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CHEM1612 - Pharmacy Week 11: Kinetics - Rate Law Dr. Siegbert Schmid School of Chemistry, Rm 223 Phone: 9351 4196

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Presentation on theme: "CHEM1612 - Pharmacy Week 11: Kinetics - Rate Law Dr. Siegbert Schmid School of Chemistry, Rm 223 Phone: 9351 4196"— Presentation transcript:

1 CHEM Pharmacy Week 11: Kinetics - Rate Law Dr. Siegbert Schmid School of Chemistry, Rm 223 Phone:

2 Unless otherwise stated, all images in this file have been reproduced from: Blackman, Bottle, Schmid, Mocerino and Wille, Chemistry, John Wiley & Sons Australia, Ltd ISBN:

3 Lecture Blackman, Bottle, Schmid, Mocerino & Wille: Chapter 14 KINETICS: the study of REACTION RATES and their relation to the way the reaction proceeds, i.e., its MECHANISM. Thermodynamics tells whether a reaction favours products or reactants (i.e. relative stabilities), but gives us no information on HOW FAST the reaction goes from reactants to products, e.g. Chemical Kinetics H 2 should react with O 2 (ΔH° = –286 kJ mol -1 ) At RT the reaction is spontaneous and K = 3.6 x !  But no reaction occurs!!!! 2 H 2 + O 2 2 H 2 O

4 Lecture Factors affecting reaction rate Rate is proportional to collision rate which is proportional to Figure from Silberberg, “Chemistry”, McGraw Hill,  Concentration of some or all of the molecules present  Physical state: reactants need to mix to collide  Temperature: the higher T, the more energetic the collisions, the faster the reaction  Pressure (similar to concentration)  Presence of a catalyst

5 Lecture Rate of a Reaction The rate of a reaction is the change in concentration of one of the reactants that occurs during a given period of time. -Δ[A] Δt Rate = = Figure from Silberberg, “Chemistry”, McGraw Hill, Δ[B] Δt

6 Lecture Average reaction rate = – Δ[A] Δ t Time (s) [A] (mol L -1 ) Ave. Rate (mol L -1 s -1 ) The reaction rate varies with time as the reaction proceeds. Average rate is not constant. Rate of a Reaction

7 Lecture An infinitesimally small change in the concentration, d[A], that occurs over the infinitesimally short period of time, dt, gives the instantaneous rate of reaction. You can work out that rate for any moment in time by determining the slope of a tangent drawn to the concentration-time curve at that exact moment. Rate of a Reaction - d[A] dt Rate 50s =

8 Lecture Expressing Reaction Rates For a generic chemical reaction the reaction rate is defined as: A + C → 2 B (1) (2) Expression 2 is just a rearrangement of 1, but its numerical value for the rate is double that of (1). The expression and its numerical value depend on the reactant taken as reference.

9 Lecture Express the rate in terms of the change in concentration with time of each substance for the reaction: 2 N 2 O 5 → 4 NO 2 + O 2 Rate of production of O 2 = 2.6·10 -6 M s -1. Rate of production of NO 2 = 4 × 2.6·10 -6 = 1.0·10 -5 M s -1 Rate of consumption of N 2 O 5 = - 2 × 2.6 · = - 5.2·10 -6 M s -1 Expressing Reaction Rates

10 Lecture Expressing Reaction Rates a A +b B → c D + d D In practice, you will commonly choose as a reference the species that appears with stoichiometric coefficient of 1.

11 Lecture Expressing Reaction Rates Express the rate of reaction in terms of concentration of reactants and products for the reaction: 4 NH 3 (g) + 5 O 2 (g)  4 NO (g) + 6 H 2 O (g) Solution: Rate of reaction

12 Lecture The concentrations of N 2 O 5 are 1.24 ·10 -2 and 0.93 · M at 600 s and 1200 s after the reactants are mixed at the appropriate temperature. Calculate the reaction rates for 2 N 2 O 5 → 4 NO 2 + O 2 What is the rate of formation of the products? rate of formation of NO 2 = (2 × rate N 2 O 5 ) =1.0 · M s -1. rate of formation of O 2 = (0.5 × rate N 2 O 5 ) = 2.6 x M s -1. Solution: Rate of decomposition of N 2 O 5 = Example 1

13 Lecture Example 2 Express the rate in terms of the change in concentration with time of each substance for the reaction: 2 O 3 → 3 O 2 Answer: If the rate at which O 2 appears is 6·10 -5 Ms -1, at what rate is O 3 disappearing at the same time?

