Factors affecting reaction rate Knowledge-Based Study

Study notes

1. Chemical kinetics investigates the speeds of chemical processes and the pathways through which reactants convert into products. According to collision theory, for a chemical reaction to occur, reactant particles must collide with one another.

   However, not every collision results in a chemical reaction. A collision that leads to the formation of products is called an effective collision. For a collision to be effective, two fundamental criteria must be met simultaneously:

   • Particles must collide with a proper spatial orientation so that the relevant bonds can break and new bonds can form.
   • Particles must possess a minimum amount of kinetic energy, known as the activation energy ($E_a$), to overcome the electrostatic repulsion of their electron clouds and reach the transition state.

   Consequently, the overall rate of a chemical reaction is directly proportional to the frequency of effective collisions occurring per unit of time.

   The chemical composition, structure, and physical state of the reacting species strongly influence the rate of a reaction.

   • Bond breaking and bond forming: Reactions involving the rearrangement of strong covalent bonds or the cleavage of many chemical bonds generally proceed slowly at room temperature because they require high activation energies. In contrast, reactions between simple, oppositely charged aqueous ions (precipitation or neutralization reactions) require minimal bond reorganization and occur almost instantaneously.
   • Physical state: Homogeneous mixtures of gases or aqueous solutions generally react faster than heterogeneous mixtures involving solids, primarily because particles in fluids mix uniformly and collide more frequently.

   The concentration of reactants in solution and the partial pressure of gaseous reactants dictate how crowded the particles are within a given volume.

   Concentration in Solutions

   Increasing the concentration of a dissolved reactant increases the number of solute particles per unit volume. As particle density increases, the probability of collisions between reacting particles rises. Because the frequency of total collisions increases while the fraction of collisions with sufficient energy remains constant, the frequency of effective collisions increases, thereby raising the reaction rate.

   Pressure in Gaseous Systems

   For reactions involving gases, increasing the pressure (for instance, by decreasing the volume of the reaction vessel) decreases the mean free path between gas molecules. This concentration increase leads to more frequent collisions and a correspondingly higher reaction rate. Adding an inert gas at constant volume does not alter reactant concentrations and therefore does not change the reaction rate.

   In heterogeneous reactions—where reactants exist in different phases, such as a solid reacting with a liquid or gas—collisions can occur only at the interface between the phases.

   Subdividing a solid into smaller particles (e.g., grinding a marble chip into a fine powder) exposes a significantly larger total surface area to the surrounding fluid. With more reactant atoms or molecules exposed simultaneously:

   • The total number of collision sites increases.
   • The frequency of collisions per second increases proportionally.
   • The reaction rate increases without changing the activation energy or temperature of the system.

   Temperature is a measure of the average kinetic energy of the particles in a substance. Increasing the temperature of a reaction mixture has two main effects on collision dynamics:

   1. Collision Frequency: Molecules move faster at higher temperatures, causing them to collide slightly more often.
   2. Energy of Collisions: Far more significantly, a rise in temperature broadens the Maxwell–Boltzmann distribution curve toward higher kinetic energies. As a result, a much greater fraction of particles possesses energy equal to or exceeding the activation energy ($E \ge E_a$).

   Because the proportion of high-energy collisions increases exponentially with temperature, even a modest temperature increase (e.g., 10 °C) typically causes a substantial increase in the reaction rate for both endothermic and exothermic reactions.

   A catalyst is a substance that increases the rate of a chemical reaction without undergoing a permanent net chemical change itself.

   Catalysts operate by providing an alternative reaction pathway or mechanism that features a lower activation energy ($E_a$). Key characteristics of catalysts include:

   • By lowering $E_a$, a substantially greater fraction of collisions possesses sufficient energy to react at a given temperature.
   • A catalyst does not alter the enthalpy change ($\Delta H$) or the initial and final states of the reaction; it only changes the energy profile of the transition state.
   • In reversible processes, a catalyst lowers the activation energy of both the forward and reverse reactions by the exact same amount, meaning it accelerates both directions equally and helps the system reach equilibrium faster without shifting the equilibrium position.