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    1. 1 rate of reaction 1.3 The Collision Theory

    2. 1.3 The Collision Theory 1 According to the kinetic theory of matter, all matter is made up of tiny, discrete particles. These particles are continually moving and so have kinetic energy.

    3. 1.3 The Collision Theory 2 Based on the assumption that the particles in matter are moving all the time and collide with each other, the collision theory was introduced to explain (a) how chemical reactions occur, and (b) the factors affecting the rates of reactions.

    4. 1.3 The Collision Theory 3. When the reactants are mixed, the reactant particles (atoms, molecules or ions) collide with each other. However, not all collisions will result in a chemical reaction to form the products of the reaction. It is likely that the particles collide but are bounced back without producing any changes. The collisions that are successful in producing a chemical reaction are called the effective collisions.

    5. 1.3 The Collision Theory 4 Collisions of particles that are unsuccessful in producing a chemical reaction are called ineffective collisions.

    6. 1.3 The Collision Theory 5 The collision theory states that for a chemical reaction to occur, the reacting particles must: (a) collide with each other so that the breaking and formation of chemical bonds can occur.

    7. 1.3 The Collision Theory 5 The collision theory states that for a chemical reaction to occur, the reacting particles must: (a) collide with each other so that the breaking and formation of chemical bonds can occur. (b) possess energy that is equal to or more than the minimum energy called the activation energy.

    8. 1.3 The Collision Theory 5 The collision theory states that for a chemical reaction to occur, the reacting particles must: (a) collide with each other so that the breaking and formation of chemical bonds can occur. (b) possess energy that is equal to or more than the minimum energy called the activation energy. (c) collide in the correct orientation.

    9. Activation Energy 1 The activation energy is the minimum energy that the reactant particles must possess at the time of collision in order for a chemical reaction to take place.

    10. Activation Energy 2 The activation energy can also be considered as an energy barrier that must he overcome by the colliding particles in order that collision will result in the formation of product molecules.

    11. Activation Energy 3 (a) If the two molecules with sufficient energy (that is, energy equal to or more than the activation energy) collide in the correct orientation, the chemical bonds in the reactant molecules will break and reaction will occur to form new bonds in the product molecules (Figure 1.35). For example,

    12. Activation Energy 3 (b) However, if two molecules, with energy equal to or more than the activation energy, do not collide with each other in the correct orientation, then reaction will not occur.

    13. Activation Energy 3 (c) If two reactant molecules, with energy less than the activation energy, collide in the correct orientation, then reaction will also not occur. The colliding molecules will simply rebound and move away from each other.

    14. Relating the frequency of effective collisions with factors influencing the rate of reaction According to the collision theory, two important factors that determine the rate of a chemical reaction are: (a) the frequency of effective collisions and (b) the magnitude of the activation energy

    15. Relating the frequency of effective collisions with factors influencing the rate of reaction 1 Frequency of effective collisions For a given reaction, if the frequency of collisions between reactant molecules is high, it follows that the frequency of effective collisions that causes a reaction to occur will also be high. As a result, the rate of reaction increases.

    16. Relating the frequency of effective collisions with factors influencing the rate of reaction 2 Magnitude of activation energy Reactions that have high activation energy will occur at a slow rate. This is because only a small fraction of the molecules possess sufficient energy to overcome the activation energy for the reaction to occur. In contrast, reactions that possess low activation energy will occur at a fast rate. This is because most of the molecules have sufficient energy to overcome the activation energy and this enables the reaction to occur.

    17. Relating the frequency of effective collisions with factors influencing the rate of reaction 3 In general, any factor that increases the rate of effective collisions will also increase the rate of reaction.

    18. Effect of surface area on the reaction rate 1 The sodium chloride crystal as shown in Figure 1.29 (a) has a surface area of 16 cm2. When this crystal is divided into smaller crystals as shown in Figure 1.29 (b), the total surface area is increased to 24 cm2. Surface area of the NaCl crystal =(1 x 2) x 4 + (2 x 2) x 2 =16cm2

    19. Effect of surface area on the reaction rate Total surface area

    20. Effect of surface area on the reaction rate 2. The smaller the particle size, the greater the total surface area exposed for reaction to occur.

    21. Effect of surface area on the reaction rate 3 In the reaction between hydrochloric acid and marble chips (calcium carbonate), CaCO3(s) + 2HCl(aq) ? CaCl2(aq) + H2O(l) + CO2(g) the rate of reaction will increase if powdered marble is used instead of marble chips.

