Chemistry
Chemistry
Problem
The reversible reaction A(g) ⇌ B(g) has ΔH = –75 kJ mol⁻¹. At 298 K the system is at equilibrium with K = 4.0 and concentrations [A] = 1.0 M and [B] = 4.0 M in a 1.0-L vessel. The temperature is rapidly raised to 350 K (where K = 1.0), and at that same instant 1.0 mol of B is injected into the vessel, raising [B] to 5.0 M. Which statement best predicts the direction in which the reaction will shift as it re-establishes equilibrium? a) No net shift will occur because, after the changes, already equals the new . b) The reaction will shift toward reactants (formation of ). c) The reaction will shift toward products (formation of ) because adding product drives the reaction forward. d) The reaction will shift toward products until returns to 4.0, the value of the original equilibrium constant at 298 K.Intuitive explanation
This problem tests whether you can predict equilibrium shift by comparing the reaction quotient to the equilibrium constant at the new temperature. The key insight is that tells you the system's current state relative to equilibrium: if , there's too much product and the reaction shifts left (toward reactants); if , there's too little product and it shifts right. Calculate using the concentrations immediately after the temperature jump and B injection, then compare it to the new value at 350 K to determine the direction.
Intuitive explanation
This problem tests whether you can predict equilibrium shift by comparing the reaction quotient to the equilibrium constant at the new temperature. The key insight is that tells you the system's current state relative to equilibrium: if , there's too much product and the reaction shifts left (toward reactants); if , there's too little product and it shifts right. Calculate using the concentrations immediately after the temperature jump and B injection, then compare it to the new value at 350 K to determine the direction.