Reaction Rate

1. Reaction Rate • The rate of reaction is how fast the reaction is. • To measure rate we must take into consideration 2 measurable quantities : • one of these must be time • the other could be the decrease in mass of the products or the increase in mass of the reactants.

2. Calculating Rate Average rate = Change in reactant or product / time taken for change.

3. Example Example – calculate the rate of reaction in the first 20 sec Average rate = Change in mass /time taken = 50 – 20/ 20 – 0 = 30/20 =1 .5g s -1

4. Collision Theory! • For a chemical reaction to take place the reactant particles must collide with each other. • The more frequently they collide – the greater chance they will be in the correct “ geometry” and so the collision will be successful and a reaction can take place!

5. Factors that affect rate • Temperature • Concentration • Particle size • Catalyst

6. Concentration • When we increase the concentration of reactants the rate increases. This is because the more particles - The greater chance they will collide - The greater chance the collision will be successful i.e. correct geometry - The greater chance of a reaction! ( PPA)

7. Particle Size • In order for particles to collide their surfaces must come into contact. • The bigger the surface – the bigger the chance of collision. • The greater chance of successful collisions • The greater chance of a reaction. • The smaller the particle – the bigger the surface area.

8. Activation Energy • This is the minimum amount of energy particles must have in order to react with each other. • The lower the activation energy, of a reaction, the greater chance the reaction will happen successfully.

9. Temperature • When we heat up particles – they move! (kinetic energy) • Temperature is a measure of the average kinetic energy. • If we increase the temperature more particles will have equal to or greater than Activation Energy. • More collisions – more successful collisions – increase in reaction rate! ( PPA)

10. Catalysts • Catalyst lower the activation energy of a reaction. • Therefore there is more chance of collisions – more successful collisions – increase in rate! • Catalysts are not used in the reaction – they can be used again! • Catalyst inhibitors will slow down rate.

11. Types of Catalysts • Homogeneous – Catalyst and reactants are in same state. • Heterogeneous – Catalyst and substrate are in a different state. • Biological - Enzymes

12. Homogeneous Catalysts • A homogenous catalyst takes part in the reaction and is then reformed at the end. • The reactants and catalyst are in the same physical state! • Example – Cobalt ions are the catalyst in the reaction between Hydrogen peroxide and Potassium sodium tartrate.

13. Heterogeneous catalysts • The catalyst provides a surface for the reaction. • The reactant particles are absorbed onto “active sites” on the catalyst surface. • The reactant bonds weaken. • Particles are in the correct position for “ collisions” • Collision occur – new product bonds form. • The new product particles leave catalyst surface.

14. Biological catalysts • Enzymes are catalysts inside living tissue. • They are very specific – usually only catalyse one reaction. • They work at “ optimum conditions” – i.e. pH, temperature. • E.g – amylase helps the break down of starch to sugar.

15. Denaturing Catalysts • This is when a catalyst is destroyed. • It can be caused by the “ active site” being blocked. • E.g. lead petrol denatures the active site in a catalytic converter. Or • Too high a temperature or the wrong pH would denature an enzyme.

16. Exothermic Reactions • An exothermic reaction is one where heat is released. This means that the energy released when the product bonds were made is greater then the energy put in to break the reactant bonds. • E.g. combustion, neutralisation.

17. Endothermic reaction • This is when overall energy is taken in. • The energy released when the product bonds were made is less than the energy put in to break the reactant bonds. • E.g. Dissolving some salts – ammonium carbonate.

18. Enthalpy Change • The difference between potential energy of reactants and products is called – ENTHALPY CHANGE. • It is denoted by the symbol ΔH • ΔH = Hp – Hr • Hp = enthalpy of products • Hr = enthalpy of reactants.

19. ΔH • The ΔH of exothermic reaction is always a negative value. • The ΔH of endothermic reactions is always a positive value. ( graphs)

20. Calculations using equations! • Example • What mass of water is produced when 1g of Methane is burned. • Step 1 – Balanced equation • CH4 + 2O2 —> CO2 + 2H2O • Step 2 – Mole ratio • 1 mole of CH4 produces 2 moles of H2O.

21. Step 3 – Mass • 1 mole = 16g of CH4 —>2moles of H2O = 2 x 18 36g 1g of CH4 —> 1 / 16 x 36 =2.25g

22. Excess! • Example • 8g of Methane reacts with 16g of Oxygen. Which reactant is in excess? • Step 1 – Balanced Equation. • CH4 + 2O2 —>2H2O + CO2 • Step 2 – Mole ratio. • 1 mole of CH4 reacts with 2 moles of O2

23. Step 3 – Put in Mass • CH4:O2 • 1 : 2 • 16g —>2 x 16g = 32g • Step 4 – Put in actual moles • 8g of CH4 = 0.5 moles • 16g of O2 = 0.5 moles • Since we only need 0.25 moles of CH4 to react with 16g of O2 – CH4 is in excess.