Unit 1. Reaction Kinetics

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1 Unit 1. Reaction Kinetics Given: That butane takes less energy input to burn than a nacho chip; draw the graph of the reaction for both items. Reaction kinetics is the study of the rates and the factors, which affect the rates. Definition Where means the change in. REACTION RATE = amount TIME 1

2 1.2 Methods of measuring reaction rates Given the following reaction: Cu(s) + 4 HNO 3 (aq) Cu(NO 3 ) 2 (aq) + 2 H 2 O(l) + 2 NO 2 (g) + heat red-brown colorless blue brown The rate of reaction can be measured by a) Color change: using a spectrophotometer, the intensity of the characteristic blue color of Cu(NO 3 ) 2 whereby Rate = color intensity time b) Temperature change: since the reaction produces heat, using a graph of temperature vs. Time graph whereby Rate = temperature time 2

3 c) Pressure change: since a gas is liberated in the reaction, using a sealed vessel and a pressure gauge, and a pressure vs. Time graph whereby Rate = pressure time d) Mass change: since copper is the only solid in the reaction, the rate at which copper is used up can be measures and plotted on a graph of mass vs time whereby Rate = mass time 3

4 1.3 Factors Affecting Reaction Rates Reaction rates are determined by the rate at which molecules of two (or more) reactants collide. For bonds to be broken, reactants must approach each other at not only the right orientation, but also the right energy (force/speed). This constitutes an effective collision. Any change of conditions that leads to an increase in collision rates will increase the rate of reaction. THINK THIS THROUGH FOR EACH OF THE FOLLOWING FACTORS. a) Temperature An increase in temperature increases the rate of collisions. HOW SO? Increase in energy of collisions (molecules are moving faster) Increase in the number of collision (an overall increase in collisions per unit time implies that there will also be more at the right orientation) What will happen to the speed at which the reaction proceeds? b) Concentration 4

5 An increase in the concentration increases the rate of collisions. HOW SO? What will happen to the speed at which the reaction proceeds? c) Pressure An increase in the pressure increases the rate of collisions. HOW SO? What will happen to the speed at which the reaction proceeds? The volume of a system is inversely proportional to the pressure applied to the system. When the pressure on a system is increased, the volume of the system decreases, and visa versa. 5

6 d) The nature*** of the reactants ***This does not refer to the phase, but to the strength of the bonds being broken. The more and the stronger the bonds: the slower the reaction The transfer of electrons is fast because NO bonds are broken. Exercises 10, 11 6

7 e) Surface area The greater the surface area available for reaction, the greater the reaction rate. HOW SO? ***Although similar to concentration effects, surface area effects are separate from concentration effects. Homogeneous reaction - when reactants are in the same phase two gases two dissolved substances two miscible liquids Heterogeneous reaction when reactants are in different phases solid/liquid liquid/gas solid/gas two immiscible liquids FASTEST SLOWEST Aqueous ions > (gases or liquids) > solids 7

8 f) Catalysts and inhibitors catalysts - increase reaction rates without themselves being involved in the reaction inhibitors - decreases reaction rates by either combining with a catalyst or with a reactant to prevent reaction. 1.4 Experimental Measurement of Reaction Rates Refer to p.11 slope of the line (m) = y 2 - y 1 x 2 -x 1 Exercises 8

9 1.5 Reaction Rates and Collision Theory (a.k.a. Kinetic Molecular Theory - KMT) Molecules act as small, hard spheres which bounce off each other and transfer energy among themselves during collisions. Before two molecules can react, they must collide. a) The effect of concentration on KMT and why? b) The effect of temperature and why? 1.6 Enthalpy (Heat) Changes in Chemical Reactions a) Bond energies 9

10 Potential energy (PE) - energy due to an object s position in space, plus the sum of all attractive and repulsive forces amongst the particles making up the object. PE of a chemical system = the energy of the electrons in the chemical bonds + number and type of atoms in the molecules. Kinetic energy (KE) - energy due to movement in the system (either the movement of the entire system, or the molecules within the system) Bond energy - the amount of energy required to break a bond between two atoms E.g. Cl 2 (g) kj 2 Cl(g) bond broken <energy added> 2Cl(g) Cl 2 (g) kj bond formation <energy released> Exercises. b) Reaction heats ENTHALPY = H = KE + PE of a system at constant atmospheric pressure Change in enthalpy = H = H PROD - H REACT 10

