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The chemical molecule ${\mathbf{\text{SO}}}_{\mathbf{\text{3}}}$ stands for sulphur trioxide. Sulfur oxide has been called "unquestionably the most important economically." As a precursor of sulfuric acid, it is manufactured on a large scale. Gaseous monomer, crystalline trimer, and solid polymer are all examples of sulphur trioxide.

At normal temperature, the oxidation of${\text{SO}}_2{\text{to SO}}_{\text{3}}$is a slow and exothermic reaction.

$\begin{array}{rcl}{\mathrm{SO}}_2\left(\mathrm{g}\right)+\frac{1}{2}{\mathrm{O}}_2\left(\mathrm{g}\right)& ↔& {\mathrm{SO}}_3\left(\mathrm{g}\right)\\ {\mathrm{Δ\pm á}}^{\mathrm{o}}& =& -99\mathrm{kJ}\end{array}$

We can improve the yield by altering the temperature, raising the pressure (since there are more moles of gas on the left than on the right), and manipulating the concentrations (adding more ${\text{O}}_2$, eliminating ${\text{SO}}_{\text{3}}$).

Increasing the temperature would increase the frequency of ${\text{SO}}_2-{\text{O}}_2$collisions, and thus the rate of creation of ${\text{SO}}_{\text{3}}$. However, because the reaction enthalpy is exothermic, lowering the temperature of the mixture would cause the equilibrium to shift to the right, increasing the yield of ${\text{SO}}_{\text{3}}$.

As a result, a catalyst is required in this case - a material that aids in the establishment of equilibrium by lowering the activation energy. The equilibrium is attained faster and at a lower temperature when ${\text{V}}_{\text{2}}{\text{O}}_{\text{5}}$or inert silica is used as a catalyst.