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Rates of reactions

Rates of reactions . Topic 6. 6.1 Determining the rate of a chemical rxn. What is the rate of a chemical rxn ? The speed at which reactants are used up and products are formed. Experiments to measure rate of rxn. Consider this rxn :

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Rates of reactions

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  1. Rates of reactions Topic 6

  2. 6.1 Determining the rate of a chemical rxn • What is the rate of a chemical rxn? • The speed at which reactants are used up and products are formed

  3. Experiments to measure rate of rxn • Consider this rxn: • CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2(g) + H2O • We can measure the rate of rxn in 2 ways • Measure rate that CO2 is produced • Measure the rate at which the mass decreases

  4. Measuring CO2 production CO2(g) Delivery Tube Measuring Cylinder HCl(aq) CaCO3(s) Water

  5. Measurement of the rate mass decreases Mass decreases as CO2 is given off Measure mass decrease every 10 sec Plot the data Cotton Wool HCl(aq) CaCO3(s) Balance

  6. Rate of rxn defined • Change in concentration of reactants or products per unit of time • (Time could be 1 sec, 1 min, etc) • Average rate= (change in concentration)/(time) • Unit are dm-3 s-1, dm-3 min-1, etc • If a given amount of reactant is used up, the same amount will be produced

  7. 6.2 collision theory • What is the product of particles colliding? • Reactions • For collisions to result, there are two conditions that must be fulfilled • 1. Collision must involve more than a certain minimum amount of energy • 2. Molecules must collide with the correct orientations

  8. Collision must involve a certain min. amount of energy • To react, particles must collide with sufficient energy • The min. energy must result in the activation energy (Ea) • Ea is the minimum amount of energy that colliding particles must posses for a collision to result in a rxn • If particles collide and do not overcome Ea, they will just bounce off each other • Collisions resulting in a reaction are called successful or effective collisions • Not all rxns that overcome Eawill result in rxn

  9. Molecules collide with correct orientation • If molecules do not collide with correct orientation they will not react • Main factors affecting rate of rxn • Concentration of reactants • Pressure for reactions involving gases • Surface area of solid reactants • Temperature • Catalysis

  10. Effects of conc. • Does higher reactant concentrations, how will it effect the rate? • With more particles in a given volume, there will be more collisions • Increases the chances of successful collisions

  11. Effects of pressure • How will increased pressure effect rate? • Similar to concentration increase • Increases collision frequency • Only reactions involving gases are significantly affected by changing pressure

  12. Effect of surface area of solid reactants • Reactions generally only occur on the surface of solids • How can surface area of a solid be increase? • Finely divide the solid to open up more areas for reactions • Allows for more potential collision opportunities

  13. Relationship between temp and the energy of particles in a gas • How does temp effect the movement of particles? • Ideal gas: kinetic energy of the particles in a gas is proportional to its temp in K • If temp is doubled, average energy is generally doubled • Does not depend on identity of gas, however lighter molecules will be traveling faster than heavier ones

  14. Effects of temp on rate • How does temp effect rate? • The rate will generally double for every 10 K increase in temp • Why? • Collision frequency increases due to particles moving faster • The also collide harder • Increases the chance for collision

  15. Maxwell-boltzmann • Distribution of molecular kinetic energies at a particular temp • Only a few particles with high energy and only a few low energy • Most particles have average energy

  16. catalysis • Catalyst A substance that increases the rate of a chemical rxn without itself being used up in the rxn • Often written above the yield arrow in the equation • Also allow the rxn to proceed by an alternative pathway of lower activation energy • May be homogeneous (same state as reactants) or heterogeneous

  17. HL2

  18. 6.3 the rate expression • Rate equation/ expression • Consider A B • Rate is directly proportional to [A] • Rate= k[A] • k is the rate constant • The rate expression is experimentally determined equation relating to the rate of rxn to the concentration of substances in the rxn mixture

  19. Rate expression • General equation • xA + yB C + D • Rate= k[A]m[B]n • Rate constant a constant of proportionality relating the concentrations in the experimentally determined rate expression to the rate of the rxn • Only constant at a particular temp

  20. Order of rxn • In respect to a particular reactant is the power of the reactant’s concentration in the experimentally determined rate equation • m and n are the orders of the reactants • Overall order is m+ n • Rate expression can be determined experimentally

  21. Experimental determination of rate expression • Consider, A + B C • The initial rate can be taken because we know the initial concentrations of A and B • An experiment with a fixed amount of B and varied concentrations of A is performed • Then do the same with fixed A and varied B • Use data to determine orders of A and B

  22. Determining order of reaction and rate expression from experimental data • Given the reaction 2A B

  23. cont • We want to determine: • The order with respect to A • The rate expression • The value of the rate constant (w/ units) • The rate of reaction when [A]= 1.3 mol dm-3

  24. Zero-order rxns • Rate independent of concentration • Rate equation is: rate=k • Units of k are conc. time-1 • Could be mol dm-3 s-1 • moldm-3h-1 • etc.

