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6: Kinetics

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    • 6.0: Introduction
      Reaction Rates to Medical Breakthroughs: The Kinetics Legacy of Maud Menten
    • 6.1: Reaction Rates
      The rate of a reaction can be expressed as the decrease in reactant concentration or the increase in product concentration per unit time. The relationships between different rate expressions are determined by the stoichiometric coefficients of the balanced equation.  The average rate is calculated as the change in concentration over a given time interval. The instantaneous rate refers to the rate at a specific moment in time and is determined by the slope of a tangent line on a concentration vs.
    • 6.2: Factors Affecting Reaction Rates
      Several factors affect the rate of a chemical reaction. Surface area is particularly important for reactions involving solids—greater surface area leads to faster reactions by increasing contact between reactants. Factors like temperature, concentration, and catalysts also influence reaction rates and will be discussed in later sections.
    • 6.3: Rate Laws
      Rate laws describe how reaction rates depend on reactant concentrations and must be determined experimentally rather than predicted from stoichiometry. The reaction order indicates how reactant concentration changes affect the reaction rate. While most reactions are zero, first, or second order, fractional and negative orders are possible. The method of initial rates can be used to determine reaction orders by varying reactant concentrations and observing the effect on the initial rate.
    • 6.4: Integrated Rate Laws
      Rate laws describe how the reaction rate depends on reactant concentration and can be determined using methods such as the method of initial rates or graphical analysis of concentration vs. time data.  Integrated rate laws describe how reactant concentrations change over time and are derived by integrating the rate laws.  The half-life of a reaction is the time required to decrease the amount of a given reactant by one-half.
    • 6.5: Collision Theory, Activation Energy, and the Arrhenius Equation
      Collision theory states that molecules must collide with the correct orientation and enough energy to react. The activation energy is the minimum energy needed to reach the transition state, a very unstable species between reactants and products.  Reaction energy diagrams illustrate how energy changes during a reaction.  The Arrhenius equation quantitatively describes how temperature and activation energy affect the rate constant and can be used to find the activation energy.
    • 6.6: Reaction Mechanisms
      A reaction mechanism is the sequence of elementary steps by which reactants are converted to products. The overall rate of a reaction is determined by the rate of the slowest step, called the rate-determining step.   Each elementary step has its own transition state, shown as a maximum on a reaction energy diagram.  Intermediates are species formed in one step and consumed in a later step.
    • 6.7: Catalysis
      Catalysts speed up reactions by providing a lower-energy pathway, but they are not consumed in the reaction.  In heterogeneous catalysis, catalysts provide a solid surface where reactants bind (adsorption).  In homogeneous catalysis, catalysts and reactants exist in the same phase.  Enzymes are biological catalysts that dramatically increase reaction rates and are highly specific to their substrates.  In enzyme-catalyzed reactions, the substrate binds to the enzyme’s active site, forming a tempo
    • 6.E: Kinetics (Exercises)
      Working through the end-of-chapter problems is an important part of learning kinetics. These questions help you assess your understanding of fundamental concepts, practise applying them to new situations, and strengthen your problem-solving skills. Regular practice will deepen your grasp of reaction rates and mechanisms and prepare you to tackle more complex chemical problems.

    Thumbnail: Molecular collisions frequency. (Public Domain; Sadi Carnot via Wikipedia)

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