REACTION KINETICS 1 Chemical Kinetics Kinetics – how fast does a reaction proceed? Will the reaction produce significant change within milliseconds or thousands of years? Reaction rate is the change in the concentration of a reactant or a product with time (M/s). A B Δ[A] rate = - Δ[A] = change in concentration of A over Δ t time period Δ t Δ[B] Δ[B] = change in concentration of B over rate = Δ t time period Δ t Because [A] decreases with time, Δ[A] is negative. 2 A B D[A] rate = - Dt D[B] rate = Dt 3 Reaction Rates and Stoichiometry 2 N2O5 (g) 4 NO2 + O2 4 Consider the decomposition of N2O5 to give NO2 and O2: 2 N2O5 (g) 4 NO2 + O2 Time Concentration (M) (s) N2O5 NO2 O2 0 0,02 0 0 100 0,0169 0,0063 0,0016 200 0,0142 0,0115 0,0029 300 0,0120 0,0160 0,0040 400 0,0101 0,0197 0,0049 500 0,0086 0,0229 0,0057 600 0,0072 0,0256 0,0064 700 0,0061 0,0278 0,0070 Reactants decrease Products increase with time with time 5 0,03 0,02 N2O5 NO2 O2 Concentration (M) Concentration 0,01 0 0 100 200 300 400 500 600 700 800 Time (sec) 6 From the graph looking at t = 300 to 400 sec 0.0009 M −6 −1 Rate O2 = = 9 x 10 M푠 100 s 0.0037 M −5 −1 Rate NO2 = = 3.7 x 10 M푠 100 s −0,0019 M −5 −1 Rate N2O5 = = -1.9 x 10 M푠 100 s Why do they differ? Recall : 2 N2O5 (g) 4 NO2 + O2 7 To compare the rates one must account for the stoichiometry. 1 −6 −1 −6 −1 Rate O2 = x 9 x 10 M푠 = 9 x 10 M푠 1 1 −5 −1 −6 −1 Rate NO2 = x 3.7 x 10 M푠 = 9.2 x 10 M푠 4 Now they 1 −5 −1 −6 −1 agree ! Rate N2O5 = - x 1.9 x 10 M푠 = 9.5 x 10 M푠 2 8 Reaction Rates and Stoichiometry 2A B Two moles of A disappear for each mole of B that is formed. D[A] D[B] rate = - 1 rate = 2 Dt Dt aA + bB cC + dD 1 D[A] D[B] D[C] D[D] rate = - = - 1 = 1 = 1 a Dt b Dt c Dt d Dt 9 Example : rA rB rC rD aA + bB cC + dD Rate = a b c d 3A 2B + C st If the overall rate of rxn is 1 order wrt to A , then r = k [A] ra rb rc rate = 3 2 1 rA = rxn. rate wrt to rA = -3 k [A] rB = rxn. rate wrt to rB = 2 k [A] rC = rxn. rate wrt to rC = k [A] 10 The Rate Law and Reaction Order The rate law expresses the relationship of the rate of a reaction to the rate constant and the concentrations of the reactants raised to some powers. aA + bB cC + dD Rate = k [A]x[B]y reaction is xth order with respect to reactant A reaction is yth order with respect to reactant B reaction is (x +y)th order overall 11 @ constant temperature, the rate of rxn. (disappearance of a reactant or formation of a product) is some function of the concentration of the reactants aA + bB cC + dD order of the rxn. respect to individual reactants A and B (0, 1, 2 may be fractional) Overall order of rxn = r= k [A]α [B]β (mol/L.t) molar concentrations of reactants (mol/ L) reaction rate constant Overall order of rxn = α+β Order of rxn with respect to reactant A =α 12 Order of rxn with respect to reactant B = β F2 (g) + 2ClO2 (g) 2FClO2 (g) x y rate = k [F2] [ClO2] Double [F2] with [ClO2] constant Rate doubles x = 1 rate = k [F2][ClO2] Quadruple [ClO2] with [F2] constant Rate quadruples 13 y = 1 Rate Laws • Rate laws are always determined experimentally. • Reaction order is always defined in terms of reactant (not product) concentrations. • The order of a reactant is not related to the stoichiometric coefficient of the reactant in the balanced chemical equation. F2 (g) + 2ClO2 (g) 2FClO2 (g) 1 rate = k [F2][ClO2] 14 Unit of k (reaction rate constant) k = M1-n t-1 mol For 0 order rxn. k = rate L.t r = k[A]⁰ r = k mol rate 1 For 1st order rxn. k = L . t = mol mol t r = k [A] L L mol nd rate L.t L For 2 order rxn. k = 2 2 = mol mol mol.t r = k *A+² L L 15 Example : r = k *A+² *B+ Order of overall rxn. = 2 + 1 = 3 Order of rxn. with respect to reactant A = 2 Order of rxn. with respect to reactant B = 1 2 L 1 Unit of k = M-2 t-1 = mol t 16 Classes of Reactions That Occur In Nature Homogenous Reactions