14 Lecture The Iodine Clock Mix different amounts of HIO 3 + NaHSO 3 + starch. Concentration of reactants is: [beaker I] > [beaker II] >[beaker III]. The following reactions take place consecutively in each beaker: Starch forms a blackish blue complex with iodine. As the final reaction is the fastest, the colour of the elemental iodine only becomes apparent once the sulphite is fully consumed. The reaction is slowest in the solution with the lowest concentration, as the reaction time is dependent on the concentration.

15 Lecture Rate Law Expresses the rate as a function of reactant concentrations and T. For a generic reaction: aA + bB + …→ cC + dD + …. The rate law has the form: rate = k [A] m [B] n ….. k = rate constant, is independent of conc. but increases with T m,n,… reaction orders; if the rate doubles for doubling of [A], m = 1 In general m, n,… ≠ a, b, c, …

16 Lecture Rate Law Rate of hydrolysis of cis-platin is proportional to [Pt(NH 3 ) 2 Cl 2 ] We express this as a RATE LAW Rate laws can be determined ONLY experimentally, they cannot be deduced by reaction stoichiometry. Rate of reaction = k [Pt(NH 3 ) 2 Cl 2 ] Hydrolysis of cisplatin [Pt(NH 3 ) 2 Cl 2 ](aq) +H 2 O(l)  [Pt(NH 3 ) 2 (H 2 O)Cl](aq) + Cl - (aq)

17 Lecture Experimental Tools Many methods are available to monitor reaction rates, e.g.: Spectrometric Methods (measure light adsorbed by a reactant or product) Conductometric Methods (measure change in conductivity during reaction) Manometric Methods (Monitor the change in pressure over time, at constant V, T) Direct Chemical Methods (a small aliquot of reaction mixture is sampled, cooled down, and titrated)

18 Lecture For the general reaction: a A + b B + c C …  d D + e E …. rate = k [A] m [B] n [C] o … m is the order of the reaction with respect to A (or “in” A), n is the order of the reaction with respect to B… Overall order of the reaction is = m + n + o +…. e.g. if rate = k [A] 2 [B], then the reaction is second order with respect to A, first order with respect to B, and overall third order. Reaction orders cannot be deduced from the balanced reaction. Reaction Orders

19 Lecture Reaction Orders For most reactions the order is a small positive integer or zero, but also: Fractional number: CHCl 3 (g) + Cl 2 (g) → CCl 4 (g) + HCl (g) Rate = k [CHCl 3 ] [Cl 2 ] ½ Negative number: 2 O 3 (g) → 3 O 2 (g) Rate = k [O 3 ] 2 [O 2 ] -1 = k [O 3 ] 2 / [O 2 ]

20 Lecture What is the order of reaction with respect to NO, O 3, and the overall order of reaction for the reaction: NO (g) + O 3 (g)  NO 2 (g) + O 2 (g) Rate = k [NO] [O 3 ] Answer: First order with respect to NO and O 3, overall second order (1+1). Reaction Orders

21 Lecture What order is the following reaction? H 2 (g) + 2 ICl (g)  2 HCl (g) + I 2 (s) The reaction order can be determined ONLY by experiment. Rate= -d[H 2 ] / dt= k [H 2 ][ICl] = k [H 2 ] 1 [ICl] 1 This reaction is first order with respect to H 2, first order with respect to ICl and second order overall. Reaction Orders

22 Lecture Reaction Orders Express the rate in terms of the change in concentration with time of each substance for the reaction: 2 NO (g) + 2 H 2 (g)  N 2 (g) + 2 H 2 O (g) What is the order of reaction with respect to NO, H 2 and the overall order of reaction for the reaction: Rate = k [NO] 2 [H 2 ] The reaction is second order with respect to NO, first order with respect of H 2, overall third order (2+1).


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