Reaction Rates

1. Topic 6 Reaction Rates

2. IB Core Objective 6.1.1 Define the term rate of reaction.

3. 6.1.1 Define the term rate of reaction. • Different chemical reactions occur at different rates (ie. speeds) • Rapid reactions, eg. The neutralisation of a strong acid by a strong base in aqueous solution • Slow reactions, eg. Rusting of iron • Rate of a chemical reaction is a measure of the speed at which products are formed, measured as the change in concentration divided by the change in time (units = mol dm-3s-1)

4. IB Core Objective 6.1.2 Describe suitable experimental procedures for measuring rates of reactions.

5. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions. • Different changes can be measured vs. time to give a rate: • Rate = Change/time • Rate = Δ[Concentration]/ Δ[time] • Rate = Δ[Pressure]/ Δ[time] • Rate = Δ[Absorbance/ transmittance]/ Δ[time] • Rate = Δ[pH]/ Δ[time] • Reaction rate describes how fast a reaction will take place.

6. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions. • What factors will determine rate? • Surface area (usually refers to solid) • Concentration or pressure (if gas) • Temperature • Catalyst • How is it found? • Through experimentation only • Rate of decreasing reactants • −Δ[Concentration]/ Δ[time] • Rate of increasing products • Δ[Concentration]/ Δ[time] Notice the negative sign

7. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions. • pH changes (pH meter or acid/base titration) • Volume, pressure or mass changes • Conductivity (aq) (Conductivity meter or titration) • Spectrometer or colorimeter for colour change

8. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions. • Data loggers and probes • Conductivity • Acid/base • Temperature • Titrations • Acid/ base • Concentration

9. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions. TITRATION • Removal of small samples from the reaction mixture at different times and then titrating the sample to determine the concentration of either one of the reactants or one of the products at this time. • Results can then be used directly to generate a graph of concentration against time. • Best for quite slow reactions

10. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions. • Spectrometer • Absorbance and transmittance • Measures concentration

11. Collecting gas

12. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions. COLLECTION OF EVOLVED GAS • Gas produced in the reaction is collected either in a gas syringe or in a graduated vessel over water. • The volume of gas collected at different times can be recorded. • This technique is limited to reactions that produce a gas (obviously!) PLUS if the gas is to be collected over water, the gas must not be water soluble. • An alternative technique is to carry out the reaction in a vessel of fixed volume and monitor the increase in the gas pressure

13. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions. COLLECTION OF EVOLVED GAS • Example: Measuring the rate of reaction between a moderately reactive metal (such as zinc) and an acid (such as hydrochloric acid). Zn (s) + 2 H+(aq) → Zn2+(aq) + H2 (g)

14. 6.1.2 Describe suitable experimental procedures for measuring rates of reactions. MEASUREMENT OF THE MASS OF REACTION MIXTURE • The total mass of the reaction mixture will only vary if a gas is evolved. • The gas should have a high molar mass (ie. not hydrogen) so there is a significant change in mass • Example: measuring the rate of reaction between a metal carbonate (such as calcium carbonate, marble chips) and an acid (such as hydrochloric acid) by measuring the rate of mass loss resulting from the evolution of carbon dioxide: CaCO3 (s) + 2H+(aq) → Ca2+(aq) + H2O (l) + CO2 (g)

15. IB Core Objective • 6.1.3 Analyse data from rate experiments. • Students should be familiar with graphs of changes in concentration, volume, and mass against time.

16. 6.1.3 Analyse data from rate experiments. • Numerical value will vary according to amount of the substance involved in the stoichiometric equation MnO4-(aq) + 8 H+(aq) + 5 Fe2+(aq) → Mn2+ (aq) 4 H2O (l) + 5 Fe3+(aq) • The rate of appearance of Fe3+ is five times as great as the rate at which MnO4- is consumed Rate = - ∆ [MnO4-] = 1∆ [Fe3+] ∆ t 5 ∆ t Units: mol dm-3 s-1

17. 6.1.3 Analyse data from rate experiments. • Or more simply for a reaction: a A → b B, then Rate = - 1∆ [A]= 1∆ [B] a ∆ t b ∆ t • Any property that differs between the reactants and products can be used to measure the rate of the reaction General Formula

18. 6.1.3 Analyse data from rate experiments. • To find rate at an instant in time (instantaneous rate) we calculate the slope of a tangent on the experimentally obtained graph. • Here, Δx is the change in time and Δy is the change in concentration.

19. IB Core Objective How does kinetic energy and contact differ for solids when compared to liquids and gases? • 6.2.1 Describe the kinetic theory in terms of the movement of particles whose average energy is proportional to temperature in kelvins. • As heat is supplied to a substance, the velocity (kinetic energy) of the particles will increase. • When the velocity increases, so does the temperature. • Therefore, the absolute temperature in kelvin is proportional to the average kinetic energy of all the particles. Solids are fixed and only vibrate, so kinetic energy is limited

20. IB Core Objective • 6.2.2 Define the term activation energy, Ea. • The minimum amount of energy required for reaction is the activation energy, Ea.

21. 6.2.2 Define the term activation energy, Ea. • Just as a ball cannot get over a hill if it does not roll up the hill with enough energy, a reaction cannot occur unless the molecules possess sufficient energy to get over the activation energy barrier.

