Chemistry 40S Chemical Kinetics (This unit has been adapted from

Transcription

1 Chemistry 40S Chemical Kinetics (This unit has been adapted from Name: 1

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3 Lesson 1: Introduction to Kinetics Goals: Identify variables used to monitor reaction rate. Formulate an operational definition for reaction rate in terms of rate units. Measure the average rate of a chemical reaction. Determine the instantaneous rate of a reaction using graphical analysis. Rate of reaction Measuring Rate The reaction below shows that A and B combine to form C and D and heat is given off. Also, substance B has a green colour in solution and D has a red colour in solution. A (s) + B (aq) C (g) + D (aq) + heat In this reaction, rate can be measured in terms of several different properties: 1. Temperature change over time ( C/min). As the reaction proceeds, the system will increase in temperature. The faster heat is produced, the higher the rate. 2. Pressure change over time (kpa/s or mmhg/s). As the reaction proceeds, more gaseous C is produced. This increases the pressure of the system. The faster the pressure increases, the greater the rate. 3

4 3. Mass change over time (g of C/min). As the reaction proceeds, reactants are used up and converted to gas. This results in a loss of mass of reactants, or an increase in mass of product, C. 4. Colour change over time. Colour change is usually measured in terms of how much light of a specific wavelength can be absorbed. The amount of absorbance can be directly related to the concentration. The greater the concentration of D, the deeper the red colour and the more light that is absorbed. Calculating Average Rate The rate of a reaction is often described in terms of the change in concentration of the reactants or the products. Average rate of a reaction 4

5 Example 1: According to the reaction A B, the following data was collected: Time (s) Concentration of B (mol/l) a) What is the average rate over the entire 50 s? b) What is the average rate for the interval 20 s to 40 s? Example 2: Calculate the average reaction rate over this time period expressed as moles of C4H9Cl consumed per liter per second. 5

6 Example 3: The decomposition of nitrogen dioxide produces nitrogen monoxide and oxygen according to the reaction: 2NO2 (g) 2NO (g) + O2 (g) The following data has been collected: Time (s) [NO2] (mol/l [NO] (mol/l) [O2] (mol/l) Calculate the average rate of decomposition of NO2 over 400 s. Instantaneous Rate Reactions often start quickly, but slow down over time. If we graph the change in concentration of reactants or products versus time we find that the graphs are not linear. 6

7 Instantaneous rate 7

8 Practice: Introduction to Kinetics 1. For the reaction A products, the following data was collected: Time (min) Mass of A (g) a) What is the average rate, in g of A/min, over the entire 5 minutes? b) What is the average rate for the interval between 2.0 and 4.0 minutes? c) What is the average rate, in g of A/min, over the entire 5 minutes? d) What is the average rate for the interval between 2.0 and 4.0 minutes? 8

9 2. The decomposition of acetaldehyde to methane and carbon dioxide occurs according to the following equation: CH3CHO (g) CH4 (g) + CO (g) The results of an experiment are given below: Time (s) [CH3CHO] (mol/l) a) What is the rate of decomposition of acetaldehyde between 42 s and 105 s? b) What is the rate of decomposition in the interval 190 s to 480 s? 9

10 3. Below is the data from an experiment that studied the following reaction: 2HCl (aq) + CaCO3 (s) CaCl2 (aq) + H2O (l) + CO2 (g) HCl was placed in a beaker and massed immediately after adding CaCO3 chips (time = 0). The mass of the beaker was recorded at 1.0 minute intervals for a total of 15 minutes. We will assume the loss of mass is the amount of carbon dioxide gas that escapes from the beaker. Time (min) Mass of beaker and contents (g) a) Complete the table. Mass loss (CO2 produced) (g) b) Determine the average rate, in g of CO2/min, over the entire 10 minutes. c) Determine the average rate for the first 3 minutes and the last 3 minutes. 10

