Chemistry: Chapter 14

Chemical Reaction Basics

• Activation energy is the minimum energy required for a chemical reaction to occur.
• High activation barrier means only a few molecules have sufficient kinetic energy, resulting in a slower reaction.
• Temperature increase leads to an increase in reaction rate.

Transition State and Arrhenius Equation

• The transition state is a hypothetical species between reactants and products on a reaction profile.
• The Arrhenius equation shows that the rate constant increases with increasing temperature and decreasing activation energy.
• The equation is: k = Ae^-Ea/RT, and ln k = (-Ea/RT) + ln A.

Reaction Mechanism

• A reaction mechanism is a step-by-step description of a chemical reaction.
• Each step is an elementary process, and the mechanism must be consistent with stoichiometry and experimentally determined rate law.
• Elementary processes can be unimolecular or bimolecular, and are reversible.

Catalysis

• Catalysis is an alternative reaction pathway with lower energy.
• It increases the rate of reaction, reduces activation energy, and completes the reaction in less time.

Rate of Reaction

• Rate of reaction = ΔM / Δt, where ΔM is the change in concentration of products.
• Rate of reaction = rate of disappearance of reactants = rate of appearance of products.

Rate Law

• The rate law is Rate = k [A]^m [B]^n, where k is the rate constant.
• The overall order of reaction = m + n +…, and m and n are not related to a and b.

Types of Rate Laws

• Zero-Order Rate Law: Rate = k, where the sum of exponents is 0.
• First-Order Rate Law: Rate = k [A], where the sum of exponents is 1.
• Second-Order Rate Law: Rate = k [A]^2, where the sum of exponents is 2.

Half-Life

• The half-life (t½) is the time taken for one-half of a reactant to be consumed.
• For a first-order reaction, the half-life is constant and independent of initial concentration.