Question

Write the balanced equation for the reaction of nitrogen monoxide with oxygen to produce nitrogen dioxide:

If the following experimental data were collected for this reaction:

Run #	(NO)(M)	[O2] (M)	Initial Rate of NO (M/s)
1	0.0126	0.0125	1.40 × 10^-2 (= 0.0140)
2	0.0252	0.0250	1.14 × 10^-1 (= 0.114)
3	0.0252	0.0125	5.62 × 10^-2 (= 0.0562)
4	0.0504	0.00625	1.13 × 10^-1 (= 0.113)

(a) What is the order with respect to each reactant ?

(b) What is the rate constant ?

(c) What is the complete rate law for this reaction ?

(d) Most reactions exhibit different rates for individual reactants and products based on stoichiometry. Which of those multiple rates does the term “rate of the reaction” refer to?

(e) What is the rate of disappearance of nitrogen monoxide when [NO] = 0.100 M and [O2] = 0.200 M ?

(f) What is the rate of disappearance of oxygen at the concentrations given in Part (e)?

Answer 1

Let rate law

r = k [NO]^m [O2]ⁿ

From trial 1 and 3

r1/r3 = ( [NO]1^m [O2]1ⁿ ) / ( [NO]3^m [O2]3ⁿ )

0.0140/0.0562 = (0.0126/0.0252)^m

0.25 = 0.5^m

0.5² = 0.5^m

m = 2

Similarly from trial 2 and 3

r2/r3 = ( [NO]2^m [O2]2ⁿ ) /  ( [NO]3^m [O2]3ⁿ )

0.114/0.0562 = (0.0250/0.0125)ⁿ

2 = 2ⁿ

n = 1

rate law

r = k [NO]² [O2] .............. Eq1

Order with respect to NO = 2

Order with respect to O2 = 1

From trial 1

r1 = k [NO]1² [O2]1

0.0140 M/s = k (0.0126)² (0.0125) M3

k = 7.05 x 10^3 M-2 s-1

Every reactant or product represent its own rate with respect to other reactants or products.

Rate of reaction is represented by the eq1

Rate of disappearance of NO

r_NO = - 7.05 x 10^3 x (0.1*0.1) x (0.2)

= - 14.11 M/s

Negative sign shows the disappearance of NO at the same concentration

Rate of disappearance of oxygen = - 14.11 M/s