl3, and AlCl3 under a fixed temperature and total pressure (which then fixes the argon partial pressure). This fixes the system at equilibrium. It is true that the phase rule can be satisfied for this reaction system by less restrictive experimental conditions. Our statement at the beginning of page 458 of the original article should have said, “According to the Gibbs phase rule, in a fivecomponent system, when total pressure, temperature, argon partial pressure, and Mg/Cl ratio are fixed, . . . .” (Fixing the argon pressure together with Mg/Cl in the input is tantamount to fixing the argon pressure at equilibrium, although the two values may be different.) It is noted, however, that different input conditions satisfying the same phase rule (i.e., the same Mg/Cl ratio with different Mg/Ti or Al/Ti ratios) will yield different relative amounts of the condensed-phase products at equilibrium. Thus, the satisfaction of the Gibbs phase rule does not uniquely fix the experimental conditions to produce a solid mixture of a specific equilibrium composition. Therefore, the choice of experimental conditions solely based on the degree of freedom according to the Gibbs phase rule, as Ghosh suggests, does not ensure the equilibrium solid product to have a desired overall composition. As to the possible formation of liquid MgCl2, while the thermodynamic calculation of Ghosh is a worthy exercise, it must be remembered that the experimental work was not designed to necessarily create equilibrium. Even for the experimental run in which the conversion of the limiting reactant TiCl3 vapor reached the highest value at 94 pct, the attainment of equilibrium would have entailed essentially complete consumption (about 99.94 pct) of TiCl3 vapor, over 90 pct of AlCl3, and some 99.95 pct of Mg, according to Ghosh’s own calculation. In addition to the lower than equilibrium conversion of TiCl3, the experimental conversion of AlCl3 in this run was even greatly lower at 8 pct. This indicates that the experimental values of TiCl3 and AlCl3 partial pressures were greatly higher than the calculated equilibrium values and that the MgCl2 partial pressure was much lower than the equilibrium value. In most other runs, conversion was significantly less than in this run. Thus, it is highly unlikely that liquid MgCl2 was present in our experiment. Additionally, the reliability of the thermodynamic data going as far back as 1967, especially involving those for titanium aluminides, must be questioned for the purpose of ascertaining the difference in the calculated MgCl2 partial pressure and its vapor pressure. Aside from the presence or absence of liquid MgCl2, another likely reason
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