2. The collision theory of how chemical reactions occur

Doc Brown's Chemistry KS4 science GCSE/IGCSE/O level Revision Notes - Factors affecting the Speed-Rates of Chemical Reactions - Doc Brown's chemistry revision notes: basic school chemistry science GCSE chemistry, IGCSE  chemistry, O level & ~US grades 8, 9 and 10 school science courses or equivalent for ~14-16 year old science students for national examinations in chemistry

Rates of reaction notes INDEX

2. The theory of how reactions happen

MORE COLLISIONS INCREASE THE RATE OF A REACTION

(the more the particles hit each other the greater the probability of reaction!)

MORE ENERGETIC COLLISIONS INCREASE THE RATE OF A REACTION

(the more kinetic energy the particles have, the more likely they are to break bonds and form products, so the faster the reaction goes!)

• WHAT CAUSES A CHEMICAL REACTION?

• WHAT MUST HAPPEN FOR A CHEMICAL REACTION TO TAKE PLACE?

• CAN WE MAKE PREDICTIONS ABOUT HOW THE SPEED OF A REACTION MAY CHANGE IF THE REACTION CONDITIONS ARE CHANGED?

PARTICLE COLLISION THEORY

• Reactions can only happen when the reactant particles collide, but most collisions are NOT successful in forming product molecules despite the incredible high rate of collisions between ALL the particles in ANY liquid or gas.

  • The collision frequency is about 10⁹ per second between air molecules at room temperature!

  • It means even in the air around you, although no chemical reactions are usually taking place, each oxygen, nitrogen and any other molecule is undergoing around a 1000 million collisions are second! scary!

  • So, if there are so many collisions, even in a reacting mixture, why doesn't every reaction go at an explosive rate!

• The reason is that particles have a wide range of kinetic energy BUT only a small fraction of particles have enough kinetic energy to break bonds and bring about chemical change.

  • The diagram above tries to give you an idea about the concepts of fruitful collisions (minority) leading to products and the vast majority of collisions are unfruitful, producing no product, the molecules just bounce of each other.

• The minimum kinetic energy required for a reaction to take place is known as the activation energy (shown in the diagrams below).

  • (i)

    • Reaction profile (i) An activation energy diagram for an exothermic reaction.

  • (ii)

    •  Reaction profile (ii) An activation energy diagram for an endothermic reaction.

• This 'activation' kinetic energy is needed and to be sufficient to break bonds in the reactant molecules so new bonds are created when the reaction products are formed.

• The majority of particle collisions do NOT form products - otherwise all reactions would be superfast!

• The minority high kinetic energy collisions between particles which do produce a chemical change are called 'fruitful collisions', those that don't produce products are called 'unfruitful' collisions.

• The reactant molecules must collide with enough kinetic energy to break the original bonds to enable new bonds to form in the product molecules.

• Basically reaction rates are controlled by the frequency of collision of reactant particles AND the kinetic energy the particles have.

  • The more collisions there are in a given time AND the greater the kinetic energy the particles have, the faster the reaction goes, and each rates factor requires a particular interpretation of these concepts and ideas.

  • Collision frequency means the 'rate of particle collision' or 'the number of collisions in a given time'.

• ALL the rate-controlling factors described in sections 3a to 3e are to do with either ...

  • (a) [sections 3a, 3b and 3c] the collision frequency (chance of collision) to give a fruitful collision and products,

    • so increasing the reactant concentration of solutions, increasing gaseous reactant pressure or reducing particle size of a solid reactant (increasing surface area) all favour increasing the rate of fruitful collisions,

  • OR,

  • (b) [section 3d and 3e] the combined kinetic energy of reactant particle collision (>= activation energy) to give a fruitful collision and products,

    • so, increasing temperature increases the KE of particles giving more fruitful energetic collisions,

    • AND, using a catalyst to decrease the activation energy means more molecules already have enough kinetic energy to overcome the activation energy and react without having to increase the temperature.

  • both these explanations are all about the 'chance of a fruitful collision' leading to reactant bonds breaking product formation via new bonds forming.

• 'Concept picture'

• In the case of temperature, the energy of the collision is even more important than the frequency effect.

• In each of the sections 3a to 3e the collision theory is applied in more detail to that particular factor affecting the speed/rate of a reaction, so read on!

• The particle theory of gases and liquids and the particle diagrams and their explanation will also help you understand or describe in your coursework what is going on.

• For more details on activation energy see GCSE/IGCSE/O Level notes on Chemical Energetics

• For A Level students Advanced Level Chemistry Theory pages on "CHEMICAL KINETICS" covers all advanced theoretical aspects of rates of reaction.

• More details of laboratory investigations ('labs') involving 'rates of reaction' i.e. experimental methods for observing the speed of a reaction and including the effect of a catalyst are given in the INTRODUCTION

Concept picture for a solid catalyst

Concept picture for a solid reactant and solution

Rates of reaction notes INDEX

GCSE/IGCSE MULTIPLE CHOICE QUIZ on RATES of reaction

Website content © Dr Phil Brown 2000+. All copyrights reserved on revision notes, images, quizzes, worksheets etc. Copying of website material is NOT permitted. Exam revision summaries & references to science course specifications are unofficial.