PHA 111 Chemical Kinetics PDF

PHA 111 Chemical Kinetics: Rate & Order of Reaction Dr Stephen Childs Senior Lecturer in Pharmaceutical Chemistry [email protected] Slide 1 PHA 111: Chemical Kinetics  Recommended R...

PHA 111 Chemical Kinetics: Rate & Order of Reaction Dr Stephen Childs Senior Lecturer in Pharmaceutical Chemistry [email protected] Slide 1 PHA 111: Chemical Kinetics  Recommended Reading Atkins’ Physical Chemistry Chapters 21.2 – 21.5 (9th Edition) Slide 2 PHA 111: Chemical Kinetics  Why study rates of reaction? To better understand reaction mechanisms eg. SN1 & SN2 To optimise reaction rates Improve ! side-products ↑ yield / reduce To minimise drug degradation Predict / improve shelf-life eg. pH dependant hydrolysis of aspirin to salicylic acid Understand what drives a reaction forward H2 + ½ O2 → H2O ΔH = -286 kJ/mol Slide 3 PHA 111: Chemical Kinetics  What affects reaction rate? Physical state Many pharmaceutically relevant reaction occur in solution Concentration Molecules must come in contact to react Increased concentration = increased frequency of collisions Temperature Increases frequency of collision Increases the (vibrational) energy of the bonds Catalysts increase the rate of reaction Alternative reaction mechanism Slide 4 PHA 111: Chemical Kinetics  Collision Theory Rate of reaction will depend on how often reacting molecules collide with sufficient energy to react. Collision rate depends on: – Concentration: chances of collision increased, depends on reaction order! – Pressure : for a gas, rate increases with pressure (same mechanism as concentration) 𝒏 𝒑= 𝒙 𝑹𝑻 𝑽 – Surface area and molecular orientation – Temperature and molecular speed Slide 5 PHA 111: Chemical Kinetics  Effect of Surface Area Collision rate depends available surface area Mg(s) + 2H+ (aq) -> Mg2+(aq) + H2(g) al dispersed - much more surface outside crea only available to react wr Slide 6 PHA 111: Chemical Kinetics  Effect of Molecular Orientation Ethere ecoudea O other ↑ replech electroetive na - Big ex enzyme 1" Must have correct orientation – Not all collisions will produce a reaction, – Consider HCl and ethene, (electrophilic addition) – Less likely the more complex the reactants – Only occurs 1 in 105 collisions for complex reactions Slide 7 PHA 111: Chemical Kinetics  Effect of Molecular Speed Collision rate depends on molecular speed – Temperature affects speed – Based on kinetic theory of gases 3RT Vrms = - root mean M squared – At 5°C Vrms = 240 m/s (120 amu) ↓ slightly ↑ as ↑ temperature ↑ temperature ↑ 4 volume speed => – At 20°C Vrms = 247 m/s (120 amu) (90s) Slide 8 PHA 111: Chemical Kinetics  Activation Energy Favourable ( entropy enthalpy not energetically bonds - too many Energy barrier associated with the transition state – Distortion of molecular shape some put energy kick Sufficient energy to = (normal) thea – Energy barrier usually 50 – 100 kJ/mol ↓ – Average energy at 20°C approx. 