11 4. The formation of nitrogen dioxide from nitrogen dioxide and oxygen gas was studied. The balanced equation for the reaction is: O2 (g) + 2NO (g) 2NO2 (g) The chemist measured the concentration of the three gases at various time intervals. The data is in the table below. Construct a well-labeled graph, labeled axis with units, a title, etc to represent this data. Along the y-axis plot gas concentration and time on the x-axis. Time (min) Concentration (mol/l) O2 NO NO Once the graph is completed, answer the following questions: a) What is the average rate of consumption of nitrogen oxide and oxygen over the entire 71-minute interval? Determine the average rate for each. b) What is the average rate of formation of nitrogen dioxide over the entire 71 minute interval? c) What is the average rate of the consumption of NO and O2 and the production of NO2 each? 11

12 Lesson 2: Rate and Stoichiometry Goal: Relate the rate of formation of a product to the rate of disappearance of a reactant given experimental data and reaction stoichiometry. The graph below is from Lesson 1 s practice question 4: The equation for this reaction is O2 (g) + 2NO (g) 2NO2 (g). If we examine the graph, the concentration of the NO2 increases about the same amount as the NO decreases and the concentrations of the NO and NO2 change about twice the amount that the O2 changes. 12

13 We can use the equation below to determine the rate Example 1: The decomposition of nitrogen dioxide occurs according to the equation 2NO2 O2 + 2NO. If the rate of decomposition of NO2 is determined to be 0.50 mol/ls at a certain temperature, predict the rate of creation of both products. Example 2: For the reaction 2A + B 3C, what is the rate of production of C and the rate of disappearance of B if A is used up at a rate of 0.60 mol/ls? 13

14 Practice: Rate and Stoichiometry 1. If NOCl (g) is decomposing at a rate of mol/l/min in the following reaction: 2NOCl (g) 2NO (g) + Cl2 (g) a) What is the rate of formation of NO (g)? b) What is the rate of formation of Cl2 (g)? 2. Thiosulfate ion is oxidized by iodine according to the following reaction: 2S2O3 2 (aq) + I2 (aq) S4O6 2 (aq) + 2I (aq) If, in a certain experiment, mol of S2O3 2- is consumed in 1.0 L of solution each second, what is the rate of consumption of I2? At what rates are S4O6 2 and I produced in this solution? 14

15 3. If the decomposition of N2O5 gas occurs at a rate of 0.20 mol/l/s, what would be the rate of formation of NO2 gas and O2 for the reaction of 2N2O5 (g) 4NO2 (g) + O2 (g)? 4. If ammonia gas, NH3, reacts at a rate of mol/l/s according to the reaction 4NH3 (g) + 5O2 (g) 4NO (g) + 6H2O (g). a) At what rate does oxygen gas react under the same conditions? b) What is the rate of formation of water? c) What is the rate of production of nitrogen monoxide? 15

16 Lesson 3: Explaining Why Reactions Occur Goals: Define the Collision Theory for chemical reactions. Interpret Kinetic Energy Distribution curves Describe the role of activation energy in chemical reactions. Interpret reaction coordinate diagrams for endothermic and exothermic reactions. The Collision Theory Collision theory In the atmosphere, ozone is converted to oxygen gas and nitrogen dioxide by reacting with nitrogen monoxide, according to the following reaction NO (g) + O3 (g) NO2 (g) + O2 (g). If the oxygen atoms collide in orientation a), no reaction occurs. But if the nitrogen atom collides with the oxygen atom (b), a reaction occurs. 16

17 Activation Energy Not only do particles need collide with the correct orientation or reaction geometry, they must also collide with enough energy. Chemical reactions involve the making and breaking of bonds, which requires energy. If colliding particles do not have sufficient velocity (Kinetic Energy) they will not produce a reaction. The minimum amount of energy required for colliding particles to produce a chemical reaction is called the Activation Energy (EA) of that reaction. The greater the activation energy, the slower the reaction rate and the longer the reaction takes. For example, hydrogen and oxygen can be kept in the same container at room temperature without a reaction occurring. Even though the molecules collide, they do not possess activation energy. If the mixture is heated to 800 C, or a flame or spark is introduced, an explosive exothermic reaction occurs. The heat, flame or spark provides enough energy for the particles to reach activation energy. They are able to strike each other with enough energy to break the hydrogen-hydrogen and oxygen-oxygen bonds. Kinetic Energy Distribution 17

