(c) doc bDoc Brown's Chemistry KS4 science GCSE/IGCSE/O level Revision Notes

Factors affecting the Speed-Rates of Chemical Reactions

2. The collision theory of how chemical reactions occur

What must happen for two chemicals to react together? How do reactions take place? Why do reactions take place? What is collision theory and can it be used to explain factors affecting the speed of a reaction? What is the activation energy of a reaction? Why is the activation energy of a reaction so important? These revision notes are suitable for GCSE IGCSE O Level KS4 science chemistry students studying the theory of how chemical reactions happen. The descriptions of experiments to do with rates of reaction provide data that require theoretical explanations using particle models, and the explanations below (and the other rates pages) should help with homework, coursework assignments and interpreting laboratory experiments ('labs') on the factors affecting the rate (speed) of a chemical reaction. These notes on the collision theory of chemical reaction and the factors affecting reaction rate and activation energy are designed to meet the highest standards of knowledge and understanding required for students/pupils doing GCSE chemistry, IGCSE chemistry, O Level chemistry, KS4 science courses and can be useful primer for A Level chemistry courses. These revision notes on the particle collision theory of how chemical reactions take place, should prove useful for the new AQA GCSE chemistry, Edexcel GCSE chemistry & OCR GCSE chemistry (Gateway & 21st Century) GCSE (91), (9-5) & (5-1) science courses.


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 faster the reaction!)

MORE ENERGETIC COLLISIONS INCREASE THE RATE OF A REACTION

(the more kinetic energy the particles have the faster the reaction!)

  • 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?

  • 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)(c) doc b

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

    • (ii)(c) doc b

      • (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 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 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.

  • ALL the rate-controlling factors described in section pages 3a and 3c to 3e are to do with either ...

    • (a) 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) 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.

  • 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!


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GCSE/IGCSE MULTIPLE CHOICE QUIZ on RATES of reaction


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