The document discusses the rate of chemical reactions and factors that affect it. It provides examples of reactions that occur at different rates and how rate is calculated. The average rate and instantaneous rate are defined. Experiments are described to determine the effect of surface area, concentration, temperature, catalyst, and pressure on the reaction rate. The concept of effective collision is introduced, where particles must collide with sufficient energy and correct orientation for a reaction to occur. Factors that increase collision frequency or lower activation energy can increase the reaction rate.

1. CHAPTER 1 RATE OF REACTION

2. 1.1 RATE OF REACTION Why a tall building can be destroyed with explosives within seconds??

3. Why food has to take several hours to be decay ??

4. Why do coal mines face a high risk of explosion ??

5. All these are because : difference of rate of reaction

7. What is the rate of reaction ? The speed at which reactants are converted into products in a chemical reaction For examples:

8. 1. Fireworks display occurs in just a few seconds Rate of reaction HIGH

9. 2. The photosynthesis process takes a few hours Rate of reaction LOWER

10. 3. The erosion process of stones takes many years Rate of reaction VERY LOW

11. Some reactions are fast & some reactions are very slow So do the chemical reactions in laboratory

12. How we determine & calculate the rate of chemical reactions in laboratory ??

13. It just same with the way we calculate the speed of driving How we calculate the speed of driving ??

14. Yes, the formulae is : Speed = distance (km) time (hour)

15. In the other hand, it is : Rate of = physical changes reaction time

16. What are the changes that we could measure?? Mass before and after experiment (gs -1 ) Colour changes Temperature changes (°Cs -1 ) Pressure changes (Ps -1 ) Volume of gas liberated (cm 3 s -1 ) (only if the reaction produce gases)

18. A + B C + D (reactants) (products) During the chemical reaction, reactants will become less but products become more

19. Decomposition of hydrogen peroxide to form water and oxygen gas H 2 O 2 (aq) H 2 O (l) + O 2 (g)

20. Take 5

21. For any chemical reaction, we can measure its average rate of reaction & rate of reaction at a given time

22. What’s that ??

23. Rate of = physical changes reaction time average rate of reaction rate of reaction at a given time Average rate = total changes of reaction total time taken Rate of reaction at a given time = gradient of the curve at that instant = y 2 – y 1 / X 2 – X 1

24. Total changes Total time taken Average rate = total changes of reaction total time taken Volume of hydrogen gas ( cm 3 ) Time ( s )

25. Changes = y 2 – y 1 time taken = X 2 – X 1 Rate of reaction at 90 second = gradient of the curve at that instant = y 2 – y 1 / x 2 – x 1 find the given time draw a line to curve draw a tangent Measure y 2 – y 1 Measure x 2 – x 1 Volume of hydrogen gas ( cm 3 ) Time ( s )

27. Let’s see an example:

28. CaCO 3 + HCl  CaCl 2 + H 2 O + H 2 burette water Hydrochloric acid Marble chips Hydrogen gas

29. Volume of hydrogen gas was recorded every 30 seconds

30. . . . . . . . . . . . . . . . . after 10 minutes . . .

31. Here’s the result : Time / s Reading of burette / cm 3 Volume of gas / cm 3

32. Then we plot the graph: Volume of hydrogen gas ( cm 3 ) Time ( s ) 5 10 0 35 30 25 20 15 60 30 90 180 210 240 270 300 150 120

33. Get the answers: Calculate the average rate of reaction between marble chips and hydrochloric acid. Solution: refer to table Time / s Reading of burette / cm 3 Volume of gas / cm 3

34. Average rate of reaction = total changes / total time taken Average rate of reaction = 34 cm 3 300 s = 0.113 cm 3 s -1 ** you’ll get the same answer by using graph

36. b) Calculate the average rate of reaction in the first 90 seconds . Solution: this is accumulated !! Time / s Reading of burette / cm 3 Volume of gas / cm 3

