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.