Chem 402 Lecture 26: L26 3:27:17

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27 March 2017 L26: Thermodynamics vs. Kinetics I. Intro to Kinetics A. Kinetics 1. Beginning Example 0 0 a. H 2(g) ½ O 2(g) H O2∆𝐺(l)∆𝐺𝑟(𝐻 𝑂 ) 𝑓 −227 (𝑙)𝑚𝑜𝑙 b. C (s, diamond) (s, graphit𝑟)−2.9 𝑘𝐽/𝑚𝑜𝑙 • Diamonds are not really forever, will thermodynamically favorably turn into graphite • However, thermodynamics tells us how things should go, but nothing about rates or spontaneity 2. Kinetics a. The study of reaction rates and mechanisms of reactions 3. Quantifying a rate of reaction a. For a reaction: aA + bB  cC + dD (lower case = stoichiometric coefficients) b. 𝑅𝑎𝑡𝑒 = − 1 𝑑 𝐴, where d[A] is the change in concentration of reaction species 𝑎 𝑑𝑡 • Negative because we are using up a, so 𝑑 𝐴]< 0, so the negative in front 𝑑𝑡 makes the rate positive • Can be put in terms of B, C, or D 𝑑 𝑝𝑟𝑜𝑑𝑢𝑐𝑡] • Note: if in terms of a product, > 0, so no negative 𝑑𝑡 4. Units of Rate of Reaction a. Rate is always in concentration per time b. Ex: molL s -1 -1 th B. Rate Laws (in terms of basics and 0 order) 1. Differential Rate Law a. Rate ∝ 𝐴 [ ] [ 𝐵 𝑛𝐵 • Rate is proportional to the concentration of reactants • nAand n Bre the orders of reaction with respect to the reactants, aka the partial orders b. In most cases there’s no relation between the reaction order and stoichiometric coefficient (different concepts) c. In special cases, for elementary reactions, the stoichiometric coefficients equal the orders of reaction • However, if the stoichiometric coefficients equal the orders of reaction, it doesn’t necessarily mean that it is an elementary reaction d. The proportionality constant (k) is the rate constant • Rate = 𝑘 𝐴 ] [ 𝐵 𝑛𝐵 • “k” tells us that only a fraction of collisions lead to reaction e. Factors that contribute to reaction: orientation of reactants, energy requirement for reaction, temperature, mechanism (how reaction goes) f. The units of “k” depends on the orders of reaction with respect to reactants • Overall order = ∑ 𝑖𝑛𝑖(𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠) 2. 0 Order Differential Rate Law a. If ∑ 𝑛