Question

Consider the following system at equilibrium at 698 k: 2hi(g) 2.49 kcalH₂(g) i2(g) if the temperature on the equilibrium system is suddenly increased: the reaction must: a. run in the forward direction to restablish equilibrium. b. run in the reverse direction to restablish equilibrium. c. remain the same. already at equilibrium. the concentration of i2 will: a. increase. b. decrease. c. remain the same.

1) run in the forward direction to restablish equilibrium.
2) run in the reverse direction to restablish equilibrium.
3) remain the same.
4) a. increase.
5) b. decrease.
6) c. remain the same.

Answer 1

Increasing the temperature of the equilibrium system of hydrogen iodide, hydrogen gas, and iodine gas causes the reaction to shift in the reverse direction, increasing H₂ and I₂ concentrations while decreasing HI concentration.

Step-by-step explanation:

When a system at equilibrium has its temperature suddenly increased, it responds according to Le Chatelier's Principle, which predicts how the system will shift to re-establish equilibrium. In the case of the equilibrium system of hydrogen iodide (HI) breaking down into hydrogen gas (H₂) and iodine gas (I₂), which is an exothermic reaction, increasing the temperature will decrease the equilibrium constant.

As a result, to compensate for this increase in temperature, the system will shift in the reverse direction favoring the endothermic process—thus increasing the concentrations of H₂ and I₂ and decreasing the concentration of HI. To put it simply, at higher temperatures, the reaction shifts to the left in the equation H₂(g) + I₂(g) <=> 2 HI(g) to absorb the added heat. Consequently, the correct answer for the effects of increasing temperature on the equilibrium is that the reaction must (b) run in the reverse direction to reestablish equilibrium, and (a) the concentration of I₂ will increase.