Rates of Chemical Reaction

The rate of a chemical reaction is defined as the change in the concentration of reactants or products per unit time. It indicates how fast or slow a reaction proceeds. Understanding reaction rates is crucial in industrial processes where time is a factor in profitability.

1. Factors Affecting the Rate of Reaction

Several factors influence how quickly a chemical reaction occurs:

• Temperature: Increasing temperature increases the kinetic energy of the particles. This leads to more frequent and more energetic collisions. For many reactions, a 10°C rise in temperature approximately doubles the reaction rate. Example: The reaction between Sodium Thiosulphate (Na₂S₂O₃) and Hydrochloric Acid (HCl) proceeds faster at higher temperatures.
• Concentration and Pressure: Increasing the concentration of reactants increases the number of particles in a given volume, leading to a higher frequency of effective collisions. For gaseous systems, increasing the pressure is equivalent to increasing concentration. Example: The iodine clock reaction is sensitive to changes in concentration.
• Surface Area: For reactions involving solids, increasing the surface area (e.g., using powder instead of lumps) increases the number of exposed particles available for collision. Example: Powdered marble chips (CaCO₃) react faster with HCl than marble lumps of the same mass.
• Catalysts: A catalyst is a substance that increases the rate of a reaction without being consumed. It works by providing an alternative pathway with a lower Activation Energy (Ea). Example: Manganese (IV) Oxide (MnO₂) speeds up the decomposition of Hydrogen Peroxide (H₂O₂) or Potassium Chlorate (KClO₃).
• Light: Some reactions are photochemical and are triggered by light energy. Example: The chlorination of methane requires UV light.

2. Reaction Rate Curves

A reaction rate curve typically plots the concentration of a product or reactant against time. The gradient (slope) of the curve at any point represents the rate of reaction at that instant. A steeper slope indicates a faster reaction.

3. Collision Theory and Activation Energy

According to the Collision Theory, for a reaction to occur, reactant particles must collide with:

• The correct orientation.
• Minimum energy known as the Activation Energy (Ea).

The Activated Complex is the unstable intermediate state formed during a successful collision. Activation energy is the 'energy barrier' that must be overcome for reactants to transform into products.