Heterogenous Reactions Reactants and products are in the same phase. Reactions that occur at surfaces between (i.e., liquid, solid or gas) phases. Reactants are distributed continuously(but Examples: not necessarily uniformly) throughout the Reactions on the surface of ion exchange resin fluid. Reactions that require the presence of a solid phase catalyst 1a. Single irreversible reaction: A B A+A B aA + bB C 1b. Multiple irreversible reaction: B A (Parallel reaction) A B C C (consecutive or series reaction) 1c. Reversible reaction: 17 A B A+B C+D Types of Reactions and Reaction Rates → irreversible rxns → reversible rxns → saturation rxns → autocatalytic rxns 18 Irreversible Reactions a)Single Irreversible Rxns: A P A + A P aA + bB P 19 Irreversible Reactions a)Single Irreversible Rxns: r r r A B r a b P Example: aA pP (Rxn. rate 1st order) r = k [A] rA = r x (-a) rA = -a k [A] rP = p k [A] 20 b)Multiple Irreversible Rxns: b -1) Parallel rxns: r r r A B k1 bB 1 a b aA r r k cC r A C 2 2 a c r r r Overall rate of rxn: r + r 2 A B C 1 2 a b c Rate of rxn with respect to A = rA = -ar1 – ar2 Rate of rxn with respect to B = rB = b r1 21 Rate of rxn with respect to C = rC = cr2 k1 bB aA k2 cC If both of the rates of rxn 1st order r1= k1[A] r2= k2[A] rA= -ak1[A] - ak2[A] rB= bk1[A] rC= ck2[A] 22 b-2) Consequtive rxns aAk1 bBk2 cC Slowest step limits the overall rxn. rate k1 r r r = - a r aA bB r A B A 1 1 a b rB = b r1 – b r2 k2 r r r B C bB cC 2 b c rC = c r2 If both of the rates of rxn 1st order r = -ak [A] r1= k1[A] A 1 r2= k2[A] rB= bk1[A] - bk2[B] rC= ck2[A] 23 Reversible Reactions k1 aA bB k 2 r r r A B k1 1 aA bB a b rA = ar1 + ar2 bB aA r = br + br r r B 1 2 k2 r B A 2 b a Note: for reversible rxns signs of stoichiometric coefficients (both reactants & products) are always positive (+) If both of the rates of rxn 1st order r = ak [A] + ak [B] r1= k1[A] a 1 2 r = bk [A]+ bk [B] r2= k2[A] b 1 2 24 Saturation Type Reactions Saturation type reactions have a maximum rate that is a point at which the rate becomes independent of concentration A. k [A] For the reaction aA bB r K [A] r = rate of rxn, mol /L.t [A] = conc. of reactant, mol / L k = rxn rate constant, mol/ L.t K = half saturation constant, mol/ L Fig.1.Graphical representation of Saturation –Type Reaction Ref: Tchobanoglous and Schroeder, 1985, Addison-Wesley Publishing Company 25 Saturation Type Reactions (continue) k[A] r K [A] • When K << [A] r = k [A] ≈ k (zero order) K + [A] Negligible Ref: Tchobanoglous and Scroeder, 1985, Addison-Wesley Publishing Company • When K >> [A] r = k [A] ≈ k [A] (first order) K + [A] K 27 Negligible Saturation Type Reactions (continued) k For the reaction A + B P k [A] [B] r = k [A] [B] K [A]K [B] K [A] 1 2 28 Autocatalytic Reactions Autocatalytic reaction rates are functions of the product concentration. Example: Bacterial growth (rate of increase in bacterial number is proportional to the number present) Autocatalytic rxns can be 1st order 2nd order Saturation Type Partially Autocatalytic (function of reactant & product) 29 1st order Autocatalytic Rxn. r r aA bB r = k [B] = A B a b rA = a k [B] rB = b k [B] 2nd order Autocatalytic Rxn. r r aA Bb r = k [B] [A] = A B a b rA = a k [A][B] rB = b k [A] [B] 30 Effect of Temperature On Reaction Rate Coefficients The temperature dependence of the rate constant is given by Van’t Hoff- Arhenius relationship k _ 2 (T2 T1) k1 For BOD θ = 1.047 for temperatures between 20o - 30 oC Theta remains constant for only a small range of temperatures. 31 Ea ln k1 ln A RT1 Ea ln k2 ln A RT2 Subtracting ln k2 from ln k1 gives Ea 1 1 ln k1 ln k2 ( ) R T2 T1 32 k E 1 1 ln 1 a ( ) k2 R T2 T1 k _ 2 (T2 T1) kk E T T ln 1 a ( 1 2 ) k2 R T1T2 Ea 1 1 ln k1 ln k2 ( ) R T2 T1 33 EXAMPLE: A first order reaction rate constant, k1 at 20°C = 0.23/d.