Chapter 16. Rate Laws. The rate law describes the way in which reactant concentration affects reaction rate.

Transcription

1 Rate Laws The rate law describes the way in which reactant concentration affects reaction rate. A rate law is the expression that shows how the rate of formation of product depends on the concentration of all species other than the solvent that take part in a reaction. A rate law may be simple or very complicated.

2 Rate Laws, continued By studying rate laws, chemists learn how a reaction takes place and can often make an informed guess about how atoms move in rearranging themselves from reactants into products. The reaction mechanism is the way in which a chemical reaction takes place. A reaction mechanism is expressed by a series of chemical equations.

3 Rate Laws, continued Determining a General Rate Law Equation For a reaction that involves a single reactant, the rate is often proportional to the concentration of the reactant raised to some power. rate = k[reactant] n The exponent, n, is called the order of the reaction. It can be a whole number, a fraction, or zero. The term k is the rate constant, a proportionality constant that varies with temperature.

4 Determining a Rate Law Sample Problem B Three experiments were performed to measure the initial rate of the reaction: 2HI(g) H 2 (g) + I 2 (g) Conditions were identical in the three experiments, except that the hydrogen iodide concentrations varied. The results are shown on the following slide.

5 Determining a Rate Law, continued Sample Problem B Solution Find the ratio of the reactant concentrations between experiments 1 and 2, [HI] [HI] Then see how this affects the ratio reaction rates ( rate) ( rate) 2 1 of the

6 Determining a Rate Law, continued Sample Problem B Solution, continued When the concentration changes by a factor of 2, the rate changes by 4, or 2 2. Hence the reaction order, n, is 2.

7 Rate Laws, continued Rate Laws for Several Reactants When a reaction has more than one reactant, a term in the rate law corresponds to each. There are 3 concentration terms in the rate law for: 2Br (aq) + H 2 O 2 (aq) + 2H 3 O + (aq) Br 2 (aq) + 4H 2 O(l) There is an order associated with each term: rate = k[br ] n 1[H 2 O 2 ] n 2[H 3 O + ] n 3

8 Rate Laws, continued Rate Laws for Several Reactants, continued

9 Rate Laws, continued Rate Laws for Several Reactants, continued The reaction on the previous slide participates in the destruction of the ozone layer high in the atmosphere. NO(g) + O 3 (g) NO 2 (g) + O 2 (g) There are two terms in the rate law for this reaction. rate = k[no] n 1[O 3 ] n 2 In this case, it turns out that n 1 = n 2 = 1, so this reaction has a simple one-step mechanism.

10 Rate Laws, continued Rate-Determining Step Controls Reaction Rate Although a chemical equation can be written for the overall reaction, it does not usually show how the reaction actually takes place. For example, the reaction shown below is believed to take place in four steps. 2Br (aq) + H 2 O 2 (aq) + 2H 3 O + (aq) Br 2 (aq) + 4H 2 O(l) The order with respect to each of the three reactants was found to be 1.

11 Rate Laws, continued Rate-Determining Step Controls Reaction Rate, continued These four steps add up to the overall reaction : Br ( aq ) H3O ( aq ) HBr( aq ) H2O( l) HBr( aq ) H O ( aq ) HOBr( aq ) H O( l) Br ( aq ) HOBr( aq ) Br 2( aq) OH ( aq) OH ( aq ) H3O ( aq ) 2H 2O( l)

12 Rate Laws, continued Rate-Determining Step Controls Reaction Rate, continued Three of the steps are shown as equilibria. These are fast reactions. Step 2, however, is slow. If one step is slower, it will control the overall rate. A reaction cannot go faster than its slowest step. Such a step is known as the rate-determining step. Species such as HOBr that form during a reaction but are then consumed are called intermediates.

13 Reaction Pathways and Activation For two particles to react, they must collide violently enough to overcome the repulsion of their electrons. The kinetic energies of gas particles vary widely. Only particles with high kinetic energy are likely to react. The minimum energy that two colliding particles need to have before a chemical change is possible is called the of the activation energy, E a, of the reaction. No reaction is possible if the colliding pair has less energy than E a.

14 Reaction Pathways and Activation Energy, continued Activation-Energy Diagrams Model Reaction Progress With a combined kinetic energy equal to the E a, the molecules reach a state where there is a 50:50 chance of either returning to the initial state without reacting, or of rearranging to become products. This point is called the activation complex or transition state of the reaction.

15 Be Explained Activation Energies

16 Reaction Pathways and Activation Energy, continued Activation-Energy Diagrams Model Reaction Progress, continued 2HI H 2 I 2 H 2 + I 2 initial state activated complex final state (reactants) (products) In the initial state, the bonds are between the H and I. In the activated complex, four weak bonds link the four atoms into a deformed square. In the final state, the bonds link H to H and I to I.