18 Heat of Reaction The rate of a reaction is determined by the height of the barrier that the particles must cross in order that they are converted into products. The activation energy is the barrier colliding particles must overcome. As the particles collide, they form an unstable intermediate particle called the activated complex or transition state. The energy required to produce the activated complex is the activation energy. The activated complex has a maximum amount of potential energy but exists for a very small instant in time. Enthalpy (H) 18

19 Potential Energy Diagrams A reaction coordinate diagram, or potential energy (EP) diagram, or reaction progress diagram represents the energy change that occurs during a chemical reaction. In the exothermic reaction, the products possess less energy than the reactants. During the reaction, heat is lost from the system and ΔH is a negative value. In an endothermic reaction, the products possess more potential energy than the reactants. This energy is absorbed from the surroundings, increasing the systems energy content, giving a positive ΔH value. 19

20 Potential energy diagrams provide a picture of the energy changes, which occur as a chemical reaction proceeds. The reaction coordinate axis can be time or any other means of measuring reaction progress. Let's look at the potential energy diagram for the formation of water from hydrogen and oxygen: 2H2 + O2 2H2O Hydrogen and oxygen can exist together without exploding or burning due to the need to overcome activation energy. The activation energy is high because of the need to break covalent hydrogen-hydrogen bonds in hydrogen molecules and oxygen-oxygen bonds in oxygen molecules. 20

21 Example 2: Let s sketch and analyze the reaction progress diagram for the reaction of CH3CH2Br + OH CH3CH2OH + Br. 21

22 Practice: Explain Why Reactions Occur 1. Given the following reaction coordinate diagram: a) What is the activation energy of the reaction shown by the diagram? b) What is the enthalpy change for this reaction? c) Is this reaction endothermic or exothermic? 2. What is the activated complex or transition state and how is it related to reaction rates? 22

23 3. Given the following reaction coordinate diagram: a) What is the activation energy of the reaction in the diagram to the left? b) What is the enthalpy change for this reaction? c) Is this reaction endothermic or exothermic? d) What would be the activation energy of the reverse reaction? 4. Does every collision between reactant particles produce a reaction? Explain. 23

24 5. Given the following reaction coordinate diagram: a) What is the activation energy of the diagram to the left? b) What is the enthalpy change for this reaction? c) Is this reaction endothermic or exothermic? d) What would be the activation energy of the reverse reaction? 6. Explain why the enthalpy change for an exothermic reaction is negative, even though the container gets warmer. 24

25 Lesson 4: Factors Effecting Rate Goals: Identify the factors that could affect the rate of a chemical reaction. Describe qualitatively the relationship between factors that affect the rate of chemical reactions and the relative rate of a reaction. Use the Collision Theory to explain the factors influencing the rate of a reaction. Explain the effect of temperature increases on the shape of a kinetic energy distribution curve. Explain the effect of concentration changes on the rate of a chemical reaction using a kinetic energy distribution curve. Explain the effect of a catalyst on the rate of a chemical reaction using a reaction coordinate diagram as well as a kinetic energy distribution curve. The Nature of Reactants Chemical reactions involve the breaking of bonds between atoms in the reactants and reforming bonds to make the products. The breaking of bonds requires energy and the forming of bonds releases energy. It makes sense then that the strength and number of bonds to be broken will affect the rate of a reaction. The weaker the bonds to be broken the faster the reaction 25

26 Example 1: Between the two reactions below, which would be the fastest? Reaction 1: H2 (g) + I2 (g) 2HI (g) Reaction 2: Pb(NO3)2 (aq) + 2NaI (aq) PbI2 (s) + 2 NaNO3 (aq) When comparing reactions with similar bonds, the more bonds that must be broken the slower the reaction Example 2: Between the two reactions below, which would be the fastest? 2NO (g) + O2 (g) 2NO2 (g) 2C8H18 (g) + 25O2 (g) 16CO2 (g) + 18H2O (g) 26

