Kinetics is the study of of chemical reactions and the by which they occur.

Transcription

1 Kinetics is the study of of chemical reactions and the by which they occur.

2 tells us if a reaction can occur while tells us how quickly the reaction occurs some reactions that are thermodynamically feasible are kinetically so slow that they are hardly noticeable The changing of a diamond into either graphite or carbon dioxide gas is thermodynamically favorable, but so slow that it is not detectable even in a lifetime. Why? It has an enormous activation energy that must be overcome before the reaction may proceed. C O CO diamond H + + OH 2 - g H g aq aq 2 l HAPPENS 2 VERY SLOW O G G o 298 o 298 INSTANTLY 396kJ = -79 kj 2

3 1. if the temp. increases, then the rxn tends to speed up. Generally a raise of 10 o C doubles the rate. Heat em up & speed em up! ; the faster molecules move, the more likely they are to collide and the more energetic the collisions become. In other words, the more likely molecular damage will occur. Bonds will be broken and new bonds will form at a faster rate.

4 2. the more concentrated the reactants are the quicker the rxn. The more molecules present, the more collisions occur, the faster reaction proceeds, the greater its rate. It s like the crowded halls during the passing period

5 Simplified representation of effect of different numbers of molecules in the same volume. Increase in concentration, increases the chance of a rxn occurring. A(g) + B(g) Products A A B B A A B B B A A A B B B 4 different possible A- B collisions 6 different possible A- B collisions 9 different possible A- B collisions 5

6 3. Chemical identity: some chemicals react very quickly, some are very slow 2Na + 2H 2 O 2Na + + 2OH - + H 2 very fast 2Al + 6H 2 O 2Al(OH) 3 + 3H 2 very slow Physical state also matters: solids will react much more slowly than liquids or aqueous solutions

7 Nature of Reactants; Broad category that includes the different reacting properties of substances. For example: ~ Sodium reacts with water explosively at room temperature to liberate hydrogen and form sodium hydroxide. ~ Calcium reacts with water slowly at room temperature to liberate hydrogen and form calcium hydroxide. 7

8 ~ The reaction of magnesium with water at room temperature is so slow that that the evolution of hydrogen is not observable to the human eye. ~ However, Mg reacts with steam rapidly to liberate H 2 and form magnesium oxide. Differences due to nature of the reactants

9 4. - exposed surfaces affect speed. Reactions occur at the exposed surface of the substance. Therefore, the greater surface area exposed, the greater chance of collisions between particles, so the reaction should proceed at a much faster rate. Ex. coal dust is very explosive as opposed to a piece of charcoal. (Exception: substances in the gaseous state or solution) Aqueous solutions are considered ultimate exposure

10 5. chemical that speeds up a rxn w/o being used up. It increases the rate of the rxn, but it remains unchanged when the rxn is complete A catalyst may participate in multiple chemical transformations. The effect of a catalyst may vary due to the presence of other substances known as inhibitors (which reduce the catalytic activity) or promoters (which increase the activity).

11 Step 1: SiO 2 + 2Cl 2 + Ag AgSiO 2 Cl 4 Step 2: AgSiO 2 Cl 4 + heat SiCl 4 + O 2 + Ag An is a substance that is produced and then used up during a reaction A is a substance that is there at the beginning of the reaction and also there at the end of reaction. Intermediate = Catalyst =

12 Catalysts change reaction rates by a faster alternative pathway where a different, lesser amount of is needed. Note: Since a catalyst lowers the activation energy barrier, the forward and reverse reactions are both accelerated to the same degree. 12

13 catalysts exist in same phase as the reactants. catalysts exist in different phases than the reactants. Homogeneous catalysts can actually appear in the rate law because their concentration affects the reaction. Ex: H 2 O 2 decomposes relatively slowly into H 2 O and O 2 ; however; exposure to light accelerates this process (hence stored in brown bottles) AND with the help of MnO 2, it goes extremely FAST!! 13

14 A catalytic reaction may be. It s is an acceleration of a chemical reaction by the addition of an acid or a base. Below is the decomposition of sugar (sucrose) into glucose and fructose in sulfuric acid. This is also an example of a dehydration reaction.

15 Biological catalysts are called. They are very shape specific to their substrate and have an active site that attracts the substrate using intermolecular forces. Since their shape is mostly determined by IMFs (Intermolecular forces) as opposed to actual chemical bonds. IMF s are very temperature sensitive. Altering the ph or heating the enzyme easily disrupts the IMFs and causes the enzyme to denature (lose its three dimensional shape).

