Unit 11- Chemical Kinetics

The change in concentration of a reactant or product per unit time Reaction Rate The change in concentration of a reactant or product per unit time

2 NO2  2 NO + O2

COLLISION THEORY

One of the simplest models to explain reaction rates is collision theory According to collision theory, the rate of reaction is proportional to the number of effective collisions per second among the reacting molecules An effective collision is one that actually gives product molecules

There are two factors that determine whether a collision will result in products being formed: Molecular orientation is important because a collision on the “wrong side” of a reacting species cannot produce any product. Particles must collide with enough energy. The minimum energy required is called the activation energy, Ea.

This increases the rate of reaction for two reasons. The minimum kinetic energy the colliding particles must have is called the activation energy, Ea Activation energies can be large, so only a small fraction of the well-orientated, colliding molecules collide with enough energy. At higher temperatures, the average kinetic energy of the reacting particles is higher. This increases the rate of reaction for two reasons. Collisions happen more often. Collisions happen with more energy.

Kinetic energy distribution for a reaction at two different temperatures. At the higher temperature, a larger fraction of the collisions have sufficient energy for reaction to occur.

Transition state theory explains what happens when reactant particles come together Potential-energy diagrams are used to help visualize the relationship between the activation energy and the development of total potential energy The potential energy is plotted against reaction coordinate or reaction progress

The potential-energy diagram for an exothermic reaction The potential-energy diagram for an exothermic reaction. The extent of reaction is represented as the reaction coordinate.

Potential-energy diagram for an endothermic reaction Potential-energy diagram for an endothermic reaction. The heat of reaction and activation energy are labeled.

Reactions generally have different activation energies in the forward and reverse direction

The brief moment during a successful collision that the reactant bonds are partially broken and the product bonds are partially formed is called the transition state The potential energy of the transition state is a maximum of the potential-energy diagram The unstable chemical species that “exists” momentarily is called the activated complex

Formation of the activated complex in the reaction between NO2Cl and Cl. NO2Cl + Cl  NO2 + Cl2

Experimental Rate Laws

Generally, chemical reactions happen faster when the concentrations of the reactants are higher. More concentrated particles, more collisions. More collisions, more effective collisions.

Each reactant could have a different effect on the reaction rate. For some reactants, doubling the concentration could have no effect on the rate of the reaction. (Zero order) For some reactants, doubling the concentration could cause the reaction rate to double. (First order) For some reactants, doubling the concentration could cause the reaction rate to quadruple. (Second order) The only way to know is by experiment!

Reaction Mechanisms

Chemical reactions do not usually occur in a single step. Instead, chemical reactions occur in a series of simple steps called a reaction mechanism.

Facts about Reaction Mechanisms The chemical equation written to describe a reaction is the sum of all of the steps in the mechanism. It is not possible to determine the mechanism simply by looking at the overall reaction. Each step in a reaction mechanism involves one particle or two particles colliding (very rarely, there are occasions when three particles must collide simultaneously).

Facts about Reaction Mechanisms If a step in a mechanism involves just one reactant particle, it is called unimolecular. If the step involves two reactant particles, it is bimolecular. Reaction mechanisms are educated guesses based on theory and experimentation. Mechanisms often involve the formation of products in early steps that are used up in later steps. These substances are referred to as intermediates and do not appear in the overall balanced equation.

Facts about Reaction Mechanisms Often, a substance is used as a reactant in an early step and given off as a product. This is one example of how a catalyst might work in a chemical reaction. The steps in a mechanism occur at different rates. The overall rate of the reaction is a result of the slowest step in the mechanism. Only the concentrations of the reactants that appear in the slow step will affect the rate of the reaction. The orders in the experimental rate law relate to the coefficients in the slowest step.

A catalyst is a substance that changes the rate of a chemical reaction without itself being used up. Homogeneous catalysts speed reactions by providing a different reaction mechanism with a lower activation energy. The lower activation energy results in a larger fraction of effective collisions

(a) The catalyst provides an alternate, low-energy path from the reactants to the products. (b) A larger fraction of molecules have sufficient energy to react when the catalyzed path is available.

What is the overall reaction? What is the catalyst? Are there any intermediates?

Heterogeneous catalysts are typically solids. Instead of participating in the reaction, heterogeneous catalysts attract the reactant particles to its surface. This increases the frequency of collisions. The Haber Process, which is used in the production of ammonia, utilizes a heterogeneous catalyst.

The Haber process. Catalytic formation of ammonia molecules from hydrogen and nitrogen on the surface of a catalyst.