Question

For the reaction A+B+C -> D+E, the initial reaction rate was measured for various initial concentrations of reactants. The following data were collected:

Trial [A] (M) [B[ (M) [C] (M) Initial rate (M/s)

1 0.20 0.20 0.20 6.0 x 10^-5

2 0.20 0.20 0.60 1.8 x 10^-4

3 0.40 0.20 0.20 2.4 x 10^-4

4 0.40 0.40 0.20 2.4 x 10^-4



Reaction order respect to A = 2

Reaction order in respect to B = 0

Reaction order in respect to C = 1



The value of the rate constant k for this reaction = 7.5*10^-3 M^-2 * s^-1



Given the data calculated in Parts A, B, C, and D, determine the initial rate for a reaction that starts with 0.75M of reagent A and 0.90M of reagents B and C?

Answer 1

Answer: The initial rate of the reaction is $3.8\times 10^{-3}$

Explanation:

Rate law is defined as the expression which expresses the rate of the reaction in terms of molar concentration of the reactants with each term raised to the power their stoichiometric coefficient of that reactant in the balanced chemical equation.

For the given chemical equation:

$A+B+C\rightarrow D+E$

Rate law expression for the reaction:

$\text{Rate}=k[A]^a[B]^b[C]^c$

where,

a = order with respect to A = 2

b = order with respect to B = 0

c = order with respect to C = 1

k = rate constant = $7.5\times 10^{-3}M^{-2}s^{-1}$

Calculating the rate of the reaction:

$[A]=0.75M$

$[B]=0.90M$

$[C]=0.90M$

Putting values in above equation, we get:

$\text{Initial rate}=7.5\times 10^{-3}\times (0.75)^2\times (0.90)^0\times (0.90)^1\\\\\text{Initial rate}=3.8\times 10^{-3}$

Hence, the initial rate of the reaction is $3.8\times 10^{-3}$