Chemical Kinetics

Schematic Profile (Potential Energy curve for the reaction)

  • Chemical reactions involve the breaking and forming of bonds between atoms.
  • The reaction progress can be represented by a potential energy diagram.
  • It shows the changes in potential energy as the reaction proceeds.
  • The vertical axis represents potential energy and the horizontal axis represents reaction progress.
  • Let’s analyze the different regions of the potential energy diagram.

Chemical Kinetics

Reactants

  • The reactants start with a certain amount of potential energy.
  • They have sufficient energy to undergo the reaction, but the specific arrangement of atoms makes it unstable.
  • The reactants are represented on the left side of the potential energy diagram.
  • They are usually higher in potential energy compared to the products.
  • The potential energy of the reactants is denoted as Ea (activation energy).

Chemical Kinetics

Activation Energy

  • Activation energy (Ea) is the minimum amount of energy required to initiate a chemical reaction.
  • It is the energy barrier that needs to be overcome for the reaction to occur.
  • The higher the activation energy, the slower the reaction rate.
  • Activation energy depends on factors such as temperature, concentration, and presence of catalysts.
  • Catalysts lower the activation energy, therefore increasing the rate of the reaction.

Chemical Kinetics

Transition State

  • As the reactants gain energy, they reach a transition state or activated complex.
  • This is an intermediate state where old bonds are breaking and new bonds are forming.
  • The transition state is represented as the highest point on the potential energy diagram.
  • It is a short-lived species that exists at the peak of the energy barrier.
  • The energy of the transition state is denoted as ΔG‡ (free energy of activation).

Chemical Kinetics

Products

  • Once the transition state is crossed, the reaction proceeds towards the formation of products.
  • The products have a lower potential energy compared to the reactants.
  • The potential energy of the products is denoted as ΔG (free energy change).
  • The difference in potential energy between reactants and products determines the overall energy change of the reaction.
  • Exothermic reactions release energy and have negative ΔG, while endothermic reactions absorb energy and have positive ΔG.

Chemical Kinetics

Reaction Rates

  • The rate of a chemical reaction is the speed at which reactants are converted into products.
  • It is measured in terms of the change in concentration of a reactant or product per unit time.
  • Rate = Δ[A] / Δt, where [A] is the concentration of a reactant or product and Δt is the time interval.
  • Reaction rates depend on factors such as temperature, concentration, surface area, and presence of catalysts.
  • Increasing temperature, concentration, and surface area, or adding a catalyst, generally increases the reaction rate.

Chemical Kinetics

Rate Law

  • The rate of a reaction can be expressed using a rate law equation.
  • The rate law relates the rate of the reaction to the concentrations of the reactants.
  • It is determined experimentally and can be represented as Rate = k[A]^m[B]^n, where k is the rate constant, [A] and [B] are the concentrations of reactants, and m and n are the orders of reaction with respect to A and B, respectively.
  • The overall order of the reaction is the sum of m and n.
  • The rate constant (k) depends on temperature and provides information about the reaction’s speed.

Chemical Kinetics

Rate Determining Step

  • In complex reactions, the rate of the reaction is determined by a slowest step called the rate-determining step.
  • The rate-determining step limits the overall rate of the reaction.
  • It involves the highest activation energy and determines the order of the reaction.
  • By identifying the rate-determining step, we can focus on that step to control or enhance the reaction rate.
  • Catalysts often work by providing an alternative reaction pathway with a lower activation energy.

Chemical Kinetics

Collision Theory

  • The collision theory explains how chemical reactions occur at the molecular level.

  • According to this theory, for a reaction to take place, molecules must collide with sufficient energy (activation energy) and proper orientation.

  • The collision frequency and energy of collisions determine the reaction rate.

  • Increasing the collision frequency or the energy of collisions increases the reaction rate.

  • Temperature, concentration, pressure, and surface area affect collision frequency and energy. Chemical Kinetics - Schematic Profile (Potential Energy curve for the reaction)

  • Chemical reactions involve the breaking and forming of bonds between atoms.

  • The reaction progress can be represented by a potential energy diagram.

  • It shows the changes in potential energy as the reaction proceeds.

  • The vertical axis represents potential energy and the horizontal axis represents reaction progress.

  • Let’s analyze the different regions of the potential energy diagram.

Chemical Kinetics - Reactants

  • The reactants start with a certain amount of potential energy.

