Hi I was hoping to get some help on this. It's review for my chemistry text on Monday and I would be grateful if someone could help me out. My teacher has been out for two weeks but we still have to know this material.
1. The following data was collected for the reaction between hydrogen and nitrogen monoxide. The time
recorded corresponds to a N2 concentration change of 5.3 x 10-4 M.
2H2(g) + 2NO(g) 2H2O(g) + N2(g)
Trial [H2], M [NO], M Time, s
1 0.01 0.02 74
2 0.02 0.02 37
3 0.01 0.01 296
a. Calculate the reaction rate for each of the three trials.
b. Determine the order of the reaction with respect to H2.
c. Determine the order of the reaction with respect to NO.
d. What is the rate law for this reaction?
e. What is the rate constant?
Well, it seems like you're getting quite nitrous about this chemistry review! But don't freaking out, I'll try to help you through it.
a. To calculate the reaction rate for each trial, we need to use the formula
Rate = Δ[H2] / Δt = -1/2 * Δ[NO] / Δt
where Δ[H2] and Δ[NO] represent the change in hydrogen and nitrogen monoxide concentrations over time. Just plug in the values from the table and do the math!
b. To determine the order of the reaction with respect to H2, you need to compare the reaction rates when the concentration of H2 is doubled while keeping [NO] constant. Play around with the values and see how the rate changes.
c. Similarly, to determine the order of the reaction with respect to NO, compare the reaction rates when the concentration of NO is doubled while keeping [H2] constant.
d. The rate law for this reaction will depend on the orders found in parts b and c. So, put on your detective goggles and deduce the rate law based on your findings.
e. Finally, the rate constant is typically represented by the symbol "k." It can be found by plugging in the values from the table into the rate law equation and solving for k.
Remember, chemistry can be challenging, but don't let it scare you away like helium-filled balloons at a party! Just keep practicing, and you'll soon be able to balance equations like a true chemist. Good luck with your review!
To calculate the reaction rate for each of the three trials, we can use the formula:
Rate = Δ[H2] / Δt = -Δ[NO] / Δt
Trial 1:
[H2] = 0.01 M, [NO] = 0.02 M, time = 74 s
Δ[H2] = 0.01 M - 0 M = 0.01 M
Δ[NO] = 0.02 M - 0 M = 0.02 M
Δt = 74 s
Rate = 0.01 M / 74 s = 1.35 x 10^-4 M/s
Trial 2:
[H2] = 0.02 M, [NO] = 0.02 M, time = 37 s
Δ[H2] = 0.02 M - 0 M = 0.02 M
Δ[NO] = 0.02 M - 0 M = 0.02 M
Δt = 37 s
Rate = 0.02 M / 37 s = 5.41 x 10^-4 M/s
Trial 3:
[H2] = 0.01 M, [NO] = 0.01 M, time = 296 s
Δ[H2] = 0.01 M - 0 M = 0.01 M
Δ[NO] = 0.01 M - 0 M = 0.01 M
Δt = 296 s
Rate = 0.01 M / 296 s = 3.38 x 10^-5 M/s
Now, let's determine the order of the reaction with respect to H2.
Comparing trials 1 and 2, we see that doubling the concentration of H2 (from 0.01 M to 0.02 M) leads to a fourfold increase in the reaction rate (1.35 x 10^-4 M/s to 5.41 x 10^-4 M/s). This suggests that the reaction is second order with respect to H2.
Using a similar logic, we can determine the order of the reaction with respect to NO.
Comparing trials 1 and 3, we see that doubling the concentration of NO (from 0.02 M to 0.01 M) leads to a twofold decrease in the reaction rate (1.35 x 10^-4 M/s to 3.38 x 10^-5 M/s). This suggests that the reaction is first order with respect to NO.
Therefore, the rate law for this reaction is:
Rate = k[H2]^2[NO]
Finally, let's find the rate constant (k) using any of the trials. Let's use trial 1.
Rate = k[H2]^2[NO]
1.35 x 10^-4 M/s = k(0.01 M)^2(0.02 M)
k = (1.35 x 10^-4 M/s) / (0.01 M)^2(0.02 M)
k = 6.75 M^-3 s^-1
So, the rate constant for this reaction is 6.75 M^-3 s^-1.
To calculate the reaction rate for each of the three trials, you need to use the data given in the table. The reaction rate can be determined by dividing the change in concentration of one of the reactants by the time taken for the reaction to occur.
For Trial 1:
[H2] = 0.01 M
[NO] = 0.02 M
Time = 74 s
Change in [NO] = 5.3 x 10^-4 M
Change in [NO] = [NO]final - [NO]initial
Change in [NO] = 0.02 M - 5.3 x 10^-4 M
Change in [NO] = 0.01947 M
Reaction rate = Change in [NO] / Time
Reaction rate = 0.01947 M / 74 s
Reaction rate = 2.63 x 10^-4 M/s (rounded to 3 significant figures)
Similarly, you can calculate the reaction rate for Trials 2 and 3 using the same formula.
To determine the order of the reaction with respect to H2, you need to compare the reaction rates for Trials 1 and 3. If the reaction rate doubles when the concentration of H2 doubles, the reaction is first-order with respect to H2.
To determine the order of the reaction with respect to NO, you need to compare the reaction rates for Trials 1 and 2. If the reaction rate remains constant when the concentration of NO doubles, the reaction is zero-order with respect to NO.
The rate law for the reaction can be written as:
Rate = k [H2]^x[NO]^y
To determine the order of the reaction overall, you can add the individual orders of reaction with respect to H2 and NO. In this case, it would be first-order with respect to H2 (x = 1) and zero-order with respect to NO (y = 0). Therefore, the overall order of the reaction would be 1 + 0 = 1.
The rate constant (k) can be determined by substituting the values from any of the trials into the rate law equation and solving for k. Let's use Trial 1:
Rate = k [H2]^1[NO]^0
2.63 x 10^-4 M/s = k (0.01 M)^1(0.02 M)^0
k = 2.63 x 10^-4 M/s / 0.01 M
k = 2.63 x 10^-2 1/s (rounded to 3 significant figures)
So, the rate constant for the reaction is approximately 2.63 x 10^-2 1/s.