Fundamentals of Chemical Kinetics
Chemical kinetics studies how fast reactions proceed and what factors influence that speed. Understanding the rate of a reaction provides insight into the underlying molecular events and helps control processes in industry, biology, and the laboratory.
What a Reaction Rate Measures
The rate of a chemical reaction quantifies how fast a reaction occurs. It is expressed as the change in concentration of reactants or products per unit time. Typical units are mol L⁻¹ s⁻¹, reflecting concentration (moles per litre) divided by seconds.
Expressing Rate in Terms of Concentration
Mathematically, the rate can be written as the change in concentration per unit time. For a reactant A that disappears, the rate is –d[A]/dt; for a product B that forms, it is +d[B]/dt. Both expressions share the same units and describe the same overall speed of the process.
Rate Law and Reaction Order
The rate law is the equation that links the reaction rate to the molar concentrations of the reactants. It has the general form:
rate = k [A]^m [B]^n …
Here, k is the rate constant, and the exponents m and n are the order of the reaction with respect to each reactant. These orders are determined experimentally, not from the stoichiometric coefficients in the balanced equation.
- Order → exponent in the rate law
- Coefficient → number from the balanced equation
Because the order is experimental, it can differ from the stoichiometric coefficient; the statement that they are always equal is false.
Rate Constant (k)
The rate constant k is a proportionality factor that connects concentrations to the observed rate. When all reactant concentrations are 1 mol L⁻¹, the rate equals the value of k. Importantly, k does not depend on the initial concentrations of the reactants; it is a property of the reaction at a given temperature.
Temperature Effects on the Rate Constant
Temperature has a profound impact on k. Raising the temperature generally increases k because more molecules possess the required activation energy to form the activated complex. A common approximation is that the reaction rate roughly doubles for every 10 °C increase in temperature, though this rule does not hold for every reaction.
Activation Energy and Collision Theory
According to collision theory, a reaction occurs only when reactant particles collide. For a collision to be effective, two conditions must be satisfied:
- Sufficient energy – the kinetic energy must exceed the activation energy, the minimum energy needed to form the activated complex.
- Proper orientation – molecules must align correctly to break old bonds and form new ones.
Catalysts modify the activation energy. A positive catalyst lowers the activation energy, providing an alternative pathway and speeding up the reaction. Conversely, a negative catalyst raises the activation energy, slowing the reaction.
Physical State, Surface Area, and Pressure
Reaction rates vary with the physical state of the reactants. Generally, gases react faster than liquids, which react faster than solids because of molecular mobility:
Gases > Liquids > Solids
For solid reactants, increasing the surface area raises the rate by offering more sites for collisions. In gaseous reactions, raising the pressure increases the concentration of molecules, which similarly boosts the rate. In solution, increasing the concentration of reactants has the same effect.
Key Takeaways
- The reaction rate measures the speed of concentration change, typically in mol L⁻¹ s⁻¹.
- The rate law relates rate to reactant concentrations; exponents are the reaction orders.
- Rate constants are temperature‑dependent but independent of initial concentrations.
- Higher temperature, larger surface area, and greater pressure or concentration all increase the rate.
- Catalysts alter the activation energy, thereby influencing the rate.
