## Catalysts in Zero Order Reactions

$\frac{d[R]}{dt}=-k; [R]=-kt + [R]_{0}; t_{\frac{1}{2}}=\frac{[R]_{0}}{2k}$

Emily Warda 2L
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### Catalysts in Zero Order Reactions

I know that the rate of a zero-order reaction is independent of [reactant]. To strengthen this concept, Dr. Lavelle explained the use of a catalyst. Does anyone recall exactly how this related to zero-order reactions? Must a catalyst be present to make the first statement true?

Naveed Zaman 1C
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### Re: Catalysts in Zero Order Reactions

Basically, catalysts can turn some reactions into zero order reactions because of their very nature to facilitate collisions. A reaction that is second order, such as the decomposition of NO2 to N2 and O2, can be turned into a zero order reaction in the presence of a hot platinum wire (which acts as a catalyst for the reaction). Normally, the second order form of this reaction requires high temperatures and is second order because it requires two NO2 molecules to collide effectively, so this rate is dependent on the number of pairs of molecules that successfully collide; this means increasing the concentration of NO2 should increase the number of pairs of molecules that collide (hence the rate law).

However, a hot platinum wire turns the reaction into zero order because at some point, even if you keep increasing the concentration of NO2, all the spots available for the reaction to occur on the catalyst will be taken up. So if you have 10 available spots on the catalyst to run the reaction (simple theoretical example), the reaction won't run faster or slower if you change reactant concentration but all 10 spots of the catalyst are already filled. The reaction proceeds at the same rate whether you have enough molecules to fit 10 spots or 100000 spots on the catalyst; since there are only 10 spots on the catalyst, only 10 reactions can occur at a time.

This means the rate is independent of the concentration, or Rate = k[NO2]0 = k

Other ways we can get (pseudo)zero order is if we maybe have two reactants, but we make the concentration of 1 reactant so high that it keeps getting replenished so much faster than it gets used up. Another way is if only a small fraction of the reactant molecules are in a location or state in which they are able to react, and this fraction is continually replenished from the larger pool.

Sorry if this doesn't make sense. Here is a link that says it more concisely than I do:
https://chem.libretexts.org/Core/Physic ... _Reactions

Hope this helps!