Rates of Reaction

1. Rates of Reaction Chapter16

2. C CONCENTRATION A B TIME RATE CHANGE DURING A REACTION Reactions are fastest at the start and get slower as the reactants concentration drops. In a reaction such asA + 2B ——> Cthe concentrations might change as shown • Reactants (A and B) • Concentration decreases with time • Product (C) • Concentration increases with time • the steeper the curve the faster the • rate of the reaction • reactions start off quickly because • of the greater likelihood of collisions • reactions slow down with time as • there are fewer reactants to collide

3. gradient = y x MEASURING THE RATE RATEHow much concentration changes with time. It is the equivalent of velocity. THE SLOPE OF THE GRADIENT OF THE CURVE GETS LESS AS THE REACTION SLOWS DOWN WITH TIME CONCENTRATION y x TIME • the rate of change of concentration is found from the slope (gradient) of the curve • the slope at the start of the reaction will give the INITIAL RATE • the slope gets less (showing the rate is slowing down) as the reaction proceeds

4. Rate of Reaction The rate of a reaction is defined as the change in concentration per unit time in any one reactant or product.

5. Rate of Reaction The rate of a reaction is defined as the change in concentration per unit time in any one reactant or product. The rate of reaction depends on 5 factors. • Nature of reactants • Particle size • Concentration • Temperature • Catalysts.

6. Nature of reactants • The reaction between acidified sodium dichromate and Ammonium Iron (11) sulphate is instantaneous. (Ionic) • The reaction between acidified sodium dichromate and ethanal occurs much more slowly. (Covalent)

7. Particle Size. Large Marble chips Small marble chips CaCO3 +2HCl  CaCl2 +CO2 + H2O Powdered Marble The Rate of the reaction increases as the particle size decreases. The smaller the particle size the greater the surface area. Therefore the greater the number of collisions, the greater the number of successful collisions.

8. INCREASING SURFACE AREA • Increasing surface area increases chances of a collision - more particles are exposed • Powdered solids react quicker than larger lumps • Catalysts (e.g. in catalytic converters) are in a finely divided form for this reason • + • In many organic reactions there are two liquid layers, one aqueous, the other non-aqueous. Shaking the mixture improves the reaction rate as an emulsion is often formed and the area of the boundary layers is increased giving more collisions. 1 1 CUT THE SHAPE INTO SMALLER PIECES 1 1 3 3 SURFACE AREA 9+9+3+3+3+3 = 30 sq units SURFACE AREA 9 x (1+1+1+1+1+1) = 54 sq units

9. Coal Dust Explosion

10. Concentration The greater the concentration the greater the rate of reaction This reaction is studied using the reaction between sodium thiosulfate and Hydrochloric acid. Na2S2O3 + 2HCl -> S + 2NaCl + SO2 + H2O If the concentration of the reactants is increased, the number of collisions will also be increased. If the number of collisions is increased then the number of effective collisions will be increased.

11. INCREASING CONCENTRATION Increasing concentration = more frequent collisions = increased rate of reaction Low concentration = fewer collisions Higher concentration = more collisions However, increasing the concentration of some reactants can have a greater effect than increasing others

12. Temperature The greater the temperature the greater the rate of reaction This reaction is studied using the reaction between sodium thiosulfate and Hydrochloric acid. Na2S2O3 + 2HCl -> S + 2NaCl + SO2 + H2O The rate of a reaction increases as the temperature increases because more of the colliding molecules have the minimum activation energy needed to react.

13. INCREASING TEMPERATURE Effectincreasing the temperature increases the rate of a reaction particles get more energy so they can overcome the energy barrier particle speeds also increase so collisions are more frequent ENERGY CHANGES DURING A REACTION As a reaction takes place the enthalpy of the system rises to a maximum, then falls A minimum amount of energy is required to overcome the ACTIVATION ENERGY (Ea). Only those reactants with energy equal to, or greater than, this value will react. If more energy is given to the reactants then they are more likely to react. Typical energy profile diagram for an exothermic reaction

14. INCREASING TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY NUMBER OF MOLECUES WITH A PARTICULARENERGY MOLECULAR ENERGY Because of the many collisions taking place between molecules, there is a spread of molecular energies and velocities. This has been demonstrated by experiment. It indicated that ... no particles have zero energy/velocity some have very low and some have very high energies/velocities most have intermediate velocities.

15. INCREASING TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T1 NUMBER OF MOLECUES WITH A PARTICULARENERGY T2 TEMPERATURE T2 > T1 MOLECULAR ENERGY • 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

16. INCREASING TEMPERATURE MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY NUMBER OF MOLECUES WITH A PARTICULARENERGY NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER Ea MOLECULAR ENERGY ACTIVATION ENERGY - Ea The Activation Energy is the minimum energy required for a reaction to take place The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react.

