1. 6.1 Rates of Reaction

2. Mrs Smart likes a cup of tea at break time but she doesn’t have long to make and drink it. She likes her tea with 2 sugars in. Which would you recommend that she add to her tea ? sugar cubes or granulated sugar?

3. Objectives 6.1.1 Define the term RATE OF REACTION 6.1.2 Describe suitable experimental procedures for measuring rates of reaction. 6.1.3 Analyse data from rate experiments

4. Put these chemical reactions in order of how fast they happen from fastest to slowest: An egg frying A firework exploding A car rusting Coal forming Potassium reacting with water

5. Chemical Kinetics The study of rates of chemical reactions and the mechanisms (or steps) by which a chemical reaction takes place. Reaction rates vary greatly – some are very fast (burning) and some are very slow (disintegration of a plastic bottle in sunlight).

6. Rate of Reaction The rate of a reaction is the change in concentration of a product per unit of time (rate of formation of product). Also viewed as the negative of the change in concentration of a reactant per unit of time (- rate of disappearance of reactant). Rxn Rate (avg) = Δ [reactant or product] Δ time [square brackets] = mol/L

7. Reaction Rates Rates of reactions can be determined by monitoring the change in concentration of either reactants or products as a function of time.  [A] vs  t 7

8. Rate of Reaction A measure of the speed at which the products are formed. Measured as a change in concentration divided by the change in time. R  P Rate = ∆[P]/∆t = - ∆[R]/∆t Rate is always positive. The minus sign is shown here to realise that the reactants are decreasing.

9. Important to Remember Measuring any of the products or reactants will give you the same shape curve. The actual concentrations produced will be different based on the reactions stoichiometry.

10. Measuring Reaction Rates Decomposition of H 2 O 2 Time, s Accumulated mass O 2, g[H 2 O 2 ], M 0 0 0.882 60 2.9600.697 1205.0560.566 1806.7840.458 2408.1600.372 3009.3440.298 36010.3360.236 42011.1040.188 48011.6800.152 54012.1920.120 60012.6080.094

11. Kinetic data for the reaction: H 2 O 2 (aq)  H 2 O (l) + ½ O 2 (g) Rates are obtained from the slopes of the straight lines: an average rate from the purple dotted line the instantaneous rate at t = 300 s from the red line the initial rate from the blue line. Some Kinetic Data

12. Understanding the graph Concentration of reactant decreases with time, as it is being used up The rate is fastest when concentration of reactants is greatest; and slows when concentration of reactants are less. Decreasing numbers of reactants means less collisions producing new product.

13. Reaction Rates In this reaction, the concentration of butyl chloride, C 4 H 9 Cl, was measured at various times, t. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) [C 4 H 9 Cl] M 13

14. Reaction Rates The average rate of the reaction over each interval is the change in concentration divided by the change in time: C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) Average Rate, M/s 14

15. Reaction Rates Note that the average rate decreases as the reaction proceeds. This is because as the reaction goes forward, there are fewer collisions between reactant molecules. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) 15

16. Reaction Rates A plot of concentration vs. time for this reaction yields a curve like this. The slope of a line tangent to the curve at any point is the instantaneous rate at that time. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) 16

17. Reaction Rates The reaction slows down with time because the concentration of the reactants decreases. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) 17

18. Reaction Rates and Stoichiometry In this reaction, the ratio of C 4 H 9 Cl to C 4 H 9 OH is 1:1. Thus, the rate of disappearance of C 4 H 9 Cl is the same as the rate of appearance of C 4 H 9 OH. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) Rate = -  [C 4 H 9 Cl]  t =  [C 4 H 9 OH]  t 18

19. Reaction Rates and Stoichiometry What if the ratio is not 1:1? H 2 (g) + I 2 (g)  2 HI (g) Only 1/2 HI is made for each H 2 used. 19

20. Reaction Rates and Stoichiometry To generalize, for the reaction aA + bBcC + dD Reactants (decrease) Products (increase) 20

21. Measuring rate of reactions Any property that changes between the start and end of the reaction can in principle be used. The larger the changes the better the accuracy of its measurement. Any ideas of what we could measure?

