Chemistry 12 Unit I: Reaction Kinetics Introduction: Kinetics Definition: All reactions occur at different rates Examples: Slow Reactions Fast Reactions Chemists need to understand kinetics because sometimes it is important to slow down a reaction (food spoilage, burning of fossil fuels) and sometimes it is important to speed a reaction up (breakdown of PCB s or colouring of hair). Understanding reaction rates allows chemist s (or technologists) to work on these problems. Measuring Reaction Rate: * How do we measure the rate of something?
* We often use the molar concentration of something to measure reaction rate since it is a change in amount. The molar concentration of something is given the symbol [ ]. For example, the molar concentration of HCl would be written [HCl]. * What are some of the ways that we could measure a change in amount in the lab? Ex: CaCO 3(s) + 2HCl (aq) CaCl 2(aq) + CO 2(g) + H 2 O A Graphical Representation of a Reaction: 2H 2 O 2 2H 2 O + O 2 * If we assume that all of the H 2 O 2 is converted to products, all reaction rates should be the same since H 2 O 2 consumed = H 2 O & O 2 produced (taking stoichiometry into account). Some Experimental data for the above reaction: Time (min) [H 2 O 2 ] (mol/l) [O 2 ] (mol/l) 0.0 0.0200 0.0 200.0 0.0162 0.0019 400.0 0.0131 0.0034 600.0 0.0106 0.0047 800.0 0.0086 0.0057 1000.0 0.0069 0.0065 1200.0 0.0056 0.0072 1600.0 0.0037 0.0082
2000.0 0.0024 0.0088 Calculating Average Reaction Rates: Calculate the average rate of reaction of O 2 formation during the time interval 200.0 to 400.0 minutes. Calculate the rate of H 2 O 2 consumption for the same time interval. * Important: even though the rate for consumption should always be negative, the provincial exam (and Hebden) makes it positive. For correctness, however, a rate of consumption should be negative, but an OVERALL rate should always be positive. 2 things that come out of the examples above: 1/ 2/ Ex: If the rate of consumption of NaOH is 30g/min when reacted with H 2 SO 4, what is the rate of production of Na 2 SO 4 in g/min?
Ex: If the rate of consumption of CaCO 3 is 30 g/min in the following reaction: CaCO 3 + 2HCl CaCl 2 + CO 2 + H 2 O what is the rate of consumption of HCl in mol/min? * The above examples all involved the AVERAGE rate of reaction, but the rate changes within a reaction. In the beginning, it is really fast because of the availability of reactants. If we want to know the speed of a reaction at a certain period in time, called the instantaneous rate, we need to graph, draw a line that is tangent to the curve at that point, and calculate the slope of the tangent. This slope is the instantaneous rate (rate for that point in time). Assignment: Read pages 1-5, & 11 and do questions 1-4, 6, 8, 18 & 19 Factors Affecting Reaction Rate: What are the things that will speed up (or slow down) a reaction? 1) 2) 3) 4) 5)
Energy in Reactions: There are 2 types of energy involved in a chemical reaction: 1) 2) * The total energy is always conserved. * When 2 molecules/compounds approach one another, they both will have kinetic energy because they are moving and potential energy stored in the bonds. When they get close to each other, they start to slow down because their electrons start to repel (K.E. lowers). The potential energy starts to increase because the bonds start to deform (remember bonds are electrons and they are repelling each other). The potential energy continues to increase as bonds stretch and finally break. When the bonds reform, the potential energy drops and kinetic energy increases as the molecules/compounds push away from one another. The Potential Energy Diagram: * ΔH is the change in heat energy. It is the difference between the energy of the products and the energy of the reactants (E products E reactants ). This value will be negative for exothermic reactions and positive for endothermic reactions. * Ea is the activation energy. It is the energy it takes for reactants to form products.
Exothermic vs. Endothermic Assignment: Read pages 13-16 and do questions # 23-26 Kinetic Energy Diagrams:
* Temperature is: The Collision Theory: 1) 2) 3) * Successful reactions (ones that form products) need:
Collision Theory and the Factors Affecting Rate: 1/ Concentration: A + B C * A and B collide. If it is a successful collision, C will form. If it is not successful, A and B will re-form. * If we double the concentrations of A or B, we will increase the number of collisions. * If we have more collisions, we increase the probability that one of those collisions will have: 1) the correct orientation AND 2) enough energy to get to products When we increase the [reactants], we increase the reaction rate because we have a greater number of collisions, leading to an increased chance that these collisions will have the correct orientation and sufficient energy. 2/ Pressure: * same argument as concentration * since volume has decreased, the concentration is greater When we increase the pressure, there are more collisions therefore a greater probability of successful collisions.
