Combustion[edit]
Ozone can be used for combustion reactions and combustible gases; ozone provides higher
temperatures than burning in dioxygen (O2). The following is a reaction for the combustion
of carbon subnitride which can also cause higher temperatures:
3 C
4N
2 + 4 O
3 → 12 CO + 3 N
Ozone can react at cryogenic temperatures. At 77 K (−196.2 °C; −321.1 °F),
atomic hydrogen reacts with liquid ozone to form a hydrogen superoxide radical,
which dimerizes:[33]
H + O
3 → HO2 + O
2 HO2 → H
2O
Ozone decomposition[edit]
Types of ozone decomposition[edit]
Ozone is a toxic substance commonly found or generated in human environments
(aircraft cabins, offices with photocopiers, laser printers, sterilizers…) and its
catalytic decomposition is very important to reduce pollution. This type of
decomposition is the most widely used, especially with solid catalysts, and it has
many advantages such as a higher conversion with a lower temperature.
Furthermore, the product and the catalyst can be instantaneously separated, and
this way the catalyst can be easily recovered without using any separation operation.
Moreover, the most used materials in the catalytic decomposition of ozone in the gas
phase are noble metals like Pt, Rh or Pd and transition metals such as Mn, Co, Cu,
Fe, Ni or Ag.
There are two other possibilities for the ozone decomposition in gas phase:
The first one is a thermal decomposition where the ozone can be decomposed using
only the action of heat. The problem is that this type of decomposition is very slow
with temperatures below 250 °C. However, the decomposition rate can be increased
working with higher temperatures but this would involve a high energy cost.
The second one is a photochemical decomposition, which consists of radiating
ozone with ultraviolet radiation (UV) and it gives rise to oxygen and radical peroxide.
[34]
Kinetics of ozone decomposition into molecular oxygen[edit]
The process of ozone decomposition is a complex reaction involving two elementary
reactions that finally lead to molecular oxygen, and this means that the reaction
order and the rate law cannot be determined by the stoichiometry of the fitted
equation.
Overall reaction: 2 O3 → 3 O2
Rate law (observed): V = K · [O3]2 · [O2]−1
It has been determined that the ozone decomposition follows a first order kinetics,
and from the rate law above it can be determined that the partial order respect to
molecular oxygen is -1 and respect to ozone is 2, therefore the global reaction order
is 1.
The ozone decomposition consists of two elementary steps: The first one
corresponds to a unimolecular reaction because one only molecule of ozone
decomposes into two products (molecular oxygen and oxygen). Then, the oxygen
from the first step is an intermediate because it participates as a reactant in the
second step, which is a bimolecular reaction because there are two different
reactants (ozone and oxygen) that give rise to one product, that corresponds to
molecular oxygen in the gas phase.
Step 1: Unimolecular reaction O3 → O2 + O
Step 2: Bimolecular reaction O3 + O → 2 O2
These two steps have different reaction rates, the first one is reversible and faster
than the second reaction, which is slower, so this means that the determining step is
the second reaction and this is used to determine the observed reaction rate. The
reaction rate laws for every step are the ones that follow:
V1 = K1 · [O3] V2 = K2 · [O] · [O3]
The following mechanism allows to explain the rate law of the ozone decomposition
observed experimentally, and also it allows to determine the reaction orders with
respect to ozone and oxygen, with which the overall reaction order will be
determined. The slower step, the bimolecular reaction, is the one that determines
the rate of product formation, and considering that this step gives rise to two oxygen
molecules the rate law has this form:
V = 2 K2 · [O] · [O3]
However, this equation depends on the concentration of oxygen (intermediate),
which can be determined considering the first step. Since the first step is faster and
reversible and the second step is slower, the reactants and products from the first
step are in equilibrium, so the concentration of the intermediate can be determined
as follows:
Then using these equations, the formation rate of molecular oxygen is as shown
below:
