1. REACTION KINETICS (AS) 1.Rate of reaction = change in concentration of reactant or product over time Rate of reaction =  [reactant]/  time OR  [product]/  time

2. 2.Concentration –time graphs time Conc of a reactant Conc of reactant decreases with time

3. time Conc of product After certain time,conc of products becomes constant Conc of product increases with time

4. a. Rate of reaction at time, t : (instantaneous rate) draw a tangent to the concentration vs time curve at time t the gradient of tangent = rate of reaction

5. Example time [reactant ] t y x Gradient = y/x = rate of reaction at time, t Unit : mol dm -3 s -1 or mol dm -3 min -1

6. Note : i)Average rate : rate measured over a period of time Eg : rate = change in [reactant]/ t 2 – t 1 ii)Initial rate : rate at almost t=0 b. Rate of rxn is proportional to concentration of most reactants Concentration increases, rate increases

7. Note : Rate is independent of concentration of a reactant Concentration changes but rate is constant Zero order reaction time Conc of reactant Conc decreases with time Constant gradient Rate is constant

8. THEORIES OF REACTION RATES 1. Collision theory : a. reactions occur due to collision of reactant particles b. not all collisions results in reaction effective collisions : collisions between reacting particles that results in a reaction

9. c.Characteristics of effective collisions : i) have favourable orientation eg C – C – C – C –Br + OH -  C – C – C – C –OH + Br - collision of an OH - with the bromoethane molecule is unlikely to result in a reaction if it hits the end of the molecule away from the Br

10. ii) possess a minimum energy = E a (1) Definition : Activation energy,E a, is the minimum energy required for a reaction to take place High E a  slow reaction (2) E a is used to enable bonds in the reactants to stretch and break as new bonds form in the products

11. 2. Transition state theory : a. reactions takes place via transition state in which reactants come together b. bond making and breaking occur continuously and simultaneously In the transition state, bonds are in the process of making and breaking.

12. A-B + C  A + B-C A B C transition state Bond formingBond breaking

13. c. reaction profile/enthalpy diagram : Note : (1) Transition state is the highest point in the reaction profile (2) Energy gap between reactants and transition state = E a (3) E a forward rxn ≠ E a reverse rxn

14. Reaction profile or energy / enthalpy diagram for uncatalysed reactions exothermic reversible reaction Extent of reaction Energy Products Reactants Transition state E a forward rxn E a reverse rxn HH

15. endothermic reversible reaction Extent of rxn Energy Reactants Products Transition state E a reverse rxn E a forward rxn HH

16. d. Multi step reaction Reaction that takes place via an intermediate Mechanism of rxn involves a multi step reaction The intermediate will occur at a minimum on the graph One minimum = one intermediate

17. Eg : Step 1 : Reactants  Intermediate,  H = positive Step 2 : Intermediate  Products,  H = negative Overall : Reactants  Products,  H = negative

18. Energy Extent of rxn Reactants Products Transition state 1 Transition state 2 Intermediate E a (1) E a (2) Overall  H

19. e. Reacting particles must possess energy greater than or equal to the E a before they can react

20. FACTORS AFFECTING RATE OF REACTION Concentration Temperature Catalyst

21. I. Concentration of reactants 1. conc increases, rate normally increases ( exception : zero order ) 2. as concentration increases : frequency of collisions increases no of effective collisions increases rate of reaction increases

22. 3. Expt to show effect of concentration on rate of reaction : Eg: Na 2 S 2 O 3 (aq) + 2HCl(aq)  2 NaCl(aq) + H 2 O(l) + SO 2 (g) + S(s) H 2 O(l) + SO 2 (g) + S(s) a. Effect of [S 2 O 3 2- ] on rate of reaction b. Sulphur appears as particles of solid c. Measure time taken to block view of cross/words under conical flask

23. Experiment to show effect of concentration on rate of reaction : Eg Na 2 S 2 O 3 (aq) + 2HCl (aq)  2 NaCl(aq) + H 2 O(l) + SO 2 (g) + S(s) a. Effect of conc of S 2 O 3 2- on rate of rxn b. Sulphur appears as small particles of solid c. Measure time taken for enough sulphur to form to block view of the cross/words under conical flask

24. d. Use different volumes of S 2 O 3 2- but keep volume of HCl constant e. H 2 O used to keep total volume of all mixtures constant Hence volume of S 2 O 3 2- used  conc S 2 O 3 2- eg : volume doubles, conc doubles

25. MixtureVolume of S 2 O 3 2- /cm 3 Volume of HCl/cm 3 Volume of H 2 O/cm 3 Time/s 1 10 20 30 2 20 3 40 20 0

26. Rate of reaction α 1/time From expt, As volume of S 2 O 3 2- increases, [S 2 O 3 2- ] increases, time taken decreases Rate of reaction increases

27. [S 2 O 3 2- ] 1 / time Rate of reaction α [S 2 O 3 2- ]

28. II.Temperature 1. When temperature increases : average speed of reacting particles increases particles collide more frequently and with greater energy no of particles with energy ≥ E a increases no of effective collisions increases rate of reaction increases

29. 2. Why does rate increase with temperature? Molecules in a gas does not all have the same speed. Their speeds and therefore their energies are distributed according to the Maxwell Boltzmann distribution curve

30. Maxwell Boltzmann distribution curve Energy/speed Fraction or no of molecules with energy E Most probable energy

31. a. Shape : at a temp T, molecules in a sample of gas have different speed/energy Most probable speed/energy corresponds to the maximum of the curve. b. Area under the curve = total no of molecules in the sample

32. c. As temp increases, curve flattens ( have a lower peak ) more spread out ( moves to the right ) however total no of molecules = areas under the curves remains the same

33. Effect on Maxwell Boltzmann distribution curve Energy/speed No of molecules with energy E Lower T Higher T EaEa

34. d. Shaded area = no of molecules with energy ≥ E a As temp increases, Size of shaded area increases More molecules with energy ≥ E a No of effective collisions increases Rate of reaction increases

35. Note : At temp T and ( T + 10 K ), Size of shaded area doubles No of molecules with energy ≥ E a doubles Rate of reaction doubles

36. e. Reactions with larger E a are slower but rise in temp has more significant increase on the rate of reaction with a higher E a

37. III.Effect of catalyst ( catalysis ) 1.Catalysts are substances that affects the rate of a chemical reaction without being chemically changed themselves They are not consumed and are regenerated at the end of the reaction

38. Properties of catalyst: increase the rate of reaction amount of catalyst used affects the rate which is proportional to the amount used required in small amount chemically unchanged after the reaction do not affect  H

39. 2. Two types of catalyst : a. positive catalyst : increases rate of reaction eg ferum in Haber process b. negative catalyst / inhibitor : slows down a reaction eg glycerine or phosphoric acid inhibits decomposition of hydrogen peroxide

40. 3. Action of positive catalyst Provides alternative pathway with a lower E a More molecules with energy ≥ E a No of effective collisions increases Rate of reaction increases Note : different catalyst can affect a similar reaction differently

41. 4. Diagrams : a. Enthalpy diagram or energy profile : eg exothermic rxn Reaction pathway Energy Reactants Products E a catalysed rxn(lower) E a uncatalysed rxn

42. b. Maxwell Boltzmann distribution curve ( at a certain temp T ) Energy No of molecules with energy E E a uncatalysed E a catalysed (lower)

43. For catalysed reaction : Size of shaded area increases No of molecules with energy ≥ E a increases No of effective collisions increases Rate of reaction increases Note : another factor affecting rate is surface area ( higher surface area, faster reaction )

44. 5. Types of catalyst : 3 types a. Heterogeneous catalyst : catalyst is in a different phase compared to reactants. Examples : Reaction Catalyst N 2 (g) + 3H 2 (g)  2NH 3 (g) ferum (s) ( Haber process ) 2SO 2 (g) + O 2 (g)  2SO 3 (g) V 2 O 5 (s) ( Contact process ) C 2 H 4 (g) + H 2 (g)  C 2 H 6 (g) Ni (s) ( Hydrogenation of alkenes in manufacture of margarine )

45. b. Homogeneous catalyst : catalyst is present in the same phase as the reactants. Examples: Reaction Catalyst CH 3 COOH(aq) + C 2 H 5 OH(aq) H + (aq)  CH 3 COOC 2 H 5 (l) + H 2 O (l) S 2 O 8 2- (aq) + 2I - (aq) Fe 2+ (aq)  2SO 4 2- (aq) + I 2 (aq) or Fe 3+ (aq)

46. c. Biological catalyst ( enzymes ): Proteins which catalyses chemical reactions in living systems Are extremely specific, one enzyme normally catalyses one reaction Example: amylase found in saliva. It is used to break carbohydrates into simpler molecules.

47. Autocatalysis 1. One of the product is a catalyst for the reaction 2. Reaction proceeds slowly at first at uncatalysed rate until a significant amount of the product ( also the catalyst ) is established 3. Then reaction will speed up to catalysed rate Reaction will stop when reactants are exhausted

48. Eg : 2 MnO 4 - + 16 H + + 5 C 2 O 4 2-  2 Mn 2+ + 8 H 2 O + 10 CO 2 catalyst

49. time [ MnO 4 - ] Fast decrease in conc Faster reaction Catalysed rate Slow decrease in conc Slow reaction Uncatalysed rate

50. time rate Slow Uncatalysed rate Fast Catalysed rate