37. Average rate of reaction = total changes / total time taken Average rate of reaction = 20 cm 3 90 s = 0.222 cm 3 s -1 ** you’ll get the same answer by using graph

39. c) Calculate the rate of reaction at a) 60 second b) 210 second

40. Solution (a) : 22.00 – 10.50 = 11.50 cm 3 90-30 = 60 s Rate of reaction at 60 second = y 2 – y 1 x 2 - x 1 = 22.00 -10.50 90 -30 = 11.50 cm 3 60 s = 0.192 cm 3 s -1 ** 5 10 35 30 25 20 15 Time ( s ) 90 Volume of hydrogen gas ( cm 3 ) 0 60 30 180 210 240 270 300 150 120

41. Solution (b) : 35.50 – 30.00 = 5.50 cm 3 255-160 = 95 s Rate of reaction at 210 second = 35.50 -30.00 255 -160 = 5.50 cm 3 95 s = 0.058 cm 3 s -1 ** 5 10 35 30 25 20 15 Time ( s ) 90 Volume of hydrogen gas ( cm 3 ) 0 60 30 180 210 240 270 300 150 120

43. d) Compare the rate of reaction at 60 second and 210 second Solution: Rate of reaction at 60 second = 0.192 cm 3 s -1 Rate of reaction at 210 second = 0.058 cm 3 s -1 The rate of reaction at 60 second is higher than 210 second. This means the reaction occurs faster at the moment of 60 second

44. e) Why the rate of reaction at 60 second is higher than 210 second ? Solution: Because the concentration of reactants is decreasing during the chemical reaction.

45. f) There are no more hydrogen gas liberated after 300 second, why? Solution: Because the chemical reaction is stopped All the marble chips had been reacted

47. Conclusion Rate of reaction very high Rate of reaction lower Curve becomes horizontal line No more reaction occur Rate of reaction = 0 cm 3 s -1 Rate of reaction More lower Total volume of gas Liberated (cm 3 ) Time (s / min)

48. Learning Task 1.2 page 8 WORKSHOP HOUR

50. 1.2 Factors affecting the rate of reaction

51. Some reactions are fast & some reactions are very slow So do the chemical reactions in laboratory

52. What are the factors affect the rate of reaction ??

53. 2. T emperature 1. T otal surface area 4. C atalyst 5. P ressure 3. C oncentration 2T 2C 1P

54. How does the factors of Total surface area, Temperature, Concentration, Catalyst & Pressure affect the reaction rate?? 2T 2C 1P

55. Let’s do the experiment to investigate that: 1.Surface area smaller , rate of reaction 2.Concentration higher, rate of reaction 3.Temperature higher, rate of reaction 4.Catalyst added, rate of reaction

58. Surface area smaller , rate of reaction Concentration higher, rate of reaction Temperature higher, rate of reaction Catalyst added, rate of reaction Pressure higher, rate of reaction Conclusion

60. Problem solving

61. A + B C + D (reactants) (products) During the chemical reaction, reactants will become less but products become more The chemical reaction will stop once one of the reactants is used up .

62. Pattern of graphs Reactant A + Reactant B  gases Less Less More

65. More examples of interpreting graphical results

66. Mg (s) + HCl (aq)  MgCl (aq) + H 2 (g) excess Draw out the curves of graph for: 1. Magnesium mass / 2. Hydrochloric acid concentration/ 3. Magnesium chloride concentration / 4. Hydrogen gas volume against the time

67. 1. the curves of graph for Magnesium mass against the time Magnesium Excess >>>

68. 2. the curves of graph for hydrochloric acid concentration against the time Hydrochloric acid Limited reactant

69. 3. the curves of graph for magnesium chloride concentration against the time Hydrochloric acid has been used up & reaction stopped

70. 4. the curves of graph for hydrogen gas volume against the time Hydrochloric acid has been used up & reaction stopped

71. Q1. Compare these: Set 1 : 1g of zinc plate (excess) reacts with sulphuric acid 0.1 mol dm -3 to form zinc sulphate and hydrogen gas Set 2 : 1g of zinc powder (excess) reacts with sulphuric acid 0.1 mol dm -3 to form zinc sulphate and hydrogen gas Find the difference !!

