A chemical mechanism is a theoretical conjecture that tries to describe in detail what takes place at each stage of an overall chemical reaction. The detailed steps of a reaction are not observable in most cases. The conjectured mechanism is chosen because it is thermodynamically feasible, and has experimental support in isolated intermediates (see next section) or other quantitative and qualitative characteristics of the reaction. It also describes each reactive intermediate, activated complex, and transition state, and which bonds are broken (and in what order), and which bonds are formed (and in what order). A complete mechanism must also explain the reason for the reactants and catalyst used, the stereochemistry observed in reactants and products, all products formed and the amount of each.
A reaction mechanism must also account for the order in which molecules react. Often what appears to be a single-step conversion is in fact a multistep reaction.
Reaction intermediates are chemical species, often unstable and short-lived (however sometimes can be isolated), which are not reactants or products of the overall chemical reaction, but are temporary products and/or reactants in the mechanism's reaction steps. Reaction intermediates are often free radicals or ions.
The kinetics (relative rates of the reaction steps and the rate equation for the overall reaction) are explained in terms of the energy needed for the conversion of the reactants to the proposed transition states (molecular states that corresponds to maxima on the reaction coordinates, and to saddle points on the potential energy surface for the reaction).
Consider the following reaction for example:
In this case, experiments have determined that this reaction takes place according to the rate law . This form suggests that the rate-determining step is a reaction between two molecules of NO2. A possible mechanism for the overall reaction that explains the rate law is:
Each step is called an elementary step, and each has its own rate law and molecularity. The elementary steps should add up to the original reaction. (Meaning, if we were to cancel out all the molecules that appear on both sides of the reaction, we would be left with the original reaction.)
When determining the overall rate law for a reaction, the slowest step is the step that determines the reaction rate. Because the first step (in the above reaction) is the slowest step, it is the rate-determining step. Because it involves the collision of two NO2 molecules, it is a bimolecular reaction with a rate which obeys the rate law .
Other reactions may have mechanisms of several consecutive steps. In organic chemistry, the reaction mechanism for the benzoin condensation, put forward in 1903 by A. J. Lapworth, was one of the first proposed reaction mechanisms.
A chain reaction is an example of a complex mechanism, in which the propagation steps form a closed cycle. In a chain reaction, the intermediate produced in one step generates an intermediate in another step. Intermediates are called chain carriers. Sometimes, the chain carriers are radicals, they can be ions as well. In nuclear fission they are neutrons.
Chain reactions have several steps, which may include:
Even though all these steps can appear in one chain reaction, the minimum necessary ones are: Initiation, propagation and termination.
An example of a simple chain reaction is the thermal decomposition of acetaldehyde (CH3CHO) to methane (CH4) and carbon monoxide (CO). The experimental reaction order is 3/2, which can be explained by a Rice-Herzfeld mechanism.
This reaction mechanism for acetaldehyde has 4 steps with rate equations for each step :
For the overall reaction, the rates of change of the concentration of the intermediates oCH3 and CH3COo are zero, according to the steady-state approximation, which is used to account for the rate laws of chain reactions.
d[oCH3]/dt = k1[CH3CHO] - k2[oCH3][CH3CHO] + k3[CH3COo] - 2k4[oCH3]2 = 0
and d[CH3COo]/dt = k2[oCH3][CH3CHO] - k3[CH3COo] = 0
The sum of these two equations is k1[CH3CHO] - 2 k4[oCH3]2 = 0. This may be solved to find the steady-state concentration of oCH3 radicals as [oCH3] = (k1 / 2k4)1/2 [CH3CHO]1/2.
It follows that the rate of formation of CH4 is d[CH4]/dt = k2[oCH3][CH3CHO] = k2 (k1 / 2k4)1/2 [CH3CHO]3/2
Thus the mechanism explains the observed rate expression, for the principal products CH4 and CO. The exact rate law may be even more complicated, there are also minor products such as acetone (CH3COCH3) and propanal (CH3CH2CHO).
Many experiments that suggest the possible sequence of steps in a reaction mechanism have been designed, including:
A correct reaction mechanism is an important part of accurate predictive modeling. For many combustion and plasma systems, detailed mechanisms are not available or require development.
Even when information is available, identifying and assembling the relevant data from a variety of sources, reconciling discrepant values and extrapolating to different conditions can be a difficult process without expert help. Rate constants or thermochemical data are often not available in the literature, so computational chemistry techniques or group additivity methods must be used to obtain the required parameters.
In general, reaction steps involving more than three molecular entities do not occur, because is statistically improbable in terms of Maxwell distribution to find such transition state.
L.G.WADE,ORGANIC CHEMISTRY 7TH ED,2010