1. AH Chemistry – Unit 1 Kinetics

2. How fast does it go? Thermodynamics Is the reaction feasible? How far will the reaction go? Thermodynamics is about start and finish Kinetics is about what happens along the way.

3. Kinetics Studies the rate at which a chemical process occurs. Besides information about the speed at which reactions occur, kinetics also sheds light on the reaction mechanism (exactly how the reaction occurs).

4. Measuring reaction rates Properties suitable for measuring rate The mass of the apparatus (when a gas is being released) The volume of a gaseous product The pH of a solution (when H + or OH - ions are used or produced) The conductivity of a solution (when the number or nature of the ions present changes) The colour (when products and reactants have different colours)

5. Measuring reaction rates Time/s Concentration of reactants mol l -1

6. Measuring reaction rates Time/s Concentration of products mol l -1

7. Reaction Rate C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) Note that the average rate decreases as the reaction proceeds. This is because as the reaction goes forward, there are fewer collisions between reactant molecules. We will study 2 examples – one on the PPT and one in your notes

8. Reaction Rates A plot of concentration vs. time for this reaction yields a curve like this. The slope of a line tangent to the curve at any point is the instantaneous rate at that time. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq)

9. Reaction Rates and Stoichiometry In this reaction, the ratio of C 4 H 9 Cl to C 4 H 9 OH is 1:1. Thus, the rate of disappearance of C 4 H 9 Cl is the same as the rate of appearance of C 4 H 9 OH. C 4 H 9 Cl (aq) + H 2 O (l)  C 4 H 9 OH (aq) + HCl (aq) Rate = -  [C 4 H 9 Cl]  t =  [C 4 H 9 OH]  t

10. Reaction Rates and Stoichiometry What if the ratio is not 1:1? H 2 (g) + I 2 (g)  2 HI (g) Only 1/2 HI is made for each H 2 used.

11. Reaction Rates and Stoichiometry To generalize, for the reaction aA + bBcC + dD Reactants (decrease) Products (increase)

12. The Rate Law aA + bB → cC + dD Rate  [A] m.[B] n Rate = k[A] m.[B] n Where… k = rate constant m and n = orders of reaction with respect to A and B

13. Rate Laws A rate law shows the relationship between the reaction rate and the concentrations of reactants. – For gas-phase reactants use P A instead of [A]. k is a constant that has a specific value for each reaction. The value of k is determined experimentally. “Constant” is relative here- k is unique for each reaction k changes with temperature.

14. order of a reaction The power to which the concentration of a particular reactant is raised in the rate equation overall order of a reaction The sum of the powers to which the concentrations of all reactants in the rate equation are raised in the rate equation

15. Order for the reaction is the power to which the conc. of that reactant is raised in the rate equation Rate equation Order of reaction Rate = k [ A ] 00 Rate = k [ A ] 11 Rate = k [ A ] 22 Rate = k [ A ] 1 [ B ] 12 Rate = k [ A ] 1 [ B ] 23 Units for k depend on the overall rate of the reaction.

16. Small integral values refer to the number of molecules in the single step of the reaction which controls the overall reaction rate – the RATE DETERMINING STEP. These do not necessarily refer to the number of molecules in the balanced equation. Order cannot be deduced from the balanced equation. It can only be determined experimentally

17. Rate equation (from experiment) Overall order of reaction = H 2 O 2 + 2HI  H 2 O + I 2 Rate [H 2 O 2 ][HI] 2 NOTE: Balanced equation has 3 reactant molecules Order of reaction has been experimentally determined

18. ZERO ORDER Changing the concentration has no effect on the rate of the reaction. Rate = k [reactant] 0  rate = k

19. FIRST ORDER Doubling the concentration doubles the rate. Rate = k [reactant] 1  rate = k [R] Tripling the concentration triples the rate.

20. SECOND ORDER Doubling the concentration quadruples (4x) the rate. Rate = k [reactant] 2  rate = k [R] 2 k – the value for k is found experimentally Its value depends on the experimental conditions e.g. varies with temperature

21. Determining Order A + B + C  D (g) Expt. no [A] [B] [C] (mol l -1 s -1 ) Initial rate (ml of D s -1 ) 1 1.0 1.0 1.020 2 2.0 1.0 1.040 3 1.0 2.0 1.020 4 1.0 1.0 2.080 Rate Law/equation  Rate = k [A] [B] 0 [C] 2 Rate Law/equation  Rate = k [A] [C] 2

22. Orders of reaction Reactions usually occur in a series of steps. The integers in the rate equation refer to the actual number of particles involved in the slowest step of the reaction. This is the rate determining step.

23. Calculating the Rate Constant Can be calculated from a series of experiments in which the starting concentrations of reactants are varied…

24. Example

28. Reaction rates and reaction mechanisms