17. INCREASING TEMPERATURE TEMPERATURE T2 > T1 MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T1 T2 NUMBER OF MOLECUES WITH A PARTICULARENERGY EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER Ea MOLECULAR ENERGY Explanation increasing the temperature gives more particles an energy greater than Ea more reactants are able to overcome the energy barrier and form products a small rise in temperature can lead to a large increase in rate

18. Catalysts

19. Types of Catalysis Homogeneous catalysis. This is catalysis in which both the reactants and the catalyst are in the same phase. (Iodine snake experiment) Heterogeneous catalysis. This is catalysis in which the reactants and catalyst are in different phases. (Hydrogen peroxide (liquid) and Manganese dioxide (solid)) Auto catalysis.One of the products in the reaction catalyses the reaction. (Permanganate ions and Fe2+ ions.)

20. Mechanisms of catalysis Intermediate compound theory. A + B  AB SLOW A + C  AC FAST AC + B AB + C FAST The decomposition of hydrogen peroxide catalysed by the presence of I ions (iodine snake reaction) illustrates the formation of an intermediate. Overall Reaction 2H2O2 2H2O + O2 Step 1 H2O2 + I- H2O + IO- Step 2 H2O2 + IO- H2O + O2 +I- Also Sodium Hydrogen tartrate + Hydrogen peroxide + Co2+ ions (pink) Pink Blue/Green Pink Cobalt Intermediate Cobalt

21. Surface Adsorption Theory Methanol Methanal CH3OH HCHO Platinum 2H2 + O2  2H2O A good example is the reaction of Hydrogen and Oxygen to form water using finely divided Platinum as the catalyst. The Hydrogen and oxygen molecules settle on the surface of the catalyst. The adsorbed atoms form weak bonds with the metal atoms. Transition metals can act as catalysts because they have vacant d orbitals. The hydrogen and oxygen molecules then react to form water. The products leave the surface of the catalyst. (desorption)

22. Catalytic converters. Exhaust fumes contain carbon monoxide (CO), nitrogen monoxide (NO), nitrogen dioxide (NO2) and unburnt hydrocarbons. a catalytic converter converts these gases to environmentally friendly gases. The catalytic converter consists of a thin coating of platinum, palladium, and rhodium on a ceramic or metal honeycomb inside a stainless steel case. Pt/Pd/Rh (Temp =300 C) 2CO + 2NO  2CO2 + N2 The un-burnt hydrocarbons react with oxides of nitrogen to form carbon dioxide nitrogen and water.

23. Catalytic converter

25. Collision Theory For a reaction to occur, the reacting particles must collide with each other. For the formation of product a certain minimum energy is required in the collision. Such a collision is called an effective collision.

26. COLLISION THEORY • Collision theory states that... • particles must COLLIDE before a reaction can take place • not all collisions lead to a reaction • reactants must possess at least a minimum amount of energy - ACTIVATION ENERGY • plus • particles must approach each other in a certain relative way - the STERIC EFFECT • According to collision theory, to increase the rate of reaction you therefore need... • more frequent collisions increase particle speed or • have more particles present • more successful collisions give particles more energy or • lower the activation energy

27. The Activation energy • The Activation energy is the minimum energy which colliding particles must have for a reaction to occur.

28. Energy profile diagram • A catalyst works by reducing the activation energy.

29. ADDING A CATALYST • Catalysts provide an alternative reaction pathway with a lower Activation Energy (Ea) • Decreasing the Activation Energy means that more particles will have sufficient • energy to overcome the energy barrier and react • Catalysts remain chemically unchanged at the end of the reaction. WITHOUT A CATALYST WITH A CATALYST

30. ADDING A CATALYST • Catalysts provide an alternative reaction pathway with a lower Activation Energy (Ea) • Decreasing the Activation Energy means that more particles will have sufficient • energy to overcome the energy barrier and react • Catalysts remain chemically unchanged at the end of the reaction. WITHOUT A CATALYST WITH A CATALYST

31. ADDING A CATALYST MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY NUMBER OF MOLECUES WITH A PARTICULARENERGY NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER Ea MOLECULAR ENERGY The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react. If a catalyst is added, the Activation Energy is lowered - Ea will move to the left.

32. ADDING A CATALYST MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY NUMBER OF MOLECUES WITH A PARTICULARENERGY EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER Ea MOLECULAR ENERGY The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react. Lowering the Activation Energy, Ea, results in a greater area under the curveafterEashowing that more molecules have energies in excess of the Activation Energy

33. CATALYSTS - A REVIEW • work by providing an alternative reaction pathway with a lower Activation Energy • using catalysts avoids the need to supply extra heat - safer and cheaper • catalysts remain chemically unchanged at the end of the reaction. • TypesHomogeneous CatalystsHeterogeneous Catalysts • same phase as reactants different phase to reactants • e.g. CFC’s and ozone e.g. Fe in Haber process • Usesused in industry especially where an increase in temperature results in • a lower yield due to a shift in equilibrium (Haber and Contact Processes) • CATALYSTS DO NOT AFFECT THE POSITION OF ANY EQUILIBRIUM • but they do affect the rate at which equilibrium is attained • a lot is spent on research into more effective catalysts - the savings can be dramatic • catalysts need to be changed regularly as they get ‘poisoned’ by other chemicals • catalysts are used in a finely divided state to increase the surface area

34. Monitoring the rate of production of oxygen from hydrogen peroxide using manganese dioxide as a catalyst Hydrogen peroxide decomposes into water and oxygen as follows: H2O2(l) → H2O(l) + 1/2 O2(g) This occurs much too slowly to be monitored. However, manganese dioxide acts as a suitable catalyst, and the reaction occurs at a measurable rate.

