Be sure to answer all parts. The rate law for 2 NO(g) + O2(g) → 2 NO2(g) is rate = k[NO]2[O2]. The following mechanisms have been proposed: I. 2 NO(g) + O2(g) → 2 NO2(g) II.2 NO(g) ⇌ N2O2(g)[fast] N2O2(g) + O2(g) → 2 NO2(g) [slow] III.2 NO(g) ⇌ N2(g) + O2(g)[fast] N2(g) + 2 O2(g) → 2 NO2(g) [slow] (a) Which of these mechanisms is consistent with the rate law? Check all that apply. I. II. III. None of the above (b) Which of these mechanisms is most reasonable? I. II. III.

a. I

b. I

Explanation:

A rate law is an equation that relates the rate of a reaction to the concentration of reactants (and catalyst) raised to various powers.

In the equation above

2 NO(g) + O2(g) → 2 NO2(g)

rate = k[NO]²[O2]

where k = rate constant

From the mechanism above we see that in;

1: Rate law= k[NO]² [O2]

2: Rate law= k[N2O2][O2] [slow eq determines rate law]

3: Rate law= k[N2][O2]²

We can observe that the resembling equation is 1.

The rate of a chemical reaction is determined by the slowest step.

Rate = k[Concentration of reactants individually raised to their stoichiometric co-efficients]

In mechanism I,

Overall reaction occurs in a single step. Therefore,

rate= k[NO]²[O2]

This is this consistent with the observed rate law.

In mechanism II,

The overall reaction occurs in two steps, through the involvement of an intermediate, N2O2.

Rate of the slowest step should be the overall reaction rate.

Therefore, overall rate= k[N2O2][O2]

Again considering non-accumulation of intermediate, N2O2 in the overall reaction.

Its rate of production will be equal to its rate of decomposition.

Thus, k1[NO]²= k[N2O2][O2]

➡ [N2O2]= (k1/k). [NO]²/[O2]

Overall rate= k(k1/k).([NO]²[O2])/[O2]

=k1[NO]²

So, this is not consistent with the rate law.

Mechanism III,

the overall rate =k[NO]².

Therefore, we see that only mechanism I is is most appropriate , reasonable and consistent with the observed rate law.

if you mean hydrogen and oxygen then water is the product