Reaction Rates

AP chapter 14.3

Reaction Rates

• Describe how quickly concentration of reactants or products are changing

• Units typically M/t for aqueous reactants and products

• Could be units of P/t for gaseous products

• Effective concentration solids and liquids does not change over the course of a reaction, so it will be more difficult to model these changes

Reaction equations

• Rate=k[A]m[B]n

• [A] is a symbol roughly meaning “concentration of” (in units of molarity or partial pressure)

• More precisely, it means “Activity of” or “effective concentration”. Remember, that even in an aqueous solution, not all of the ions are available to react.

• This is given only as one example• The equation for each type of reaction must be

tested experimentally. The values of m and n are determined experimentally

Reaction orders

• Rate=k[A]m[B]n

• The reaction above is “m” order for reactant A, “n” order for reactant B, and “m+n” order for the reaction overall. “Reaction order” describes the influence which increasing concentration has on the reaction rate

You Try

• Describe the reaction orders for the following rate equation:

• Rate=k[A]2[B]1

• Sketch the shape of the graph showing the relationship between – [A] and rate– [B] and rate– If [A] and [B] are measured in molar units,

what is the unit for the rate constant?

What’s the relationship between concentration and time?• Note that as time passes, the

concentrations change, so the rate changes.

• Based on this analysis will rate increase or decrease as time passes?

• The general solutions for these relationships require some calculus, but even without calculus, you can learn the equations for some simple cases.

Integrated rate laws

• These show the relationship between [A] and time over the course of a single experiment

• First order: Rate=k[A]

• First order: ln [A] = -kt + ln[A]0

• Other integrated rate laws

• Sketch the graph of ln[A] against t

Collision model

• Reactions occur when molecules collide

• Even for unimolecular reactions, collisions are necessary to change the kinetic energy of colliding particles

• Example: 3 O2 + h 2 O3

• Example: Cl2 + F2 2 ClF

Distribution of kinetic energies• Temperature is proportional to average

kinetic energy for a collection of molecules

• However, the kinetic energy of any individual molecule could be a little less or a little more.

• Distribution of molecular energies is a predictable function

• http://intro.chem.okstate.edu/1314F00/Laboratory/GLP.htm

Reactions at a molecular level

• Molecules only react if they possess enough combined energy to overcome the activation energy.

• The orientations of the molecules also matters.

• http://www.mp-docker.demon.co.uk/chains_and_rings/mechanisms/index.html

Arrhenius equation

• Relates Temperature, reaction rate and activation energy.

• Various forms of the Arrhenius equation

Reaction pathway diagrams

• Show the relative potential energy of reactants, products, intermediates and transition states.

• “intermediates” are the various molecular forms which appear as reactant becomes product.

• Depending on context, “intermediate” could mean only the stable intermediates, or could also include short lived (transient) “transition states”

• Energy diagram

Catalysts

• Catalysts are substances which speed up a reaction, without being altered or consumed in the process

• Catalysts may be temporarily altered as part of one of the reaction intermediates.

• How Catalysts work

How catalysts work at a chemical level• Catalysts lower activation energy, either

by bringing reactants closer together, or otherwise stabilizing the reaction transition state.

• A catalyst speeds up both forward and reverse reactions, so the mixture comes to equilibrium more rapidly. A catalyst can not change the equilibrium concentrations

• What do you think would be the result of adding a catalyst to a mixture already at equilibrium?

Enzymes

• “Enzyme” is a term for a catalyst found in, or obtained from a biological system.

• Enzymes are primarily made of protein, but may also include metal ions, nucleic acids, or other structural materials.

• How enzymes work

Non-enzyme catalysts

• Both enzyme and non-enzyme catalysts are important in manufacturing processes, such as the synthesis of ammonia

• Catalysts can be homogeneous (in the same phase as the reaction) or heterogeneous (a finely divided metal in solution, for example.

Heterogeneous catalysts

• Why do you think heterogeneous catalysts must be finally divided?

• What other methods do chemical engineers to increase the surface area of a catalyst?