72. Hydrogen Gas / cm 3 Time / s Set 2 Set 1 The number of moles of chemical subtances are same Factor : total of surface area / Particle size

73. Q2. Compare these: Set 3 : 1 g of zinc powder (excess) reacts with 500cm 3 sulphuric acid 0.1 mol dm -3 to form zinc sulphate and hydrogen gas Set 4 : 1 g of zinc powder (excess) reacts with 250cm 3 sulphuric acid 0.2 mol dm -3 to form zinc sulphate and hydrogen gas Find the difference !!

74. Hydrogen Gas / cm 3 Time / s Set 4 Set 3 The number of moles of chemical subtances are same Factor : Concentration

75. Q3. Compare these: Set 5 : 1 g of zinc powder (excess) reacts with 250cm 3 sulphuric acid 0.1 mol dm -3 to form zinc sulphate and hydrogen gas Set 6 : 1 g of zinc powder (excess) reacts with 250cm 3 sulphuric acid 0.2 mol dm -3 to form zinc sulphate and hydrogen gas Find the difference !!

76. Hydrogen Gas / cm 3 Time / s Set 6 Set 5 The number of moles of chemical subtances are same Factor : Concentration

77. Q4. Compare these: Set 7 : 1 g of zinc powder (excess) reacts with 25cm 3 sulphuric acid 0.15 mol dm -3 to form zinc sulphate and hydrogen gas Set 8 : 1 g of zinc powder (excess) reacts with 50cm 3 sulphuric acid 0.1 mol dm -3 to form zinc sulphate and hydrogen gas Find the difference !!

78. Hydrogen Gas / cm 3 Time / s Set 8 Set 7 The number of moles of chemical subtances in Set 8 is higher, so the volume of gas realeased is higher Factor : Concentration

79. 1.3 Collision Theory

80. Find out what they means!! Collision Correct orientation Activation energy Effective collision Collision frequency Effective collision frequency Energy profile diagram Achieve effective collision

81. Base on kinetic theory, particles moves constantly and collide each other all the time However, majority of collisions do not lead to a reaction,why ?

82. Only those in which the colliding species have: Achieve minimum amount of energy, called the ACTIVATION ENERGY, E a And with CORRECT ORIENTATION EFFECTIVE COLLISIONS

83. For a reaction to take place, the particles of the substances that are reacting have to collide . If they collide, with enough energy then they will react. The minimum amount of energy that particles need to react is called the ACTIVATION ENERGY

85. Activation energy

87. therefore some main ways of increasing the rate of a reaction: 1) increase the number of collisions 2) increase the amount of kinetic energy so that more collisions lead to a reaction 3) decrease the energy activation so that more reactants could be reacted what factors could cause 1, 2, 3?

88. Size of solid reactant smaller Total surface area exposed to collision with other reactant particle is bigger The frequency of collision between reactant particles increase Frequency of effective collision increase Rate of reaction higher Surface Area

89. Higher temperature Kinetic energy of reactant particles increase Particles move faster and collide more often The frequency of collision between reactant particles increase Frequency of effective collision increase Rate of reaction higher Temperature

90. Higher concentration More number of particle per unit volume of solution The frequency of collision between reactant particles increase Frequency of effective collision increase Rate of reaction higher Concentration

91. Positive catalyst used enables the reaction occur through an alternative path which requires lower activation energy More colliding particles are able to achieve the lower activation energy Frequency of effective collision increase Rate of reaction higher Catalyst

92. Activation energy Lower Activation energy

93. Pressure Higher pressure More number of particle per unit volume of solution The frequency of collision between reactant particles increase Frequency of effective collision increase Rate of reaction higher

94. Conclusion