4 kJ/mol ge – Just Most molecules have insufficient energy – Only about 1 in 109 reaction ? Slide 9 PHA 111: Chemical Kinetics  Maxwell-Boltzmann Distribution For gases, the number of molecules with sufficient energy to react can be represented as below: Many particles · don't have enough energy to react ! James Clerk Maxwell 4 50k]/mol Ludwig Boltzmann Slide 10 PHA 111: Chemical Kinetics  Effect of Temperature Increasing Temperature disproportionally increased the number of particles with sufficient activation energy [ ↑ temp. ↑ the number of moves to the right Particles can - No. of 293 ° K react 303 ° K - Particles smal 4 gives large fraction of particles average available for the - energy reaction over the system& Activation Energy Energy Slide 11 PHA 111: Chemical Kinetics  Arrhenius Equation Increasing temperature affects rate by increasing k From a small number of experiments we can deduce k at varying temperatures, as well as determining t 50 and t90 haf shelf - life - life 𝐸𝑎 Arrhenius equation: ln 𝑘𝑟 = ln 𝐴 − 𝑅𝑇 Alternative form of Arrhenius equation relates k to temperature: * k = Ae-Ea/RT – k: kinetic rate constant - indio ) ? – A: frequency factor (a.k.a. pre-exponential factor) – Ea: activation energy (probability collision will result in reaction) – R: gas constant (8.3145 J mol-1 deg-1) – T: temperature (in Kelvin) Slide 12 PHA 111: Chemical Kinetics  Effect of Temperature k = Ae-Ea/RT A is rate constant if energy barrier is absent – Also if T is infinite (or very high) e-Ea/RT is the fraction of molecules with sufficient energy to react Assuming an activation energy or 50,000 J/mol, and temperatures of 293 and 303°K respectively, the fraction of molecules able to react (e) 20 ° C 20 C · is almost doubled by a 10°C increase in temperature: most doubled with 10c Increase Slide 13 PHA 111: Chemical Kinetics  Catalysts Catalysts increase the rate of reaction without being consumed – Acid catalysed hydrolysis of esters – Chlorine radicals catalyse breakdown of ozone Catalysts provide a lower activation energy pathway for the reaction half-way ex) up the hill > - Tunnel New AE Original AE X 104 Increase in rote I Energy Slide 14 PHA 111: Chemical Kinetics  Catalysts Catalysts do not affect the position of equilibrium of a - reaction Catalysts do not make unfavourable reactions favourable - – Overall G does not change only speeding up ! Enzymes are proteins which act as biological catalysts Acetylcholinesterase hydrolyses acetylcholine in 100 s L transmission of signals in the brain bindstrongly Slide 15 PHA 111: Chemical Kinetics  Rate of Reaction Following concentration as a function of time A+B→C - Decrease in reactant concentration (A or B) Increase in product concentration (C) −Δ𝐴 −Δ𝐵 Δ𝐶 𝑅𝑎𝑡𝑒 = = = Δ𝑡 Δ𝑡 Δ𝑡 Difference over a finite time Want change over infinitesimally small time: −𝑑𝐴 −𝑑𝐵 𝑑𝐶 𝑅𝑎𝑡𝑒 = = = 𝑑𝑡 𝑑𝑡 𝑑𝑡 Slide 16 PHA 111: Chemical Kinetics  Measuring Rate of Reaction Slide 17 PHA 111: Chemical Kinetics  Chemical Reactions Most reactions are multi-step Molecularity Number of molecules involved in each elementary reaction step Can be obtained from the stoichiometry 1 species = unimolecular 2 species = bimolecular 3 species = termolecular Molecularity must be an integer Not the same as reaction order! Slide 18 PHA 111: Chemical Kinetics  Rate of Reaction What happens when we increase [A] ? 