22. IB Core Objective • 6.2.3 Describe the collision theory • Just because chemicals collide /interact does not mean they will react! • Reaction rate depends on: • Collision frequency • Number of particles with E ≥ Ea. • Appropriate collision geometry or orientation.

23. IB Core Objective • 6.2.4 Predict and explain, using collision theory, the qualitative effects of particle size, temperature, concentration and pressure on the rate of a reaction.

24. 6.2.4 Surface Area • By increasing the surface area we increase the contact area. • Collision rate will increase. Marble Marble chip

25. 6.2.4 Concentration & Pressure • Increasing the concentration or pressure will increase opportunity for molecules to react • Collision rate will increase • Pressure only affects gases

26. 6.2.1 and 6.2.4Temperature and Rate • Temperature is directly related to kinetic energy or particle speed. • Faster particles increases probability of molecule interaction • Will increase # of particles with sufficient energy to over come activation energy.

27. Orientation

28. IB Core Objective • 6.2.6 Describe the effect of a catalyst on a chemical reaction. • What do you know about catalysts? How would you define one in your own terms? • Have you heard the expression “catalyst for change”?

29. Catalysts 6.2.6 • Catalysts provide an alternate reaction pathway. • One that takes less energy. Similar to a tunnel through a mountain side, whether you go up and over, or go through you still end up at the same place. • Catalysts are just a facilitator • They are not used up • They lower the activation energy allowing for more particles to have the correct energy requirement

30. Analogy • Catalysts are like a dating service • They bring compatible people together • They are not involved in what happens after the pairs are together • They can be reused again and again. • This is not to say that people will not get together on their own, however it lowers the energy required to find a match

31. 6.2.6 Describe the effect of a catalyst on a chemical reaction. One way a catalyst can speed up a reaction is by holding the reactants together and helping bonds to break.

32. Draw an Energy Diagram for a catalyzed reaction

33. IB Core Objectives • 6.2.5 Sketch and explain qualitatively the Maxwell-Boltzmann energy distribution curve for a fixed amount of gas at different temperatures and its consequences for changes in reaction rate. • Students should be able to explain why the area under the curve is constant and does not change with temperature.

34. Maxwell–Boltzmann Distributions 6.2.5 Activation Energy

35. Maxwell–Boltzmann Distributions 6.2.5 • As the temperature increases, the curve flattens and broadens. • At higher temperatures, a larger number of molecules has higher energy. This system has a fixed number of particles

36. Maxwell–Boltzmann Distributions • If the dotted line represents the activation energy, as the temperature increases, so does the fraction of molecules that can overcome the activation energy barrier. • As a result, the reaction rate increases. Be sure to draw the higher temp. Curve shorter and wider than the original.

37. IB Core Objective • 6.2.7 Sketch and explain Maxwell-Boltzmann curves for reactions with and without catalysts.

38. Catalysts and Activation Energy 6.2.7 Is this a Maxwell-Boltzmann distribution curve? Why or why not? Sketch what a Maxwell-Boltzmann curve would look like.

39. HL

40. Return Rate Law

41. HL Objective • 16.1.1 Distinguish between the terms rate constant, overall order of reaction and order of reaction with respect to a particular reactant.

42. 16.1.1 Distinguish between the terms rate constant, overall order of reaction and order of reaction with respect to a particular reactant. • Rate Rxn = Δ[A]/ Δt Rate = k[A]m [B]n • k = rate constant • A and B are reactants. • m and n represent the order of reaction for each reactant. • m + n = overall order of reaction • ‘k’ represents a constant and does not change EXCEPT • Temperature!! • Particle size (solids)

43. HL Objective • 16.1.2 Deduce the rate expression for a reaction from experimental data.

44. 16.1.2 Deduce the rate expression for a reaction from experimental data. • 1) If concentration of [A] is doubled yet no change in rate, than order = zero. [A]0 • 2) If doubling conc. of [A] = a doubling of rate, than order = one. [A]1 • 3) If doubling conc. of [A] = a quadrupling of rate, than order = 2. [A]2

45. 16.1.2 Deduce the rate expression for a reaction from experimental data. For the reaction A + 2B → C Deduce the order of reaction in respect to reactant A Deduce the order of reaction in respect to reactant B State what the overall order of the reaction is. Deduce the rate expression for this reaction A: 1. 2nd order. 2. 1st order 3. 2+1=3 3rd order 4. rate = k[A]2[B]

46. IB HL Objective • 16.1.3 Solve problems involving the rate expression.

47. 16.1.3 Solve problems involving the rate expression. • From the previous problem, rate = k[A]2[B]. • To find the rate constant, plug in numbers from the results: Notice the units! A: k= 4 x 103 mol-2 dm6 s-1

48. 16.1.3 Solve problems involving the rate expression. Units for rate constant Zero order overall: mol dm-3 s-1 First order overall: s-1 Second order overall: mol-1 dm3 s-1 Third order overall: mol-2 dm6 s-1

49. IB HL Objective • 16.1.4 Sketch, identify and analyse graphical representation for zero-, first- and second-order reactions. • Note: Students should be familiar with both concentration-time and rate-concentration graphs.

50. 16.1.4 Sketch, identify and analyse graphical representation for zero-, first- and second-order reactions. • Zero order reaction • 1st order reaction • 2nd order reaction Concentration Rate Time Time