27 If the reactions have similar bonds and similar numbers of bonds, the state of the reactants is important Temperature Changes According to the collision theory, the rate of a reaction is determined by the frequency or number of successful or effective collisions. An effective or successful collision is a collision between reactant particles with enough energy and the correct orientation. At any given temperature there is a fixed number of particles that possess enough energy to produce an effective collision, that is, activation energy. Let s take a look at the diagram below. 27

28 Concentration Changes Increasing the concentration of reactants increases the total number of particles in a container. If the number of particles in a container increases, so do the number of particles with activation energy. More particles with activation energy increase the frequency of effective collisions. Increasing the number of particles in a container will also increase the chances of a collision. 28

29 Let s look at the diagram below. In a), with two particles of each, there exists four possible collisions which could produce a reaction. In b), by doubling the number of red particles, the number of possible collisions increases to eight. This will increase collision frequency. Catalysts Catalyst 29

30 The catalyst does NOT affect the reaction products or the enthalpy change for the reaction. Both remain the same. Catalysts will only provide an easier path for the reaction to proceed. Lowering the activation energy of a reaction increases the number of particles with enough energy to produce an effective collision, as shown by the kinetic energy distribution below. More particles with activation energy mean more frequent effective collisions and an increased reaction rate. Enzymes Enzymes are capable of increasing the rate of biological reactions by over one million times. Enzymes accomplish this by bringing the reactants, called substrates, together. The reactants bind to a certain part of the enzyme, called the active site. The reactants are then placed into a position, which favours a reaction. The active sites can be very specific, binding to only one set of reactants, or they can catalyze several different types of reactions. The reaction is a two-step mechanism: Enzyme + Substrate Enzyme-substrate complex Enzyme-substrate complex Enzyme + Product The enzyme remains unchanged after the reaction and is able to catalyze another reaction. An enzyme can remain active for an indefinite number of reactions. 30

31 The control of biological reactions is accomplished by the enzymes. The body will either begin producing enzymes when the reaction is required, or the enzymes can be "activated" by the presence of certain chemicals or ions. Pressure Changes Changing the pressure on a system usually only affects the reaction rates of gaseous reactions. Pressure is the force of particles upon the walls of their container. If the number of particles in a container increases without changing volume, the pressure increases. There are 3 ways to change pressure: 1. Add more product and/or reactant particles to the container. 2. Increase or decrease the volume of the container. 3. Add an inert or unreactive gas. If adding more reactant particles increases pressure, the concentration increases causing an increased rate. If removing reactant reduces the pressure, the rate decreases due to decreased concentration. If you decrease the volume of a container without changing the number of particles in the container, as in figure c) in the diagram below, the concentration of the reactants increases. The spaces between the particles decrease, increasing the chances of a collision. If the concentration of the reactants increases, the reaction rate increases. 31

32 If you increase the volume of a container without changing the number of particles, as in figure a in the diagram above, the concentration of the reactants decreases. The spaces between the particles increase, decreasing the chances of a collision. Decreasing the concentration of reactants decreases the reaction rate. Surface Area The size of the reactants, or surface area in contact, is usually only a factor in heterogeneous reactions. Increasing the surface area of the reactants by crushing, grinding or other means, increases the number of particles of reactants in contact. The rate of reactions increases when surface area in contact also increases. Increasing surface area increases the frequency of collisions, increasing reaction rate. Practice: Factors Affecting Rate 1. In general, what effect does an increase in the concentration of the reactants have on the rate of the reaction? Explain using the collision theory. 32

33 2. How do changes each of the following factors affect the rate of a chemical reaction? Use diagrams to clarify your explanations. a) Temperature b) Particle size 33

34 c) Pressure 3. Which equation of the following pairs of equations would occur the fastest at under the same conditions? Explain your answers. a) Zn 2+ (aq) + S 2 (aq) ZnS (s) Zn(s) + S(s) ZnS(s) b) 2H2O2 (aq) 2H2O (l) + O2 (l) Cu (s) + 2AgNO3 (aq) 2Ag + (aq) + Cu(NO3)2 (aq) c) Pb(NO3)2 (aq) + 2KI (aq) PbI2 (aq) + 2KNO3 (aq) C3H8 (g) + 5O2 (g) 3CO2 (g) + 4H2O (g) d) 2Fe (s) + 3O2 (g) 2Fe2O3 (s) 2NO (g) + O2 (g) 2NO2 (g) 34