  25. First –order rxns • Rate is directly proportional to the concentration • Half- life is first-order • Half-life is related to rate constant: • rate constant= 0.693/ half-life • Rate equation: rate=k[A] • Units of k are time-1

  26. Second-order rns • Rate of rxn is proportional to concentration squared • Rate expression: rate=k[A]2

  27. Units of rate constant

  28. 6.4 the arrhenius equation • Shows the variation of the rate constant with temperature • As temp. increases, the rate constant increases exponentially • Equation: k=Ae-Ea/RT • A pre-exponential factor, A-factor or frequency factor • Relates to frequency and orientation of collision • A constant that varies slightly with temperature • e-Ea/RT  the fraction of collision where E> Ea (Energy is greater than activation energy) • Not all reactions where E> Eawill result in collision

  29. The other Arrhenious equation • May also be written as lnk= (-Ea/R) x (1/T) + lnA • This form is used to solve for activation energy, if these procedures are followed • Conduct a series of temperature-varied experiments • Calculate rate constant for each temp. • Plot a graph of lnk (y-axis) vs. 1/T (x-axis) • Temp in K • Slope = -Ea/R (R= gas constant)

  30. Effects of a catalyst on rate constant • If the rate equation is: rate= k[A][B] • Catalyst increases rate constant

  31. 6.5 Mechanisms of reactions • Consider the reaction, 2NO2(g) +F2(g)  2NO2F(g) • If this rxn were to occur in a single step, all 3 molecules would have to collide in correct orientation at the same time • We could assume that if the concentration was increased the chances of proper collision would increase • This makes the rate dependent upon reactant concentrations • Rate equation will be: rate= [NO2]2[F2] • The 2 superscript comes from the coefficient in the balanced equation

  32. Con’t • The rate derived from experimentation was found to be • rate= [NO2][F2] • This suggests that the rxn does not occur in a single step • Thus, this reaction (as many more) must occur in multiple steps • The chances of molecules colliding in perfect orientation simultaneously is quite low

  33. suggestedMechanisms of this rxn • NO2 + F2 NO2F +F Step 1 • NO2 + F  NO2F Step 2 • 2NO2 + F2  2NO2F Overall Equation • Step 1: rate= k1[NO2][F2] • Step 2: rate= k2[NO2][F]

  34. Con’t • Step 1 is the same as the overall equation, thus must be the step determining the rate for the overall rxn • Called the rate-determining step (the slow step) • Step 2 is the fast step and does not influence the overall rate of rxn to a great extent • thus, the concentrations of the step species are not included in the rate equation

  35. Generic example • A + 2B  C • B + B  Q Step 1 rate- determining step • Q + A  C Step 2 fast • Mechanism • B + B  Q • Q+ A  C • 2B + A  C • Thus, rate= k[B]2 since it is the only species remaining from the first step

  36. Rate determining step as second step • B + B  Q Step 1 fast • Q + A  C Step 2 rate-determining step • Process of determining overall rate is basically the same as if the first step as the slow step • Rate= k[Q][A] • Because Q is produced by 2 B molecules in the first step, we can replace [Q] with [B]2 • Rate= k[B]2[A]

  37. Another mechanism • A + B  S step 1 fast • S + B  C step 2 rate-determining step • Reactants involved up to and including the rate determining step are included in rate equation • Rate= k[B]2[A]

  38. Reaction involving a catalyst • CH3COCH3(aq) + I2(aq)  CH3COCH2I(aq) + HI(aq) • The rxn is acid (H+) catalysed • Experimental rate expression is • Rate= k[CH3COCH3][H+] • Does not include I2, so it is only involved after the rate-determining step

  39. Con’t • Proposed mechanism: • CH3COCH3 + H+ X rate-determining step • X + I2  CH3COCH2I + HI + H+ fast • Catalyst is involved in the rate- determining step but is regenerated in second step and does not appear in the overall chemical equation • X is an intermediate • H+ cancels out

  40. Sn1 vs sn2 mechanisms • (CH3)3CBr + OH- (CH3)3COH + Br– • This is a nucleophilic substitution • Experimental rate expression: rate= k[(CH3)3CBr] • Since OH- is not included, it is in the fast step • Suggested mechanism: • (CH3)3CBr  (CH3)3C+ + Br – rate- determining step • (CH3)3C++ OH-  (CH3)3COH fast

  41. Con’t • This is considered SN1 • S= substitution • N= nucelophilic • 1= molecularity of rate- determining step • Molecularity # of ‘molecules’ that react in a particular step (usually rate-determining)

  42. Summary of mechanism rules • Mechanism must agree with overall stoichiometric equation • Maximum of 2 particles can react in any one step • All species in rate equ. Must appear in mechanism in or before the rate-determining step • The power of a particular reactant’s concentration in the rate equ. Indicates the # of times it appears in the mechanism up to and including the rate determining step

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