Question

Consider the following equilibrium system:

CO(g) + H₂O(g) + CO₂(g) + H₂O(9)
Which of the following actions would cause the concentration of H₂O(g) to decrease?

A. Adding CO(9)
B. Removing CO(9)
C. Adding H (9)
D. Adding O(9)

Answer 1

The concentration of H₂O(g) in the equilibrium system CO(g) + H₂O(g) = CO₂(g) + H₂(g) will decrease if additional H️(g) is added due to Le Châtelier's Principle, which states that the system will adjust to minimize the change by shifting the equilibrium to the right.

Step-by-step explanation:

Considering the equilibrium system CO(g) + H₂O(g) = CO₂(g) + H₂(g), adding H️(g) (hydrogen gas) to the system will cause the concentration of H₂O(g) (water vapor) to decrease. This is due to Le Châtelier's Principle, which states that if a change is made to a system at equilibrium, the system will adjust to counteract the change and re-establish equilibrium.


Adding H️(g) will shift the equilibrium to the right according to the equation, resulting in more CO(g) and H₂O(g) reacting to form additional CO₂(g) and H₂(g), thus reducing the concentration of H₂O(g) in the system.

As for the different scenarios given where substances are added or removed from various equilibrium systems:

• Adding more H₂ to the system would shift the equilibrium toward the formation of more CH3OH(g), increasing its concentration while decreasing the concentration of H₂.
  
• Removing CO would shift the equilibrium to the left, increasing the concentrations of H₂ and CO while decreasing the concentration of CH3OH.
  
• Adding CH3OH would shift the equilibrium to the left, increasing the concentrations of H₂ and CO and decreasing the concentration of CH3OH as it's used up to re-establish equilibrium.
  
• Increasing the temperature of the system would depend on the exothermic or endothermic nature of the reaction. If the reaction is exothermic, increasing the temperature would favor the formation of reactants, while if it's endothermic, it would favor the formation of products.

For the problem involving the equilibrium concentrations of CO, H₂O, H₂, and CO₂ at equilibrium with a given K value of 1.34, we would use the equilibrium constant expression to calculate the unknown concentrations.

Finally, when manipulating the volume of a container with a gaseous equilibrium mixture, according to the ideal gas law (PV=nRT), reducing the volume while maintaining the same temperature will increase the pressure of the gases involved. So for the SbCl5 decomposition example, the decrease in volume will cause an increase in the partial pressures of the gases, potentially shifting the equilibrium position.