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Chemical Kinetics. A 2(g) + B 2(g) 2 AB (g). Fundamental questions: Will it take place? Thermodynamics If it does, how long will it take to reach completion or equilibrium? Chemical kinetics: is the study of the speeds, or rates, of chemical reactions. Outline.
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A2(g) + B2(g) 2 AB(g) • Fundamental questions: • Will it take place? Thermodynamics • If it does, how long will it take to reach completion or equilibrium? • Chemical kinetics: is the study of the speeds, or rates, of chemical reactions
Outline • Why study kinetics? • Factors affecting rates • Measuring rates • reactant concentration • temperature • action of catalysts • surface area. • Concentration vs. Rate (Rate Laws) • Concentration vs. Time (Integrated Rate Laws) • Theories about Rxn Rates • Activation Energy • Mechanisms • Catalysts
D[AB] Rate of appearance of AB = Dt -D[A2] Rate of disappearance of A2 = Dt mol/liter Unit: = mol / L s = mol dm-3 s-1 S The rate of the reaction is a measure of how fast the changes are taking place.
The nature of the reactants • 2 NO+ O2 2 NO2 fast reaction at 25°C • 2 CO + O2 2 CO2 very slow at 25°C Rate constant -D[A2] = k [A] mol dm-3 s-1 Dt (in the simplest case: A products) Factors affecting reaction rates 2. The concentration of reactants Rate equation or rate law The rate of reaction is proportional to the rate of disappearance of reactants:
Rate Laws: Basic Assumptions • aA + bB cC + dD • Simple rate laws depend only on the concentrations of the reactants, not the products. • Rate = k[A]x[B]y • k = rate constant, units depend on the values of x and y. • x and y = the orders of A and B respectively • x + y = overall order of the rxn • 3 unknowns (x, y, k) 3 experiments
Slope (Rate) Changes w/ Time • Rate is a function of concentration. • When the concentration is high, the rate is large. • Concentration is a function of time. • When the reaction starts the concentration changes the most. 2HI H2 + I2
Zero order reaction: 2 N2O+ O2 2 N2 + O2 rate = k Au First order reaction: 2 N2O5 4 NO2 + O2 rate = k [A] Factors affecting reaction rates rate = k [N2O5]
Second order reaction: 2 A products A+B products rate = k [A]2 rate = k [A] [B] Third order reaction: 3 A products rate = k [A]3 rate = k [A] [B] [C] A + B + C products rate = k [A]2 [B] 2 A + B products
The order of a reaction is given by the sum of the exponents of the conc. terms in the rate equation. rate = k [A]m [B]n [C]p …. order = m + n + p + …
Temperature and Rate Constants (k) • Since the rate law has no temperature term in it, the rate constant must depend on temperature. • The size of k depends on T. • Greater T gives a larger k • Consider the reaction H2(g) + I2(g) 2HI(g).
effective collisions most probable energy at t1 t2 No of molecules t2 t1 Minimum energy required for reaction Ea 3. The temperature of the reaction. Collision theory of reaction rates
-Ea/RT n = n0e -Ea/RT k = Ae Ea ln k = lnA - RT Ea log k = log A - 2.303 RT Maxwell-Boltzmann distribution law Arrhenius equation Ea = activation energy A = frequency factor K = rate constant
Arrhenius Equation • K = rate constant • A = frequency factor • Ea = activation energy • T = temperature • R = 8.314 J/mol K
Frequency Factor • Solve for A Effective collisions / 6.02 x 1023
Temperature Increases Rate • Most reactions speed up as temperature increases. (E.g. food spoils when not refrigerated.) • When two light sticks are placed in water: one at room temperature and one in ice, the one at room temperature is brighter than the one in ice. • The chemical reaction responsible for chemi-luminescence is dependent on temperature: the higher the temperature, the faster the reaction and the brighter the light.
Half-Life • Half-lives are typically reported for radioactive materials, medications, toxins…etc. • 238U has a t½ = 5 x 109 y • 234P has a t½ = 7 h • Amount of time required for half of the compound to react. • Half life equations are found from the integrated rate laws. • Find the time required for the concentration to be one half the initial concentration.
The Collision Model • Observations: rates of reactions are affected by concentration and temperature. • Concentration terms already accounted for in rate laws • R = k[A]x • Goal: develop a model that explains why rates of reactions increase as temperature increases. Low T Collision High T Collision
The Rules of Reaction Mechanisms • Elementary step: any process that occurs in a single step. • Molecularity: the number of molecules present in an elementary step. • Unimolecular: one molecule in the elementary step, • Bimolecular: two molecules in the elementary step, and • Termolecular: three molecules in the elementary step. • It is not common to see termolecular processes (statistically improbable).
