- Lindemann mechanism
In
chemical kinetics , the Lindemann mechanism, sometimes called the Lindemann-Hinshelwood mechanism, is a schematicreaction mechanism .Frederick Lindemann discovered the concept in 1921 andCyril Hinshelwood developed it. [ [http://www.chem.iitm.ac.in/pradeep/webpage/teaching/02L15.pdf] Access date 8 December 2007]It breaks down a
stepwise reaction into two or moreelementary step s, then it gives arate constant for each elementary step. Therate law andrate equation for the entire reaction can be derived from this information.Lindemann mechanisms have been used to model
gas phase decomposition reactions. Although the net formula for a decomposition may appear to be first-order (unimolecular ) in the reactant, a Lindemann mechanism may show that the reaction is actually second-order (bimolecular ). [ [http://links.jstor.org/sici?sici=0036-1399(199112)51%3A6%3C1489%3AGPDBTL%3E2.0.CO%3B2-X] "Gas phase decomposition by the Lindemann mechanism" by S. L. Cole and J. W. Wilder. "SIAM Journal on Applied Mathematics ", Vol. 51, No. 6 (Dec., 1991), pp. 1489-1497.]Activated reaction intermediates
A Lindemann mechanism typically includes an activated
reaction intermediate , labeled A* (where A can be any element or compound). The activated intermediate is produced from thereactant s only after a sufficientactivation energy is applied. It then decomposes or reacts with another reactant in order to form the products or a second intermediate.The steady-state approximation
In some cases, one of the elementary steps is much slower than the other steps. This slow step is called the
rate-determining step because it is the only step that affects the rate. In layman's terms, a rate-determining step could be compared to traveling through atraffic jam : the time it takes to complete a journey is most severely affected by the time spent waiting in the traffic jam, which is the slow step of the journey.In the
steady-state approximation , it is assumed that each of the elementary steps influences the rate, so there is no "fast" or "slow" step. Therefore, all of the steps must be accounted for in calculating the rate equation. It is also assumed that the concentration of intermediate A* remains constant over time because the concentration of A* builds up very quickly but decays very slowly over the course of a reaction, and the concentration of A* never becomes large. This assumption simplifies the calculation of the rate equation.Examples
General schematic example
The schematic reaction A + M → P is assumed to consist of two elementary steps:
# A + M → A* + M (forward reaction rate = k1; reverse reaction rate = k-1)
# A* + M → P (forward reaction rate = k2)Assuming that the
concentration of intermediate A* is held constant according to the steady-state approximation, what is the rate of formation of product P?First, find the rates of production and consumption of intermediate A*. The rate of production of A* in the first elementary step is simply:
:d [A*] /dt = k1 [A] [M] (forward first step)
A* is consumed both in the reverse first step and in the forward second step. The respective rates of consumption of A* are:
:-d [A*] /dt = k-1 [A*] [M] (reverse first step):-d [A*] /dt = k2 [A*] (forward second step)
According to the steady-state approximation, the rate of production of A* equals the rate of consumption. Therefore:
:k1 [A] [M] = k-1 [A*] [M] + k2 [A*]
Solving for [A*] , it is found that
: [A*] = (k1 [A] [M] ) / (k-1 [M] + k2)
The overall reaction rate is
:d [P] /dt = k2 [A*]
Now, by substituting the calculate value for [A*] , the overall reaction rate can be expressed in terms of the original reactants A and M as follows:
:d [P] /dt = (k1k2 [A] [M] ) / (k-1 [M] + k2) [The question is taken from GRE Chemistry Test Practice Book, based on the 2000 exam, question 98.]
Specific practical example
The decomposition of
dinitrogen pentoxide tonitrogen dioxide andnitrogen trioxide :N2O5 → NO2 + NO3
is postulated to take place via two elementary steps, which are similar in form to the schematic example given above:
# N2O5 + N2O5 → N2O5* + N2O5
# N2O5* → NO2 + NO3Using the steady-state approximation, the rate equation is calculated to be
:Rate = k2 [N2O5] * = k1k2 [N2O5] 2 / (k-1 [N2O5] + k2)
Experiment has shown that the rate is observed as first-order in the original concentration of N2O5 sometimes, and second order at other times.
If k2 >> k-1 (>> means "much larger than"), then the rate equation may be simplified by assuming that k-1 ~= 0. Then the rate equation is
:Rate = k1 [N2O5] 2
which is second order. In contrast, if k2 << k-1 (<< means "much less than"), then the rate equation may be simplified by assuming k2 ~= 0. Then the rate equation is
:Rate = k1k2 [N2O5] / k-1
which is first order. [ [http://www.chem.arizona.edu/~salzmanr/480a/480ants/lindeman/lindeman.html "Lindemann Mechanism"] by W. R. Salzman at the
University of Arizona , 2004. Access date 8 December 2007.]References
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