  • They have sufficient energy to undergo the reaction, but the specific arrangement of atoms makes it unstable.

  • The reactants are represented on the left side of the potential energy diagram.

  • They are usually higher in potential energy compared to the products.

  • The potential energy of the reactants is denoted as Ea (activation energy).

Chemical Kinetics - Activation Energy

  • Activation energy (Ea) is the minimum amount of energy required to initiate a chemical reaction.

  • It is the energy barrier that needs to be overcome for the reaction to occur.

  • The higher the activation energy, the slower the reaction rate.

  • Activation energy depends on factors such as temperature, concentration, and presence of catalysts.

  • Catalysts lower the activation energy, therefore increasing the rate of the reaction.

Chemical Kinetics - Transition State

  • As the reactants gain energy, they reach a transition state or activated complex.

  • This is an intermediate state where old bonds are breaking and new bonds are forming.

  • The transition state is represented as the highest point on the potential energy diagram.

  • It is a short-lived species that exists at the peak of the energy barrier.

  • The energy of the transition state is denoted as ΔG‡ (free energy of activation).

Chemical Kinetics - Products

  • Once the transition state is crossed, the reaction proceeds towards the formation of products.

  • The products have a lower potential energy compared to the reactants.

  • The potential energy of the products is denoted as ΔG (free energy change).

  • The difference in potential energy between reactants and products determines the overall energy change of the reaction.

  • Exothermic reactions release energy and have negative ΔG, while endothermic reactions absorb energy and have positive ΔG.

Chemical Kinetics - Reaction Rates

  • The rate of a chemical reaction is the speed at which reactants are converted into products.

  • It is measured in terms of the change in concentration of a reactant or product per unit time.

  • Rate = Δ[A] / Δt, where [A] is the concentration of a reactant or product and Δt is the time interval.

  • Reaction rates depend on factors such as temperature, concentration, surface area, and presence of catalysts.

  • Increasing temperature, concentration, and surface area, or adding a catalyst, generally increases the reaction rate.

Chemical Kinetics - Rate Law

  • The rate of a reaction can be expressed using a rate law equation.

  • The rate law relates the rate of the reaction to the concentrations of the reactants.

  • It is determined experimentally and can be represented as Rate = k[A]^m[B]^n,

    where k is the rate constant, [A] and [B] are the concentrations of reactants, and m and n are the orders of reaction with respect to A and B, respectively.

  • The overall order of the reaction is the sum of m and n.

  • The rate constant (k) depends on temperature and provides information about the reaction’s speed.

Chemical Kinetics - Rate Determining Step

  • In complex reactions, the rate of the reaction is determined by a slowest step called the rate-determining step.

  • The rate-determining step limits the overall rate of the reaction.

  • It involves the highest activation energy and determines the order of the reaction.

  • By identifying the rate-determining step, we can focus on that step to control or enhance the reaction rate.

  • Catalysts often work by providing an alternative reaction pathway with a lower activation energy.

Chemical Kinetics - Collision Theory

  • The collision theory explains how chemical reactions occur at the molecular level.

  • According to this theory, for a reaction to take place, molecules must collide with sufficient energy (activation energy) and proper orientation.

  • The collision frequency and energy of collisions determine the reaction rate.

  • Increasing the collision frequency or the energy of collisions increases the reaction rate.

  • Temperature, concentration, pressure, and surface area affect collision frequency and energy.

Chemical Kinetics - Reaction Mechanisms

  • Complex reactions often occur through a series of elementary steps.
  • The reaction mechanism is the sequence of these elementary steps that lead to the overall reaction.
  • Each elementary step involves the collision of molecules and the formation or breaking of chemical bonds.
  • The molecularity of a step refers to the number of molecules involved in that step.
  • Unimolecular steps involve one molecule, bimolecular steps involve two molecules, and termolecular steps involve three molecules.
  • The overall reaction rate depends on the slowest step in the mechanism, known as the rate-determining step. Chemical Kinetics - Elementary Reactions
  • Elementary reactions are the individual steps in a reaction mechanism.
  • They often involve the formation or breaking of chemical bonds.
  • Elementary reactions are characterized by their molecularity and reaction rate expressions.
  • Examples of elementary reactions:
    • A + B → C (unimolecular)

    • 2A → C (unimolecular)

    • A + B → C + D (bimolecular)

    • 2A + B → C (termolecular)

