1 / 35

Physical Chemistry II

Physical Chemistry II. CHEM 3320. List Required Textbooks "Physical Chemistry" , 4 th edition, by kieth Laidler , Meisser and Sanctury . "Physical Chemistry", 8 th Edition, By Peter Atkins and Julio de Paula. Chapter I. Chemical reaction Mechanism 1

buttrey
Download Presentation

Physical Chemistry II

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Physical Chemistry II CHEM 3320

  2. List Required Textbooks • "Physical Chemistry", 4th edition, by kiethLaidler, Meisser and Sanctury. • "Physical Chemistry", 8th Edition, By Peter Atkins and Julio de Paula.

  3. Chapter I • Chemical reaction Mechanism 1 • 1. Rate law and order of reaction • 2. Differential rate laws • 3. integrated rate laws • 4.Molecularity • 5.Reaction mechanisms and elementary processes • 6. Collision Theory, activation energy • 7. Arrhenius Equation • 8. Activated Complex Theory

  4. UNIT 3 Chapter 6: Rates of Reaction Section 6.2 Factors Affecting Reaction Rate • 1. Nature of reactants • Ions faster than molecules • 2. Concentration • Higher concentration a greater number of effective collisions • 3. Temperature • Higher temperature, higher sufficient energy needed for a reaction (energy is ≥ Ea)

  5. UNIT 3 Chapter 6: Rates of Reaction Section 6.2 Factors Affecting Reaction Rate • 4. Pressure for gases • increased pressure , increased the number of collisions . • 5. Surface area • a greater surface area of solid reactant  a greater chance of effective collisions • 6. Presence of a catalyst • a catalyst is a substance that increases a reaction rate without being consumed by the reaction

  6. UNIT 3 Chapter 6: Rates of Reaction Section 6.3 The Rate Law The rate law shows the relationship between reaction rates k and concentration of reactants [] for the overall reaction. rate = k[A]m[B]n m: order of the reaction for reactant A n: order of the reaction for reactant B k: rate constant overall order of the reaction : m+ n

  7. UNIT 3 Chapter 6: Rates of Reaction Section 6.3 First-order Reactions A  Product • the rate law equation: rate = k [A] k: rate constant, [A]: concentration of reactant A

  8. UNIT 3 Chapter 6: Rates of Reaction Section 6.3 Second-order Reactions A + B  Product • the rate law equation: rate = k [A][B] OR rate = k [A]2 Rate of reaction is first order with respect to reactant A Rate of reaction is first order with respect to reactant B Overall order of chemical reaction is 2nd order reaction

  9. Reaction Mechanisms • Reaction Mechanism: explains how the overall reaction proceeds. • Reaction Mechanisms: is the sequence of several steps that describes the actual process by which reactants become products. • Each of these step is known as an elementary reaction or elementary process.

  10. 2NO (g) + O2 (g) 2NO2 (g) Elementary step 1: NO + NO N2O2 + Elementary step 2: N2O2 + O2 2NO2 Overall reaction: 2NO + O2 2NO2 Reaction Mechanisms Elementary step: any process that occurs in a single step For example, oxygen and nitrogen are not formed directly from the decomposition of nitrogen dioxide: N2O2 is detected during the reaction!

  11. For Example: • Now we will examine what path the reactants took in order to become the products. • The reaction mechanism gives the path of the reaction. • Mechanisms provide a very detailed picture of which bonds are broken and formed during the course of a reaction. A  B Reaction Mechanisms

  12. Elementary Steps & Molecularity • Molecularity: 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 uncommon to see termolecularProcesses…statistically improbable for an • effective collision to occur.)

  13. Bimolecular reaction Bimolecular reaction A + A  products A + B  products Rate Laws and Molecularity rate = k [A] Unimolecular reaction A  products rate = k [A][B] rate = k [A]2

  14. Rate Laws of Elementary Steps • Since this process occurs in one single step, the stoichiometry can be used to determine the rate law! • The rate law for an elementary step is written • directly from that step

  15. Collision Theory Molecules of reactants must collide each other . Not all collisions are effective (i.e. leads to chemical reaction). Conditions of occurring chemical reaction according to collision theory:

  16. Effective Collision Criteria 1.The Correct Orientation of Reactants For a chemical reaction to occur, reactant molecules must collide with the correct orientation relative to each other, which called (collision geometry). Five of many possible ways that NO(g) can collide with NO3(g) are shown. Only one has the correct collision geometry for reaction to occur.