    22. Effect of surface area on the reaction rate

    23. Effect of surface area on the reaction rate 4. The total surface area is larger in powdered marble. With greater surface area exposed for reaction with the acid (Figure 1.30), collisions between reacting molecules become more frequent.

    24. Effect of surface area on the reaction rate 5. This means that the frequency of effective collisions (that is, collisions with the correct orientation and with energy equal to or greater than activation energy) between reacting particles will increase. As a result, the rate of reaction also increases.

    25. Effect of concentration on the rate of reaction 1 Magnesium reacts with dilute hydrochloric acid according to the equation: Mg(s) + 2HCl(aq) ? MgCl2(aq) + H2(g) When the concentration of hydrochloric acid increases, the rate of reaction also increases.

    26. Effect of concentration on the rate of reaction 2 Figure 1.37 shows the arrangement of particles in 1 mol dm-3 hydrochloric acid and 2 mol dm-3 (more concentrated) hydrochloric acid.

    27. Effect of concentration on the rate of reaction When the concentration of hydrochloric acid increases, the number of particles per unit volume also increases and the particles are closer together.

    28. Effect of concentration on the rate of reaction 3 When the number of particles increases, the frequency of collisions also increases. As a result, the frequency of effective collisions increases. This causes the rate of reaction to increase.

    29. Effect of pressure on the rate of reaction 1 In chemical reactions involving gases, increasing the pressure increases the rate of reaction. Conversely, decreasing the pressure decreases the rate of reaction. For example, the rate of reaction between nitrogen and oxygen to produce nitrogen monoxide can be increased by increasing the pressure. N2(g) + O2(g) ? 2NO(g)

    30. Effect of pressure on the rate of reaction 2 At low pressures, the gaseous molecules are spread far apart (Figure 1.32 (a)). When the pressure is increased, the volume of the gas is reduced (Figure 1.32(b)).

    31. Effect of pressure on the rate of reaction This means that at high pressures, (a) the number of gaseous molecules per unit volume is increased, and (b) the gaseous molecules are packed closer together.

    32. Effect of pressure on the rate of reaction 3 As a result, increasing the pressure causes the gaseous molecules to collide more frequently. Consequently, the frequency of effective collisions increases and the rate of reaction also increases.

    33. Effect of temperature on the rate of reaction 1 Calcium carbonate reacts with hydrochloric acid to form carbon dioxide according to the following equation: CaCO3(s) + 2HCl(aq) ? CaCl2(aq) + H2O(l) + CO2(g)

    34. Effect of temperature on the rate of reaction Figure 1.33 shows the graphs of total volume of carbon dioxide given off against the time taken for the reaction between calcium carbonate and dilute hydrochloric acid at 25 C and 30 C.

    35. Effect of temperature on the rate of reaction 2. At low temperatures, particles of reactants move at a slower speed. However, when the temperature is increased, the particles absorb the heat energy. As a result, the kinetic energy of the particles increases. Hence, (a) the reacting particles move faster, and (b) the number of reacting particles with the activation energy required increases.

    36. Effect of temperature on the rate of reaction 3 Consequently, the frequency of effective collisions increases and hence, the rate of reaction also increases.

    37. Effect of temperature on the rate of reaction 4 Temperature has a great effect on the rate of reaction. For most reactions, the rate of reaction approximately doubles when the temperature of reaction increases by 10 C.

    38. Effect of catalysts on reaction rates 1 The decomposition of hydrogen peroxide to water and oxygen occurs very slowly at room temperature. 2H2O2(aq) ? 2H2O(l) +O2(g). In the presence of a catalyst, the decomposition of hydrogen peroxide occurs rapidly.

    39. Effect of catalysts on reaction rates 2 Figure 1.34 shows the rate of evolution of oxygen for the decomposition of hydrogen peroxide without a catalyst and in the presence of a catalyst such as manganese(IV) oxide or iron(III) oxide.

    40. Effect of catalysts on reaction rates 3. A catalyst provides an alternative reaction route (or pathway) for the reaction to occur. In the presence of a positive catalyst, an alternative route has a lower activation energy.

    41. Effect of catalysts on reaction rates In other words, a positive catalyst lowers the activation energy required for the reaction (Figure 1.34). As a result, more reacting particles possess sufficient energy to overcome the lower activation energy required for effective collisions. Hence, the frequency of effective collisions increases and the rate of reaction increases