CHEM 106 1nd Edition Lecture 11Outline of Last Lecture I. Free EnergyII. Calculating ΔG°fIII. Temperature and SpontaneityIV. Coupled ReactionsOutline of Current Lecture II. Chemical KineticsIII. Reaction RatesIV. Effect of ConcentrationCurrent LectureChapter 15: Chemical KineticsThermodynamics tells us whether a reaction should or should not occur, but it tells us nothing about rate or how it occurs.Chemical Kinetics tells us how fast, and the mechanism or pathway of the reaction. Photochemical Smog:N2(g) + O2(g) → 2NO(g) ΔH°rxn = 180.6 kJReaction is endothermic But ΔG < 0 at very high temperatureSo, ΔS°rxn has to be very positive to overcome the endothermic You cant really tell this by just looking at the reaction in this case (because same numberof particles and no phase changes)Fig. 15.2 page 70315.2 Reaction RatesReaction rate: how rapidly a reaction occurs, rate of change of concentrationRate = Δ[ ]Δtt = time [ ] = concentrationWhere Δ[ ] = [ ]final – [ ]initial Δt = tfinal – tinitial Back to reaction between N2 and O2:Δ[ NO]Δt = [NO]final−[NO]initialt final−t initialunits are M/s[NO] ↑ with time, so rate is positiveRules for rates of reactions: - Always positive (e.g. take absolute values, if reactant disappearing)- Can be based on any reactant or product (for a given reaction)- Value of rate must be identical for any reactant or product (adjust for stoichiometry) So we can write a rate based on consumption of N2-actual value would be negative, but we consider absolute value Rate of N2 consumption: −Δ[N2]Δt or ¿ Δ[N2]∨¿Δt¿ = [N2]f −[N2]itf −tiRemember that when 1 mol O2 + 1 mol N2 consumed, they produce 2 mols NOHave to indicate this in the rate expression:N2(g) + O2(g) → 2NO(g) Rate: −Δ[N2]Δt = −Δ[O2]Δt = 12 Δ[ NO]Δt(written by the reactants point of view) Must include stoichiometry!Fig. 15.3 page 708NO ↑ faster than N2 and O2 ↓Ex: P4 + 6H2 → 4PH3Suppose Δ[ PH3]Δt = 2.4 M/s what is Δ[ P4]Δt ?Δ[ P4]Δt = -14 Δ[ PH3]Δt = - 0.6 M/sRate of reaction would be + 0.6 M/sWhat is Δ[ H2]Δt ?−16 Δ[ H2]Δt = 14 ∆[ P H3]∆ t ∆[ H2]∆t = −64 ∆[ PH3]∆ t = -3.6 M/sEx: 2 NH3(aq) + 3O2(g) → 2H+(aq) + 2NO-(aq) + 2H2O(l)How are rates of formation of H+ and NO2- related to rate of consumption of NH3?12 +¿H¿¿Δ¿¿ = 12 2−¿NO¿¿Δ¿¿ = −12 Δ[ N H3]Δt Rates of formation = rates of consumptionRelate formation rate of NO2- to the consumption rate of O2:12 −¿NO2¿¿Δ¿¿ = −13 Δ[O2]Δt so, −¿NO2¿¿Δ ¿¿ = −23 Δ[O2]Δ tRate of NO2- formation is 2/3 the rate of O2 consumptionFig. 15.3 Instantaneous rate vs. Average rateAs a reaction proceeds, average rate decreasesThe rate at a specific time is sometimes more usefulFig. 15.4 page 709Instantaneous rate is the tangent to the concentration vs. time (1st derivative)The average rate is the slope of the line connecting 2 points on a graph15.3 Effect of ConcentrationFig. 15.6 page 711Reaction rate ↓ as t ↑Eventually, reaction ends, rate = 0 M/sReaction progresses more quickly when [ ] of reactants are highest – more reactants tocollide with each other.We expect the reaction rate to be dependent on concentration, but the way a reaction depends on [ ] is not the same. To be