CHE1031 Lecture 10: Reaction kinetics Lecture 10 topics Brown chapter 14 1. Reaction rates Factors that effect reaction rates 14.1 Visualizing rates & units 14.2 Average reaction rates Instantaneous reaction rates Stoichiometry & reaction rates 2. Concentration & reaction rates 14.3 Rate laws Reaction orders 3. Change in concentration with time 14.4 First- & second-order reactions Half-life 4. Temperature & reaction rate 14.5 Collision, orientation & Ea 5. Reaction mechanisms 14.6 Elementary Multistep 6. Catalysis ` 14.7

Concentration affects kinetics Orders of reactions Rate laws The general process of advancing scientific knowledge by making experimental observations and by formulating hypotheses, theories, and laws. It’s a systematic problems solving process AND it’s hands-on….. Experiments must be done, data generated, conclusions made. This method is “iterative”; it requires looping back and starting over if needed. [Why do you think they call it REsearch?] Often years, decades or more of experiments are required to prove a theory. While it’s possible to prove a hypothesis wrong, it’s actually NOT possible to absolutely prove a hypothesis correct as the outcome may have had a cause that the scientist hasn’t considered.

Rates laws Rate laws – mathematical equations that describe how reaction rate is proportional to concentration of reactants Where: k is the rate constant – determines how temp. affects rate m & n are reaction orders – small whole numbers aA + bB  cC + dD Rate = k[A]m[B]n NH4+1 + NO2-1  N2 + 2H2O Rate = k[NH4+1][NO2-1] So: As the conc of NH4 doubles so does the reaction rate As the conc of NO2 doubles so does the reaction rate Exp’t Initial [NH4+1] [NO2-1] M/s (x10-7) 1 0.0100 0.200 5.4 2 0.0200 10.8 3 0.0400 21.5 4 0.0202 5 0.0404 21.6 6 0.0808 43.3 Given the data in the table, calculate k. 5.4x10-7 M/s = k(0.0100)(0.200) k = 2.7x10-4/M/s …now any conc can be calc. p. 563 - 5

Reaction orders NH4+1 + NO2-1  N2 + 2H2O Rate = k[NH4+1][NO2-1] For this reaction the exponents m & n are both 1 and so not explicitly shown. We therefore say that the reaction is first order for both reactants. The overall order of reaction is the sum of the individual orders, so 2 in this case. Reaction orders must be experimentally determined & are hard to predict. CHCl3 + Cl2  CCl4 + HCl Rate = k[CHCl3][Cl2]1/2 2NO + 2H2  N2 + 2H2O Rate = k[NO]2[H2] So reaction orders can be either whole numbers or fractions. Looking at the second reaction: What are the reaction orders of each reactant? What’s the overall reaction order? How does doubling the concentration of NO change the rate? How does doubling the concentration of H2 change the rate? What are units of rate constants? Rate = M/s concentration = M k = M/s = s-1 M p. 565 - 6

Rate laws & initial data A + B  C Exp’t [A] [B] Initial Rate (M/s) 1 0.100 4.0x10-5 2 0.200 3 16.0x10-5 Determine the rate law for the reaction. Determine the rate constant. Calculate rate when [A] = 0.050 M and [B] = 0.100 M Rate = k[A]m[B]n Reactant A – the rate increases by 4X as concentration of A doubles, so m = 2. Reactant B – the rate doesn’t change when the conc of B doubles, so n = 0. Rate = k[A]2[B]0 b) k = rate/[A]2 = 4.0x10-5/(0.100)2 = 4.0x10-3 M-1 s-1 c) Rate = (4.0x10-3 M-1s-1)(0.050)2(0.100)0 = 1.0x10-5 M/s p. 568 - 9