1. Kinetics How fast does a reaction (event) occur? Reaction rates are controlled by: Nature of reactants Ability of reactants to meet Concentration of reactants Temperature Presence of a catalyst Rate of pay = €10/hour UNITS: mol/L x 1/s =mol.L -1.s -1 or M.s -1

2. Kinetics Change of reaction rate with timeConcentration and rate A + B  products In general it is found that: rate  [A] m [B] n The values of the exponents, m and n, must be determined empirically (by experiment). We can replace  by = if we introduce a rate constant, k. Rate = k [A] m [B] n This expression is the rate law

3. Rate Laws Example: H 2 SeO 3 + 6I - + 4H +  Se + 2I 3 - + 3H 2 O Rate = k[H 2 SeO 3 ] x [I - ] y [H + ] z Experimentally found that x=1, y=3, z=2 Rate = k[H 2 SeO 3 ][I - ] 3 [H + ] 2 At 0C, k=5.0 x 10 5 L 5 mol -5 s -1 (units of rate constant are such that the rate has units of mol.L -1.s -1 ) Notice that exponents in rate law frequently are unrelated to reaction stoichiometry. Sometimes they are the same, but we cannot predict this without experimental data! Exponents in the rate law are used to describe the order of the reaction with respect to each reactant. The overall order of a reaction is the sum of the orders with respect to each reactant.

4. Determining exponents in a rate law One way to do this is to study how changes in initial concentrations affect the initial rate of the reaction Initial Concs [A] [B] Initial rate (mol L -1 s -1 ) 0.10 0.20 0.100.40 0.300.100.60 0.300.202.40 0.30 5.40 A + B  products Rate = k [A] m [B] n 1-3: [B] is constant. Rate changes due only to [A] m must be 1 3-5:[A] is constant. When [B] is doubled, rate increases by factor of 4 (=2 2 ). When [B] is tripled, rate increases by factor of 9 (=3 2 ). n must be 2

5. Concentration and Time-1 st order reactions Rate = -k[A] Integrated rate law Half-life: time required for half of initial concentration of reactant to disappear. t 1/2 = ln2/k

6. Concentration and Time-2nd order reactions Simplest 2 nd order: 2A  B Rate = k[A] 2 Integrated rate law Half-life t 1/2 = 1/k[A] 0 Half-life depends on initial concentration

7. Temperature dependence of reaction rates Transition-state theory is used to explain what happens when reactants collide. Most often a change in momentum or direction simply occurs. Sometimes a reaction occurs. Only combining molecules that have kinetic energies at least as large as the activation energy can surmount the barrier and produce products. The difference in potential energy of products to reactants is the heat of reaction (exothermic in this case). PE decreases, therefore KE and T increase! The Arrhenius equation relates the activation energy to the rate constant

8. Catalysis

9. Ozone depletion

10. Radio-activity Unstable atomic nuclei may decay by emitting particles that are detected with special counters. Alpha, beta, and gamma emission are common types of radioactivity. In beta decay the emitted particles are electrons; in alpha decay they are helium nuclei, and in gamma decay they are high energy photons. Counters can be sensitive to either alpha, beta, or gamma-ray particles. The rubidium isotope 37 Rb 87 decays by beta emission to 38Sr87, a stable strontium nucleus: 37 Rb 87  38 Sr 87 + . From the following experimental data, calculate (a) the rate constant and (b) the half-life of the rubidium isotope. From a 1.00 g sample of RbCl which is 27.85% 37 Rb 87, an activity of 478 beta counts per second was found. The molecular weight of RbCl is 120.9 g mole -1.