Chapter 6 - Rates of Chemical Reactions

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