Kinetics

collision theory: molecules must collide in order to react. Collision converts kinetic energy of molecules (temperature) into enthalpy (heat energy), which is used to start the reaction by breaking chemical bonds in the reactants.

reactants: the compounds consumed in the chemical reaction

products: the compounds created by the chemical reaction

intermediates: compounds produced by a reaction and consumed by a later reaction in a multi-step process.

effective collision: leads to formation of products

ineffective collision: does not lead to products

mechanism: the details of which molecules collide with which and in what order and orientation in order for the reaction to take place.

enthalpy: the sum of the kinetic and potential energy of the molecules. (Kinetic energy is the energy due to motion and is measured by temperature. Potential energy is the energy of the chemical bonds.) Enthalpy can be converted to heat, and vice-versa.

reaction coordinate: a diagram that shows energyvs. progress of a chemical reaction.

enthalpy of reaction (ΔH_rxn): the difference between the enthalpy of the products and the enthalpy of the reactants. The enthalpy change is what produces or consumes heat in a chemical reaction.

activation energy (ΔE_act): minimum kinetic energy needed for a collision to produce the transition state (and therefore proceed to form products).

catalyst: a substance that speeds up a reaction by lowering the activation energy of the rate-limiting step.

transition state (or activated complex): the configuration of all entities at the instant an effective collision happens. In the transition state, all of the energy needed by the reaction has been converted to enthalpy.

reaction rate: the rate of formation of products in a chemical reaction, usually expressed in or (where M = molarity = )

rate-limiting step (or rate-determining step): the step that determines the overall rate of the reaction. In a multi-step reaction, the rate-limiting step is the slowest step.

For example, in the multi-step reaction:

A  B  C  D

the rate of B à C will determine the overall rate of the reaction A à D.

rate law: an equation describing the relationship between the concentrations of reactants and the reaction rate. Rate laws are always determined by experiment. The rate law cannot be predicted from the chemical equation.

A rate law has the form:

Rate = k [A]^x [B]^y [C]^z ...

where:

[A], [B], and [C] are the concentrations of A, B, and C, in (molarity)

k is a constant called the rate constant

x, y, and z are whole-number exponents determined experimentally

rate constant (k): the factor (constant) that relates the rate of a reaction to the concentrations of zero or more reactants. The units of k will be whatever is necessary for the rate to end up with units of .

Example:

Consider the following reaction:

A + 2B à C + 2D

The rate law for this reaction could be any of the following:

Rate = k [A]
Rate = k [B]
Rate = k [B]²
Rate = k [A][B]²
Rate = k

The only way to determine the rate law and rate constant for a reaction is to perform experiments using different concentrations of reactants, and measure the rate of formation of the products.

order: the order of a reaction is the number of separate molecules that react to form the transition state. The order of a reaction equals the sum of the exponents in the rate law. Some examples:

Rate Law	Order
Rate = k [A]      = k [A]^1	1^st
Rate = k [A][B]      = k [A]^1 [B]^1	2^nd
Rate = k [A]^2 [B]     = k [A]^2 [B]^1	3^rd
Rate = k      = k [A]^0 [B]^0	0

Factors that can
Affect Reaction Rates

· concentration of reactants: higher concentration means more frequent collisions = faster rate. (Only applies to molecules involved in the rate-determining step.) For gases, higher pressure = higher concentration.

· surface area of reactants: more surface area means higher probability of a collision = faster rate.

· temperature: higher temperature = faster because faster-moving molecules collide more often, and because faster-moving molecules have more kinetic energy to overcome the activation energy.

· nature of the reactants: weak bonds are easier to break than strong bonds. Reactions involving dissolved ions are very fast, because bonds are already broken.

· catalysts: catalysts speed up reactions in any of several ways:

o bring molecules into the correct orientation for an effective collision (equivalent to increasing the concentration and/or surface area)

o assist in breaking of bonds in the reactant(s) and/or formation of bonds in the products (equivalent to changing the nature of the reactants and/or lowering the activation energy)

Catalysts are not reactants; they are not consumed by the reaction.