Group Four
Reaction Mechanisms
The flame of an oxyacetylene torch is used in many industries and businesses, including
garages and on the farm, to cut through metals by quickly melting them. The temperature
of 3300 C necessary to melt most metals is produced by the combustion of acetylene to
produce carbon dioxide and water:
2 C2H2 (g) + 5 O2 (g)  4 CO2 (g) + 2 H2O (g)
The balanced equation might seem to imply that two molecules of acetylene collide with
five molecules of oxygen to yield the products. Chemists know that for two acetylene
molecules and five oxygen molecules to reach the same point at the same time is an
extremely unlikely event (Figure 10). Collisions of even three particles is rare. If a seven
particle collision were necessary for this reaction to take place, in the gas phase, the
reaction would never be seen. Since, in fact, oxygen and acetylene do react rapidly, the
reaction must proceed through a series of simple steps, where each step probably only
involves a collision of two particles. Chemists believe that most chemical reactions take
place by means of series of steps. This series of steps is known as the reaction
mechanism. The mechanism can be considered to be the pathway that a reaction takes.
Figure 10
Let’s consider another example. Earlier you read that nitrogen monoxide was a pollutant
in automobile exhaust. The mechanism whereby nitrogen monoxide reacts with oxygen
to form nitrogen dioxide (another pollutant) is believed to consist of the following single
collisions or elementary steps:
Step 1:
Step 2
2 NO (g)  N2O2 (g)
N2O2 (g) + O2 (g)  2 NO2 (g)____________________
2 NO (g) + N2O2 (g) + O2 (g)  N2O2 (g) + 2 NO2 (g)
Net Eq:
2 NO (g) + O2 (g)  2 NO2 (g)
Notice that if we add the two elementary steps we get an equation that has N2O2 on both
sides. It cancels out and does not appear as part of the overall or net equation. It is
produced in the first elementary step and used up in the second. Species such as N2O2 are
known as reaction intermediates (Figures 11). They can be atoms, molecules, or ions.
They are very short lived and are usually difficult to isolate.
Step 1
Reaction
Intermediate
Reactants
Products
Step 2
Figure 11
Note: It is incorrect to interpret the net equation above as meaning that the
product NO2 is produced from a three molecule collision between two
NO molecules and one molecule of O2. When the equation for a reaction
is written, the net equation is usually given, not the steps that add up to
the net equation. In fact, the steps that add up to the net equation are not
known for most reactions. Unless otherwise indicated, the equations
we will study are net equations.
The reaction mechanism for the formation of nitrogen monoxide can also be viewed in
terms of a potential energy diagram (Figure 12). Notice it shows the relative potential
energies of the reactants and products, indicating that overall the reaction is exothermic
(H is negative). During the path from reactants to products, however, a reaction
intermediate is formed. The formation requires some energy in order to cause bond
breaking and bond formation, and thus Step 1 is slightly endothermic. The reaction
intermediate then reacts with oxygen gas in Step 2. Again this reaction requires energy as
is indicated by the ‘hill’ in Step 2. However, the production of nitrogen dioxide from the
reaction intermediate releases a lot of energy and so the overall the amount of energy
used up in the reaction to break bonds is more than compensated by the energy released
in the end.
Ea Overall
Ea Step 2
Reaction
Intermediate
Ea Step 1
H is negative
(exothermic)
Figure 12
Reaction mechanisms do not necessarily involve only two elementary steps. Consider the
reaction between hydrogen and iodine:
H2 (g) + I2 (g)  2 HI (g)
Originally it was thought to be a bimolecular reaction involving the collision of one
hydrogen molecule and one iodine molecule, as the above equation predicts. Recent
research has shown that it a three-step mechanism involving two intermediates (I and
H2I):
uv light
Step 1:
Step 2:
Step 3:
I2  I + I
I + H2  H2I
H2I + I  2 HI_________________________
I2 + I + H2 + H2I + I  I + I + H2I + 2 HI
Net Eq:
H2 + I2  2 HI
Ultraviolet light is known to increase the dissociation of iodine molecules (Step 1) and, in
turn, to increase the concentration of the reaction intermediate, the I atom. This produces
the second reaction intermediate, H2I, which reacts with the I atoms present to produce
the HI product (Figure 13).
Figure 13
Demonstration:
Materials: Lego or tinker toys
Analogy : Lego or tinker toy race.
- Have students form teams and build a model step by step
- Provide students with the step by step procedure of forming a model of your
choice
- Each student on the team may perform one step.
Sample Questions
1. Define the following: reaction mechanism, reaction intermediate, elementary step.
2. Why is it more likely for a reaction to take place in a series of steps rather that in one
step?
3. The following mechanism has been proposed for the reaction between iodobutane and
chloride ion:
Step 1:
Step 2:
C4H9I  C4H9+1 + I-1
C4H9+1 + Cl-1  C4H9Cl
(a) Give the net equation for the reaction.
(b) Identify the reaction intermediates.
4. The following mechanism has been proposed for the reaction between nitrogen
dioxide and methane (CH4):
Step 1:
NO2 + NO2  NO3 + NO
Step 2:
Step 3:
NO3 + CH4  HNO3 + CH3
CH3 + NO2  CH3NO2
(a) Give the net equation for the reaction.
(b) Identify the reaction intermediates.