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36 Lesson 5: Mechanism of Reactions Goals: Explain the concept of a reaction mechanism. Propose an analogy for a reaction mechanism. Draw reaction coordinate diagrams for reaction mechanisms. Predict the relative rate of a chemical reaction or step of a mechanism based on reaction coordinate diagrams. Collision Theory and Mechanisms Consider the reaction 2NO (g) + O2 (g) 2NO2 (g). According to the collision theory, for this reaction to occur in one step, 3 particles must collide. Reaction Intermediates 36

37 Let s take a look at the complex reaction of 2NO (g) + O2 (g) 2NO2 (g). Net Reactions Example 1: For the reaction of NO2 + CO NO + CO2, determine the reaction s mechanism. 37

38 Catalysts, as well as intermediates, do not appear in the overall reaction. Example 2: The decomposition of ozone using chlorine as a catalyst, 2O3 (g) 3O2 (g), determine the reaction s mechanism. Rate Determining Step 38

39 Example 3: Given the following mechanism: P + Q X + T (slow) X + P Y + R (fast) Y + S T (moderate) a) What is the net reaction? b) What are the reaction intermediates? c) Which is the rate determining step? d) What would be the effect of increasing the concentration of P? e) What would be the effect of decreasing the concentration of Q? f) What would be the effect of increasing the concentration of S? 39

40 Practice: Mechanism of Reactions 1. Given the following reaction mechanism: H2O2 + I H2O + IO H2O2 + IO H2O + O2 + I a) Write the balanced net reaction. b) Identify the reaction intermediate(s). c) Identify the catalyst(s). 2. Examine the following reaction mechanism: P + Q I1 + R (slow) I1 + P I2 + W (moderate) I2 + S T (fast) a) Write out the net reaction. b) Identify the overall rate of the net reaction. c) Increasing [P] increases the rate of the net reaction, increasing [Q] increases the rate of the net reaction, increasing [S] has no effect of the rate. Explain why this is possible. 40

41 3. Write the net reaction for the mechanism. A + B I1 + C (very slow) I1 + A C + D (fast) 4. A proposed mechanism for the preparation of the poisonous liquid nitrobenzene (C6H6NO2) is: C6H6 + NO2 + C6H6NO2 + (slow) H2SO4 H + + HSO4 very fast) C6H6NO2 + + HSO4 C6H5NO2 + H2SO4 (fast) a) What is the RDS? Why? b) What is the net reaction? c) Without H2SO4 this is a very slow reaction. Explain. 41

42 Lesson 6: Quantitative Concentration Changes Goals: Determine the rate law of a chemical reaction from experimental data. Explain the effect of concentration on reaction rates in terms experimental data and rate law. Derive rate law form a reaction mechanism. Predict a reaction mechanism from rate law. Rate Law Rate Law For the reaction: A products The rate of consumption of A is directly proportional to its concentration. That is, the faster A is consumed, the lower its concentration. The Rate Law is represented by the equation: Rate = k[a] x The order of a reaction indicates how concentration of reactants affects the rate of a reaction. 42

43 For a reaction with more than one reactant, such as A + B products The rate law would be: Rate = k[a] x [B] y Example 1: What is the rate law for the following reaction, given the experimental data below? H2O2 + 2HI 2H2O + I2 Trial [H2O2] (mol/l) [HI] (mol/l) Initial Rate (mol/ls)

44 Example 2: For the reaction A + B products, the following data was collected: Trial [A] (mol/l) [B] (mol/l) Initial Rate (mol/ls) Determine the rate law. Example 3: Determine the value of the specific rate constant, k, for the reaction in example 2. 44