  17. Effective Collision 2.Sufficient Activation Energy: For a chemical reaction, reactant molecules must also collide with sufficient energy. Activation energy, Ea, is the minimum amount of collision energy required to initiate a chemical reaction. Collision energy depends on the kinetic energy of the colliding particles.

  18. Potential Energy Hill • Ea: is the minimum energy that reactants must have to form products. • the height of thepotential barrier  (sometimes called the energy barrier). Activation Energy Curve called : 1. Potential Energy Hill Or 2. Potential Energy Barrier

  19. Activation energy, Ea The shaded part of the Maxwell-Boltzmann distribution curve represents number of particles ( i.e. number of collisions) that have enough collision energy for a reaction (i.e. the energy is ≥ Ea). • Suppose: Number of collisions 100 • Where, Ea= 70 j/mole • 25 collisions have 20 j/mole • 40 collisions have 45 j/mole • 20 collisions have 50 j/mole • 10 collisions have 60 j/mole • 5 collisions have 70 j/mole Number of collisions Energy

  20. Maxwell–Boltzmann Distributions • Temperature is defined as a measure of the average kinetic energy of the molecules in a sample. • At any temperature there is a wide distribution of kinetic energies.

  21. Maxwell–Boltzmann Distributions • As the temperature increases, the curve flattens and broadens. • Thus at higher temperatures, a larger number of molecules has higher energy.

  22. Maxwell–Boltzmann Distributions • If the dotted line represents the activation energy, as the temperature increases, so does the number of molecules ( i.e. number of collisions) that can overcome the activation energy barrier. • As a result, the reaction rate increases.

  23. Arrhenius Equation A mathematical relationship between k : ( rate constant of chemical reaction) and Ea: activation energy. where A : “Frequency Factor”-- a constant indicating how many collisions have the correct orientation to form products.

  24. Arrhenius Equation: Temperature Dependence of the Rate Constant Ea= the activation energy (J/mol) R= the gas constant (8.314 J/K•mol) T= is the absolute temperature ( in Kelvin) A= is the frequency factor

  25. Arrhenius Equation • Taking the natural logarithm (ln) of both sides, the equation becomes, • Ln (natural logarithm): is inverse function of exponential function. eln(x) = x ln(ex) = x

  26. Arrhenius Equation ln(k) = - Ea/R(1/T) + ln(A) • When k is determined experimentally at several temperatures,Eacan be calculated from the slope of a plot , Slope= -Ea/R Straight Line Equation y= mx +b

  27. Exothermic Reaction A → B + Heat • Potential Energy of reactant = Energy of chemical bond = Heat content = H • H product < H reactant • Enthalpy (∆H) • ∆H = H product - H reactant < 0 • ∆H = negative value (-) • Activation Energy (Ea) = H transition state- H reactant • Exothermic reaction has LowEa Exothermic

  28. Endothermic Reaction A + Heat → B • Potential Energy of reactant = Energy of chemical bond = Heat content = H • H product > H reactant • Enthalpy (∆H) • ∆H = H product - H reactant > 0 • ∆H = positive value (+) • Activation Energy (Ea) = H transition state- H reactant • Endothermic reaction has HighEa Endothermic

  29. Activation Energy and Enthalpy • The Ea for a reaction cannot be predicted from ∆H. • ∆H is determined only by the difference in potential • energy between reactants and products. • △H has no effect on the rate of reaction. • The rate depends on the size of the activation energy Ea • Reactions with lowEaoccur quickly. Reactions with • highEa occur slowly. Potential energy diagram for the combustion of octane.

  30. Activation Energy for Reversible Reactions • Potential energy diagrams  both forward and reverse reactions. • follow left to right for the forward reaction • follow right to left for the reverse reaction

  31. Activated Complex (Transition State) Activated complex is unstable compound and can break to form product. Activated complex: The arrangement of atoms found at the top of potential energy hill or barrier.

  32. Activated Complex (Transition State) 1. The collision must provide at least the minimum energy necessary to produce the activated complex. 2. It takes energy to initiate the reaction by converting the reactants into the activated complex. 3.If the collision does not provide this energy, products cannot form.

  33. Analyzing Reactions Using Potential Energy Diagrams • Forward Reaction is Exothermic Reaction • Reversible Reaction is Endothermic Reaction • BrCH3 molecule and OH- must collide with the correct orientation and sufficient energy and an activated complex forms. • 2. When chemical bonds reform, • potential energy decreases and • kinetic energy increases as the • particles move apart. Ea(rev) is greater than Ea(fwd)

More Related