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Other gas mixtures

The oxygen required to cause oxidation in the gas phase need not to come from oxygen gas. Consider the following reaction:

2CO (g) + O2 (g) = 2CO2 (g)

For this reaction,

$${K_{{{CO} \over {C{O_2}}}}} = {{p_{C{O_2}}^2} \over {p_{CO}^2.{p_{{O_2}}}}}$$

or $${p_{{O_2}}} = {{p_{C{O_2}}^2} \over {p_{C{O_{}}}^2.{K_{{{CO} \over {C{O_2}}}}}}}$$

and hence

$$\displaylines{ \ln {1 \over {{p_{{O_2}}}}} = \ln {K_{{{CO} \over {C{O_2}}}}} + 2\ln {{p_{C{O_{}}}^{}} \over {p_{C{O_2}}^{}}} \cr = {{ - \Delta G} \over {RT}} \cr} $$

We see that pO2 is equivalent to a ratio: $${{p_{C{O_2}}^{}} \over {p_{CO}^{}}}$$ .

Another nomographic scale may be added to the diagram, with a new origin, C where the CO/CO2 line crosses the y-axis.

Similarly for the reaction 2H2 + O2 = 2H2O; pO2 is equivalent to $${{p_{{H_2}O}^{}} \over {p_{{H_2}}^{}}}$$ Adding a further nomographic scale to the diagram, we see that the equilibrium pressure ratios of CO and CO2 or H2 and H2O for a given oxidation of metal, or reduction of an oxide, can be deduced at a given temperature from the diagram.

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