17 Reaction Pathways and Activation Energy, continued Hydrogen Bromide Requires a Different Diagram The graph for HBr is similar to that of HI. One difference is that the E a is lower in the case of HBr. Because the activation energy of HBr is lower than that of HI, a larger fraction of the HBr molecules have enough energy to react than in the HI case. As a result, hydrogen bromide decomposes more quickly than hydrogen iodide does.

18 Reaction Pathways and Activation Energy, continued Hydrogen Bromide Requires a Different Diagram, continued In both graphs, the initial states are not at the same energy as the final states. The products have a lower energy than the reactants in the case of the HI decomposition reaction. The opposite is true for hydrogen bromide decomposition.

19 Reaction Pathways and Activation Energy, continued Hydrogen Bromide Requires a Different Diagram, continued The decomposition of hydrogen iodide is exothermic. 2HI(g) H 2 (g) + I 2 (g) H = 53 kj The decomposition of hydrogen bromide is endothermic. 2HBr(g) H 2 (g) + Br 2 (g) H = 73 kj

20 Reaction Pathways and Activation Energy, continued Not All Collisions Result in Reaction In order to react, molecules must collide with enough energy to overcome the activation energy barrier. Correct orientation in a collision is also important. New bonds are formed between specific atoms in colliding molecules. Molecules will not react without the proper orientation, no matter how much kinetic energy they have.

21 Reaction Pathways and Activation Energy, continued Not All Collisions Result in Reaction, continued For example, if a chlorine molecule collides with the oxygen end of the nitrogen monoxide molecule, the following reaction may occur. NO(g) + Cl 2 (g) NOCl(g) + Cl(g) This reaction will not occur if the chlorine molecule strikes the nitrogen end of the molecule.

22 Be Explained Particle Collisions

23 Catalysts Increase Reaction Rate Often, adding a catalyst to a reaction mixture will increase the reaction rate, even though the catalyst is not changed or used up at the end of the reaction. The process is called catalysis. For example, hydrogen peroxide decomposes slowly. 2H 2 O 2 (aq) 2H 2 O(l) + O 2 (g) Adding a drop of potassium iodide solution speeds up the reaction, because it acts as a catalyst.

24 Catalysts Increase Reaction Rate, continued Carbon monoxide is a poisonous gas in car exhaust. A car s catalytic converter causes the reaction of carbon monoxide with oxygen to take place much more quickly than it would alone: 2CO(g) + O 2 (g) 2CO 2 (g) Catalysis does not change the overall reaction at all. The stoichiometry and thermodynamics are the same. The changes affect only the path the reaction takes from reactant to product.

25 Catalysts Increase Reaction Rate, continued Catalysts Lower the Activation Energy Barrier Catalysis works by making a different pathway available between the reactants and the products. The new pathway has a different mechanism and a rate law from that of the uncatalyzed reaction. The catalyzed pathway may involve a surface reaction. Or, the catalytic mechanism may take place in the same phase as the uncatalyzed reaction.

26 Catalysts Increase Reaction Rate, continued Catalysts Lower the Activation Energy Barrier, continued The iodide-catalyzed decomposition of hydrogen peroxide is catalysis that does not involve a surface: 1) I (aq) + H 2 O 2 (aq) IO (aq) + H 2 O(l) 2) IO (aq) + H 2 O 2 (aq) I (aq) + O 2 (g) + H 2 O(l) The I ion, used up in step 1, reforms in step 2, and the IO ion, reformed in step 1, is used up in step 2.

27 Catalysts Increase Reaction Rate, continued Catalysts Lower the Activation Energy Barrier, continued In principle, a single iodide ion could break down an unlimited amount of hydrogen peroxide. This is the characteristic of all catalytic pathways the catalyst is never used up. Each pathway corresponds to a different mechanism, a different rate law, and a different activation energy.

28 Catalysts Increase Reaction Rate, continued Comparison of Pathways for the Decomposition of H 2 O 2

29 Catalysts Increase Reaction Rate, continued Enzymes Are Catalysts Found in Nature Enzymes are large protein molecules whose biological role is to catalyze processes that otherwise would happen too slowly to help an organism. For example, the enzyme lactase catalyzes the reaction of water with lactose, a sugar present in milk. People whose bodies lack the ability to produce lactase have what is known as lactose intolerance.

30 Catalysts Increase Reaction Rate, continued Enzymes Are Catalysts Found in Nature, continued Enzymes are very specific and catalyze only one reaction. This is because the surface of an enzyme molecule has a detailed arrangement of atoms that interacts with the target molecule. The enzyme site and the target molecule are said to have a lock and key relationship to each other.

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