1. Chpt.16: Rates of Reactions

2. Chemical reactions occur at different speeds: Lighted taper put in test tube of H 2 gas – immediate reaction and a loud bang FAST REACTION Rusting of iron – takes several months – SLOW REACTION Important for chemists to know how quickly a reaction occurs and how to change the speed of a reaction – i.e. Industry

3. The term used to describe how quickly a reaction occurs is known as the rate of a reaction. Chemists measure the rate of a chemical reaction by measuring how the concentration of any one reactant or product changes with time.

4. Definition: The rate of reaction is defined as the change in concentration per unit time of any one reactant or product Question: Why does the rate of chemical reactions decrease with time??? Answer: Concentration reactant decreases/The reactants are used up

5. Definition: Average rate of reaction: is the amount of product divided by the time taken

7. Definition: Instantaneous Rate: the rate of reaction at a particular instant in time is called the instantaneous rate. To measure instantaneous rate at given time from graph: - find the slope of the tangent to the curve at the given time i.e. the slope of the tangent tells you the rate of production of oxygen at that instant in time Slope = Rise Run

8. Measuring Rates of Reactions Mandatory Experiment: To monitor the rate of production of oxygen from hydrogen peroxide using manganese dioxide as a catalyst To calculate the average rate of evolution of oxygen: Average = Total volume of oxygen produced (cm 3 ) Rate Total time for reaction to go to completion (secs)

9. Method 1

10. Method 2

11. Question: Manganese dioxide was added to hydrogen peroxide in a flask, and the oxygen produced was collected in a gas syringe. The total volume of oxygen was recorded every 30 seconds, and the data is shown in the table: 1.Draw a graph of volume Vs. Time 2.How long did it take for 55cm 3 of oxygen to be produced 3.Calculate the rate of reaction after 90 seconds (i.e. Instantaneous rate) Time0306090120150180210240 Volume of O 2 (cm 3 ) 0203546535760

12. Factors Affecting Rates of Reaction Rate of reaction depends on five factors: A. Nature of Reactants B. Particle Size C. Concentration D. Temperature E. Presence of a catalyst

13. A. Nature of Reactants: In general – - ionic reactions which only involve the coming together of oppositely charged species in solution are fast at room temperature - covalent reactions which involve the breaking of bonds are usually slow at room temperature. (This is because it takes time for the covalent bonds to be broken and new bonds to be formed) *Demo

15. B. Particle Size: In reactions where one of the reactants is a solid, and the other reactant is a liquid, the particle size of the solid has an important effect on the rate of reaction The more finely divided the solid, the faster the reaction – when the solid is finely divided, i.e. powdered, the solid has a greater surface area. The number of particles at the surface and therefore available to react with the liquid is greatly is increased, so reaction rate is greater

16. The area of one face of the cube will be 2 x 2 = 4cm 2 The cube has six faces, so the total surface area is 4cm 2 x 6 = 24cm 2 Each of the small cubes has a face area of 1cm x 1cm = 1cm 2 The six faces give a total surface area for each smaller cube of 6cm 2 There are eight cubes so the total surface area is 6cm 2 x 8 = 48cm 2 We could cut that cube horizontally and vertically along each face so that we have eight smaller cubes.

17. Reaction of calcium carbonate (marble chips) with dilute hydrochloric acid solution: CaCO 3(s) + 2HCl (aq) CaCl 2 + H 2 O (l) + CO 2(g) Powdered marble reacts much more quickly than marble chips

18. Demonstration: To quantify the effect of particle size on the rate of reaction

21. Dust Explosion!!!! First recorded dust explosion occurred in an Italian Flour Mill in 1785. Dust explosions remain a hazard in coalmines, grain silos and other industrial situations.

23. Very finely divided particles may cause a dust explosion Any solid material that can burn in air will do so at a rate that increases with increased surface area. If the heat produced is sufficiently great, then an explosion will occur.