22. Measuring Rate of Reaction Find a property that we can measure that changes over time. What things in a chemical reaction that can change over time?????

23. Ways of Measuring Rate of Reaction Temperature Thermometer Colour Spectrophotometer Mass Balance Electrical Conductivity Resistance Acidity pH meter Light Absorption Eye, Colorimeter Gas Measuring cylinder, Gas syringe

24. Taking Initial Rate Measuring initial rate from the graph.

25. Notes to Remember Property verses time graphs is usual. If wanting units in mol dm -3 s -1 need to produce graphs of concentration vs time. Important to keep the temperature constant in a water bath (as we know that temperature affects rate). Common practice to place reactants in a water bath before mixing.

26. Measuring reaction rates Rate implies we are measuring how things change over a period of time. To measure the rate of a reaction we have to track the manner in which the amount of product (or reactant) changes over time. Rate of gas formation can be measured using a syringe. For a reaction in which sulfur is precipitated we can time how long the solution takes to go cloudy.

27. Collecting an evolved gas The gas produced is collected in a syringe or in a graduated vessel over water (gas can’t be water soluble) The volume can be collected a different times and recorded

28. Slower and slower Reactions do not proceed at a steady rate. They start fast and get slower and slower. This is not surprising because the reactant concentration (and the chance of collision) gets lower and lower as time progresses. Percentage completion of reaction 0 fast 100 stopped 25 slower 75 very slow reactants product

29. Rates and Graphs These show the increasing amount of product or the decreasing amount of reactant. Amount of product Time Amount of reactant Time Steep gradient Fast reaction Shallow gradient Slow reaction Steep gradient Fast reaction Shallow gradient Slow reaction

30. Rate graphs and reactant concentrations Amount of product Time reactants product Reactant Concentration falls Rate of Reaction falls All product All reactant Mix of reactant And product Gradient of graph decreases

33. The following slides describe the four chemical reactions that are commonly used as examples.

34. Acid and marble Marble chips are calcium carbonate. They react with acid to evolve a gas. calcium carbonate + hydrochloric acid  calcium chloride + water + carbon dioxide CaCO 3 (s) + 2HCl(aq)  CaCl 2 (aq) + H 2 O(l) + CO 2 (g) The gas given off can be collected in a syringe and readings taken every 30 seconds or so.

35. Acid and marble 1. Measure the agreed mass of marble chips 2. Set up the syringe, flask and connector 3. Measure the acid / water. 4. Add the marble chips and quickly insert the bung and start stop clock. 5. Take syringe readings at 30 second intervals. TimeReading 0 s0 cm 3

36. Acid and marble

37. Mg HCl Acid and metal Reactive metals (eg. Magnesium) react with acid to evolve hydrogen gas. magnesium + hydrochloric acid  magnesium chloride + hydrogen As the gas given off leaves the flask the total mass of the flask and its contents decreases slightly. Readings of the mass(g) can be taken. Typically at 1 minute intervals. Mg(s) + 2HCl(aq)  MgCl 2 (aq) + H 2 (g) 11.80 11.77 11.74 11.73 11.72 11.71

38. Mg HCl Cotton wool Acid and metal 1. Measure the agreed volume of acid / water into the conical flask. 2. Have a loose plug of cotton wool to prevent “spitting” of droplets of liquid. 3. Have a piece of magnesium of known mass ready. 4. Add the magnesium, place the cotton wool in the neck and start taking mass readings immediately. TimeReading 0 s0 cm 3 60 120 11.80 11.77 11.74 11.73 11.72 11.71

39. Acid and metal

40. Decomposition of Hydrogen Peroxide Hydrogen peroxide decomposes into water and oxygen. Hydrogen peroxide  water + oxygen Oxygen gas is given off and can be measured using a gas syringe or a balance. The reaction is catalysed by a wide range of solids. 2H 2 O 2 (aq)  2H 2 O(l) + O 2 (g) Remember the catalyst NEVER produces more product - just quicker