35. Method 2

38. Questions 1. Why is the slope of the graph steepest in the early stages of the reaction? Since rate is proportional to concentration, the greatest rate, indicated by the steepest slope, is evident in the early stages when the concentration of hydrogen peroxide is at a maximum. 2. At what stage is the reaction complete? When the graph becomes horizontal. 3. What would be the effect on the graph of doubling the amount of manganese(IV) oxide? The increased surface area of catalyst would speed up the reaction, giving a steeper slope and an earlier completion. The volume of oxygen produced would be unchanged.

39. 4. Would doubling the manganese(IV) oxide create a practical difficulty? Explain your answer. Yes. The production of oxygen could become too quick for accurate monitoring. 5. What would be the effect on the graph of doubling the concentration of hydrogen peroxide? Increasing the concentration of a reactant would speed up the rate, as indicated by a steeper slope. Doubling the concentration would produce double the final volume of oxygen. 6. Would doubling the concentration of hydrogen peroxide create a practical difficulty? Explain your answer. Yes. The capacity of the collection vessel could be exceeded.

40. Studying the effects on reaction rate of (i) concentration and (ii) temperature The reaction used is that between a sodium thiosulfate solution and hydrochloric acid: 2HCl(aq) + Na2S2O3(aq) 2NaCl(aq) + SO2(aq) + S(s)↓ + H2O(l) The precipitate of sulfur formed gradually obscures a cross marked on paper and placed beneath the reaction flask. The rate of reaction, and consequently the time taken to obscure the cross, depends on a number of variables such as temperature, concentration and volume. By varying one of these and keeping the others constant, the effect on rate can be studied. The inverse of the time taken to obscure the cross is the measure of reaction rate used in this experiment.

41. Effect of concentration 1. Place 100 cm3 of the sodium thiosulfate solution into a conical flask. 2. Add 10 cm3 of 3 M hydrochloric acid to the flask, while starting the stop clock at the same time. 3. Swirl the flask and place it on a piece of white paper marked with a cross. 4. Record the time taken for the cross to disappear. 5. Repeat the experiment using 80, 60, 40 and 20 cm3 of. sodium thiosulfate solution respectively. In each case, add water to make the volume up to 100 cm3 and mix before adding HCl. 6. If the initial sodium thiosulfate concentration is 0.1 M, subsequent concentrations will be 0.08 M, 0.06 M, 0.04 M and 0.02 M respectively.

42. 7.Record the results in a table similar to the following: Concentration of thiosulfate 0.l M 0.08 M 0.06 M 0.04 M 0.02 M Reaction time (s) 1/time 8. Draw a graph of 1/time against concentration. This is effectively a graph of reaction rate against concentration.

44. Effect of temperature Procedure NB: Wear your safety glasses. 1 . Place 100 cm3 of 0.05 M sodium thiosulfate solution into a conical flask. 2. Warm the flask gently until the temperature is about 20 0C. 3. Add 5 cm3 of 3 M HCl, starting a stop clock at the same time, before proceeding. 4. Without delay, swirl the flask, place it on a piece of white paper marked with a cross, and record the exact temperature of the contents of the flask. 5. Record the time taken for the cross to disappear 6. Repeat the experiment, heating the thiosulfate to temperatures of approximately 30 0C, 40 0C, 50 0C and 60 0C respectively (before adding the HCl). 7. Record the results in a table similar to the following

46. Questions Suggested Answers to Student Questions 1. What is the effect of increasing the concentration on the reaction time? The reaction time is decreased. 2. What is the effect of increasing the concentration on the reaction rate? The rate is increased. 3. What is meant by saying that two quantities are directly proportional? If one of the quantities is increased/decreased by a certain factor, the other changes in exactly the same way. 4. What is the effect of raising the temperature on the reaction time? The reaction time is decreased. 5. What is the effect of raising the temperature on the reaction rate? Suggest two factors responsible for the result observed. The rate is increased. The higher temperature results in greater kinetic energy of the particles present. This causes: (i) more collisions per unit time, and (ii) a greater proportion of the collisions to have the activation energy needed for products to form. Both (i) and (ii) result in a rate increase.

47. 6. Suggest a reason why it is not recommended to carry out the experiment at temperatures higher than about 60 0C. The reaction occurs so quickly that it is not possible to measure the time accurately. 7. Which is the limiting reactant in the temperature experiment? 100 cm3 of 0.05 M Na2S2O3 contains: 100/1000 x 0.05 = 0.005 moles Na2S2O3 5 cm3 of 3 M HCl contains: 5/1000 x 3 = 0.015 moles HCl According to the balanced equation, the reacting ratio is Na2S2O3 : HCl = 1:2 The amounts used are in the ratio Na2S2O3 : HCl = 0.005: 0.015 = 1 : 3 Clearly Na2S2O3 is the limiting reactant.