1) Nothing happens: Rate independent of [A] ~ Kineticrate a 2) Rate = k[A] Double [A], k is double Triple [A], k is tripled 3) Rate = k[A]2 [A] is double, k is quadrupled (22 = 4) [A] is tripled, k is increased nine-fold (32 = 9) Slide 19 PHA 111: Chemical Kinetics  Reaction Order (i) Rate Multiplication Reaction Order Rate constant (k) units 1 0 M/s 2 1 1/s 4 2 1/(M*s) Units of k must give M/s when combined with concentration units (M) conc. /s The rate law is the sum of each of the reactant orders. If aA +bB + cC → dD + eE rate = k[A]x[B]y[C]z where x is the order of A, not the stoichiometry! Slide 20 PHA 111: Chemical Kinetics  Reaction Order -linear Gurva Slide 21 PHA 111: Chemical Kinetics  Reaction Order (ii) Calculate the reaction order for each reactant, and the overall reaction order for the following reaction: F₂ + 2ClO₂ → 2ClO₂F Experiment No. [Fl₂] (M) [ClO₂] (M) Initial Rate (M/s) 1 0.10 0.010 0.0012 2 3 0.10 0.20 ( X2 0.040 0.010 2x4 0.0048 0.0024 W Doubling [F₂] doubles the rate – therefore must be 1st order m Doubling [ClO₂] doubles the rate – therefore must also be 1st order - Overall order is 1+1 = 2nd order - Therefore the rate law is k [F₂] [ClO₂] Slide 22 PHA 111: Chemical Kinetics  Reaction Order Example Calculate the reaction order for each reactant, and the overall reaction order for the following reaction: BrO3- + 5Br- + 6H+ → 3 Br2 + 3H2O Experiment [BrO3-] (M) [Br-] (M) [H+ ] (M) Initial Rate (M/s) No. 1 0.10 0.10 0.10 8.0 x 10-4 (x2)x 2 0.20 (xz 0.10 0.10 1.6 x 10-3 3 0.20 0.20 0.10 3.2 x 10-3 4 0.10 0.10 0.20 3.2 x 10-3 ist ist and Doubling [BrO3-] doubles the rate – therefore must be 1st order Doubling [Br-] doubles the rate – therefore must also be 1st order Doubling [H+] quadruples the rate – therefore must also be 2nd order Therefore the rate law is k [BrO3-]1 [Br-]1 [H+]2 (overall order = 4) Slide 23 PHA 111: Chemical Kinetics  Rate Law Examples The reaction between A + B is measured in the lab: rate = k[A][B] Order or reaction w.r.t. A is first Order or reaction w.r.t. B is first Overall order is 2 rate = k[B]2 Order or reaction w.r.t. A is zero Order or reaction w.r.t. B is second Overall order is 2 rate = k[A] Order or reaction w.r.t. A is first Order or reaction w.r.t. B is zero Overall order is 1 Slide 24 PHA 111: Chemical Kinetics  Reaction Mechanisms SN1 reaction: tert-Butyl chloride and hydroxyl anion Rate = k[(CH3)3CBr] Rate order of (CH3)3CBr? Rate order of OH? effect ~ no of reaction on rate (knowing these might let us determine the mechanism!) Step 1 - (rate determining) unimolecular ↳ Step 2 - Fast (does not affect overall rate of reaction) = So what is the effect of increasing [OH]? Slide 25 PHA 111: Chemical Kinetics  Reaction Mechanisms SN2 reaction: Chloromethane and hydroxyl anion Rate = k[(CH3)3CBr][OH-] Both components effect the rate of run Rate order of (CH3)3CBr? Rate order of OH? Bimolecular Slide 26 PHA 111: Chemical Kinetics  Determining Mechanisms For a reaction between A + B, rate =k[A][B] L means both must Which mechanism is correct? involved in step-contribute slow to the rote of rxn O If the slow step is first in the mechanism the orders tell you what is taking part in that step. First order w.r.t. A and B, so both must take part. Must be mechanism 2 (Cant be 1, as B isn’t in the slow step!) Slide 27 PHA 111: Chemical Kinetics  Determining Mechanisms What if the slow step is not first in the mechanism? Rate = k[A][X] But we don’t start with X so we can’t measure it! We need to derive a new rate equation based on the equilibrium: [𝑋] 𝐾𝑐 = 𝐴 [𝐵] Re-arrange and substitute into initial rate expression: 𝑅𝑎𝑡𝑒 = 𝑘 𝐴. 𝐾𝑐 [A][B] Combine the constants to get the new rate expression: 𝑅𝑎𝑡𝑒 = 𝐾1 [𝐴]2 [B] > - when slow step is not first during the process Slide 28 PHA 111: Chemical Kinetics