Question

Nitrogen monoxide, $NO$, reacts with hydrogen, $H _{2}$, according to the following equation:
$2NO(g) + 2H_2(g) \rightarrow N_2(g) + 2H_2O(g)$
If the mechanism for this reaction were,
$2 NO ( g )+ H _{2}( g ) \rightarrow N _{2}( g )+ H _{2} O _{2}( g ) ; \text { slow } $
$H _{2} O _{2}( g )+ H _{2}( g ) \rightarrow 2 H _{2} O ( g ) ; \text { fast }$
Which of the following rate laws would we expect to obtain experimentally?

• A. Rate $=k\left[ H _{2} O _{2}\right]\left[ H _{2}\right]$
• B. Rate $=k\left[ NO ^{2}\left[ H _{2}\right]\right.$
• C. Rate $=k\left[ NO ^{2}\left[ H _{2}\right]^{2}\right.$
• D. Rate $=k[ NO ]\left[ H _{2}\right]$

Answer 1

Answer: (B)

Slow reaction is the rate determining step.