At a certain temperature the following data was collected for the reaction:
2 ICl + H2 🡪 I2 + 2HCl
Initial Concentration (mol/L)
Initial Rate of Formation of Products (mol/L•s)
[ICl][H2]
0.10
0.10
0.0015
0.20
0.10
0.0030
0.10
0.050
0.00075
Determine the order of reaction for both reactants (show your work and indicate which trials were used to determine order).
Determine k (show your work).
Write the rate law for this set of data, including the k value found in part b.
Here is a better chart for the question. Please answer all the questions.
2 ICl + H2 🡪 I2 + 2HCl
Initial Concentration (mol/L)
[ICl] [H2] Initial Rate of Formation of Products (mol/L•s)
0.10 0.10 0.0015
0.20 0.10 0.0030
0.10 0.050 0.00075
Spacing is a real problem with html. I use periods to space. It isn't perfect but it makes it a little easier to read. I've a trial column.
Trial........[ICl]........ [H2]........... Initial Rate of Formation of Products (mol/L•s)
...1........0.10....... 0.10............... 0.0015
...2........0.20....... 0.10............... 0.0030
....3........0.10....... 0.050............. 0.00075
You determine the order by picking trial 1 and 2. You want the same concentration for one of the components. That give the order to ICl. Then you pick trial 1 and 3 which gives you the order for H2. Trial 1 and 2 are handled this way. This is hard to do on the computer but stay with me. a and be are order numbers. Divide rate 2 by rate 1
......rate = k(ICl)^a(H2)^b
.... 0.0030 = k(0.20)^a(0.10)^b
-------------------------------------------------
......0.0015 = k(0.10)^a(0.10)^b
You see (0.010)^b cancels along with k leaving
2 = 2^a which makes a = 1 so the order for ICl is 1
Do the same for the order for H2 by dividing trial 1 by trial 3.
0.0015 = k(0.10)^a (0.10)^b
------------------------------------------
0.00075 = k(0.10)^a (0.05)^b
Note that k and 0.10)^a cancels leaving
2 = 2^b which makes b = 1 so the order for H2 is 1.
Determine k:
Pick any trial and calculate k
rate = k(ICl)^1*(H2)^1
Plug in the values from the chart and solve for k.
Then you can write the rate law. You should confirm the calculations above. It's late and I could have made a typo.
To determine the order of reaction for each reactant, we will use the method of initial rates.
Step 1: Calculate the initial rate of formation of products for each trial.
Trial 1:
Initial rate = 0.0015 mol/L•s
Trial 2:
Initial rate = 0.0030 mol/L•s
Trial 3:
Initial rate = 0.00075 mol/L•s
Step 2: Determine the order of reaction for the reactant ICl.
We can see that the initial concentration of ICl remains constant at 0.10 mol/L, while the initial rate of formation of products changes. This indicates that the concentration of ICl does not affect the rate of reaction. Therefore, the order of reaction for ICl is zero.
Step 3: Determine the order of reaction for the reactant H2.
We can see that the initial rate of formation of products doubles when the initial concentration of H2 doubles, while keeping the concentration of ICl constant. This indicates that the rate of reaction is directly proportional to the concentration of H2. Therefore, the order of reaction for H2 is one.
Step 4: Determine the rate constant, k.
Using any of the trials, we can substitute the concentration of reactants and the initial rate into the rate law equation:
Rate = k[H2]^1[ICl]^0
For example, using Trial 1:
0.0015 mol/L•s = k(0.10 mol/L)^1(0.10 mol/L)^0
0.0015 mol/L•s = k(0.10 mol/L)
Simplifying, we find:
k = 0.0015 mol/L•s / 0.10 mol/L = 0.015 s^(-1)
So, the rate constant, k, is 0.015 s^(-1).
Step 5: Write the rate law for this set of data, including the k value found in step 4.
Rate = k[H2]^1[ICl]^0
Simplifying, we get:
Rate = 0.015[H2]
Therefore, the rate law for this reaction is Rate = 0.015[H2].
To determine the order of reaction for each reactant, we need to compare the initial rates of formation of products when the initial concentrations of each reactant are changed while keeping the concentrations of the other reactants constant. The order of the reaction with respect to a specific reactant is determined by the change in the rate of the reaction when the concentration of that reactant is changed.
Let's analyze the data provided to find the order of reaction for each reactant:
Trial 1:
[ICl] = 0.10 M
[H2] = 0.10 M
Rate = 0.0015 M/s
Trial 2:
[ICl] = 0.20 M
[H2] = 0.10 M
Rate = 0.0030 M/s
Trial 3:
[ICl] = 0.10 M
[H2] = 0.50 M
Rate = 0.00075 M/s
To find the order of reaction with respect to ICl, we compare Trials 1 and 2 where [ICl] doubles while [H2] remains constant. We can see that the concentration of ICl doubles, but the rate of formation of products also doubles. This indicates that the rate is directly proportional to the concentration of ICl. Therefore, the order of reaction with respect to ICl is 1.
Now let's determine the order of reaction with respect to H2. We compare Trials 1 and 3 where [H2] increases by a factor of 5 while [ICl] remains constant. The concentration of H2 increases by a factor of 5, but the rate of formation of products remains the same. This suggests that the concentration of H2 does not affect the rate of reaction. Therefore, the order of reaction with respect to H2 is 0.
Now let's determine the overall order of reaction by summing up the individual orders of reaction for each reactant. In this case, the overall order of reaction would be 1 + 0 = 1.
To find the rate constant (k), we need to use one of the trials and substitute the values into the rate equation. Let's use Trial 1:
Rate = k[ICl]^1[H2]^0
Rate = k[ICl]
Rearranging the equation, we have:
k = Rate / [ICl]
Using the values from Trial 1: k = 0.0015 M/s / 0.10 M = 0.015 s^-1
Therefore, the rate constant (k) is 0.015 s^-1.
Finally, let's write the rate law for this set of data using the order of reaction for each reactant and the value of k:
Rate = k[ICl][H2]^0
Simplifying further, we have:
Rate = k[ICl]
So, the rate law for this set of data is Rate = 0.015[ICl].