2 List of TopicsNo. ofWeeksContact HoursKinetics of elementary reactions311Composite reaction mechanismsSolution of electrolytes &Debye-Huckel theoryElectrochemical cells28Surface chemistryTransport properties1A comprehensive review
3 Proportion of Total Assessment By the end of the semester Assessment taskWeek DueProportion of Total AssessmentMid term I6th week10%Mid term II11th weekParticipation and attendance5%QuizzespresentationAssignmentFrom 2nd to 14th weekFinal practical exam15th week20%Final written Exam.By the end of the semester40%
4 List Required Textbooks "Physical Chemistry", 4th edition, by kieth Laidler, Meisser and Sanctury."Physical Chemistry", 8th Edition, By Peter Atkins and Julio de Paula.
5 Chapter I Chemical reaction Mechanism 1 1. Rate law and order of reaction2. Differential rate laws3. integrated rate laws4.Molecularity5.Reaction mechanisms and elementary processes6. Collision Theory, activation energy7. Arrhenius Equation8. Activated Complex Theory
6 Factors Affecting Reaction Rate UNIT 3Chapter 6: Rates of ReactionSection 6.2Factors Affecting Reaction Rate1. Nature of reactantsIons faster than molecules2. ConcentrationHigher concentration a greater number of effective collisions3. TemperatureHigher temperature, higher sufficient energy needed for a reaction (energy is ≥ Ea)
7 Factors Affecting Reaction Rate UNIT 3Chapter 6: Rates of ReactionSection 6.2Factors Affecting Reaction Rate4. Pressure for gasesincreased pressure , increased the number of collisions .5. Surface areaa greater surface area of solid reactant a greater chance of effective collisions6. Presence of a catalysta catalyst is a substance that increases a reaction rate without being consumed by the reaction
8 The Rate Law rate = k[A]m[B]n UNIT 3 Chapter 6: Rates of ReactionSection 6.3The Rate LawThe rate law shows the relationship between reaction rates k and concentration of reactants  for the overall reaction.rate = k[A]m[B]nm: order of the reaction for reactant An: order of the reaction for reactant Bk: rate constantoverall order of the reaction : m + n
9 First-order Reactions UNIT 3Chapter 6: Rates of ReactionSection 6.3First-order ReactionsA Productthe rate law equation:rate = k [A]k: rate constant, [A]: concentration of reactant A
10 Second-order Reactions UNIT 3Chapter 6: Rates of ReactionSection 6.3Second-order ReactionsA + B Productthe rate law equation:rate = k [A][B] ORrate = k [A]2Rate of reaction is first order with respect to reactant ARate of reaction is first order with respect to reactant BOverall order of chemical reaction is 2nd order reaction
11 Reaction MechanismsReaction 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.
12 N2O2 is detected during the reaction! Reaction MechanismsElementary step: any process that occurs in a single stepFor example, oxygen and nitrogen are not formed directly from the decomposition of nitrogen dioxide:2NO (g) + O2 (g) NO2 (g)N2O2 is detected during the reaction!Elementary step 1:NO + NO N2O2Overall reaction:2NO + O NO2+Elementary step 2:N2O2 + O NO2
13 Reaction Mechanisms For Example: A B 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
14 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, andTermolecular: three molecules in the elementary step.(It is uncommon to see termolecular Processes…statistically improbable for aneffective collision to occur.)
15 Rate Laws and Molecularity rate = k [A]Unimolecular reactionA productsBimolecular reactionA + B productsrate = k [A][B]Bimolecular reactionA + A productsrate = k [A]2
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 writtendirectly from that step
17 Collision TheoryMolecules 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:
18 Effective Collision Criteria 1.The Correct Orientation of ReactantsFor 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.
19 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.
20 Potential Energy HillEa: is the minimum energy that reactants must have to form products.the height of the potential barrier (sometimes called the energy barrier). Activation EnergyCurve called :1. Potential Energy Hill Or2. Potential Energy Barrier
21 Activation energy, EaThe 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 100Where, Ea= 70 j/mole25 collisions have 20 j/mole40 collisions have 45 j/mole20 collisions have 50 j/mole10 collisions have 60 j/mole5 collisions have 70 j/moleNumber of collisionsEnergy
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.
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.
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.
25 Arrhenius EquationA mathematical relationship between k : ( rate constant of chemical reaction) and Ea: activation energy.whereA : “Frequency Factor”-- a constant indicating how many collisions have the correct orientation to form products.
26 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
27 Arrhenius EquationTaking the natural logarithm (ln) of both sides, the equation becomes,Ln (natural logarithm): is inverse function of exponential function.eln(x) = xln(ex) = x
28 ln(k) = - Ea/R(1/T) + ln(A) Arrhenius Equationln(k) = - Ea/R(1/T) + ln(A)Straight Line Equationy = mx + bWhen k is determined experimentally at several temperatures, Ea can be calculated from the slope of a plot , Slope= -Ea/R
29 Exothermic Reaction A → B + Heat Potential Energy of reactant = Energy of chemical bond = Heat content = HH product < H reactantEnthalpy (∆H)∆H = H product - H reactant < 0∆H = negative value (-)Activation Energy (Ea) = H transition state- H reactantExothermic reaction has Low EaExothermic
30 Endothermic Reaction A + Heat → B Endothermic reaction has High Ea Potential Energy of reactant = Energy of chemical bond = Heat content = HH product > H reactantEnthalpy (∆H)∆H = H product - H reactant > 0∆H = positive value (+)Activation Energy (Ea) = H transition state- H reactantEndothermic reaction has High EaEndothermic
31 Activation Energy and Enthalpy The Ea for a reaction cannot be predicted from ∆H.∆H is determined only by the difference in potentialenergy between reactants and products.△H has no effect on the rate of reaction.The rate depends on the size of the activation energy EaReactions with low Ea occur quickly. Reactions withhigh Ea occur slowly.Potential energy diagram for the combustion of octane.
32 Activation Energy for Reversible Reactions Potential energy diagrams both forward and reverse reactions.follow left to right for the forward reactionfollow right to left for the reverse reaction
33 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.
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.
35 Analyzing Reactions Using Potential Energy Diagrams Forward Reaction is Exothermic ReactionReversible Reaction is Endothermic ReactionBrCH3 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 andkinetic energy increases as theparticles move apart.Ea(rev) is greater than Ea(fwd)