IB Chemistry HL Topic 6

Rate of reaction: The rate at which products are formed, which is equal to the rate at which the reactants are consumed. It has units of moldm-3s-1.

• Rate expression (law) is the dependence of rate on concentration expressed mathematically.
• For the reaction: A + B + C → products, rate =k[A]m[B]n[C]p, where m, n, and p are the orders of the reaction with respect to each reactant A, B, C. The overall order = (m + n + p).

Rates can be determined by measuring the change in the concentration of a reactant or product with time. Possible methods are pressure measurements for gases using a manometer or pressure gauge, volume measurement for gases using a gas syringe, color changes using a colorimeter or spectrophotometer, heat changes using a thermometer, mass changes using a balance, pH changes using a pH meter, titrations, etc.
A graph may be plotted of concentration against time, with time on the x-axis and some measure of how far the reaction has gone (ie concentration, volume, mass loss etc) on the y-axis. This will produce a curve and the rate at any given point is the gradient of the tangent to this curve.

Students should be familiar with graphs of changes in concentration, volume and mass against time.

Collision theory -- reactions take place as a result of particles (atoms or molecules) colliding and then undergoing a reaction. Not all collisions cause reaction, however, even in a system where the reaction is spontaneous. The particles must have sufficient kinetic energy, and the correct orientation with respect to each other for them to react.KE is prop. to temp. in kelvin.

Activation energy: The minimum amount of energy required to initiate a chemical reaction.

Students should know that reaction rate depends on:

• collision frequency
• number of particles with E ≥ Ea
• appropriate collision geometry or orientation.
  Collision Theory
  • The reactant particles must collide together.
  • Particles must have the correct geometrical alignment.
  • Particles must have a minimum energy of E ≥ Ea , the activation energy.
  • Activation energy is the minimum energy required for reactants to react in order to convert into products (via a transition state or activated complex).
  • Transition state is an unstable arrangement in which the bonds are in the process of being broken and formed and presents the maximum point on a potential energy diagram; it can not be isolated.

• Increase in concentration: Greater number of particles per unit volume per time means greater frequency of collisions and faster rate.
• Increase in temperature: This causes an increase in average kinetic energy; particles collide more frequently and more forcefully. Thus number of particles with E ≥ Ea increases greatly and rate increases (more forceful collisions is a more important factor than more frequent collisions).
• Increasing the surface area of solids reactants: This increases the number of particles that can collide, i.e., increases the frequency of successful collisions.

Students should be able to explain why the area under the curve is constant and does not change with temperature.
Maxwell – boltzman energy distribution curve: A curve of probability of energy as a function of the energy of each collision. It is shown that at lower temperatures, the graph peaks at a lower energy and at a higher probability of energy than at higher temperatures, when the graph peaks at a lower probability of molecules but colliding with a higher energy, but the area under both curves are the same. From this graph it can be seen that at a higher temperature a greater number of molecules will have any energy at a higher temperature, such as the activation energy.

• Catalysts: Lower the activation energy of reaction by providing an alternate pathway and thus increase the rate of reaction. Catalysts are specific to reactions and are used in industry.

o Homogeneous catalysis: Where the catalyst is in the same phase as the reactants, e.g., esterification reaction to produce sweet smelling compounds for perfumes, and in the food industry. o Heterogeneous catalysis: Where the catalyst and reactants are in different phases e.g., in the manufacture of ammonia, sulfuric acid, hydrogenation of alkenes etc. o A catalyst does not change the position of equilibrium; rather it increases the rate of both the forward and reverse reactions to the same extent; does not change ΔH of a reaction.

Lower with catalyst.