The Kinetic Molecular Theory

Transcription

1 The Kinetic Molecular Theory KMT and the Collision Theory Unit 13a Copyright John Sayles 1

2 Rate vs. Equilibrium Do not confuse the RATE with EQUILIBRIUM If you re using LeChatelier, you d better be talking about equilibrium If you re talking about collisions, you d better be talking about rates For C (diamond) + O 2(g) -----> CO 2(g) K=3x10 69 YIKES!!! Thank goodness the rate is incredibly slow Rate and Equilibrium are SEPARATE considerations Often, what we do to enhance eq. hoses the rate Copyright John Sayles 2

3 Rate vs. Equilibrium There is a BIG connection For an elementary (one-step) process K eq = k forward /k reverse Also recall the definition of equilibrium rate of forward process = rate of reverse process Copyright John Sayles 3

4 Five Postulates of the KMT Gas molecules are in constant, random, straight-line motion Collisions between molecules are elastic Energy and momentum are conserved KE ave = 3/2kT Ideal gas molecules have negligible volume Ideal gas molecules have negligible IMF s Last two true when gas is expanded True at High Temp, Low Pressure Copyright John Sayles 4

5 Collision Theory - subset of KMT In order for molecules to react, they must collide with Sufficient energy (called the Activation Energy) and, Proper geometry Head-on collision Involving the right parts We ll call these successful collisions The collision serves to break the old bonds Copyright John Sayles 5

6 Importance of Molecular Orientation in the Reaction of NO and CI 2 13_12 Copyright Houghton Mifflin Company. All rights reserved Copyright John Sayles 6

7 The Baseball Analogy Successful collisions are like base hits To get lots of base hits Lots of at-bats (lots of collisions) High batting average (high success rate) Can be done with skillful batter High energy Can be done with lousy pitching Low E act to overcome Fast reaction has a high number of successful collisions Copyright John Sayles 7

8 Collision Theory To increase the rate of a reaction, you must increase the Number of collisions [ ], temp, pressure, surface area do this Catalysts sort of do this or, Energy of collisions Only temp does this It seems like pressure would, too, but NO! Catalysts decrease the energy needed, E act Copyright John Sayles 8

9 Concentration and Rate Increasing concentration increases the number of collisions For a gas, increasing pressure/decreasing volume are ways to increase [ ] Examples: 12 M vs 1 M acid Compression stroke in engine Pressure wave in an explosion Copyright John Sayles 9

10 Temperature and Rate The only double whammy Increase both the number and energy of collisions In biochem rxns, a 10 C increase in temp doubles the reaction rate Examples: Refrigeration Fevers Penny in nitric acid Copyright John Sayles 10

11 Catalysts and Rate Catalysts create an alternate reaction mechanism, one which has a different activated complex with a lower E act Instead of forming A---B, forms A--C--B A--C--B has a lower E act The catalyst frequently acts as a go-between, helping to guarantee proper geometry Lower E act means a higher percentage of collisions are successful Copyright John Sayles 11

12 Enzyme Action (Lock and Key Model) Substrate Products + + Copyright Houghton Mifflin Company. All rights reserved Copyright John Sayles 12

13 Energy per mol The Activation Energy Hump 13_13 + NOCl 2 E a (reverse) = 2 kj/mol NOCl + Cl Products E a (forward) = 85 kj/mol H= 83 kj/mol NO + Cl 2 Reactants Progress of reaction Copyright Houghton Mifflin Company. All rights reserved Copyright John Sayles 13

14 Catalysts and Rate Catalysts have a HUGE impact on rate Negative catalysts decrease the rate of reactions Tetraethyl lead, explosives manufacture Catalysts are not consumed in the overall reaction They are used in one step, but reproduced in a later step A little catalyst goes a long way Examples Enzymes CFC s and ozone Catalytic converter (Pt/Ir alloy) CO + NO x -----> CO 2 + N 2 Copyright John Sayles 14

15 Proposed Mechanism of Catalytic Hydrogenation of C 2 H 4 Copyright Houghton Mifflin Company. All rights reserved Copyright John Sayles 15

16 Automobile Catalytic Converter 13_20 Exhaust manifold Copyright Houghton Mifflin Company. All rights reserved Copyright John Sayles 16

17 Energy (kj) Comparison of Activation Energies in the Uncatalyzed and Catalyzed Decompositions of Ozone 13_17 20 E a (uncatalyzed reaction) 10 E a (catalyzed reaction) 0 Cl + O 3 + O ClO + O 2 + O Cl + O 2 + O Progress of reaction Copyright Houghton Mifflin Company. All rights reserved Copyright John Sayles 17

18 Surface Area and Stirring Increased surface area increases the number of collisions Dissolving is the ultimate increase in SA Stirring exposes the piled up SA Examples Chewing Carburetion Grain elevator explosions Vapor vs. Liquid vs. Solid Copyright John Sayles 18

19 Nature of the reactants and Rate Must break the old bonds for a reaction to occur If these bonds are strong, the rate is going to tend to be low It will take a major wallop to break Very low % of successful collisions We have no control over this Copyright John Sayles 19

20 Rate vs. Equilibrium Temp always increases rate, endo or exo Temp can cause a shift either left or right Higher [ ] reactant always increases rate Higher [ ] reactant always shifts right Catalyst has no effect on equil composition Increases rate of the forward and reverse equally Copyright John Sayles 20