24. Conditions necessary for a Dust Explosion to occur: Combustible dust particles Enclosed space Source of ignition Dryness Certain concentration Oxygen (2007 – Question 9) *Demo

25. C. Concentration: Changing the concentrations of reactants in a chemical reaction alters the rate of the reaction: - increase in concentration increases rate of reaction. By increasing the concentration of one or more of the reactants, more particles are available for collision, and therefore the reaction rate is faster. - increase in concentration increases amount of product produced

26. Consider the reaction of magnesium metal with a) 1.0M HCl b) 0.5M HCl Mg + 2HCl MgCl 2 + H 2

27. Note: Hydrochloric Acid is in excess i.e. reaction comes to an end when magnesium metal is used up. Magnesium Metal HCl Hydrogen Gas

28. 1.0M HCl 0.5M HCl

29. What does the graph tell us??? 1M HCl - Comes to completion first -Graph is a steep climb upwards – large amount of gas produced in short time - Fast rate of reaction 0.5M HCl: -Takes longer to reach completion - Graph is less steep climb upwards – gas given off more slowly - Slower rate of reaction *In both graphs they level off at the same volume of H 2 produced. This is because the same mass of Mg was used in each experiment and it was the limiting reactant.

30. Mandatory Experiment: To study the effect of concentration on the rate of reaction using sodium thiosulphate and hydrochloric acid.

31. D. Temperature: Changing the temperature in a chemical reaction alters the rate of the reaction: - increase in temperature increases rate of reaction i.e. shorter reaction time

32. Mandatory Experiment: To study the effect of temperature on the rate of reaction using sodium thiosulphate and hydrochloric acid.

33. Overview of effect of Particle Size, Concentration and Temperature on rate of reaction

34. E. Catalysts: A catalyst is a substance that alters the rate of a chemical reaction but is not consumed in the reaction. - Negative catalyst/Inhibitor – slows down the rate of a reaction e.g. Glycerine, Calcium Propionate

35. Catalysts areas of study: - properties of catalysts - types of catalysis - mechanisms of catalysis (Higher Level) - catalytic converters

36. Properties of catalysts: Remain unchanged Specific: – biological catalysts produced by living things are called enzymes and tend to be extremely specific (must know two): - Catalase (liver) breaks down hydrogen peroxide - Amylase (salivary glands) breaks down starch - Pepsin (stomach) breaks down proteins

37. Only needs to be present in small amounts Help reach equilibrium faster – without affecting the position of equilibrium (conc. values) or the composition of the final mixture in an equilibrium reaction. *Equilibrium is a state of dynamic balance where the rate of the forward reaction equals the rate of the reverse reaction e.g. formation of NH 3 Catalytic poisons destroy action of catalysts – lead in petrol destroys the catalysts in the catalytic converters in cars

38. 3 PHASES

39. Types of Catalysis: Homogeneous: both reactants and catalyst are in the same phase... e.g. aqueous potassium iodide catalyses the decomposition of hydrogen peroxide to water and oxygen – IODINE SNAKE *DEMO H 2 O 2 KI Both Liquids

40. Heterogeneous: reactants and catalyst are in different phases... e.g. decomposition of hydrogen peroxide (liquid) by manganese dioxide (solid), catalytic oxidation of methanol (liquid) to methanal by platinum (solid) *DEMO

42. Autocatalysis: where one of the products of the reaction catalyses the reaction... e.g. reaction between permanganate ions and Fe 2+ ions: MnO 4 + 8H + + 5Fe 2+ → Mn 2+ + 5Fe 3+ + 4H 2 O Mn 7+ → Mn 2+ - first few drops of permanganate were decolourised slowly but subsequent drops were decolourised more rapidly - reaction is catalysed by Mn 2+ ions formed in the reaction - Mn 2+ ions formed increases rate of reaction

43. Higher Level Only Mechanisms of Catalysis: Intermediate Formation Theory Surface Adsorption Theory

44. Intermediate Formation Theory Homogenous catalysts sometimes work by reacting with reactants to form unstable intermediate products. These intermediates decompose readily forming products and regenerating the catalyst W + X  Y + ZSlow – no catalyst Addition of catalyst C: W + C  [WC] Intermediate complex formed [WC] + X  Y + Z + C Fast

45. Example of Intermediate Formation Theory The decomposition of hydrogen peroxide in the presence of iodide ions illustrates the formation of an intermediate (Iodine Snake) Rxn. Eqn: 2H 2 O 2  2H 2 O + O 2 I-I-

