Fun With Chemical Kinetics Duncan, EJ, Bennah

12.1 Reaction Rates Chemical Kinetics- the speed (rate) at which chemical changes occur. Reaction Rate can be defined as the change in concentration of a reactant or product per unit time: Rate= Concentration of A at time t2 -- concentration of A at time t1 t2 -- t1 Basically, it is the change in concentration (mol or L) over the change in time. This rate will be defined as mol/L(s)

12.1 Continued Ex: 2NO2(g)= 2NO(g)+O2(g) Nitrogen Oxide decomposition - [NO2]t=50-[NO2]t=0 50s-- 0s = 4.2 x 10-5 mol/L(s) - 0.0079 mol/L -- 0.0100 mol/L 50s

Instantaneous Rate The rate of reaction is not constant but DECREASES OVER TIME!!!!!! You can find the rate at any time (instantaneous rate) by finding the slope of a tangent line to the curve at that point. slope of tangent line= change in y change in x

Reaction rate in terms of the product When finding reaction rates with the product, you must take into account coefficients. This lets you know at what relative rate things are being consumed and produced. Slope of the curve for product or product with a coefficient of 2 will be twice that of the product with a coefficient of 1.

Fun Graph

12.2 Forward and reverse reactions Chemical reactions are reversible. in the NO2 reaction, when enough O2 and 2NO are formed, causing the reverse reaction to be important. Here the rate depends on the change in concentrations in both reactions instead of one. This can be avoided by choosing certain conditions. We can say Rate= k[NO2]n We do say this, so we only deal with concentrations of reactants. This shows how rate depends on concentration, also known as Rate Law. k is the Rate constant and n is the is called the order. These must be determined by the experiment and the order will most likely be an integer.

What this means I’m going to draw this on the board but basically- We define reactions in terms of reactants. Things to know: Rate= 2 x rate’ k[NO2]n= 2k’[NO2]n k = 2 x k’ This is why we have to be careful to define what rate we are using. The book always uses rate of reactants unless otherwise specified.

Integrated Rate Law Differential Rate Law (we just call it rate law) is how the rate depends on concentration. The Integrated Rate Law describes how the concentrations depend on time. The two kind of rate laws are always related. Which one we use will be determined by how the data will be collected easiest. We can figure one out from the other. Rate law is important to understand so we can infer individual steps in the reaction from a specific rate law by working backwards.

12.3 Determining the Form of the Rate Law The initial rate of reaction is the instantaneous rate determined just after the reaction begins. Rate= -(change in concentration)/(change in time) = k[compound 1]n[compound 2]m The overall reaction order is the sum of n and m. k represents the rate constant

12.4 The Integrated Rate Law First-order reactions are reactions in which the rate of reaction depends on the concentration of the reactant. The integrated rate law expresses the concentrations of reactants as a function of time. The integrated first-order rate law is: ln[compound] = -kt + ln[compound]

12.4 The Integrated Rate Law cont. The time required for a reactant to reach half of its original concentration is called the half-life of a reactant (t1/2). Second-order rate law: Rate= k[A]2 Integrated second-order rate law: 1/[A] = kt + 1/[A]0 Zero-order rate law: Rate=k[A]0 = k(1) = k Integrated zero-order rate law: [A] = -kt + [A]0

12.4 The Integrated Rate Law cont. The pseudo-first-order rate law is used when you have more than one product. It is: k’ = k[A]a[B]b

12.5 Reaction Mechanisms Reaction Mechanisms are the steps in which a chemical reaction follows. Intermediate- species that is neither a reactant nor a product but a molecule that is formed and absorbed during a reaction. Elementary Step- a reaction whose rate can be written from its molecularity. molecularity- number of molecules/atoms that must be in a reaction to produce the reaction. Unimolar Step- reaction only involving one molecule bimolecular- two molecules termolecular- three molecules

12.5 Reaction Mechanisms NO2(g) + NO2(g) = NO3(g) + NO(g) Slow NO3(g) + CO(g) = NO2(g) + CO2(g) Fast NO2(g) + CO(g) = NO(g) + CO2(g) A reaction is only as fast as its slowest step Overall Rate= k[NO2]2

12.6 A Model for Chemical Kinetics Initial Reaction: aA+bB=Products Rate Law: k[A]n[B]m Collision Model: chart that shows the number of molecules that must collide to react.

12.6 A Model for Chemical Kinetics # of collisions with activation energy = (total # of collisions)e-Ea/RT Rate Constant: k=zpe-Ea/RT z: collision frequency, p: steric factor Arrhenius Equation: k=Ae-Ea/RT Line of Best Fit: ln(k)=-E(1)+ln(A) (y=mx+b) R(T)

12.7 Catalysis Catalysis- substance that speeds up a reaction without being consumed. Enzymes lowers activation energy Two Kinds: Homogenous Heterogenous

12.7 Catalysis Heterogeneous Catalysis Gaseous reactants that are absorbed onto the surface of a solid catalysis absorption- collection of one substance on the surface of another substance. Example- hydrogenation of unsaturated hydrocarbons catalysis- platinum, palladium, or nickel can also migrate adsorbed reactants to surface react adsorbed substances desorption of products

12.7 Catalysis Homogeneous Catalysis A homogeneous catalysis exists in the same phase as the reacting molecules. Example: Nitric oxide, it catalysis the production of ozone in the troposphere (part of the atmosphere closest to us), but in the upper atmosphere it catalysis the decomposition of ozone. Page 588 shows how nitric oxide switches from producing to destroying ozone.