Rates of Reaction

1. Rates of Reaction

2. Objectives • To understand that a chemical reaction involves collisions between particles • To be able to describe the factors which will affect the rate of a chemical reaction.

3. Some Everyday Chemical Reactions • Burning wood • Fruit ripening • Getting a tan • Cooking food

4. How Reactions Happen • A chemical reaction involves a collision between particles. • The particles collide and make new substances • The particles which react are called the reactants • The substances which are made are called the products

5. Reactions Happen at Different Speeds • There are chemical reactions that occur very slowly and others that occur very quickly RUST FORMATION FIREWORKS

6. Rates of Reaction • The rate of a reaction is how quickly a reaction proceeds • It may be defined as the change in concentration in unit time of any one reactant or product Rate = Change in concentration (molL-1) Time taken (s)

7. Calculate the rate of the following H2O2→ H2O + ½ O2 The initial concentration of the reaction is 5 molL-1, ten seconds later this has decreased 3.5 molL-1 What is the rate?

8. Calculate the rate of the following 2SO2 +O2→ 2SO3 The initial concentration of the reaction is 15 molL-1, 20 seconds later this has decreased 12.5 molL-1 What is the rate?

9. Average and Instantaneous Rate • The average rate of reaction is average rate over the course of the reaction • The instantaneous rate of reaction is the rate at a particular point in time during the reaction

10. Controlling the Rates of Reaction • Being able to control the speed (rate) of chemical reactions is important both in everyday life (cooking) and when making new materials on an industrial scale.

11. Factors Affecting Rate of Reaction • Concentration of reactants • Temperature of reaction • Particle size of solid reaction • Nature of reactants • Presence of catalyst

12. 1. Concentration of Reactants • The higher the concentration of reactants the higher the probability of collisions between reactant molecules There are less red particles in the same volume so there is less chance of a collision There are more red particles in the same volume so there is more chance of a collision so the reaction goes faster

13. 2. Temperature of Reaction • An increase in temperature brings about an increase in reaction rate. • You give more energy to the system in the form of heat

14. 3. Particle Size • Particle size (finely divided particles react faster) • Molecules can only collide at the surface. • Smaller particles bigger surface area

15. Dust Explosions • A dust explosion is the explosive combustion of a dust suspended in air in an enclosed location • Any solid material that can burn in air will do so at a rate that increases with increased surface area

16. Grain Dust Peril • In 1998, at a series of explosions occurred at a grain elevator facility in Haysville, Kansas • There were seven fatalities as a result of the explosions. • It is not the actual grain that is ignited, but the fine, thick dust which is released during the loading process when grain particles rub against each other

17. 4. Nature of Reactants • Ionic compounds react faster than covalent • In reactions bonds are broken and form • When an ionic compound is placed in water it dissociates • It takes more energy to break covalent bonds

18. Activation Energy (Eact) • For a reaction to happen, energy is required. • Activation energy is the minimum energy with which particles need to collide to cause a reaction • This is different according to the type of bonds of the reactants.

19. The activation energy may be shown on a reaction profile diagram (right) • These diagrams show energy as a barrier that needs to be overcome by the reactants before they become products

20. The difference between the energy of the reactants and the energy of the products is the heat of reaction (ΔH)

21. Endothermic reaction • A reaction in which heat is taken in. • In an endothermic reaction heat is taken in from the surroundings and the products formed have more energy than the reactants. • It is written as + ΔH

22. Exothermic Reaction • A reaction in which heat is liberated. • In an exothermic reaction heat is lost to the surroundings and the products formed have less energy than the reactants. • It is written as - ΔH

23. 5. Catalysts • A catalyst is a substance that alters the rate of a chemical reaction but is not consumed (used up) in the reaction. • In most cases it makes the reaction go faster. Some catalysts make a reaction go slower and are called negative catalysts or inhibitors. Eg Calcium propionate added to bread to make it stay fresh longer (ie. It slows down staleness)

24. The catalyst does not get used up in the reaction. • It gives the reaction the energy to get started

25. General Properties of catalysts • Catalysts are recovered chemically unchanged at the end of a reaction (Eg. Manganese dioxide used to speed up decomposition of Hydrogen peroxide has exactly the same chemical properties before and after the reaction)

26. 2. Catalysts tend to be specific, even though a catalyst may catalyse one reaction it may not have any effect on a similar reaction Enzymes in the body are examples of catalysts that are very specific Know two examples • Protease breaks down proteins such as blood stains on clothes and are used in washing powders • Catalase breaks down hydrogen peroxide in the body

