Kinetics/ collision theory

Kinetics

• Study of factors that affect the rates of reactions and the mechanisms (steps) by which reactions take place

Reaction rate

• number of particles ( atoms, ions,molecules) that react in a given time to form products.

  Reaction equation:change [R]/ change in time

Collision theory

• Particles must collide in order to react

• Collisions must be effective ( must have sufficient energy to break existing bonds and form new ones)

• must have correct orientation ( hit a specfic way)

Activation energy

The minimum energy colliding particles must have in order to react

Potential energy graph

• Potential energy of reactants or products is the difference from when they are a straight line until the ground

• Heat of reaction is the difference between the straight line of the products vs the reactants

• The reverse reaction is from the hump up until the straight line of the products

• the activation energy is from the hump up until the straight line of the reactions

• the activation complex or transition state is from the hump until the ground.

Activation complex:

• Bonds break and rearrange themselves

• Forward reaction

Factors that affect reaction rate

• Temperature

• Surface area ( particle size —> greater surface area so particles can hit easier)

• Catalyst ( substance that lowers activation energy without getting used up)

• Concentration ( increasing amount will increase collision with particles)

Calculating rate of reaction

• Look at coefficients of the balanced equation

• Multiply by the amount of moles for the substance you want ( ex oxygen) over the amount of moles you already have ex) The reaction rate of ozone is 8.76 × 10^-3 M/s over a certain interval of time. What is the rate of appearance of O2 during this interval?

  Using the balanced equation you can do 8.76 × 10³ x 3/2 since in the balanced equations its 2O₃ —> 3O₂

• Compare coefficients for each subtance

Reversible reaction

ex) A —>B

• Can become B and then through collisions also go back to A

• rate: change in [A]/ change in time

• will reach equilibrium

Equilibrium

• Rate of forward reaction is equal to the rate of the reverse reaction

• RATE OF PRODUCTION OF REACTANTS IS EQUAL TO THE RATE OF PRODUCTION OF PRODUCTS

• RATE IS THE SAME BUT YOU CAN HAVE MORE OF REACTANTS OR PRODUCTS

Equilibrium expression:

K_eq= [ products]/ [reactants]

• If k > 1 products are favored ( reaction goes to completion)

• if k < 1 reactants are favored ( mostly reactants in container)

• if k=1 neither reactants nor products are favored ( significant amount of both reactant and products in container)

• In equilibrium expressions you may only include substances ( reactants or products) that are gases or aqeuous. Liquids and solids have “ fixed” concentrations.

La Châtelier’s principle

• if a “ stress’ ( change) is applied to a reaction at equilibrium the reaction “ shifts” to relieve the stress

• Types of stress: change concentration of a reactant or product , change temp by adding or removing heat, change the pressure by increasing or decreasing the volume of container ( gases only)

• “ shifts” one way ( forward of reverse) reaction has a temporary increase in reaction rate in one direction

How these stresses apply:

• If concentration of reactants increases it shifts right

• if concentration of reactants decrease it shifts left to make more of what was removed

• If concentrations of products increase reaction shifts left

• If heat is added to the container reaction shifts left ( shifts in the endothermic reaction)

• If he is removed reaction shifts right ( in the exothermic direction).

• If pressure of container is increased reaction shifts towards the side with fewer moles of gase

• If pressure of container is decrease rection shifts where there are more moles.

• INCREASING VOLUME= REDUCING PRESSURE

• DECREASING VOLUME = INCREASING PRESSURE

Endothermic vs exothermic graphs

• Endothermic reactants are low and products are high

• Exothermic reactants are high and products are low.