Estimate the orders and rate constant K from the results observed for the Reaction. What is the rate when
[H2O2]=[I-]=[H+]=1.0M?
H2O2 + 3I- + 2H+ →I3- + 2H2O
Experiment H2O2 [I-] [H+] Initial rate MS-1
1 0.010 0.010 0.0050 1.15e-6
2 0.020 0.010 0.0050 2.30e-6
3 0.010 0.020 0.0050 2.30e-6
4 0.010 0.010 0.0100 1.15e-6
To estimate the orders and rate constant for the reaction, we can examine the changes in the initial rates with respect to the changes in the concentrations of the reactants.
Let's take Experiment 1 and Experiment 2 to analyze the effect of changing the H2O2 concentration while keeping the I- and H+ concentrations constant:
[H2O2] (M) [I-] (M) [H+] (M) Initial rate (M/s)
0.010 0.010 0.0050 1.15e-6
0.020 0.010 0.0050 2.30e-6
From these two experiments, we can see that doubling the concentration of H2O2 results in doubling the initial rate of the reaction while keeping the concentrations of I- and H+ constant. This indicates that the reaction rate is directly proportional to the concentration of H2O2.
Let's take Experiment 1 and Experiment 3 to analyze the effect of changing the I- concentration while keeping the H2O2 and H+ concentrations constant:
[H2O2] (M) [I-] (M) [H+] (M) Initial rate (M/s)
0.010 0.010 0.0050 1.15e-6
0.010 0.020 0.0050 2.30e-6
From these two experiments, we can see that doubling the concentration of I- results in doubling the initial rate of the reaction while keeping the concentrations of H2O2 and H+ constant. This indicates that the reaction rate is directly proportional to the concentration of I-.
Let's take Experiment 1 and Experiment 4 to analyze the effect of changing the H+ concentration while keeping the H2O2 and I- concentrations constant:
[H2O2] (M) [I-] (M) [H+] (M) Initial rate (M/s)
0.010 0.010 0.0050 1.15e-6
0.010 0.010 0.0100 1.15e-6
From these two experiments, we can see that doubling the concentration of H+ does not affect the initial rate of the reaction while keeping the concentrations of H2O2 and I- constant. This indicates that the reaction rate is independent of the concentration of H+.
Based on these observations, we can conclude that the reaction is first-order with respect to both H2O2 and I-, and zero-order with respect to H+.
The rate constant, K, can be determined by using the rate equation:
Rate = K[H2O2]^a[I-]^b[H+]^c
From the given data, we can substitute the concentrations of [H2O2], [I-], and [H+] from any experiment along with the corresponding initial rate. Let's use Experiment 1:
1.15e-6 = K(0.010)^1(0.010)^1(0.0050)^0
Simplifying this equation, we get:
1.15e-6 = K(0.00010)
Now, solving for K:
K = (1.15e-6) / (0.00010) = 1.15e-2 M^-2 s^-1
Therefore, the rate constant, K, is approximately 1.15e-2 M^-2 s^-1.
To determine the rate when [H2O2] = [I-] = [H+] = 1.0 M, we can substitute these values into the rate equation:
Rate = K[H2O2]^1[I-]^1[H+]^0
Rate = (1.15e-2 M^-2 s^-1)(1.0)^1(1.0)^1(1.0)^0
Rate = 1.15e-2 M^-2 s^-1
Therefore, the rate when [H2O2] = [I-] = [H+] = 1.0 M is approximately 1.15e-2 M^-2 s^-1.
To determine the orders and rate constant for the given reaction, we need to analyze the effect of changes in reactant concentrations on the initial rate.
First, let's examine the effect of changes in [H2O2] while keeping [I-] and [H+] constant in experiments 1 and 2.
Experiment 1: [H2O2] is doubled, and the initial rate is also doubled.
Experiment 2: [H2O2] is further doubled, and the initial rate is again doubled.
From these observations, we can conclude that the reaction rate is directly proportional to the concentration of H2O2. Therefore, the reaction is first order with respect to H2O2.
Next, let's analyze the effect of changes in [I-] while keeping [H2O2] and [H+] constant in experiments 1 and 3.
Experiment 1: [I-] is doubled, and the initial rate remains the same.
Experiment 3: [I-] is further doubled, and the initial rate remains the same.
From these observations, we can conclude that the reaction rate is independent of the concentration of I-. Therefore, the reaction is zero order with respect to I-.
Similarly, let's analyze the effect of changes in [H+] while keeping [H2O2] and [I-] constant in experiments 1 and 4.
Experiment 1: [H+] is doubled, and the initial rate remains the same.
Experiment 4: [H+] is further doubled, and the initial rate remains the same.
From these observations, we can conclude that the reaction rate is independent of the concentration of H+. Therefore, the reaction is zero order with respect to H+.
Now that we have determined the orders for each reactant, we can write the rate equation for the reaction:
Rate = k[H2O2]^1[I-]^0[H+]^0
Simplifying further, the rate equation becomes:
Rate = k[H2O2]
To estimate the rate constant (k), we can use any of the experimental data points. Let's consider experiment 1, where [H2O2] = 0.010 M and the initial rate = 1.15e-6 MS-1:
1.15e-6 MS-1 = k * (0.010 M)
k = (1.15e-6 MS-1) / (0.010 M)
k ≈ 1.15e-4 MS-2
To find the rate when [H2O2] = [I-] = [H+] = 1.0 M, we substitute the concentration values into the rate equation:
Rate = k * [H2O2] = (1.15e-4 MS-2) * 1.0 M
Rate ≈ 1.15e-4 MS-1
Therefore, the rate of the reaction when [H2O2] = [I-] = [H+] = 1.0 M is approximately 1.15e-4 MS-1.