One reaction that destroys O3 molecules in the stratosphere is NO + O3 → NO2 + O2 When this reaction was studied in the laboratory, it was found to be first order with respect to both NO and O3, with a rate constant of 1.9 x 104 L mol-1 s-1. If [NO] =1.2 x 10-5 mole L-1 and [O3] = 2.0 x 10-5, what is the rate of this reaction? (NB;Please this is my assignment question and i desperately need your help.Thank you)

You would plug the concentrations into the rate equation: rate= k[NO][O3] with k being the rate constant.

To calculate the rate of the reaction, we need to use the given rate constant and the concentrations of the reactants.

The rate equation for a first-order reaction can be written as:

Rate = k[NO]^a[O3]^b

Where:
- Rate is the rate of the reaction (which we need to find)
- k is the rate constant
- [NO] and [O3] are the concentrations of NO and O3, respectively
- a and b are the orders of NO and O3, respectively

From the given information, we know that the reaction is first order with respect to both NO and O3 (a = 1, b = 1). The rate constant is given as 1.9 x 10^4 L mol^-1 s^-1. The concentrations of [NO] and [O3] are 1.2 x 10^-5 mol L^-1 and 2.0 x 10^-5 mol L^-1, respectively.

Now, plug the values into the rate equation:

Rate = k[NO]^a[O3]^b
Rate = (1.9 x 10^4 L mol^-1 s^-1)(1.2 x 10^-5 mol L^-1)^1(2.0 x 10^-5 mol L^-1)^1

Simplifying the expression:

Rate = (1.9 x 10^4)(1.2 x 10^-5)(2.0 x 10^-5) L mol^-1 s^-1

Multiply the numbers:

Rate = 4.56 x 10^-6 L mol^-1 s^-1

Therefore, the rate of this reaction is 4.56 x 10^-6 L mol^-1 s^-1.