3/ Catalysts and Inhibitors: * 2 ways that catalysts help the reaction: 1) They hold the reactants together (correct orientation) * Enzymes allow reactions to occur at body temperature (normally they would require much more heat energy). Ex: Oxidation of sugar 2) Provide a lower energy pathway (usually more steps) to produce products. * Inhibitors either keep reactants apart, remove the catalyst, or do side reactions so as to lead to different products than the ones we are interested in. A catalyst increases the rate of reaction by: - Creating the correct orientation - Creating a lower energy pathway to products 4/ Nature of Reactants: a) Reactivity: some combinations are more reactive than others (Na and H 2 O) BUT charges are important and can determine whether the reaction will occur quickly or not. If one is positive and the other negative, they will be pulled together. Other mixtures may have to rely on a chance collision.
b) # of bonds (steric hindrance): The reactants need to be able to get close in order to react. Less bonds to be broken and formed leads to a faster rate. Reactants of opposite charge will lead to a faster reaction. 5/ Surface Area: * In case 1, only the atoms on the outside of the crystal are able to react. In case 2, ALL atoms are available for collisions. An increase in surface area leads to an increase in reaction rate because more atoms are able to have collisions. More possible collisions leads to an increased probability of successful collisions. 6/ Phase: * There needs to be a reasonable amount of mixing and they must be able to come close together in order to have the most chance for a successful collision.
(s) (s): (l) (l): (g) (g): (aq) (aq): (s) (l): (s) (g): (l) (g): Reaction Rate & Phase: (aq) > (l) > (g) > (s) 7/ Temperature: For SLOW reactions a 10 C increase in temperature doubles the rate.
Minimum Threshold Energy: is the energy cut-off for a reaction. Below this K.E., no products will form. Above the M.T.E., the collision will have enough K.E. to form products. * when we increase the temperature (average K.E.) more molecules will have enough energy to form products. * Remember that M.T.E. does not change, it is characteristic for each reaction. Only the average K.E. changes (the curve moves around it).
Increasing the temperature increases the reaction rate because it increases the number of collisions that have sufficient energy. Surface Area: Concentration: Pressure: Nature of Reactants: Phase: Temperature: Catalysts: Assignment: Read pages 5-9 & 12 and do questions #12-16 and 20-22 Kinetic and Potential Energy Revisited: * Minimum Threshold Energy is the Kinetic Energy equivalent to Activation Energy (P.E.) * Molecules will need to be carrying sufficient K.E. to match or better E a (need to be moving fast enough).
Activation Energy Revisited: * at the top of the hill, the activated complex (or transition state) is thought to exist. * The Activated Complex is a short-lived, high energy, and very unstable complex that is neither reactant nor product. It is usually all of the reactant atoms clumped together with weak bonds. Because of repulsions, it is very unstable.
* reactants and products have a few repulsions, but mostly attractions making them stable. * the activation energy for the reaction is the energy it takes for the reactants to get to the activated complex (start to top of hill) because its formation is what the energy barrier for the reaction actually is. * once the activated complex is formed, it can either breakdown to form reactants again, or can form products. Assignment: Read pages 17-25 and do questions 29, 32, 34-37, & 41-45 Reaction Mechanism: * It is rather difficult for 2 reactants to collide and have sufficient energy and correct orientation. Therefore, it would be virtually impossible for 3 reactants to come together and collide in the correct way. Since many reactions involve more than 2 reactants coming together, there must be some other way for the reaction to occur. * It turns out that most reactions occur in more than one step. * Reaction Mechanism: * The reactions you have seen up to now have been net reactions (overall reactions). * Each step in a reaction mechanism is called an elementary step. Ex: 2NO + F 2 2 NOF (g)
* NOF 2 is created in one step and is then used up in another step. It is called a reaction intermediate. * The potential energy diagram would look like:
Ex: HBr + O 2 H 2 O + Br 2 * Step 1 is the slowest and it will determine the rate of the reaction. If we want to increase the rate of the reaction, we have to increase the rate of THAT step * The slowest step in a reaction mechanism is called the rate determining step. What are the reaction intermediates? Where are their energies? What is the overall reaction? Where are the energies of the activated complexes? Read pages 26-30 and do questions # 46, 47, 49, 51, & 55
Catalysts: * We have already learned that catalysts increase the rate of the reaction by: offering a lower energy pathway or increasing the chance of having the correct orientation, but they may also allow bond breakage to be easier. Definition of a catalyst: Potential energy diagrams: P.E. and K.E.
* M.T.E. and Ea are lowered. More molecules now have sufficient kinetic energy to react successfully. In this case (unlike temperature), the curve doesn t move and the molecules have not gained more kinetic energy. Which is the catalyzed curve? Which step is the rds? Where are the energies of the reaction intermediates? Label Ea, Ea(catalyzed), & ΔH. Can we pick out the catalyst if we are given the reaction mechanism? Ex:
Which is it based on the definition? Reaction Intermediate? * A catalyst is present as a reactant in a beginning step and in the products of a later step. (reactant then product) * A reaction intermediate is in the products of one of the steps and in the reactants of a subsequent step. Overall reaction of example above? * Catalysts are not usually written in the overall equation as a reactant. They are usually written on top of the arrow. * There are 2 types of catalysts: 1) 2) Some examples of catalysts: 1) Haber process: 2) Alcohol production:
3) Catalytic Converters: 4) Hydrogenation of butter/margarine: 5) ALL enzymes Assignment: Read pages 30-36 and do questions 56, 58, 59, & 61-63
A Typical Provincial Exam Short Answer Question Find the missing step for the reaction mechanism below if the overall reaction is 2NO + 2H 2 N 2 + 2H 2 O: Step1: 2NO N 2 O 2 Step 2: N 2 O 2 +H 2 N 2 O + H 2 O Step 3:? Tips: 1) Write the overall reaction at the bottom 2) Cancel out all things appearing on both sides 3) Add things that need to appear at the end 4) Add things that you need to cancel (opposite side of equation) - End of Reaction Kinetics-