Question

Step 1: H2(g) + → ICI(9) HI(g) + HCl(g) (slow)
Step 2: HI(g) +ICI(g) → + HCl(g) + I2 (g) (fast)

The reaction is carried out at constant temperature inside a rigid container. Based on this mechanism, which of the following is the most likely reason for the different rates of step 1 and step 2?

a. The only factor determining the rate of step 2 is the orientation of the HI and ICl polar molecules during a collision, but it has a negligible effect when H, and ICl molecules collide.
b. The amount of energy required for a successful collision between H, and ICl is greater than the amount of energy required for a successful collision between HI and ICI.
c. The fraction of molecules with enough energy to overcome the activation energy barrier is lower for HI and ICl than for H2, and ICI.
d. The frequency of collisions between H2, and ICl is greater than the frequency of collisions between HI and ICI.

Answer 1

Answer:

B

Explanation:

Step 1 occurs at a slower rate than Step 2.

Answer 2

The fraction of molecules with enough energy to overcome the activation energy barrier is lower for HI and ICl than for H2, and ICI.

We must bear in mind that the rate determining step in a reaction is the slowest step in the chemical reaction. As such, we can clearly see that step 1 is the rate determining step in the reaction.

The fast reaction must have a minimum energy barrier between the reactants and the products. Hence, the fraction of molecules with enough energy to overcome the activation energy barrier is lower for HI and ICl than for H2, and ICI.

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