46. Step 1: one of the H 2 O 2 molecules reacts with an I - ion to form the IO - intermediate H 2 O 2 + I -  H 2 O + IO - (intermediate) Step 2: the intermediate then reacts with another H 2 O 2 molecule to form the products and regenerate the catalyst H 2 O 2 + IO -  H 2 O + O 2 + I -

47. Demonstration: of the oxidation of potassium sodium tartrate by hydrogen peroxide, catalysed by cobalt(II) salts Reactants Before During - intermediate Products After

48. Surface Adsorption Theory Most heterogeneous catalysis can be explained using Surface Adsorption Theory Absorption: moving into Adsorption: moving onto

49. Example of Surface Adsorption Theory The reaction between carbon monoxide and nitrogen monoxide to form nitrogen and carbon dioxide illustrates the surface adsorption theory Carbon monoxide and nitrogen monoxide adsorbed on a platinum surface

50. Step 1: Carbon monoxide and nitrogen monoxide molecules settle on platinum surface (catalyst) - ADSORPTION Increased concentration of these molecules means increased collision probability

51. Step 2: Carbon monoxide and nitrogen monoxide molecules held on surface of platinum by temporary bonds. Carbon monoxide and nitrogen monoxide arranged prior to reacting

52. Step 3: The molecules on the surface of the catalyst then react to form CO 2 and N 2 Carbon dioxide and nitrogen form from previous molecules

53. Step 4: Desorption occurs when the products leave the platinum surface. Carbon dioxide and nitrogen gases escape (desorb) from the platinum surface The oxidation of methanol using hot nichrome (seen previously) is an example of a) heterogeneous catalysis b) surface adsorption mechanism

54. Question: If catalyst was more finely divided what effect would it have and why? Answer: The more finely divided or porous the catalyst, the greater is the reaction rate due to the availability of catalytic sites

55. D. Catalytic Converters: Exhaust fumes (motor vehicles) contain: CO – due to incomplete combustion of hydrocarbons in fuel NO & NO 2 – N 2 in air is oxidised at high temperatures Unburned hydrocarbons Lead compounds

56. Important to decrease emissions as: CO – highly poisonous NO & NO 2 – toxic and contribute to acid rain Unburned hydrocarbons – contribute to photochemical smog

57. Catalytic converters (fitted to exhaust systems of cars) convert CO, NO and NO 2 into harmless gases. They consist of a thin coating of Rh, Pd or Pt on a ceramic or metal honeycomb inside a steel case which has a large surface area.

58. 2CO + 2NO → 2CO 2 + N 2 HC’s + NO + NO 2  CO 2 + N 2 + H 2 O Rh / Pd / Pt The mixture of hot harmful gases (300 o C) is converted into harmless products. The main reactions taking place are: A lot of research being done into developing catalytic converters with low operation temperatures All naturally present in the air

60. Catalytic Poisons When a catalyst is poisoned a lot of the surface of the catalyst is blocked off due to permanent bond formation between the metal and the substance acting as the catalytic poison. Cars fitted with catalytic converters cannot use leaded petrol because the lead atoms bond strongly to the platinum/rhodium surface poisoning it and making it ineffective. Examples of catalytic poisons arsenic, sulphur (Haber process) and lead (Catalytic converters in cars)

61. The Collision Theory (HIGHER LEVEL) This theory was developed from the kinetic theory of gases in order to account for the factors which influence the rate of reactions. Basic Assumptions: - for a reaction to occur the reacting particles must collide with each other - a collision only results in the formation of products if a certain minimum energy is exceeded in the collision. Such a collision is called an EFFECTIVE COLLISION.

62. According to Collision Theory: when an effective collision occurs bonds are broken and new bonds are formed = PRODUCTS Not just important that collisions occur but rate of reaction dependent on number that are effective collisions

63. By examination it has been shown that: - no molecules have zero energy/velocity - some molecules have very low and some molecules have very high energies/velocities - most molecules have average/intermediate velocities

64. INCREASING MOLECULAR (KINETIC) ENERGY NUMBER OF MOLECUES WITH A PARTICULAR ENERGY MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY Distribution of kinetic energies of the molecules of a gas at 15 O C Because of the many collisions taking place between molecules, there is a spread of molecular energies and velocities – GRAPH. This has been demonstrated by experiment.