11 Whereby: H PROD = combined enthalpies of all products H REACT = combined enthalpies of all reactants c) The sign of H i) endothermic - H > 0 Why? ii) exothermic - H < 0 Why? 1.7 Kinetic Energy Distribution (Boltzmann Distribution) The real picture of molecular collisions: Molecules in a system do not have identical kinetic energies (REMEMBER: kinetic energy is the energy due to movement. Simply put, molecules in a system travel at a range of different speeds). For any system there is a continuous range of speeds with a minimum, maximum and an average speed (or energy). In many ways the Boltzmann Distribution is similar to a Bell Curve except a Bell Curve is symmetrical, whereas a Boltzmann rarely is. 11

12 Ethanol decomposes to ethane and water as in the following reaction: C 2 H 5 OH C 2 H 4 + H 2 O At room temperature reaction does not detectably occur At 200 o C the reaction is very slow At 400 o C the reaction is rapid If at room temperature there are about collisions per second and yet no reaction occurs; what does that tell us about the requirement for reactivity? What factor must govern the reactivity? Since there is only a modest increase in collisions per temperature increment (1% per 10 o C rise) then: THE INCREASED REACTION RATE DUE TO THE INCREASED TEMPERATURE IS PRIMARILY DUE TO THE INCREASED NUMBER OF MOLECULE OVER THE MINIMUM OR THRESHOLD ENERGY OF REACTION, NOT THE NUMBER OF COLLISIONS Rule of thumb:. 12

13 . For slow reaction, a 10 o C increase in temperature doubles the rate 13

14 p19 figure Exercises p Activation Energies Activation energy (Ea) the minimum potential energy required to change the reactants into the activated complex Activated complex (AC) the arrangement of atoms that occurs when the reactants are in the process of rearranging to form products. (an intermediate form). Ineffective collision when the KE is less than what is required to break bonds and form an activated complex. Effective collision when the KE is sufficient to produce a reaction. *** if the KE is JUST sufficient to buy the required PE to form an AC, it may or may not occur. The molecules may come to a standstill and then separate without the breaking of bond. 14

15 ***if more than enough KE is present, the AC will move through space as a unit. The mechanism of reaction. As molecules approach each other, the outer electrons on one molecule start to repel the electrons on the other, causing a slowdown and the conversion of the KE (of motion) to PE (within bonds- a sum of the attractive and repulsive forces between two molecules. Remember?) If molecules gain enough PE in this conversion, then bonds can be broken and an activated complex can be formed. At this point the outer electrons of these new product molecules repel each other (PE). The energy with which they move apart (KE = motion energy) is a re-conversion of PE to KE. KE to PE Ratios In a Reaction. 15

16 The Importance Of Correct Molecular Alignment In A Reaction ***experimental E a values assume correct alignment. Hebden p.23 Exercises Reversibility of Reactions Reactants Products Ea(f) = energy of activation for the forward reaction Ea(r) = energy of activation for the reverse reaction AE is always endothermic, even in reactions that are exothermic overall Ea(f) = Ea(r ) + H Hebden p. 25 Exercises 16

17 1.9 Reaction Mechanisms Reaction Mechanisms the sequence of steps in an overall reaction. Because the probability of more than two atoms/molecules arriving at the same place at the same time is extremely small, complex reactions must proceed in a step-wise fashion. Elementary process individual steps in a multi-step process. Rate determining step the slowest step in a reaction. Build up of reaction intermediates depends on which elementary process is the rate determining step. Hebden p. 26 4HBr + O 2 2H 2 O + 2Br 2

18 Hebden p. 28. Exercise 1.10 Energy Diagrams Of A Multi-Step Reaction Mechanism. HBr + O 2 HOOBr (slow) HBr + HOOBr 2HOBr (fast) 2HBr + 2HOBr 2H 2 O + 2Br 2 (fast) Hebden p. 30 Exercises 1.11 The Effects of Catalysts on E a catalyst a substance which provides an overall reaction with an alternative mechanism having a lower E a. 18

19 A lower Ea means more molecules will have sufficient KE to react increased reaction rate. Increased reaction rate is for both forward and reverse reactions!!! p.32 example 1.12 Effects of Catalyst on Reaction Mechanism Read Hebden p ; Exercises on p

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