Study Notes on Equilibrium Reaction of Iron(II) Oxide and Carbon Monoxide

Reaction Overview

• Chemical Reaction: Iron(II) oxide reacts with carbon monoxide.
  • Equation:
    $ext{FeO (s) + CO (g) } \rightleftharpoons \text{ Fe (s) + CO}_2 (g)$

Equilibrium Constant

• Equilibrium Constant (Ke):
  • At a certain temperature, the equilibrium constant, denoted as $K_e$, is given as:
  • $K_e = 7.47$

Concentration Data During Reaction

• At a certain point in the reaction mixture:
  • Concentration of CO2: 10.9 M
  • Concentration of CO: 1.46 M

Analyzing Change in Concentration of Carbon Dioxide

• Possible Changes in Concentration of CO2:
  • Options for Change:
  1. The concentration of carbon dioxide will increase.
  2. The concentration of carbon dioxide will decrease and then increase.
  3. The concentration of carbon dioxide will decrease.
  4. The concentration of carbon dioxide will not change.

Determining Change in Concentration

• When analyzing this equilibrium system:
  • The reaction shifts towards the right to produce solid iron and carbon dioxide when carbon monoxide is consumed.
  • Given that the mixture currently has an excess of carbon dioxide (10.9 M), as the reaction proceeds:
  • Phase of Reaction: Initial concentrations indicate that the reaction may favor the formation of products if the reagents are depleted.
  • As CO (g) is used up to form Fe (s) and CO2 (g), the concentration of CO2 (g) would initially increase.

Conclusion on Concentration Change

• Final Analysis:
  • Therefore, based on Le Chatelier's principle, the initial increase in CO2 concentration would lead to:
  • The concentration of carbon dioxide will not change significantly throughout the progress of the reaction as long as equilibrium is maintained.
  • The final choice would depend on the specific dynamics of concentration changes but it is expected to stabilize due to equilibrium.

Answer: The concentration of carbon dioxide will not change significantly as equilibrium is reached.