CHEM 106 1nd Edition Lecture 19Outline of Last Lecture I. Measuring RadioactivityII. Radiometric DatingIII. Nuclear MedicineIV. Primordial GeochronometersOutline of Current Lecture II. EquilibriaIII. Writing Equilibrium Constant ExpressionsIV. Values and Terminology for KV. Relationship between Kc and KpCurrent LectureChapters 16: Chemical EquilibriaThis section is the heart of general chemistry 106Reactants ↔ Products→ forward reaction ← reverse reactionEquilibria, it’s dynamic!Fig. 16.1 H2O + CO ↔ CO2 + H2 Water-gas shift reactionT = 200°C, catalyzed with Cu/ZnONote Δ [reactants] and Δ [products] as f (time)“kinetic” part of curve vs. “equilibrium” partNote that at equilibrium, the reaction has not stopped, it is still occurringFig. 16.2; at equilibrium, forward rate is equal to reverse reaction rateThermodynamic Equation<we will eventually relate this to ΔG>Ex: 2NO2 ↔ N2O4dimerized NO2Assume that mechanism is the same as the stoichiometric reactionRate forward = kf [NO2]2 rate reverse = kr [N2O4]At equilibrium, rate forwards = rate reverseSo: kf [NO2]2 = kr [N2O4]N O2¿2¿¿kfkr=[ N2O4]¿K is equilibrium constant → always [ products][reactants]16.2 Writing Equilibrium Constant ExpressionsH2O + CO ↔ CO2 + H2(T = 500K)Table 16.1, pg. 767Initial [reactants], [products] vs. final [reactants],[products]Compare experiments 1 & 2:At equilibrium, comes to same concentrations regardless of whether you start with all reactants or all products.Now experiments 3 & 4:Different combinations of concentrationsDo equilibrium concentrations show same extent of the reaction?K = (0.0166)(0.0166)(0.0034)(0.0034)=24 at T =500 KWith K, you can ignore unitsCalculate K for experiments 3 & 4, you can also use values of K to estimate concentrationsLaw of Mass Action: aA + bB ↔ cC + dDKc = [C]c[ D]d[ A ]a[B ]b (products over reactants)Kc vs KpKc is based on concentrations (molarity)Kp is based on partial pressures (gases)Dalton’s Lawchapter 6 pg. 276Ptotal = P1 + P2 + P3…Kp = PPP(¿¿ A)a(¿¿ B)b¿(PC)c(¿¿ D)d¿¿Again, units are not needed, but the value of Kp ≠ value of Kc- You can’t mix partial pressures and concentration in the same equation/problemWhen K > 1 → there are more products than reactants, so reaction shifts to the rightWhen K < 1 → there are more reactant than products, so reaction shifts to the leftValues of K and Terminology:aA + bB ↔ cC + dDKc = [C]c[ D]d[ A ]a[B ]b ratio of products : reactantsIf Kc = 1, which is favored?[products] = [reactants], so neitherIf Kc > 1, which is favored?[products] favored; equilibrium lies far to the rightIf Kc < 1, which is favored?[reactants] favored; equilibrium lies far to the leftEx: What are the Kc and Kp expression for:CH4 (g) + 2H2O (g) ↔ CO2 (g) + 4H2 (g)Kc = H2(g)¿4¿H2O(g)¿2[C H4(g)]¿[CO2(g)]¿¿Kp = PH24PC O2PC H4PH2OAnother example with value for Kc:2H2 (g) + O2 (g) ↔ 2H2O (g)Kc = H2O( g)¿2¿H2(g)¿2[O2(g)]¿¿¿ = 3 e81lies far to the right!Ex: 2CO2 (g) ↔ 2CO (g) + O2 (g)Kc = CO(g)]2[O2(g)]¿C O2(g)¿2¿¿¿ = 3 e-92 lies very far to the left! (favors