Chpt.16: Rates of Reactions

1. Chpt.16: Rates of Reactions

2. Chemical reactions occur at different speeds: Lighted taper put in test tube of H2 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 (cm3) • 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: • Draw a graph of volume Vs. Time • How long did it take for 55cm3 of oxygen to be produced • Calculate the rate of reaction after 90 seconds (i.e. Instantaneous rate)

12. Factors Affecting Rates of Reaction • Rate of reaction depends on five factors: • Nature of Reactants • Particle Size • Concentration • Temperature • 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 = 4cm2 The cube has six faces, so the total surface area is 4cm2 x 6 = 24cm2 We could cut that cube horizontally and vertically along each face so that we have eight smaller cubes. Each of the small cubes has a face area of 1cm x 1cm = 1cm2 The six faces give a total surface area for each smaller cube of 6cm2 There are eight cubes so the total surface area is 6cm2 x 8 = 48cm2

17. Reaction of calcium carbonate (marble chips) with dilute hydrochloric acid solution: CaCO3(s) + 2HCl(aq) CaCl2 + H2O(l) + CO2(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 MgCl2 + H2

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

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 H2 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 NH3 • 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 H2O2 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 Fe2+ ions: MnO4 + 8H+ + 5Fe2+ → Mn2+ + 5Fe3+ + 4H2O Mn7+ → Mn2+ - first few drops of permanganate were decolourised slowly but subsequent drops were decolourised more rapidly - reaction is catalysed by Mn2+ ions formed in the reaction - Mn2+ 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 + Z Slow – 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: 2H2O2 2H2O + O2 I-

46. Step 1: one of the H2O2 molecules reacts with an I- ion to form the IO- intermediate H2O2 + I-  H2O + IO- (intermediate) Step 2: the intermediate then reacts with another H2O2 molecule to form the products and regenerate the catalyst H2O2 + IO-  H2O + O2 + I-

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

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