Collision Theory and Activation Energy Unit 3: Chemical Kinetics and Equilibrium

Review: Chemical Reaction Generic Format of reaction 5 different types of chemical reactions A + B C + D ReactantsProducts 1) Synthesis 2) Decomposition 3) Single Displacement 4) Double Displacement 5) Combustion

Review: Kinetic Molecular Theory  Used to explain many observations and chemical events! All matter is made up of microscopic-sized particles (atoms, ions, molecules) These particles are in constant motion (possess kinetic energy) There is space between the particles (speed and spacing determine the physical state of matter) Adding energy increases the speed of the moving particles (thus inc kinetic energy)

How do reactions occur? Collision Theory

Turns out…… By calculating how many collisions are taking place per second and how quickly product is being produced… chemists learned that most collisions are not successful (no product formed) To think that reactant particles collide and products are automatically produced is over simplified. There must be other requirements for a collision to be successful!

The Collision Theory is an explanation of what is necessary for a chemical reaction to occur. When a chemical reaction takes place, the reactant particles must meet two conditions (or requirements) during collision for the collisions to be successful: 1. Proper orientation 2. Particles must collide with a certain minimum amount of energy, called activation energy.

o Particles must collide with the proper geometry or orientation for atoms to come in direct contact and form the chemical bonds of the products. 1. Orientation

2. Activation Energy Particles must collide with a certain minimum amount of energy, called activation energy (E a ). This energy is required to break chemical bonds in the reactants. Note: The energy of each particle is not important, it is the energy of the collision.

2. Activation Energy

Effect of Temperature:

Potential Energy Diagram: We can represent the increase in potential energy during a chemical reaction using a potential energy diagram The kinetic energy of reactants is transferred to potential energy as the reactants collide

“The hill” The hill represents the activation energy that must be overcome for the reaction to occur “high hill”  slow rate of reaction “low hill”  fast rate of reaction Top of hill is called “change over point”  There is a chemical species that exists here is referred to as “activated complex” Activated complex Neither reactant or product Partial bonds, highly unstable

Potential Energy Diagram Handout

If both of these conditions are not met, particles will merely collide and bounce off one another without forming products.

Some collisions are successful…. Although, the percentage of successful collisions is extremely small, chemical reactions still take place at a reasonable rate because there are so many collisions per second between reactant particles!

Reaction Mechanism: Converting reactants to products often involves more then one step…as you know! Each step is called an elementary reaction Molecules formed during elementary reactions are called reaction intermediates (neither reactants nor products) Example: 2NO + O 2  2NO 2  Step 1: NO + O 2  NO 3  Step 2: NO 3 +NO  2NO 2

The Rate Determining Step: In multi-step reactions there can be 2 or more elementary reactions….  There is always one reaction that is slowest …this determines the overall rate of reaction  Thus, the slowest elementary reactions is called the rate determining reaction.

Catalysts: Works by lowering the activation energy of a reaction so that a larger fraction of reactants have sufficient energy to react They do this by providing an alternative reaction mechanism

For example:  A + B  AB (no catalyst)  With catalyst: Step 1: A + catalyst  A-cat Step 2: A-cat + B  AB + catalyst Overall: A + B  AB  Both steps are faster then the original, uncatalyzed reaction  Catalyst remains unchanged in the end

Catalysts cont’d

Real life catalyst example: Manganese dioxide (a black powder) will catalyze the breakdown of hydrogen peroxide. 2H 2 O 2 (aq) → 2H 2 O(l) + O 2 (g) (uncatalyzed) With catalyst: Car exhaust pipes use catalytic converters to get rid of some of the nasty gases from the engines.

Practice: Read Chapter 12!!!!! Pg 484 # 9, 12a,b,d Pg 486 # 1-4, 12