42. Acid and Sodium Thiosulphate In this reaction sulphur is precipitated which makes the solution turn cloudy. The effect of changing conditions such as temperature or concentrations can be studied by measuring how long it takes to produce enough sulphur to make the solution opaque (non see-through). Na 2 S 2 O 3 (aq) + 2HCl(aq)  2NaCl(aq) + H 2 O(l) + SO 2 (g) + S(s) Sodium thiosulphate + hydrochloric acid  sodium chloride + water + sulphur dioxide + sulphur

43. 1. Measure the agreed volume of thiosulphate / water into the conical flask. 2. Prepare a piece of paper with a cross drawn on it. 3. Measure the required volume of acid in a measuring cylinder. 4. Add the acid to the flask, start the clock, swirl the flask. 5. Look down through the flask until the cross disappears. Note the time. Look down here

44. Temp ( o C) Time taken (s) 25100 3060 3540 25 4515 5010 Imran studied the effect of temperature upon the time it took for the flask to go cloudy. 1.Sketch a graph of the results. 2.Using the reactions at 25 o C and 40 o C, explain how the time taken lets you work out the relative rate of reaction. Time(s) vs Temp( o C) Time Temp

47. A pupil performed an investigation into the rate of reaction between a metal and an acid. The results below where obtained. i) Plot a graph of gas volume (y-axis) against time (x-axis) ii) When was the rate of reaction fastest? iii) Use the graph to find the volume of gas produced after 35 seconds. iv) Use the graph to tell after how long the reaction stopped. v) On the graph sketch a line showing the experiment repeated at a higher temperature.

48. Experimental Results 0 50 100 150 200 250 300 0102030405060708090 Time / seconds Volume / cm3 ii) The reaction was fastest at about 25 seconds as the gradient of the line is highest at this point. iii) About 175 cm 3 iv) About 55 seconds. v) Higher temperature reaction is in red. Answer

49. A flask was connected to a gas syringe by a glass delivery tube. 30cm 3 of water and 0.5g of manganese dioxide were added to the flask. Then 5cm 3 of hydrogen peroxide was added and the stopper quickly fitted. Readings of the volume of gas produced were taken every 10 seconds. i) Plot a graph of volume of gas (y-axis) against time (x-axis). Label this curve A. ii) Without emptying the flask another 10cm 3 of water and a further 5cm 3 of hydrogen peroxide were added. Sketch the shape of the second experiment and label it B.

50. Notes: Curve B is an experiment with half the concentration of hydrogen peroxide. This should produce about half the rate as shown by a line with half the gradient of A. However, the same amount was added so 58cm 3 of gas will still be produced. A B Time (s) Answer

51. i) The results above were obtained from an experiment where the loss in mass was recorded as lumps of zinc reacted with hydrochloric acid. Plot a graph of mass loss (y-axis) against time (x-axis). ii)On the graph sketch the lines you would expect if a) the concentration of acid was reduced, b) the temperature was increased.

52. (b) (a) Answer

53. Which of these would speed up the rate at which magnesium dissolves in acid? A. Cool the acid. B. Cut up the magnesium. C. Add water. D. Coat the magnesium in oil.

54. Why does breaking up solids increase the rate of reaction? A. Makes more solid. B. Creates more energy. C. Increases surface area. D. Increases the concentration.

55. Why does temperature increase the rate of reaction? A. Acts as a catalyst. B. Increases the concentration. C. Increases number of molecules. D. Makes collisions more frequent and harder.

56. Why does a catalyst increase the rate of reaction? A. Provides a route with a lower activation energy. B. Helps provide energy for the reaction. C. Increases the speed of reactant molecules. D. Reduces the number of molecular collisions.

57. Why do most reactions start fast and get slower and slower? A. They run out of energy. B. They run out of catalyst. C. The concentration of reactant molecules gets less and less. D. The surface area increases.

58. Experiment Design Looking at designing an experiment. Investigating the effect of concentration of thiosulphate on the rate of reaction.