Chemical Kinetics Unit 10 – Chapter 12.

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Presentation transcript:

Chemical Kinetics Unit 10 – Chapter 12

Reaction Rate Change in concentration of a reactant over time 12.1 Reaction Rates Reaction Rate Change in concentration of a reactant over time [A] means “concentration of A” in mol/L.

Decomposition of NO2 into O2 and NO. 12.1 Reaction Rates Decomposition of NO2 into O2 and NO.

Instantaneous Rate Value of the rate at a particular time. 12.1 Reaction Rates Instantaneous Rate Value of the rate at a particular time. Can be obtained by computing the slope of a line tangent to the curve at that point (1st derivative).

12.2 Rate Laws: An Introduction Rate Law Shows how the rate depends on the concentration of the reactants (no equilibrium reactions…yet). For the decomposition of nitrogen dioxide: 2NO2(g) → 2NO(g) + O2(g) Rate = k[NO2]n k = rate constant n = order of the reactant

n must be determine experimentally. 12.2 Rate Laws: An Introduction Rate Law Rate = k[NO2]n The concentrations of the products do not appear in the rate law because the reaction rate is being studied under conditions where the reverse reaction does not contribute to the overall rate (it is not an equilibrium reaction and the products do not sterically inhibit the reaction). n must be determine experimentally.

12.2 Rate Law: An Introduction Types of Rate Laws Differential Rate Law (rate law): shows how the rate of a reaction depends on concentrations. Integrated Rate Law – shows how the concentrations of species in the reaction depend on time.

Method of Initial Rates 12.3 Determining the Form of the Rate Law Method of Initial Rates The value of the initial rate is determined for each experiment at the same value of t as close to t = 0 as possible. Several experiments are carried out using different initial concentrations of each of the reactants, and the initial rate is determined for each run. The results are then compared to see how the initial rate depends on the initial concentrations of each of the reactants.

Overall Reaction Order 12.3 Determining the Form of the Rate Law Overall Reaction Order The sum of the exponents in the reaction rate equation. Rate = k[A]n[B]m Overall reaction order = n + m k = rate constant [A] = concentration of reactant A [B] = concentration of reactant B

ln[A] = –kt + ln[A]o First Order Rate Law 12.4 The Integrated Rate Law First Order Rate Law Rate = k[A] Integrated: ln[A] = –kt + ln[A]o [A] = concentration of A at time t k = rate constant t = time [A]o = initial concentration of A

12.4 The Integrated Rate Law Plot of ln [N2O5] vs. Time

First Order; Half Life k = rate constant Half–Life: 12.4 The Integrated Rate Law First Order; Half Life Half–Life: k = rate constant Half–life does not depend on the concentration of reactants.

[A] = concentration of A at time t k = rate constant t = time 12.4 The Integrated Rate Law Second-Order Rate = k[A]2 Integrated: [A] = concentration of A at time t k = rate constant t = time [A]o = initial concentration of A

Second Order; Half-Life 12.4 The Integrated Rate Law Second Order; Half-Life Half–Life: k = rate constant [A]o = initial concentration of A Half–life gets longer as the reaction progresses and the concentration of reactants decrease. Each successive half–life is double the preceding one.

12.4 The Integrated Rate Law Exercise For a reaction aA  Products, [A]0 = 5.0 M, and the first two half-lives are 25 and 50 minutes, respectively. Write the rate law for this reaction. b) Calculate k. c) Calculate [A] at t = 525 minutes. rate = k[A]2 a) rate = k[A]2 We know this is second order because the second half–life is double the preceding one. b) k = 8.0 x 10-3 M–1min–1 25 min = 1 / k(5.0 M) c) [A] = 0.23 M (1 / [A]) = (8.0 x 10-3 M–1min–1)(525 min) + (1 / 5.0 M) k = 8.0 x 10-3 M–1min–1 [A] = 0.23 M

Rate = k[A]0 = k Integrated: [A] = –kt + [A]o Zero Rate Order 12.4 The Integrated Rate Law Zero Rate Order Rate = k[A]0 = k Integrated: [A] = –kt + [A]o [A] = concentration of A at time t k = rate constant t = time [A]o = initial concentration of A

[A]o = initial concentration of A 12.4 The Integrated Rate Law Zero Order; Half-Life Half–Life: k = rate constant [A]o = initial concentration of A Half–life gets shorter as the reaction progresses and the concentration of reactants decrease.

