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Physical Chemistry II CHEM 3320. List of Topics No. of Weeks Contact Hours Kinetics of elementary reactions311 Composite reaction mechanisms311 Solution.

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Presentation on theme: "Physical Chemistry II CHEM 3320. List of Topics No. of Weeks Contact Hours Kinetics of elementary reactions311 Composite reaction mechanisms311 Solution."— Presentation transcript:

1 Physical Chemistry II CHEM 3320

2 List of Topics No. of Weeks Contact Hours Kinetics of elementary reactions311 Composite reaction mechanisms311 Solution of electrolytes & Debye-Huckel theory 311 Electrochemical cells28 Surface chemistry28 Transport properties13 A comprehensive review 13

3 Assessment taskWeek DueProportion of Total Assessment Mid term I6 th week10% Mid term II11 th week10% Participation and attendance5% Quizzes5% presentation5% AssignmentFrom 2 nd to 14 th week5% Final practical exam15 th week20% Final written Exam.By the end of the semester40% Mid term I6 th week10%

4 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.

5 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 5

6 UNIT 3 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 ≥ E a ) Chapter 6: Rates of Reaction

7 UNIT 3 Section 6.2 Factors Affecting Reaction Rate Chapter 6: Rates of Reaction 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

8 UNIT 3 Section 6.3 The Rate Law The rate law shows the relationship between reaction rates k and concentration of reactants [] for the overall reaction. Chapter 6: Rates of 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

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

10 UNIT 3 Section 6.3 Second-order Reactions Chapter 6: Rates of Reaction 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 2 nd order reaction A + B  Product

11 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. 11

12 Reaction Mechanisms 2NO (g) + O 2 (g) 2NO 2 (g) N 2 O 2 is detected during the reaction! Elementary step 1:NO + NO N 2 O 2 Elementary step 2:N 2 O 2 + O 2 2NO 2 Overall reaction:2NO + O 2 2NO 2 + Elementary step: any process that occurs in a single step 12 For example, oxygen and nitrogen are not formed directly from the decomposition of nitrogen dioxide:

13 Reaction Mechanisms 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 13

14 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 termolecular Processes…statistically improbable for an effective collision to occur.) Elementary Steps & Molecularity 14

15 Unimolecular reaction rate = k [A] Bimolecular reaction A + B  products rate = k [A][B] rate = k [A] 2 Rate Laws and Molecularity A  products Bimolecular reaction A + A  products 15

16 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 16

17 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: 17

18 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 NO 3 (g) are shown. Only one has the correct collision geometry for reaction to occur. 18

19 Effective Collision 2.Sufficient Activation Energy: For a chemical reaction, reactant molecules must also collide with sufficient energy. Activation energy, E a, is the minimum amount of collision energy required to initiate a chemical reaction. Collision energy depends on the kinetic energy of the colliding particles. 19

20 Potential Energy Hill E a : is the minimum energy that reactants must have to form products. the height of the potential barrier (sometimes called the energy barrier). 20 Activation Energy Curve called : 1. Potential Energy Hill Or 2. Potential Energy Barrier

21 21 Activation energy, E a 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 ≥ E a ). 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

22 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. 22

23 Maxwell–Boltzmann Distributions As the temperature increases, the curve flattens and broadens. Thus at higher temperatures, a larger number of molecules has higher energy. 23

24 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. 24

25 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. 25

26 Arrhenius Equation: Temperature Dependence of the Rate Constant E a = the activation energy (J/mol) R = the gas constant (8.314 J/Kmol) T = is the absolute temperature ( in Kelvin) A = is the frequency factor 26

27 Arrhenius Equation Taking the natural logarithm (ln) of both sides, the equation becomes, Ln (natural logarithm): is inverse function of exponential function. 27 e ln(x) = x ln(e x ) = x

28 Arrhenius Equation 28 When k is determined experimentally at several temperatures, E a can be calculated from the slope of a plot, Slope= -E a /R Straight Line Equation y = mx + b ln(k) = - E a /R(1/T) + ln(A)

29 Exothermic Reaction 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 (E a ) = H transition state - H reactant Exothermic reaction has Low E a 29 A → B + Heat Exothermic

30 Endothermic Reaction 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 (E a ) = H transition state - H reactant Endothermic reaction has High E a 30 A + Heat → B Endothermic

31 31 Activation Energy and Enthalpy The E a 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 low E a occur quickly. Reactions with high E a occur slowly. Potential energy diagram for the combustion of octane.

32 32 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

33 Activated Complex (Transition State) 33 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.

34 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. 34

35 35 Analyzing Reactions Using Potential Energy Diagrams E a(rev) is greater than E a(fwd) Forward Reaction is Exothermic Reaction Reversible Reaction is Endothermic Reaction 1.BrCH 3 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.


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