The speed at which a chemical reaction proceeds.
Collision theory: This concept explains how the rate of a chemical reaction depends on the frequency, energy, and orientation of collisions between reactant particles.
Activation energy: This is the minimum amount of energy required to initiate a chemical reaction by breaking the bonds of the reactant molecules.
Catalysts: These are substances that increase the rate of a chemical reaction by lowering the activation energy required for the reaction to proceed.
Concentration: As the concentration of the reactants increases, the frequency of collisions between them also increases, leading to an increase in reaction rate.
Temperature: An increase in temperature leads to an increase in the kinetic energy of the reactants, which increases the frequency and energy of collisions, leading to an increase in reaction rate.
Surface area: A larger surface area exposes more reactant particles to collisions, increasing the frequency of collisions and thus the reaction rate.
Reaction mechanisms: These are a series of steps that explain how a chemical reaction proceeds from reactants to products and can affect the rate of the reaction.
Order of reaction: The order of a reaction describes how the rate of the reaction depends on the concentration of its reactants.
Rate law: The rate law is an equation that expresses the relationship between the reaction rate and the concentrations of reactants involved in the reaction.
Reaction rate constants: These are constants that specify the rate of a reaction for a given set of reactants and conditions, and can be used to predict reaction rates.
Reaction intermediates: These are species that are formed during a reaction but are not the final products, and can affect the overall reaction rate.
Enzyme kinetics: This concept involves the study of how enzymes catalyze specific chemical reactions and how their activity is affected by different factors such as temperature, pH, and substrate concentration.
Reaction equilibrium: This describes the state of a chemical reaction where the rate of the forward reaction is equal to the rate of the reverse reaction, leading to a constant concentration of reactants and products.
Kinetics of complex reactions: Some chemical reactions involve multiple steps or intermediate compounds, and the study of their kinetics involves analyzing each step separately to understand the overall reaction rate.
Chemical kinetics in everyday life: This concept focuses on the application of chemical kinetics to everyday life and real-world scenarios, such as the study of atmospheric pollution, pharmaceutical drug development, and industrial catalysis.
Instantaneous rate: Rate of reaction at a specific moment in time.
Average rate: Rate of reaction over a given period of time.
Initial rate: Rate of reaction at the start of a reaction.
Final rate: Rate of reaction at the end of a reaction.
Instantaneous reaction rate: The rate of a reaction at any point in time.
Stoichiometric reaction rate: The rate at which a given reactant is consumed or a product is formed in relation to its stoichiometric coefficient.
Rate law: A mathematical expression that relates the rate of a reaction to the concentration of reactants and/or products.
Activation energy: The amount of energy required to initiate a reaction.
Transition state: An intermediate state in a reaction with a higher energy than reactants or products.
Rate-determining step: The slowest step in a reaction that determines the overall rate.
Chain reaction: A reaction that involves a chain of reactive intermediates and propagates until consumed by termination steps.
Auto-catalysis: A reaction in which one of the reaction products acts as a catalyst to make the reaction proceed faster.
Homogeneous reaction: A reaction occurring in a single phase, e.g. all reactants and products are in the same phase (liquid or gas).
Heterogeneous reaction: A reaction occurring in multiple phases, e.g. a solid catalyst in a gaseous reaction.
Reversible reaction: A reaction that can proceed in both forward and reverse directions, depending on the concentrations of reactants and products.
Irreversible reaction: A reaction that proceeds only in one direction and is completed without any reactants present.
Photochemical reaction: A reaction initiated by light energy.
Catalytic reaction: A reaction facilitated by a catalyst, which decreases the activation energy of the reaction and increases the reaction rate.