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Unimolecular reactions involve the transformation of a single molecule. In these reactions, one molecule undergoes a change to form one or more products. For instance, consider the reaction: - \(\mathrm{CH}_3 \mathrm{NC} \rightarrow \mathrm{CH}_3 \mathrm{CN}\\)This example shows a unimolecular reaction where one molecule of methyl isocyanide (\(\mathrm{CH}_3 \mathrm{NC}\)) rearranges to form methyl cyanide (\(\mathrm{CH}_3 \mathrm{CN}\)). Such reactions are generally simple because they require no collisions with other molecules to proceed.
Unimolecular reactions are often driven by energy changes within the molecule itself, such as breaking and forming bonds, or electronic restructuring. These reactions can occur as the molecule absorbs energy from its surroundings, possibly in the form of heat or light.
One important factor in these reactions is that the reaction rate is dependent only on the concentration of the single reactant molecule involved, making it a first-order reaction. This can be represented by the rate law: - \( \text{Rate} = k \cdot [\text{Reactant}]\\)where \(k\) is the rate constant and \( [\text{Reactant}] \) is the concentration of the reactant.

Bimolecular reactions involve two reactant molecules colliding to produce one or more products. These types of reactions are common in chemical processes and follow the basic idea that molecules must collide with one another to initiate a reaction. Consider the reaction:- \(\mathrm{SO} + \mathrm{O}_2 \rightarrow \mathrm{SO}_2 + \mathrm{O}\\)In this bimolecular reaction, one molecule of sulfur monoxide (\(\mathrm{SO}\)) collides with one molecule of oxygen (\(\mathrm{O}_2\)) to produce sulfur dioxide (\(\mathrm{SO}_2\)) and an oxygen atom (\(\mathrm{O}\)).
Collision theory suggests that for the reaction to occur, the molecules must collide with sufficient energy and in the correct orientation. This is because the reacting bonds must break and new bonds must form.
The reaction rate in bimolecular reactions depends on the concentration of both reactants, and is expressed as second-order reaction: - \( \text{Rate} = k \cdot [\text{Reactant 1}] \cdot [\text{Reactant 2}]\\)where \(k\) is the rate constant, and \( [\text{Reactant 1}] \) and \( [\text{Reactant 2}] \) are the concentrations of the two reactants.

Termolecular reactions involve three reactant molecules coming together in a single step to form products. These reactions are less common due to the statistical improbability of three molecules colliding simultaneously with the required energy and orientation.
A classic example is the reaction:- \(2 \mathrm{NO} + \mathrm{Br}_2 \rightarrow 2 \mathrm{NOBr}\\)This reaction involves the collision between two molecules of nitrogen monoxide (\(\mathrm{NO}\)) and one molecule of bromine (\(\mathrm{Br}_2\)) to form two molecules of nitrosyl bromide (\(\mathrm{NOBr}\)).
Due to the complexity and rarity of such simultaneous collisions, termolecular reactions are not often seen in comparison to unimolecular or bimolecular reactions.
Termolecular reactions can be very sensitive to temperature and pressure conditions, as these factors affect the likelihood of three molecules meeting at the right place and time. The reaction rate for termolecular reactions can be described by a third-order rate law:- \( \text{Rate} = k \cdot [\text{Reactant 1}] \cdot [\text{Reactant 2}] \cdot [\text{Reactant 3}]\\)where \(k\) is the rate constant and each \( [\text{Reactant}] \) represents the concentration of one of the three participating molecules.