1. Unless otherwise stated, all images in this file have been reproduced from: Blackman, Bottle, Schmid, Mocerino and Wille, Chemistry, 2007 (John Wiley) ISBN: 9 78047081 0866

2. Slide 2/18 e CHEM1002 [Part 2] A/Prof Adam Bridgeman (Series 1) Dr Feike Dijkstra (Series 2) Weeks 8 – 13 Office Hours: Monday 2-3, Friday 1-2 Room: 543a e-mail: adam.bridgeman@sydney.edu.au e-mail: feike.dijkstra@sydney.edu.au

3. Slide 3/18 e Complexes IV The metals include many essential elements (such as Na +, K + and Ca 2+, some toxic (such as Hg 2+ and Al 3+ ) and some which are now being used in medicines (such as Pt 2+ ) The essential metals have a variety of functions in the body (such as Na + /K + in the nervous system, Fe 2+/3+ in oxygen transport and Zn 2+ in CO 2 transport) The biological function is related to the oxidation number, the coordination type and the size of the atom Summary of Last Lecture

4. Slide 4/18 e Lecture 16 Chemical Kinetics Rate of Reaction Rate Laws Reaction Order Blackman Chapter 14, Sections 14.1 - 14.3 Lecture 17 Half lives The Temperature Dependence of Reaction Rates Catalysis Blackman Chapter 14, Sections 14.4 - 14.6 Chemical Kinetics I

5. Slide 5/18 e Kinetics vs Thermodynamics Thermodynamics (ΔG, Δ univ S, E o ) tells us if a reaction favours the products or reactants It also gives the extent a reaction occurs (K eq ) Thermodynamics says nothing about how fast or slow the reaction goes It gives us the equilibrium concentrations but not how long it takes to get to equilibrium

6. Slide 6/18 e Reaction Rate The rate of a reaction is how fast the concentration of the molecules present change. Reaction rate: change in concentration of a product or a reactant per unit time. Rate is given by the gradient of concentration vs time graph d[A] = — — dt reaction rate

7. Slide 7/18 e Expressing Reaction Rates There are a number of ways to express the rate. hydrolysis of cisplatin [Pt(NH 3 ) 2 Cl 2 ] + H 2 O  [Pt(NH 3 ) 2 (H 2 O)Cl] + + Cl - d[reactant] dt rate = d[product] — + =

8. Slide 8/18 x Express the rate of reaction for each reactant and product in the reaction: 4NH 3 (g) + 5O 2 (g)  4NO(g) + 6H 2 O (g) d[NH 3 ] dt NH 3 : d[O 2 ] dt O2:O2: d[NO] dt NO: d[H 2 O] dt H 2 O: Expressing Reaction Rates - 1 4 - 1 5 + 1 4 + 1 6

9. Slide 9/18 e Rates are Determined Experimentally d[O 3 ] C 2 H 4 + O 3  C 2 H 4 O + O 2 Rate is dependent of concentration of O 3 dt rate = - fast slow

10. Slide 10/18 e k rate = k [O 3 ] [C 2 H 4 ] The Rate Law The rate of the reaction is proportional to the concentration [O 3 ] and to that of [C 2 H 4 ] This is expressed as a rate law: k increases with T is the rate constant and is independent of concentration

11. Slide 11/18 e For the general reaction: The Rate Equation aA + bB + cC …  mM + nN …. dt = k [A] x [B] y [C] z … -d[A] the rate law can only be determined by experiment, not from the stoichiometric equation x is the order of the reaction with respect to A, y is the order of the reaction with respect to B… the overall order of the reaction is given by x + y + z … there is no relationship between a and x, b and y ….

12. Slide 12/18 x Example For the reaction: 2 NO (g) + O 2 (g)  2 NO 2 (g) rate = k[NO] 2 [O 2 ] What is the order of reaction with respect to the reactants and the overall order of reaction? second order with respect to NO first order with respect to O 2 third order overall

13. Slide 13/18 e Interpreting Rate Laws If x = 1, reaction is 1 st order in A: rate α [A] 1 If [A] doubles, then rate goes up by factor of = k [A] x [B] y [C] z … rate If x = 2, reaction is 2 nd order in A: rate α [A] 2 If [A] doubles, then rate goes up by factor of If x = 3, reaction is 3 rd order in A: rate α [A] 3 If [A] doubles, then rate goes up by factor of two four eight 16

14. Slide 14/18 x Using Data to Determine Order ClO 3 - (aq) + 9 I - (aq) + 6 H + (aq)  Cl - (aq) + 3 I 3 - (aq) + 3 H 2 O(l) rate = k [ClO 3 - ] x [I - ] y [H + ] z 40.200.200.200.80 30.200.200.100.20 20.200.100.100.10 10.100.100.100.05 [ClO 3 - ] / M [I - ] / M [H+] / M initial rate / M s -1 x = 1, y = 1 and z = 2 so

15. Slide 15/18 x Using Data to Determine Order Calculate the rate constant, k: in experiment 1: rate = 0.05 M s -1, [ClO 3 - ] = 0.1 M, [I - ] = 0.1 M, [H + ] = 0.1 M ClO 3 - (aq) + 9 I - (aq) + 6 H + (aq)  Cl - (aq) + 3 I 3 - (aq) + 3 H 2 O(l) rate = k [ClO 3 - ] [I - ] [H + ] 2 so, (units of k depend on overall order) 0.05 M s -1 = k (0.10 M)(0.10 M)(0.10 M) 2 k = 5.0 x 10 2 M -3 s -1

16. Slide 16/18 x Question NO 2 (g) + CO (g)  NO (g) + CO 2 (g) [NO 2 ] / M[CO] / M Initial rate / M s -1 10.100.10 0.0050 20.400.10 0.080 30.100.20 0.0050 determine the rate equation and value of the rate constant for this reaction (zero order in [CO]) k = 0.5 M -1 s -1 rate = k[NO 2 ] 2

17. Slide 17/18 e Summary: Chemical Kinetics I Learning Outcomes - you should now be able to: Work out the order of the reaction with respect to each reactant from experimental data Work out the rate constant (including its units) from experimental data Hence, write down the rate law Answer review problems 14.66 - 14.97 in Blackman Complete the worksheet Next lecture: Half lives, temperature dependence and catalysis

18. Slide 18/18 x Practice Example The data given in the table below were obtained for the reaction between nitric oxide and chlorine at 1400 K. 2NO(g) + Cl 2 (g)  2NOCl(g) Experiment Number Initial [Cl 2 ] (mol L –1 ) Initial [NO] (mol L –1 ) Initial Reaction Rate (mol L –1 s –1 ) 10.10 0.18 20.200.100.36 30.100.200.72 Deduce the rate law for this reaction. Calculate the value of the rate constant.