Create an image showcasing a simplified visual representation of a chemistry reaction process. The image should visually demonstrate the chemical change from NO2(g) and CO(g) into CO2(g) and NO(g). This then moves onto the suggested mechanism where 2NO2(g) transforms into NO3(g) and NO(g) and subsequently, NO3(g) and CO(g) transforms back into NO2(g) and CO2(g). The image should not contain any accompanying text.

The mechanism for the reaction described by

NO2(g) + CO(g) ---> CO2(g) + NO (g)

is suggested to be

(1) 2NO2(g) --->(k1) NO3(g) + NO (g)
(2) NO3(g) +CO(g) --->(k2) NO2(g) + CO2(g)

Assuming that [NO3] is governed by steady-state conditions, derive the rate law for the production of CO2(g) and enter it in the space below.

Rate of CO2(g) production = ???

5 answers

k1[No2]^2

Well, let's put on our chemistry clown wig and try to figure this out!

Looking at the proposed mechanism, we can see that the second step involves the reaction we're interested in: NO3(g) + CO(g) ---> NO2(g) + CO2(g).

Since the first step involves the formation of NO3, and we assume that [NO3] is governed by steady-state conditions, we can say that the rate of formation of NO3 is equal to the rate of its consumption.

So, the rate of formation of NO3 can be represented as:

rate(NO3 formation) = k1 * [NO2]^2 (since it's a second order reaction)

Now, let's focus on the second step that forms CO2. According to this step, the rate at which CO2 is formed is proportional to the concentration of NO3 and CO.

rate(CO2 formation) = k2 * [NO3] * [CO]

Since we assumed steady-state conditions, the rate of NO3 formation in the first step (k1 * [NO2]^2) is equal to the rate of NO3 consumption in the second step (k2 * [NO3] * [CO]).

So, we can equate these two expressions:

k1 * [NO2]^2 = k2 * [NO3] * [CO]

Now, let's solve this expression for [NO3]:

[NO3] = (k1 * [NO2]^2) / (k2 * [CO])

Finally, we can substitute this expression for [NO3] into the rate expression for CO2 formation:

rate(CO2 formation) = k2 * [(k1 * [NO2]^2) / (k2 * [CO])] * [CO]

Simplifying this expression, we get:

rate(CO2 formation) = k1 * [NO2]^2

So, the rate law for the production of CO2(g) in this mechanism is:

Rate of CO2(g) production = k1 * [NO2]^2

There you have it! The clownish rate law for the production of CO2 in this reaction is all solved.

student

To derive the rate law for the production of CO2(g) in the given mechanism, we need to determine the rate-determining step.

In this mechanism, the second step (equation 2) involves the reactant CO(g) and the intermediate NO3(g). Assuming that the concentration of NO3 is governed by steady-state conditions, we can express the rate of formation of NO3 (d[NO3]/dt) as zero. This means that the rate of its formation is equal to the rate of its consumption.

Now, let's consider the rate of the second step (equation 2):
Rate2 = k2 * [NO3]*[CO]

Since [NO3] is governed by steady-state conditions, we can substitute it using the first step (equation 1):
[NO3] = k1 * [NO2]^2

Substituting this expression back into the rate of the second step:
Rate2 = k2 * (k1 * [NO2]^2) * [CO]

Simplifying the equation:
Rate2 = k1 * k2 * [NO2]^2 * [CO]

Therefore, the rate law for the production of CO2(g) is:
Rate of CO2(g) production = k1 * k2 * [NO2]^2 * [CO]

Please note that this rate law is provided based on the proposed mechanism and steady-state assumption for [NO3].