16 increase in concentration of a product per unit time or decrease in concentration of a reactant per unit time series of steps by which a reaction occurs the rate of rxn is always determined by the slowest step in its mechanism (rate determining step) Ways to measure rxn speed: for gases - pressure for concentration color intensity 16

17 ~ rates at which reactants disappear or products appear the change in concentration of a reactant or product. = change in concentration / (number of moles x change in time) Note: Rate = - A - B + C + D n t n t or n t n t

18 Ex. 1) 4A + B 2C + 3D

19 Reaction Mechanism: a series of steps that must agree with the determined rate law and the sum of the elementary steps must give the overall equation for the rxn. The individual steps in a reaction mechanism are called elementary steps. Assume they happen as written, wise

20 You don t want to have intermediates in a rate expression. An is a chemical that is neither a reactant nor a product. It is formed and consumed in the course of the rxn. (It is cancelled out in the overall reaction) NO + Br 2 <-> NOBr 2 NOBr 2 + NO 2NOBr NOBr 2 is the intermediate

21 the number of species that must collide to produce the reaction indicated by that step Examples of Elementary Steps Elementary Step Molecularity Rate Law A products Unimolecular rxn involving one molecule Rate = k[a] A + A products (2A products) Bimolecular rxn involving collision of two species Rate = k[a] 2 A + B products Bimolecular Rate = k[a][b] A + A + B prod. (2A + B prod.) Termolecular rxn involving collision of 3 species Rate = k[a] 2 [B] A + B + C prod. Termolecular Rate = k[a][b][c]

22 must be determined cannot be determined from balanced equations most chemical reactions are not one-step reactions Rate law expressions are also called: rate laws rate equations rate expressions 22

23 By definition, the rate law for the reaction: 2 A + 3 B 5 C is: Rate = k A x B y k = specific rate constant at a given temp. k is temperature dependent & must be evaluated by experiment x and y are called orders of rxn. x + y + = overall order of rxn Note that rate laws are written using only the reactants. The orders of reaction, x & y, are from data, the coefficients of the equation. In fact, the only time the orders equal the coefficients is when the reactions are elementary rxns.

24 Orders of a reaction are expressed in terms of either: each or reaction The order with respect to a certain reactant is simply the exponent on its concentration term in the rate expression. We must analyze concentration & rate data to make this determination. The overall order of the reaction is simply the sum of all the exponents on all the concentration terms in the expression. 24

25 1. order: The change in concentration of reactant has no effect on the rate. These are not very common. 2. order: Rate is directly proportional to the reactants concentration; doubling [rxt], doubles rate. These are very common. Nuclear decay reactions usually fit into this category. 3. order: Rate is quadrupled when [rxt] is doubled and increases by a factor of 9 when [rxt] is tripled etc. These are common, particularly in gas-phase reactions. 4. orders are rare, but do exist! 25

26 1. Zero order: General form of rate equation: Rate = k 2. First order: General form of rate equation: Rate = k [A] 1 OR Rate = k[a] 3. Second order: General form of rate equation: Rate = k [A] 2 OR Rate = k[a] 1 [B] 1 ~ which has an overall order of two (2 nd order) 4. Fractional orders: Can be any fraction: Rate = k [A] 1/2 OR Rate = k [A] 1/4 OR Rate = k [A] 7/8 OR 26

27 For the general rate expression: Rate = k[a] m [B] n If m = 0; Say, The reaction is zero order with respect to A. If m = 1 ; Say, The reaction is 1st order with respect to A. If m = 2 ; Say, The reaction is 2nd order with respect to A. 27

28 For the general rate expression: Rate = k[a] m [B] n If n = 0 ; Say, The reaction is zero order with respect to B. If n = 1 ; Say, The reaction is 1st order with respect to B. If n = 2 ; Say, The reaction is 2nd order with respect to B. Again, summing all of the orders for each reactant gives the overall order of the reaction. 28

29 For example: 2 N O 2 5 R = k first g 2 g 2 g N 2 O 5 order in N 4NO 2 O 5 + O and first order overall 29

30 Orders of a reaction Notice that some reactants do not affect the rate of the reaction: - - CH3 CBr aq OH aq CH3 COH aq Br 3 3 aq R = k[ CH CBr] first order in CH CBr, zero order in OH, and first order overall

