Consider the following for the reaction at 300 K:3 ClO - (aq)   ClO3 - (aq) + 2 Cl - (aq) ExperimentInitial [ClO-] (M) Initial Rate of Formation of ClO3 - (aq) (M/min) 10.4521.048 * 10-420.9034.183 * 10-4 (7) (4 pts) What is the order of the reaction with respect to ClO - (aq) ? A) 0B) 1C) 2 (8) (4 pts) For experiment #2, what is the initial rate of consumption of ClO-

Second order

Δ[ClO⁻]/Δt = - 4.183 x 10⁻⁴ M/min

Explanation:

Given the dа ta:

Experiment # [ClO-] (M) Initial Rate of Formation of ClO3 - (M/min)

1 10.452 1.048 x 10⁻⁴

2 20.903 4.183 x 10⁻⁴

we need to determine the order of the reaction with respect to ClO⁻.

We know the rate law for this reaction will have the form:

Rate = k [ClO⁻]^n

where n is the order of the reaction. Thus, what we need to do is to study the dependence of the initial rate on n for the experiment.

If the reaction were zeroth order the rate would not change, so we can eliminate n = 0

If the reaction were first order, doubling the concentration of [ClO-], as it was done exactly in experiment # 2, the initial rate should have doubled, which is not the case.

If the reaction were second order n: 2, doubling the concentration of [ClO-], should quadruple the initial rate of formation of ClO3-, which is what it is observed experimentally. Therefore the reaction is second order respect to ClO-.

The initial rate of consumption of ClO⁻ is the same as the rate of formation of ClO₃⁻ since:

Δ = - Δ[ClO⁻]/Δt = + Δ[ClO₃⁻]/Δt = + 1/2 [Cl⁻] / Δt

where t is the time.

from the coefficients of the balanced chemical equation.

- Δ[ClO⁻]/Δt = + Δ[ClO₃⁻]/Δt = + 1/2 [Cl⁻ ] = rate

Δ[ClO⁻]/Δt = - 4.183 x 10⁻⁴ M/min