To solve this question, we need to analyze the Gibbs energy vs temperature plot for the formation of oxides, particularly focusing on the reactions involving FeO and ZnO at 600°C.
The Ellingham diagram provided shows the standard Gibbs free energy change for the formation of various metal oxides as a function of temperature.
The key points to consider are:
Reduction Feasibility: A reducing agent will work if the line for the reducing agent (e.g., C or CO) lies below the line for the metal oxide it aims to reduce. This indicates that the reaction is thermodynamically favorable.
Given Temperature: We need to focus on the plot at 600°C and compare the positions of the lines for carbon, carbon monoxide, zinc oxide, and iron oxide.
At 600°C:
The line for C + O2 → CO2 is below the line for FeO + O2 → Fe2O3. This means that carbon can reduce FeO at this temperature as the Gibbs energy change for combustion of carbon is more negative than that for FeO formation, making the reduction reaction feasible.
The line for C + O2 → CO2 is above the line for ZnO + O2 → Zn, indicating that carbon cannot reduce ZnO at this temperature.
Similarly, the line for 2CO + O2 → 2CO2 is also above both oxide lines, particularly for FeO, ruling out CO as a reducing agent for FeO and ZnO at this temperature.
Conclusion: Based on the diagram, the correct statement is that at 600°C, carbon can reduce FeO since the reaction involving carbon is thermodynamically favorable compared to the formation of FeO.