31 2 NO + O second order in NO, first g 2 g 2 g R = k[no] 2 NO 2 order in O O 2 2, and third order overall ALL RATE EXPRESSIONS ARE DETERMINED EXPERIMENTALLY Sometimes the coefficients from the balanced equation matches the rate law but that is not normally the case 31

32 Ex. 2) Step 1: NO + Cl 2 + Pt NOCl 2 Pt Fast Step 2: NOCl 2 Pt + NO 2NOCl + Pt Slow A. What is the overall reaction?

33 Ex. 2) Step 1: NO + Cl 2 + Pt NOCl 2 Pt Fast Step 2: NOCl 2 Pt + NO 2NOCl + Pt Slow B. What is the catalyst?

34 Ex. 2) Step 1: NO + Cl 2 + Pt NOCl 2 Pt Fast Step 2: NOCl 2 Pt + NO 2NOCl + Pt Slow C. What is the intermediate? (an intermediate is a substance that is produced and then used up during a reaction)

35 Ex. 2) Step 1: NO + Cl 2 + Pt NOCl 2 Pt Fast Step 2: NOCl 2 Pt + NO 2NOCl + Pt Slow D. What is the molecularity in step 1?

36 Ex. 2) Step 1: NO + Cl 2 + Pt NOCl 2 Pt Fast Step 2: NOCl 2 Pt + NO 2NOCl + Pt Slow E. What is the rate determining step? (For a mechanism to be consistent, the steps must add up to the overall rxn, and the rate determining step must give the derived rate law.)

37 Ex. 2) Step 1: NO + Cl 2 + Pt NOCl 2 Pt Fast Step 2: NOCl 2 Pt + NO 2NOCl + Pt Slow F. If this rxn is 1 st order [NOCl 2 Pt] and 1 st order [NO] because they are the reactants in the rate determining step, what is the rate law for this expression?

38 Ex. 2) Step 1: NO + Cl 2 + Pt NOCl 2 Pt Fast Step 2: NOCl 2 Pt + NO 2NOCl + Pt Slow G. What is the overall reaction order?

39 Rate of a simple one-step reaction is directly proportional to the concentration of the reacting substance [X] is concentration of X in molarity or moles/l A(g) B(g) + C(g) R A or R = k A For a simple expression like R = k[a]: doubling the initial concentration of A doubles the initial rate of reaction 2 1 = 2 halving the initial concentration of A halves the initial rate of reaction (1/2) 1 = 1/2 39

40 Look at the following and its determined rate-law expression 3 A B C 2 g g g Rate = k A because it is a order rate-law expression doubling the [A] increases the rate of reaction by a factor of = 4 halving the [A] decreases the rate of reaction by a factor of 4 (1/2) 2 = 1/4 40

41 Ex. 3) For the reaction 3A + B C, the rate law is R = k [A] 2 A. What order is the reaction?

42 Ex. 3) For the reaction 3A + B C, the rate law is R = k [A] 2 B. How will these change the rate if you triple the concentration of A

43 Ex. 3) For the reaction 3A + B C, the rate law is R = k [A] 2 C. Double the concentration of B

44 The orders of the reactions are derives by the following mathematical equation: ratios of the rates = ratios of A x ratios of B y Or rate 1 = k[a] x 1[B] y 1[C] z 1 rate 2 = k[a] x 2[B] y 2[C] z 2 (we ll solve some examples later)

45 Reactant How [M]Δ Effect on rate Order A Doubled no change _ A Doubled doubles _ A Doubled quadruples _ A Doubled eight times _

46 Order [ΔM] order =Δrate = _ = _ = _ = _

47 Or in other words: If concentration doubles (2) and the rate stays the same: ask yourself ~ 2 to what power = 1 so 2 X = 1 ( the power is 0 so the reaction is in A) Conc. doubles & rate doubles: 2 X = 2 ( x = 1 or ) Conc. doubles & rate quadruples: 2 x = 4 ( x = 2 or )

48 Three basic events must occur for a reaction to occur the atoms, molecules or ions must:

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50 The collisions must occur with enough energy to overcome the electron/electron repulsion of the valence shell electrons of the reacting species and must have enough energy to transform translational energy into vibrational energy in order to penetrate into each other so that the electrons can rearrange and form new bonds. These collisions are referred to as.

51 This new collision product is at the peak of the activation energy hump and is called the or the transition state. At this point, the activated complex can still either fall to reactants or fall back to the products.

