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3: Reaction Kinetics

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    • 3.1: Chemical Reaction Rates
      The rate of a reaction can be expressed either in terms of the decrease in the amount of a reactant or the increase in the amount of a product per unit time. Relations between different rate expressions for a given reaction are derived directly from the stoichiometric coefficients of the equation representing the reaction.
    • 3.2: Factors Affecting Reaction Rates
      The rate of a chemical reaction is affected by several parameters. Reactions involving two phases proceed more rapidly when there is greater surface area contact. If temperature or reactant concentration is increased, the rate of a given reaction generally increases as well. A catalyst can increase the rate of a reaction by providing an alternative pathway that causes the activation energy of the reaction to decrease.
    • 3.3: Rate Laws
      Rate laws provide a mathematical description of how changes in the amount of a substance affect the rate of a chemical reaction. Rate laws are determined experimentally and cannot be predicted by reaction stoichiometry. The order of reaction describes how much a change in the amount of each substance affects the overall rate, and the overall order of a reaction is the sum of the orders for each substance present in the reaction.
    • 3.4: Integrated Rate Laws
      Differential rate laws can be determined by the method of initial rates or other methods. We measure values for the initial rates of a reaction at different concentrations of the reactants. From these measurements, we determine the order of the reaction in each reactant. Integrated rate laws are determined by integration of the corresponding differential rate laws. Rate constants for those rate laws are determined from measurements of concentration at various times during a reaction.
    • 3.5: First-Order Reactions
      In a first-order reaction, the reaction rate is directly proportional to the concentration of one of the reactants. First-order reactions often have the general form A → products.
    • 3.6: Second-Order Reactions
      The simplest kind of second-order reaction is one whose rate is proportional to the square of the concentration of one reactant. These generally have the form 2A → products. A second kind of second-order reaction has a reaction rate that is proportional to the product of the concentrations of two reactants. Such reactions generally have the form A + B → products.
    • 3.7: Zero-Order Reactions
      The rates of zero-order reactions is apparently independent of reactant concentrations. This means that the rate of the reaction is equal to the rate constant, k, of that reaction, but clearly a zero-order process cannot continue after a reactant has been exhausted. Just before this point is reached, the reaction will revert to another rate law instead of falling directly to zero  Situations which are apparently zero order occur when a reaction is catalyzed by attachment to a solid surface (hete
    • 3.8: Pseudo-1st-order reactions
      Under certain conditions, the 2nd order kinetics can be well approximated as first order kinetics. These Pseudo-1st-order reactions greatly simplify quantifying the reaction dynamics.
    • 3.9: Reaction Mechanisms
      The sequence of individual steps, or elementary reactions, by which reactants are converted into products during the course of a reaction is called the reaction mechanism. The overall rate of a reaction is determined by the rate of the slowest step, called the rate-determining step. Unimolecular elementary reactions have first-order rate laws, while bimolecular elementary reactions have second-order rate laws.
    • 3.10: The Collision Theory
      Collision theory explains why different reactions occur at different rates, and suggests ways to change the rate of a reaction. Collision theory states that for a chemical reaction to occur, the reacting particles must collide with one another. The rate of the reaction depends on the frequency of collisions. The theory also tells us that reacting particles often collide without reacting.
    • 3.11: Arrhenius Equation
    • 3.12: Kinetics of Catalysis

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