  • The forward and reverse rates of an elementary reaction are often proportional to reactant concentrations. Chemical Kinetics - Rate-Determining Step
  • The rate-determining step is the slowest step in a reaction mechanism.
  • It limits the overall rate of the reaction.
  • The rate law of the rate-determining step provides the overall rate law for the reaction.
  • The rate-determining step usually involves a high-energy transition state.
  • Identifying the rate-determining step is crucial in understanding and predicting the reaction kinetics.
  • Changing the conditions of the rate-determining step can significantly alter the reaction rate. Chemical Kinetics - Catalysts
  • Catalysts are substances that increase the rate of a chemical reaction without being consumed.
  • They provide an alternative reaction pathway with a lower activation energy.
  • Catalysts participate in the reaction but are regenerated at the end.
  • Homogeneous catalysts are in the same phase as the reactants, while heterogeneous catalysts are in a different phase.
  • Catalysts can be pure substances or mixtures of substances.
  • Examples of catalysts: enzymes, transition metals, acid-base catalysts. Chemical Kinetics - Temperature and Reaction Rate
  • Temperature has a significant effect on the reaction rate.
  • Increasing temperature generally increases the rate of a reaction.
  • Higher temperatures provide more kinetic energy, leading to more frequent and energetic collisions.
  • The relationship between temperature and rate is described by the Arrhenius equation:
    • k = A * exp(-Ea / RT) where k is the rate constant, A is the pre-exponential factor, Ea is the activation energy, R is the gas constant, and T is the temperature in Kelvin.
  • Activation energy (Ea) represents the energy barrier that must be overcome for a reaction to occur.
  • Higher activation energy leads to a slower reaction rate. Chemical Kinetics - Concentration and Reaction Rate
  • The concentration of reactants affects the rate of a chemical reaction.
  • Increasing the concentration of reactants generally increases the reaction rate.
  • Higher concentrations result in more frequent collisions, leading to a higher reaction rate.
  • The relationship between concentration and rate can be described using the rate law equation.
  • For example, the rate law for a reaction A + B → C is given as: - Rate = k[A]^m[B]^n where [A] and [B] represent the concentrations of A and B, respectively, and m and n are the reaction orders with respect to A and B.
  • The overall reaction order is the sum of the individual reaction orders. Chemical Kinetics - Surface Area and Reaction Rate
  • The surface area of a solid reactant can significantly affect the reaction rate.
  • Increasing the surface area increases the reaction rate.
  • More surface area provides more contact area for reactant molecules, leading to more frequent collisions.
  • Examples: finely powdered solids have a higher reaction rate compared to large solid pieces because they have a larger surface area.
  • Surface area also plays a role in heterogeneous catalysis, where the reactants are in a different phase than the catalyst.
  • Increasing the surface area of the catalyst can enhance the catalytic activity. Chemical Kinetics - Pressure and Reaction Rate
  • Pressure can affect the reaction rate, especially for gaseous reactions.
  • Increasing the pressure generally increases the reaction rate.
  • Higher pressure leads to a higher concentration of gas molecules in a given space, resulting in more frequent collisions.
  • This is known as the collision theory of chemical reactions.
  • Pressure does not have a significant effect on reactions involving only solids or liquids.
  • For gaseous reactions, pressure is often expressed in terms of partial pressure. Chemical Kinetics - Rate Constant
  • The rate constant (k) is a proportionality constant in the rate law equation.
  • It relates the reaction rate to the concentrations of reactants.
  • The rate constant is temperature-dependent and provides information about the reaction’s speed.
  • The units of the rate constant depend on the overall reaction order.
  • The rate constant can be determined experimentally using various techniques.
  • Activation energy (Ea) and temperature (T) influence the value of the rate constant. Chemical Kinetics - Arrhenius Equation
  • The Arrhenius equation relates the rate constant (k) to temperature (T) and activation energy (Ea).
  • It is commonly used to describe the temperature dependence of reaction rates.
  • The Arrhenius equation is given as: - k = A * exp(-Ea / RT) where k is the rate constant, A is the pre-exponential factor, Ea is the activation energy, R is the gas constant, and T is the temperature in Kelvin.
  • The Arrhenius equation shows that an increase in temperature leads to a higher rate constant and vice versa.
  • Activation energy affects the exponential term, determining the sensitivity of the reaction rate to temperature changes.