Reactant molecules combine to create new products, which are known as a chemical reaction. However, have we ever thought the true causes of molecular reactions? What causes some reactions to occur more quickly than others? A correct and rational explanation for these questions can be found in the Collision Theory of Chemical Reactions.
What is Collision Theory?
According to collision theory, reactant particles must collide with one another in order for a chemical reaction to take place. But not every collision results in a reaction. Products are only formed in collisions that satisfy specific requirements.
Conditions for Effective Collisions
The following prerequisites need to be fulfilled for a reaction to be successful:
Adequate Energy (Activation Energy): The interacting particles need to possess sufficient energy to create new connections and break old ones. Activation energy (Ea) is the bare minimum of energy required to start a reaction.
Proper Orientation: For new bonds to form in the planned way, the reactant molecules must be properly positioned during the impact. Even if there is sufficient energy, the reaction will not take place if the orientation is off.
Collision Frequency: A reaction is more likely to occur when reactant molecules collide more frequently. Only a small percentage of these encounters, or “effective collisions,” are successful.
How Collision Theory Explains Reaction Rates
The number of effective collisions per second determines the rate of a chemical reaction. This is influenced by following factors:
Temperature: As the temperature rises, reactant molecules gain kinetic energy, which causes them to move more quickly and collide more forcefully. This makes it more likely that activation energy will be overcome and accelerating the reaction.
Reactant Concentration: A higher concentration results in more particles available in a given volume, which raises the frequency of collisions and the rate of response.
Pressure (for gases): As pressure rises, gas molecules are compressed and drawn closer to one another, increasing the likelihood of collisions and quickening the process.
Surface Area: By reducing solid reactants to smaller fragments, their surface area is increased, exposing more particles to impacts and speeding up the reaction.
Catalysts: A catalyst increases the number of effective collisions without being consumed in the process by offering a different pathway with a lower activation energy.
Mathematical Expression of Collision Theory
Collision theory is quantitatively expressed by the Arrhenius equation: k = Ae-Ea /RT
where,
k = rate constant,
A = frequency factor (related to collision frequency and orientation),
Ea = activation energy,
R = universal gas constant (8.314 J/mol·K),
T = temperature in Kelvin,
e = exponential factor.
This formula demonstrates that the reaction rate grows exponentially with temperature.
Real-Life Examples of Collision Theory
Faster Food Cooking at Higher Temperatures: Food cooks more quickly at higher temperatures because molecules have more energy at these temperatures, which speeds up reactions.
Combustion Reactions: When oxygen is added, fuels burn more quickly because there are more effective collisions.
Iron Rusting: Because oxygen and water molecules collide with the metal surface more frequently in humid environments, iron rusts more quickly.
Gas Explosions: When gas molecules are compressed to a great degree, they collide more frequently, which causes quick reactions, or explosions.
Summary
We can understand the better way the components that affect chemical reactions and they happen at varying rates in collision theory. It offers a starting point for explaining reaction mechanisms and creating experiments to regulate reaction rates. We can increase the efficiency of reactions in a variety of everyday and industrial applications by controlling factors e.g; temperature, concentration, and catalysts.
The collision of reactant molecules is the basis for collision theory, which describes how chemical reactions take place. It asserts that molecules must collide with sufficient energy and in the right direction for a reaction to occur.
Not all collisions result in a reaction because:
The colliding molecules may not have sufficient energy (Activation Energy).
The molecules may not be aligned in the proper orientation to form new bonds.