Biomolecular reaction A + B products 2A products Termolecular reaction A + B + C products 2A + B products 3A products Unimolecular reaction A products A unimoleculare reaction is first order They are not common!
Energy of Collision • Chemical rxns usually occur with bonds breaking and bonds forming. • Collision energy supplies the energy needed to break the bonds. • With increased temperature, collision energy increases and the rate of rxn increases.
Effective Collisions • The right molecules must collide. • The collision must be energetic enough. • The molecules must collide in the proper orientation. • All of the above conditions must be met for a rxn to occur! • There are far more collisions between the wrong molecules or in the wrong orientation or with the wrong energy! • Some estimate that only 1 in 1018 collisions are effective.
Molecular Orientation • If H2 and I2 collide, they can collide along any possible trajectory. • However, only one trajectory leads to a rxn.
H H I I Top of the Peak ‡ • At the top of the peak, we have… • Maximum potential energy • Transition state • Point where reactants change to products • Bonds breaking and bonds forming • Intermediate structures • The [activated complex]‡ 2 HI Potential Energy H2 + I2 Progress of Rxn
Activation Energy • In order to form products, bonds must be broken in the reactants. • Bond breakage requires energy • Arrhenius: molecules must posses a minimum amount of energy to react.. • Ea is the minimum energy required to initiate a chemical reaction.
Energy of Collision Activation Energy Ea Potential Energy Energy of Rxn Enthalpy DH Progress of Rxn Progress of Reaction Diagrams • Relates the PE of the rxn to time. • For example, an exothermic rxn between A + B is shown on the right. Reactants Note: DH and Ea are unrelated! 100% of Ea is returned! Products
Energy of Collision Activation Energy Ea Potential Energy Energy of Rxn Enthalpy DH Progress of Rxn Endothermic System is Reversed • Products are higher in PE than reactants. Products Note: Less than 100% of Ea is returned. Reactants
Reaction Mechanisms Rate Laws of Multistep Mechanisms • Rate-determining step: is the slowest of the elementary steps. • Therefore, the rate-determining step governs the overall rate law for the reaction. Mechanisms with an Initial Fast Step • It is possible for an intermediate to be a reactant. • … 2NO2Cl(g) Cl2(g) + 2NO2(g)
E1 E2 Reaction mechanisms: The rate equation for a reaction must be determined by experimentation Two step mechanism Rate-determining step (intermediate)
Uncatalysed reaction Catalysed reaction reactants products 5. The catalysis A catalyst is a substance that increases the rate of a chemical reaction without being used up in the reaction 4. The surface area of a solid Increasing the surface area of a solid reactant will increase the rate of reaction (explosion of flour dust)
Catalysis • A catalyst changes the rate of a chemical reaction. • There are two types of catalyst: • homogeneous, and • heterogeneous. • Chlorine atoms are catalysts for the destruction of ozone. • Homogeneous Catalysis • The catalyst and reaction is in one phase. • Hydrogen peroxide decomposes very slowly: • 2H2O2(aq) 2H2O(l) + O2(g).
Promoters catalytic poisons homogeneous Catalyst heterogeneous positive (accelerate…) negative or inhibitor (retard…)
Catalysis • Homogeneous Catalysis • 2H2O2(aq) 2H2O(l) + O2(g). • In the presence of the bromide ion, the decomposition occurs rapidly: • 2Br –(aq) + H2O2(aq) + 2H+(aq) Br2(aq) + 2H2O(l). • Br2(aq) is brown. • Br2(aq) + H2O2(aq) 2Br –(aq) + 2H+(aq) + O2(g). • Br – is a catalyst because it can be recovered at the end of the reaction. • Generally, catalysts operate by lowering the activation energy for a reaction.
Catalysis • Homogeneous Catalysis
Catalysis • Heterogeneous Catalysis
Catalysis • Enzymes • Enzymes are biological catalysts. • Most enzymes are protein molecules with large molecular masses (10,000 to 106 amu). • Enzymes have very specific shapes. • Most enzymes catalyze very specific reactions. • Substrates undergo reaction at the active site of an enzyme. • A substrate locks into an enzyme and a fast reaction occurs.
1. S + E ES Enzyme-substrate complex Substrate (reactant) Enzyme 2. ES E + P Product Natural catalysts = enzymes (very specific) Michaelis – Menten mechanism:
First order Zero order Reaction rate Concentration of substrate At low concentrate: Rate of disappearence of S = k [S] At high concentrate:Rate of disappearence of S = k’ [S]0 = k’ The enzyme is saturated!
Catalysis • Enzymes