Arrows disappeared from below added them back to the equations.

??(a) I don't understand ln(No/N)=akt

The decomposition of IBr(g)into I2 (g) and Br2 (g) is first order in IBr with k = 0.00255/sec.

(a)Starting with [IBr] = 1.50M, what will [IBr] become after 2.50 minutes?

(b)How long, in minutes, will it take for [IBr] to go from 0.500M to 0.100M?

(c)What is the half-life for this reaction in seconds?

(d)Enough IBr is added to an evacuated container to make [IBr] = 0.350M. How long will it take [I2] = 0.100M?

(e) The following mechanism has been proposed for the reaction above:

(step 1)IBr (g) --> I (g) + Br (g)(step 2)IBr(g)+ Br(g) --> Br2(g)+ I (g)(step 3)I(g) + I (g) --> I2 (g)

Based on the rate law described above, which step is the rate-determining step? ______________

What is the molecularity of the rate-determining step? _____________________

To answer these questions, you need to understand the given rate law and the concepts of first-order reactions, half-life, and rate-determining steps. Let's break down each question and explain how to approach it.

(a) Starting with [IBr] = 1.50M, what will [IBr] become after 2.50 minutes?

To solve this, you need to use the first-order reaction equation:

ln(No/N) = akt

where No is the initial concentration of IBr, N is the concentration of IBr at a given time t, k is the rate constant, and t is time.

Rearranging the equation to solve for N, we have:

N = No * e^(-kt)

Substituting the given values: No = 1.50M, k = 0.00255/sec, and t = 2.50 minutes, convert t to seconds (1 minute = 60 seconds):

t = 2.50 minutes * 60 seconds/minute = 150 seconds

Now, plug in these values and calculate N:

N = 1.50M * e^(-0.00255/sec * 150 seconds)

Calculate the exponential term and multiply it by 1.50M to get the final concentration [IBr].

(b) How long, in minutes, will it take for [IBr] to go from 0.500M to 0.100M?

To determine the time required for the concentration to change, you can rearrange the first-order reaction equation:

N = No * e^(-kt)

Rearranging for t:

t = (ln(No) - ln(N))/k

Given No = 0.500M, N = 0.100M, and k = 0.00255/sec, plug these values into the equation and solve for t.

(c) What is the half-life for this reaction in seconds?

The half-life of a first-order reaction can be determined using the equation:

t(1/2) = ln(2)/k

Given k = 0.00255/sec, plug this value into the equation and calculate t(1/2).

(d) Enough IBr is added to an evacuated container to make [IBr] = 0.350M. How long will it take [I2] = 0.100M?

To answer this, we need to know the stoichiometry of the reaction. According to the proposed mechanism:

(step 1) IBr (g) --> I (g) + Br (g)
(step 2) IBr(g) + Br(g) --> Br2(g) + I (g)
(step 3) I(g) + I (g) --> I2 (g)

Based on the stoichiometry, we know that the concentration of I2 formed is half of the decrease in [IBr]. So, for every 1 M decrease in [IBr], [I2] decreases by 0.5 M.

Given [IBr] = 0.350M and [I2] = 0.100M, we can calculate the decrease in [IBr] and convert it to the corresponding decrease in [I2]. Then, divide this decrease by the rate of reaction to find the time required.

(e) Based on the rate law described above, which step is the rate-determining step?

The rate-determining step is the slowest step in a reaction mechanism, and its rate law determines the overall rate of the reaction. You need to analyze the proposed mechanism and identify the step that matches the given rate law. In this case, the rate law is first-order in IBr, so you need to find the step in the mechanism that includes IBr.

Compare the proposed steps with the rate law, and identify the step that includes IBr.

What is the molecularity of the rate-determining step?

Molecularity refers to the number of reactant molecules involved in an elementary reaction step. To determine the molecularity of the rate-determining step, you need to count the number of reactant molecules in that step. In this case, identify the rate-determining step and count the number of IBr molecules participating in that step.