27. 3. Catalysts need only be present in very small amounts • Increasing the amount of catalyst does not greatly affect the rate of a reaction and in cases where it does it is usually something to do with the reaction itself

28. 4. Catalysts help reactions reach equilibrium quicker but do not change what the equilibrium of a reaction is

29. 5. Action of catalysts may be destroyed by catalytic poisons for example lead in petrol can destroy the catalytic converters in cars Arsenic is a poison that inhibits the action of certain enzymes in the body

30. Types of Catalysis • Chemists have discovered 3 types of catalysis • Homogenous Catalysis • Heterogeneous Catalysis • Autocatalysis

31. Homogenous Catalysis • This describes when reactants and catalysts are in the same phase and there is no boundary between them • Eg Iodine Snake where Potassium Iodide catalyses the decomposition of H2O2 to release oxygen both catalyst + Reactant are liquids

32. Heterogeneous Catalysis • Catalysis where the catalyst and reactants are in different phases Eg. a liquid and a solid • There is a boundary between the catalyst and the reactants • Eg. MnO2catalyses the decomposition of H2O2 to release oxygen both catalyst + Reactant are liquids

33. Autocatalysis • When one of the products of the reaction catalyses the reaction • In this type of reaction it occurs slowly at first but as the reaction proceeds it gets quicker this is because the products drive the reaction forward.

34. Theories of Catalysts • Speed up a reaction by giving the reaction a new path. • The new path has a lower activation energy and more molecules have this energy. • The reaction goes faster.

35. Think of a catalyst as a tunnel through a mountain CATALYST By lowering the activation energy a catalyst makes it possible to carry out a reaction at lower temperatures (lower energy)

36. Mechanisms of Catalysis • The mechanism of catalysis tells us how the catalyst works • There are two main mechanism of catalysis for you to study • Intermediate Formation theory • Surface Adsorption theory

37. Intermediate Formation Theory of Catalysts • Homogeneous catalysts sometimes work by reacting with reactants to form unstable intermediate products • The intermediate exists for a very short time and reacts with the other reactant to give the final product and regenerate the catalyst X+A + B → C AX + B

38. See it with you own eyes • Oxidation of Potassium Sodium Tartrate using hydrogen peroxide • Catalyst in this reaction is Cobalt ions which give a pink colour • The intermediate is a green colour which appears as carbon dioxide + steam are given off • The pink colour is restored at the end indicating the Cobalt ions have not been used up

39. Surface Absorption Theory of Catalysts • Heterogenous catalysis of gas reactions by metals • The reaction happens on the surface of the metal • The reaction occurs at the active site of the catalyst • A catalyst can have multiple active sites

40. Stages of reactions of ethene CATALYST Reactants get absorbed onto catalyst surface. Bonds are weakened CATALYST

41. Bonds Break CATALYST New bonds formed CATALYST

42. Second bond forms and product diffuses away from catalyst surface, leaving it absorb fresh reactants CATALYST Catalytic Poisons • Catalysts can be poisoned they can become less efficient and sometimes they no longer work at all • In heterogeneous catalysis particles that poison the catalyst (lead / arsenic) are absorbed more strongly onto the catalyst surface than the reactant particles • Catalytic poisons block the active sites of enzymes

43. How it works (Catalytic Converter) • The catalyst in the converter speeds up reactions that reduce atmospheric pollution • The catalyst remains unchanged at the end of the reaction • The catalyst is a mix of transition metals (platinum, rhodium, palladium)

44. Reactions Catalysed • Carbon monoxide is converted to carbon dioxide by reaction with oxygen CO + ½ O2→ CO2 • Carbon monoxide can react with nitrogen monoxide to give carbon dioxide 2CO + 2NO → 2CO2 + N2 • Unburnt hydrocarbons are oxidised to carbon dioxide and water C8H18 + 12½ O2 → 8CO2 + 9H2O

45. Environmental Benefits • Reduction in emissions of toxic gases including unburnt hydrocarbons • Reductions in photochemical smog

46. Smog in Beijing

47. Mandatory Experiment 14.1 Monitoring the rate of production of oxygen from hydrogen peroxide, using manganese dioxide as a catalyst

48. Hydrogen peroxide→Oxygen + Water

49. How to layout your results

50. O2 Cm3 60 Time secs 15