2A + 3B + C ⟶ D + E
The rate constant for this reaction is 0.25, and the reaction is second order in A and is first order in B and C.
Which option correctly gives the rate of the reaction if [A] is 0.4 M, [B] is 0.1 M and [C] is 0.2 M?
5.0 × 10–4
8.0 × 10–4
2.0 × 10–3
4.0 × 10–3
rate = k(A)^2*(B)(C)
Substitute the numbers and solve for rate. Post your work if you get stuck.
To determine the rate of the reaction, we can use the rate law equation:
Rate = k[A]^m[B]^n[C]^p
where:
- Rate is the rate of the reaction,
- k is the rate constant,
- [A], [B], [C] are the concentrations of reactants A, B, and C, respectively,
- m, n, p are the reaction orders with respect to A, B, and C, respectively.
Given that the reaction is second order in A (m = 2), first order in B (n = 1), and first order in C (p = 1), the rate law equation becomes:
Rate = k[A]^2[B][C]
Plugging in the given concentrations:
[A] = 0.4 M
[B] = 0.1 M
[C] = 0.2 M
Rate = k(0.4)^2(0.1)(0.2)
= k(0.04)(0.1)(0.2)
= 0.0008k
Now, we know the rate constant for the reaction is 0.25.
Rate = 0.0008 * 0.25
= 0.0002
Therefore, the correct option is 2.0 × 10–3.