12.4 The Integrated Rate Law Summary

NO2(g) + CO(g) → NO(g) + CO2(g) 12.5 Reaction Mechanisms Reaction Mechanisms Most chemical reactions occur by a series of elementary steps. An intermediate is formed in one step and used up in a subsequent step and thus is never seen as a product in the overall balanced reaction. NO2(g) + CO(g) → NO(g) + CO2(g)

Elementary Step (Molecularity) 12.5 Reaction Mechanisms Elementary Step (Molecularity) Unimolecular – reaction involving one molecule; first order. Bimolecular – reaction involving the collision of two species; second order. Termolecular – reaction involving the collision of three species; third order.

Rate-Determining Step A reaction is only as fast as its slowest step. 12.5 Reaction Mechanisms Rate-Determining Step A reaction is only as fast as its slowest step. The rate-determining step (slowest step) determines the rate law and the molecularity of the overall reaction. The sum of the elementary steps must give the overall balanced equation for the reaction. The mechanism must agree with the experimentally determined rate law.

Decomposition of N2O5 2N2O5(g)  4NO2(g) + O2(g) 12.5 Reaction Mechanisms Decomposition of N2O5 2N2O5(g)  4NO2(g) + O2(g) Step 1: N2O5 NO2 + NO3 (fast) Step 2: NO2 + NO3 → NO + O2 + NO2 (slow) Step 3: NO3 + NO → 2NO2 (fast) 2( )

12.5 Reaction Mechanisms Reality Check The reaction A + 2B  C has the following proposed mechanism: A + B D (fast equilibrium) D + B  C (slow) Write the rate law for this mechanism. rate = k[A][B]2

12.6 A Model for Chemical Kinetics Collision Model Molecules must collide to react. Main Factors: Activation energy, Ea Temperature Molecular orientations Activation Energy: Energy that must be overcome to produce a chemical reaction.

Change in Potential Energy 12.6 A Model for Chemical Kinetics Change in Potential Energy

For Reactants to Form Products 12.6 A Model for Chemical Reactions For Reactants to Form Products Collision must involve enough energy to produce the reaction (must equal or exceed the activation energy). Relative orientation of the reactants must allow formation of any new bonds necessary to produce products.

Arrhenius Equation Linear Version Ea = activation energy 12.6 A Model for Chemical Kinetics Arrhenius Equation A = frequency factor Ea = activation energy R = gas constant (8.3145 J/K·mol) T = temperature (in K) Linear Version

A substance that speeds up a reaction without being consumed itself. 12.7 Catalysis Catalyst A substance that speeds up a reaction without being consumed itself. Provides a new pathway for the reaction with a lower activation energy.

Effect of a Catalyst on the Number of Reaction-Producing Collisions 12.7 Catalysis Effect of a Catalyst on the Number of Reaction-Producing Collisions

Heterogeneous Catalyst 12.7 Catalysis Heterogeneous Catalyst Most often involves gaseous reactants being adsorbed on the surface of a solid catalyst. Adsorption – collection of one substance on the surface of another substance. Adsorption and activation of the reactants. Migration of the adsorbed reactants on the surface. Reaction of the adsorbed substances. Escape, or desorption, of the products.

Exists in the same phase as the reacting molecules. 12.7 Catalysis Homogeneous Catalyst Exists in the same phase as the reacting molecules. Enzymes are nature’s catalysts.

12.1 Reaction Rates

12.1 Reaction Rates

12.1 Reaction Rates