Week 10.3 chemical kinetics


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Week 10.3 chemical kinetics

  1. 1. Chapter 5 Chemical Kinetics Reaction Rates Rate Laws Collision Model Prepared by: Mrs Faraziehan Senusi PA-A11-7C Reaction mechanism Catalysis Reference: Chemistry: the Molecular Nature of Matter and Change, 6th ed, 2011, Martin S. Silberberg, McGraw-Hill
  2. 2. Reaction Mechanisms • The sequence of events that describes the actual process by which reactants become products is called the reaction mechanism. • It is the step-by-step pathway by which a reaction occurs. • Reactions may occur all at once or through several discrete steps. • Each of these processes is known as an elementary reaction or elementary step. • Some reactions take place in a single step, but most reactions occur in a series of elementary steps.
  3. 3. Elementary Reaction/Step The molecularity of a process tells how many molecules are involved in the process.
  4. 4. • A reaction mechanism is defined as a proposed set of elementary steps, which account for the overall features of the reaction. • Each of the reactions that comprises the mechanism is called an elementary step. • We believe it is elementary because it takes place in a single reactive encounter between the reactants involved. • These elementary steps are the basic building blocks of a complex reaction and cannot be broken down any further.
  5. 5. Example of Reaction Mechanism
  6. 6. Rate Determining Step in Reaction Mechanism • In a reaction mechanism, one of the elementary steps will be slower than all others. • The overall reaction cannot occur faster than this slowest, rate-determining step. • Therefore, this elementary, rate-determining step establishes the rate of the overall reaction. • The speed at which the slow step occurs limits the rate at which the overall reaction occurs.
  7. 7. Slow Initial Step NO2 (g) + CO (g)  NO (g) + CO2 (g) • The rate law for this reaction is found experimentally to be Rate = k [NO2]2 • CO is necessary for this reaction to occur, but the rate of the reaction does not depend on its concentration. • This suggests the reaction occurs in two steps.
  8. 8. • A proposed mechanism for this reaction is Step 1: NO2 + NO2  NO3 + NO (slow) Step 2: NO3 + CO  NO2 + CO2 (fast) • In this proposed mechanism two molecules of NO2 collide to produce one molecule each of NO3 and NO. • The reaction intermediate NO3, then collides with one molecule of CO and reacts very rapidly to produce one molecule each of NO2 and CO2. • The NO3 intermediate is consumed in the second step. • As CO is not involved in the slow, rate-determining step, it does not appear in the rate law.
  9. 9. Fast Initial Step 2 NO (g) + Br2 (g)  2 NOBr (g) • The rate law for this reaction is found to be Rate = k [NO]2 [Br2] • Because termolecular processes are rare, this rate law suggests a two-step mechanism.
  10. 10. • A proposed mechanism is Step 1: NO + Br2 • • (fast) Step 2: NOBr2 + NO  2 NOBr • NOBr2 (slow) The first step involves the collision of one NO molecule (reactant) and one Br2 molecule (reactant) to produce the intermediate species NOBr2. The NOBr2 can react rapidly, however, to re-form NO and Br2. We say that this is an equilibrium step includes the forward and reverse reactions. Eventually another NO molecule (reactant) can collide with a short-lived NOBr2 molecule and react to produce two NOBr molecules (product).
  11. 11. • • • The rate of the overall reaction depends upon the rate of the slow step. To analyze the rate law that would be consistent with this proposed mechanism, we again start with the slow (rate-determining) step, step 2. Denoting the rate constant for this step as k2, we could express the rate of this step as Rate = k2 [NOBr2] [NO] But how can we find [NOBr2]? NOBr2 is a reaction intermediate, so its concentration at the beginning of the second step may not be easy to measure directly.
  12. 12. • NOBr2 can react two ways: – With NO to form NOBr Step 2: NOBr + NO  2 NOBr – By decomposition to reform NO and Br2 2 Step 1: NO + Br2 ↔ NOBr2 (slow) (fast) • The reactants and products of the first step are in equilibrium with each other. • Therefore, Ratef = Rater • Because Ratef = Rater , k1f [NO][Br2] = k1r [NOBr2]
  13. 13. • Solving for [NOBr2] gives us k1f [NO][Br2] = k1r [NOBr2] k1f [NO] [Br2] = [NOBr2] k1r rate law for the rate-determining step: Rate = k2 [NOBr2] [NO] • Substituting this expression for [NOBr2] in the rate law for the rate-determining step gives Rate = k2 k1f [NO] [Br2] [NO] k1r Rate = k [NO]2 [Br2]
  14. 14. Catalysts • Catalysts are substances that can be added to reacting systems to increase the rate of reaction. • They allow reactions to occur via alternative pathways that increase reaction rates by lowering activation energies. • Catalysts change the mechanism by which the process occurs.
  15. 15. • A catalyst does take part in the reaction, but all of it is re-formed in later steps. • Thus, a catalyst does not appear in the balanced equation for the reaction.
  16. 16. • How does activation energy affects rate of reaction?? Arrhenius Equation k = A e−Ea/RT When a catalyst is present, the energy barrier is lowered. Thus, more molecules possess the minimum kinetic energy necessary for reaction.
  17. 17. CATALYSIS • • A catalyst changes the rate of a chemical reaction. Two categories of catalysts: (1) homogeneous catalysts (2) heterogeneous catalysts
  18. 18. Homogeneous catalysts • A homogeneous catalyst exists in the same phase as the reactants. • Catalyst can operate by increasing the number of effective collisions. • That is, from the Arrhenius equation: catalyst increase k by increasing A or decreasing Ea. • A catalyst may add intermediates to the reaction. • Example: In the presence of Br-, Br2 (aq) is generated as an intermediate in the decomposition of H2O2. • When a catalyst adds an intermediate, the activation energies must be lower than the activation energy for the uncatalyzed reaction.
  19. 19. Heterogeneous catalysts • A heterogeneous catalyst is present in a different phase from the reactants. • Such catalysts are usually solids, and they lower activation energies by providing surfaces on which reactions can occur. • The first step in the catalytic process is usually adsorption, in which one or more of the reactants become attached to the solid surface. • Some reactant molecules may be held in particular orientations, or some bonds may be weakened; in other molecules, some bonds may be broken to form atoms or smaller molecular fragments. This causes activation of the reactants.
  20. 20. • As a result, reaction occurs more readily than would otherwise be possible. • In a final step, desorption, the product molecules leave the surface, freeing reaction sites to be used again. • Most contact catalysts are more effective as small particles, because they have relatively large surface areas.
  21. 21. A schematic representation of the catalysis of the reaction on a metallic surface (Pt, NiO) 2CO (g) + O2(g)  2CO2(g)
  22. 22. Enzymes • Enzymes are proteins that act as catalysts for specific biochemical reactions in living systems. • The reactants in enzymecatalyzed reactions are called substrates. • Most enzymes are protein molecules with large molecular masses (10,000 to 106 amu). • The substrate fits into the active site of the enzyme much like a key fits into a lock.
  23. 23. • Enzymes have very specific shapes. • Most enzymes catalyze very specific reactions. • Substrates undergo reaction at the active site of an enzyme. • A substrate locks into an enzyme and a fast reaction occurs. • The products then move away from the enzyme. • Only substrates that fit into the enzyme lock can be involved in the reaction. A space-filling model of the enzyme lysozyme. This enzyme catalyzes the hydrolysis of polysaccharides (complex carbohydrates) found in bacterial cell walls.