Chapter 6.1: Expressing and Measuring Reaction Rates

• Reaction rate: change in the number of reactants or products over time

  • Rate of reaction = ∆Quantity / ∆t

    • Moles per second

  • In a gaseous or solution units are in concentration A (moles/Liters or mol/L) over seconds

    • Rate of reaction = ∆[A] (in mol/L) / ∆t (in seconds)

  • Reaction rates are always positive

• Average rate: average change in the concentration of a reactant or product per unit time over a given time interval

  • Only give an idea of how the reaction is progressing

  • Finding slope by secant line of the graph

• Instantaneous rate: rate of the reaction at a particular time at a tangent line

• You can measure reaction rates through:

  • Mass, pH, and Conductivity

  • Pressure

  • Color: through a spectrophotometer

  • Volume

Chapter 6.2: The Rate Law: Reactant Concentration and Rate

• Rate law equation: Rate ∝ k[A]m[B]n

  • Shows relationship between the concentration of the reactants and the rate of their reaction

  • **Rate constant: the letter k (represents a proportionality constant) Depends on temperature and is constantSpeed of reaction In s−1

  • [A] and [B] are rates of reactants

  • Rate law exponents m and n do not change with temperature and have to be found experimentally

    • Exponents of 1: first order

    • Exponents of 2: second-order

• Overall reaction order: sum of exponents (m and n)

• First-order reaction: overall reaction order is equal to 1

• Second-order reaction: overall reaction order is equal to 2

• Other times reaction rates can lead it to equal to k[A]0, so the rate of reaction is constant

• Initial rates method: comparing the initial rate of each reaction

• Half-life, t1/2: reaction is the time that is needed for the reactant mass or concentration to decrease by one half of its initial value

  • In seconds

  • t1/2 = 0.693/k

Chapter 6.3: Theories of Reaction Rates

• Collision theory: for a reaction to occur, reacting particles (atoms, molecules, or ions) must collide with one another.Increasing surface area more collisions occurs Not all collisions occur in reactions, for a collision to be effective:Must be correct orientation and sufficient collision energy

• Activation energy, Ea: minimum collision energy that is required for a successful reaction

• Maxwell-Boltzmann distribution: fraction of collisions vs energy graph at a constant temperature

  • The area under the graph represents the distribution of kinetic collision

• Transition state theory: explain what happens when molecules collide in a reaction. It examines the transition, or change, from reactants to products

  • The kinetic energy of the reactants is transferred to potential energy as the reactants collide, due to the law of conservation of energy

• Potential energy diagram: a diagram that charts the potential energy of a reaction against the progress of the reaction

  • axis represents potential energy

  • *x-*axis, labeled “Reaction progress”, represent the progression of reaction through time

• There is no way to predict the activation energy of a reaction from its enthalpy change

• Transition state: top of the activation energy barrier

• Activated complex: chemical species that exist at the transition state neither product nor reactant

• Reactions rate increases at higher temperatures

Chapter 6.4: Reaction Mechanism and Catalysts

• Reaction mechanism: a series of steps that make up an overall reaction

• Elementary reaction: involves a single molecular event, such as a simple collision between atoms, molecules, or ions

• Reaction intermediates: molecules/atoms/ions that are formed in an elementary reaction and consumed in a subsequent elementary reaction

• Molecularity: refers to the number of reactant particles (molecules, atoms, or ions) that are involved in an elementary reaction

• Bimolecular: when two particles collide and react

• Unimolecular: when one molecule or ion reacts

• For an elementary reaction, the exponents in the rate law equation are the same as the stoichiometric coefficients for each reactant in the chemical equation

• Rate-determining step: slowest elementary reaction slower which becomes the overall rate-determining step

• A catalyst is a substance that increases the rate of a chemical reaction without being consumed by the reaction

  • Lowers activation energy so reactants have sufficient energy to react

  • Homogeneous catalyst: exists in the same phase as the reactants

  • Heterogeneous catalyst: exists in a phase that is different from the phase of the reaction it catalyzes

• Enzymes: biological catalysts enormous protein active site: a small part of enzyme part of the catalyst reaction Substrate: reactant molecule which binds to the active site to models:Lock and key model: the enzyme is like a lock Induced fit model: changes shape to fit the substrate