52 For reactions that have an E a 50 kj/mol, (E a = activation energy ) the rate doubles for a 10 0 C rise in temperature, near room temperature. 2 ICl(g) + H 2 (g) I 2 (g) + 2 HCl(g) Through experiment thee rate-law expression is known to be R=k[ICl][H 2 ] A t C, k = M s A t C, k = M s k a p p r o x im a te ly d o u b le s 52

53 You can mark the position of activation energy Ea, on a Maxwell-Boltzmann distribution to get a diagram like this one. If the temperature is increased, molecules speed up and collide more frequently with more energetic collisions. Only those particles represented by the area to the right of the activation energy mark will react when they collide. The great majority don't have enough energy, and will simply bounce apart. 53

54 Use experimental rate-law to The slowest step in a reaction mechanism is the Note: Experimentally determined reaction orders indicate the number of molecules involved in: the slow step only or the slow step and the equilibrium steps preceding the slow step. 54

55 Consider the iodide ion catalyzed decomposition of hydrogen peroxide to water and oxygen. 2 H O 2 H O + O 2 2 l 2 l 2 g Reaction is known to be first order in H 2 O 2, first order in IO -, and second order overall. Rate = k [H 2 O 2 ][IO - ] I - 55

56 Rate Mechanisms is thought to be: Fast step H Slow step 2 O IO I + H - 2 O 2 IO - + H H 2 2 O O + O 2 + I - Overall rxn. 2 H Experimental rate law 2 O 2 2 H 2 R = k O + O - H O IO Notice that the rate law matches the slowest step. 56

57 Important notes: one hydrogen peroxide molecule and one hypoiodite ion are involved in the rate determining step the iodide ion catalyst is consumed in step 1 and produced in step 2 in equal amounts hypoiodite ion has been detected in reaction mixture as a short-lived 57

58 Ozone, O 3, reacts very rapidly with nitrogen oxide, NO, in a reaction that is first order in each reactant and second order overall. O +NO NO +O 3 g g 2 g 2 g Experimental rate-law R= k O NO 3 58

59 A mechanism is: Slow step O + NO Fast step O + NO Overall rxn. O + NO NO + O 3 3 NO + O NO + O Notice that the rate-law matches the slowest step Rate = k[o 3 ][NO] 59

60 This mechanism is with the ratelaw expression b/c the slowest step doesn t match the rate-law found in the lab: R = k [O 3 ] [NO] Slow step O O + O 3 2 Fast step O + NO NO 2 Overall rxn. O + NO NO O Rate - law from this mechanism R = k O 3 cannot be correct 60

61 By the way, if you ve never listened to a screencast of your AP Chemistry notes, now may be the time! Go to find a quiet place and enjoy!

62 data table contains concentration and rate data. Use table logic or algebra to determine the orders of reactants and the value of the rate constant, k. data table contains concentration and time data. Use graphical methods to determine the order of a given reactant. The value of the rate constant k is equal to the absolute value of the slope of the best fit line Can be decided by performing 3 linear regressions and analyzing the regression correlation coefficient r. Not nearly as hard as it sounds! Basically, you are in search of linear data. 62

63 Integrated rate equation relates time and concentration for a chemical or nuclear reaction Why do we need this? The problem with normal rate law is that you can t figure out the rate at some later time. First Order Reactions 1st order in reactant A & 1st order overall For example: a A products common for many chemical reactions and all simple radioactive decays 63

64 ln [A] t ln [A] 0 = - akt where: [A] 0 = mol/l of A at time t=0. [A] t = mol/l of A at time t. k = specific rate constant t = time elapsed since beginning of reaction a = stoichiometric coefficient of A in balanced overall equation (not really used on the AP Exam but it is there ) 64

65 1 1_ [A] t [A] 0 = kt Same as 1 st order: [A] 0 = mol/l of A at time t=0. [A] t = mol/l of A at time t. k = specific rate constant t = time elapsed since beginning of reaction 65

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67 Reminder: To find units for k, the specific rate constant, look at overall order Overall Order Units 0 [M] time -1 1 time -1 2 [M] -1 time -1 3 [M] -2 time -1 For homework due tomorrow. Copy Table14-3, pg. 662 and Figures , pgs

68 The Chernobyl nuclear reactor accident occurred in At the time that the reactor exploded some 2.4 MCi of radioactive 137 Cs into the atmosphere. The half-life of 137 Cs is 30.1 years. In what year will the amount of 137 Cs released from Chernobyl finally decrease to 100 Ci? A Ci is a unit of radioactivity called the Curie, MCi = MegaCurie 68