The rate and extent of chemical change

Rates of reactions

• The rate of a reaction is the speed at which reactants turn into products

• Rates of reactions can be measured either through how fast reactants are used up, or by how fast products are formed

• Rate of reaction = Quantity of products formed (or reactants used) / time taken

  • This shows the average rate, where in reality reactions start quick and slow over time

Factors affecting rates of reactions

• Rates of reactions are dependent on temperature, surface area, pressure and catalysts

• Collision theory is the idea that in order for particles to react, they have to collide with sufficient energy (activation energy)

  • If they don’t have this, no reaction will occur

  • The rate depends on the amount of energy particles have, and the frequency of collisions between particles

• Temperature

  • As temperature increases, particles gain more kinetic energy

    • This causes particles to move faster, so they collide more frequently with more energy    

    • This makes the reactants more likely to exceed the activation energy, and the higher rate of successful collisions increases the rate of reaction

• Concentration/pressure

  • How many particles per unit of volume - concentration is usually solutions, pressure is usually gases

  • As concentration/pressure increases (lower volume, higher no. of particles per unit of volume), their are more particles per unit of volume

    • This causes more frequent collisions, and a higher rate of reaction

• Surface area

  • A higher surface area to volume ratio creates a larger surface for collisions, so a higher frequency and therefore a higher rate of reaction

• Catalysts

  • Catalysts are substances that speed up a reaction without being used up themselves     

    • Transition metals, enzymes, etc.

  • They lower the activation energy required by providing an alternative pathway for reactants

Measuring rates of reactions

• The mean rate in a set time is equal to to the gradient between two points

• The actual rate at a specific time is equal to the gradient at that point’s tangent

Reversible reactants and dynamic equilibrium

• In one way reactions, once products are formed they can’t change back to their reactants (break down into them)

• In reversible reactants, reactants can form products and products can break down into reactants

  • NH₄Cl ⇌ NH₃ + HCl 

  • Different directions of reactants can happen at different rates 

    • When there are lots of reactants, the forward reaction is fast until products begin to form

    • Then the forward reaction slows, and the backwards reaction speeds up 

    • Eventually, the forward and backwards reactions reach the same speeds - at equilibrium

• Equilibrium is when both reactions are still happening, but there is no overall change in concentration (effectively cancel out)

  • The concentration of reactants and products don’t change, and there are equal speeds but NOT equal masses

• If there are more products than products at equilibrium, the reaction lies to the right

• If there are more reactants at equilibrium, the reaction lies to the left

  • This can change depending on conditions and the reaction

• Equilibrium can only be reached in a closed system, where there is no energy loss to the surroundings 

• Reversible reactions are always exothermic in one direction, and endothermic in the other

La Chatelier's principle

• La Chatelier’s principle states that changing the condition of a reversible reaction causes the position of equilibrium to try and counteract the change

  • This can be used to predict the outcome of the reaction

• If equilibrium was in the centre of a reaction in a closed system, with the exothermic reaction forwards, and endothermic backwards:

  • Decreasing temperature causes the equilibrium to move in the exothermic direction to release heat 

    • Increasing temperature causes a move left (endothermic)

  • Increasing pressure makes the equilibrium to want pressure to decrease, so the equilibrium moves to the side with the less molecules being produced

    • Decreasing pressure causes a want for an increase in pressure, so more molecules being produced

  • A higher concentration of one reactant/product causes the equilibrium to shift to the opposite side to reverse the concentration change

DONE!!!