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1 of 35© Boardworks Ltd 2009. 2 of 35© Boardworks Ltd 2009.

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Presentation on theme: "1 of 35© Boardworks Ltd 2009. 2 of 35© Boardworks Ltd 2009."— Presentation transcript:

1 1 of 35© Boardworks Ltd 2009

2 2 of 35© Boardworks Ltd 2009

3 3 of 35© Boardworks Ltd 2009 What is collision theory? Collision theory states that for a reaction to occur: particles must collide with the correct orientation. particles must have sufficient energy particles must collide Most collisions do not result in reaction because they do not meet the second and third criteria.

4 4 of 35© Boardworks Ltd 2009 Activation energy

5 5 of 35© Boardworks Ltd 2009 Effect of surface area on collisions Only the particles on the surface of a solid will collide with particles of the other reactant. Surface area can be increased by decreasing the size of the reactant particles. Powders have a very large surface area. increasing surface area If the surface area is increased, more particles will be on the surface and able to collide with particles of the other reactant. This means that there will be more collisions in total and therefore more reactive collisions.

6 6 of 35© Boardworks Ltd 2009 Reaction of marble chips with acid The effect of changing surface area on the rate of reaction can be explored by reacting marble (calcium carbonate) chips and an acid such as 2 mol dm -3 hydrochloric acid. CaCO 3(s) + 2HCl (aq)  CaCl 2(s) + CO 2(g) + H 2 O (l) The carbon dioxide gas evolved can be collected and its volume measured over time. The rate at which it is produced is a measure of the rate of reaction. By repeating the experiment with marble chips of different sizes, the effect of surface area on the rate of reaction can be examined.

7 7 of 35© Boardworks Ltd 2009 Effect of surface area on rate: graph

8 8 of 35© Boardworks Ltd 2009 Effect of concentration on rate

9 9 of 35© Boardworks Ltd 2009 Effect of concentration on rate: graph

10 10 of 35© Boardworks Ltd 2009 Effect of gas pressure on rate

11 11 of 35© Boardworks Ltd 2009 Effect of pressure on rate: graph

12 12 of 35© Boardworks Ltd 2009 Effect of temperature on particles

13 13 of 35© Boardworks Ltd 2009 Effect of temperature on rate Increasing the temperature of the reaction mixture increases the rate of reaction in the following two ways: 1. At higher temperatures, the particles are moving faster, so collide more frequently. A higher number of collisions in total means a higher number of successful collisions. 2. At higher temperatures, a higher proportion of the molecules have the activation energy or more. This means that a higher proportion of collisions is successful.

14 14 of 35© Boardworks Ltd 2009 Effect of temperature on rate: graph

15 15 of 35© Boardworks Ltd 2009 Collision theory summary

16 16 of 35© Boardworks Ltd 2009 Rate of reaction summary

17 17 of 35© Boardworks Ltd 2009

18 18 of 35© Boardworks Ltd 2009 Distribution of particle speeds

19 19 of 35© Boardworks Ltd 2009 Energy distribution curves

20 20 of 35© Boardworks Ltd 2009 The effect of changing temperature

21 21 of 35© Boardworks Ltd 2009 Small temperature changes The Maxwell–Boltzmann distribution shows that for a small increase in temperature, there is a relatively large increase in the number of particles with at least the activation energy. A small increase in temperature therefore leads to a large increase in rate. The increase in collision frequency is also a factor, but its effect is small compared to the increase in energy. no. of particles energy EaEa no. particles with E a almost doubled

22 22 of 35© Boardworks Ltd 2009 Effect of temperature summary

23 23 of 35© Boardworks Ltd 2009

24 24 of 35© Boardworks Ltd 2009 Effect of catalysts on rate: graph

25 25 of 35© Boardworks Ltd 2009 What do catalysts do?

26 26 of 35© Boardworks Ltd 2009 How do catalysts work? An example is the oxidation of sulfur dioxide: SO 2(g) + V 2 O 5(s)  SO 3(g) + V 2 O 4(s) The catalyst is re-formed by reacting with oxygen: V 2 O 4(s) + ½O 2(g)  V 2 O 5(s) SO 2(g) + ½O 2(g)  SO 3(g) This is catalyzed by vanadium(V) oxide: Catalysts increase the rate of reactions without being used up during the reaction. One way in which this occurs is for the catalyst to be changed during the reaction, then changed back in a second reaction with one of the reactants or products. This is an alternative reaction pathway.

27 27 of 35© Boardworks Ltd 2009 Catalysts and energy distribution curves

28 28 of 35© Boardworks Ltd 2009 Heterogeneous catalysts There are two types of catalysts: heterogeneous and homogeneous. Heterogeneous catalysts are in a different phase to the reactants. The catalyst is usually a solid and the reactants are liquids or gases (e.g. solid catalysts for gas reactions in catalytic converters). Industrial examples of heterogeneous catalysis include the iron catalyst used in ammonia production and the Ziegler–Natta catalyst used in poly(e)thene production.

29 29 of 35© Boardworks Ltd 2009 Homogeneous catalysts Homogeneous catalysts are in the same phase as the reactants. The catalyst and the reactants are usually liquids, such as the hardener added to fibreglass resin. Another example of homogeneous catalysis is the destruction of atmospheric ozone catalyzed by chlorine free radicals. In this reaction the catalyst and reactants are in the gas phase.

30 30 of 35© Boardworks Ltd 2009 Advantages of catalysts Using a catalyst means that a reaction can take place at the same rate as the uncatalyzed reaction, but at a lower temperature and/or pressure. This has the following advantages, which are particularly important in industry: A non-industrial example is enzyme catalysis in biological washing powders, allowing efficient washing at a lower temperature. lower energy demands… …therefore less CO 2 produced… …therefore less environmental impact… …and lower production costs.

31 31 of 35© Boardworks Ltd 2009 Catalysts: true or false?

32 32 of 35© Boardworks Ltd 2009

33 33 of 35© Boardworks Ltd 2009 Glossary

34 34 of 35© Boardworks Ltd 2009 What’s the keyword?

35 35 of 35© Boardworks Ltd 2009 Multiple-choice quiz


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