45 Example 4: For the reaction 3A (g) + B(g) + 2C (g) 2D (g) + 3E (g) The following data was obtained: Trial [A] (mol/l) [B] (mol/l) [C] (mol/l) Initial Rate (mol/ls) ? 6? a) Write the rate law for this reaction. b) Calculate the value of the rate constant. c) Calculate the rate for Trial #5. d) Calculate the concentration of A in Trial #6. 45

46 For the elementary reaction such as, aa + bb cc + dd The rate law would be: Rate = k[a] a [B] b. Where a and b are the molar coefficients for the elementary reaction. Example 5: One of the reactions that result in smog is the reaction of ozone, O3 (g) and nitrogen monoxide, NO (g). This reaction is thought to occur in a single elementary step according to the equation: O3 (g) + NO (g) NO2 (g) + O2 (g) Determine the rate law for this reaction. Example 6: The mechanism for the reaction 3M + N P + 2Q is below: 2M P + X (slow) X + M Q + Y (moderate) N + Y Q (fast) a) What is the rate law for this reaction? b) What would be the effect of tripling the [M]? 46

47 c) What would be the effect of doubling the [N]? 47

48 Practice: Quantitative Concentration Changes 1. A first-order reaction initially proceeds at a rate of mol/ls. What will be the rate when half the starting material remains? When onefourth of the starting material remains? 2. Assume the N2O (g) and O2 (g) react according to the rate law Rate = k[n2o] [O2]. How does the rate change if: a) the concentration of O2 is doubled? b) the volume of the enclosing vessel is reduced by half? 48

49 3. Assume that NO (g) and H2 (g) react according to the rate law Rate = k[no] 2 [H2]. How does the rate change if: a) the concentration of H2 is tripled? b) the concentration of NO is doubled? c) the volume of the enclosing vessel is reduced by half? 4. For the reaction A + 2B 2C, find the rate law and calculate the value of the specific rate constant. [A] mol/l [B] mol/l Rate (mol/lmin)

50 5. The reaction I (aq) + OCl (aq) IO (aq) + Cl (aq) was studied and the following data were obtained: Trial [I ] mol/l [OCl ] mol/l Initial Rate (mol/l s) a) What is the rate law? b) What is the value of the rate constant? 50

51 6. For the reaction A + B + C D, find the rate law and calculate the value of the specific rate constant. Trial [A] mol/l [B] mol/l [C] mol/l Initial Rate (mol/l min)

52 7. For the reaction X + Y + Z S, find the rate law and calculate the value of the specific rate constant. Trial [X] mol/l [Y] mol/l [Z] mol/l Initial Rate (mol/lmin)

53 8. The reaction CH3COCH3 + I2 CH3COCH3 + HI is run in the presence of an excess of acid. The following data were obtained: Trial Initial [I2] (mol/l) Initial [CH3COCH3] (mol/l) Initial Rate (mol/ls) a) What is the rate law? b) What is the value of the rate constant? c) What is the rate if the concentration of CH3COCH3 is mol/l and the concentration of I2 is mol/l? d) What is the concentration of I2 if the concentration of CH3COCH3 is mol/l and the rate is mol/ls? 53

54 9. For the reaction A + 2B C + D, the following data was collected: Trial Initial [A] (mol/l) Initial [B] (mol/l) Initial Rate (moll -1 s -1 ) What is the rate law? 10. For the reaction 3A + B 2C + D, the following data was collected: Trial Initial [A] (mol/l) Initial [B] (mol/l) Initial Rate (moll -1 h -1 ) What is the rate law? 54

55 11. For the elementary reaction H2 + I2 2HI. a) Write the rate law. b) Find k if HI is produced at a rate of mol/lmin when [H2] = mol/l and [I2] = mol/l. c) What is the rate of production of HI if the concentration of both reactants is 0.10 mol/l and the temperature is the same as in b)? d) How would the rate be affected if [H2] is doubled and the [I2] is halved? 55

56 12. For the one step reaction A (g) + 2B (g) C (g). a) What is the rate law? b) How does the rate change if: i. [A] is doubled? ii. [B] is tripled? iii. the volume of the container is doubled? 56

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