65. Effect of increasing temperature on reaction rate: If the temperature is increased, the kinetic energies of the gas molecules will increase also due to the heat energy supplied. Thus increasing the temperature alters the distribution: - get a shift to higher energies/velocities - curve gets broader and flatter due to the greater spread of values - area under the curve stays constant - it corresponds to the total number of particles

66. T1T1 T2T2 TEMPERATURE T 2 > T 1 MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY MOLECULAR ENERGY NUMBER OF MOLECUES WITH A PARTICULAR ENERGY Effect of increasing temperature on reaction rate

67. Activation Energy & Collision Theory This distribution of energy among the molecules in a sample of gas explains why it is that not all the collisions between molecules bring about a reaction. A collision between molecules is not enough to ensure that a chemical reaction will take place. In a gas, even a small volume, many millions of collisions take place each second but not all collisions result in a reaction e.g. nitrogen and oxygen. Thus it can be said that molecules must collide with sufficient energy for a chemical reaction to take place. The minimum energy needed is called the activation energy, E act measured in KJ mol -1

68. Definition: Activation Energy: is the minimum energy which colliding particles must have for a reaction to occur. EaEa NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY MOLECULAR ENERGY NUMBER OF MOLECUES WITH A PARTICULAR ENERGY The area under the curve beyond E a corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react.

69. Increase in Temperature and Activation Energy The average kinetic energy of the molecules is proportional to the temperature. The higher the temperature, the greater is the energy of the molecules and the greater their average speed. This has two consequences, each of which increases the rate of the reaction: a) the number of collisions per second is increased b) each collision is more energetic, and a higher proportion of collisions have the necessary activation energy for reaction to occur.

70. Factor b) is much more significant in increasing the reaction rate than factor a), as only those collisions which have sufficiently high energy are effective. Also, a 10K rise in temperature brings about only a small increase in the number of collisions, but a relatively large increase in the number of effective collisions.

71. Increase in Temperature and Activation Energy T1T1 T2T2 TEMPERATURE T 2 > T 1 EaEa MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY NUMBER OF MOLECUES WITH A PARTICULAR ENERGY EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER MOLECULAR ENERGY

72. Reaction Profile Diagrams These diagrams show the activation energy as an energy barrier that must be overcome by the reactants before they form products

73. The size of the activation energy depends on the nature of the reactants. Reactions with SMALL activation energies – faster reaction ENERGY Progress of Reaction EaEa - Small Activation Energy – greater no of reactants have sufficient energy to change into products - Fast Rate of Reaction

74. Reactions with large activation energies – slow or may not occur at all ENERGY PROGRESS OF REACTION EaEa - Large Activation Energy – only a small number of molecules have enough energy to pass over this energy barrier every second and turn into products - Slow Rate of Reaction

75. ENERGY REACTION PATHWAY - Energy of products lower than that of reactants - ΔH is negative Exothermic & Endothermic Reactions

76. REACTION PATHWAY ENERGY - Energy of products is higher than that of reactants - ΔH is positive

77. Nature of Reactants & Activation Energy: - Covalent compounds involve bond breaking during a collision. - This breaking of bonds requires energy - These types of reactions have higher activation energies than reactions involving ionic compounds Particle Size & Activation Energy: - Smaller the particle size the more surface area exposed i.e. the greater the number of collisions that take place - Greater the number of collisions taking place = the number of effective collisions will also increase

78. Concentration & Activation Energy: - Concentration of reactants increased means number of collisions will also increase - Therefore number of effective collisions will be increased giving rise to an increased rate of reaction Catalyst & Activation Energy: - Catalyst works by providing an alternative reaction pathway with a lower activation energy - Therefore more reactant molecules possess the energy required for the effective collisions to occur - Rate of reaction increases

79. Catalyst & Activation Energy

80. According to collision theory, to increase the rate of reaction you therefore need... more frequent collisions increase particle speed have more particles present more successful collisions give particles more energy lower the activation energy