General chemistry ii chapter 14

Download General chemistry ii chapter 14

Post on 18-Nov-2014

2.462 views

Category:

Education

9 download

DESCRIPTION

 

TRANSCRIPT

  • 1. Chapter 14 Chemical Kinetics CHEMISTRY The Central Science9th Edition

2.

  • Kinetics is the study of how fast chemical reactions occur.
  • There are 4 important factors which affect the rates of chemical reactions:
    • reactant concentration,
    • temperature,
    • action of catalysts, and
    • surface area.

Kinetics 3.

  • The speed of a reaction is defined as the change that occurs per unit time.
    • It is determined by measuring the change in concentration of a reactant or product with time.
    • The speed the of the reaction is called thereaction rate .
  • For a reaction AB
  • Suppose A reacts to form B.Let us begin with 1.00 mol A.

Reaction Rates 4. Change in Concentration of Reactions 10 20 5.

    • Att= 0 (time zero) there is 1.00 mol A (100 red spheres) and no B present.
    • Att= 10 min, there is 0.54 mol A and 0.26 mol B.
    • Att= 20 min, there is 0.30 mol A and 0.70 mol B.
    • Calculating,

Calculating Reaction Rates Using Products 6.

  • For the reaction AB there are two ways of measuring rate:
    • the speed at which the products appear (i.e. change in moles of B per unit time), or
    • the speed at which the reactants disappear (i.e. the change in moles of A per unit time).
    • The equation, when calculating rates of reactants, is multiplied by -1 to compensate for the negative concentration.
      • By convention rates are expressed as positive numbers.

Calculating Reaction Rates Using Reactants 7.

  • Most useful units for rates are to look at molarity.Since volume is constant, molarity and moles are directly proportional.
  • Consider:
  • C 4 H 9 Cl( aq ) + H 2 O( l )C 4 H 9 OH( aq ) + HCl( aq )

Reaction Rates Reactants Products 8. Reaction Rates for C 4 H 9 Cl 9.

  • C 4 H 9 Cl( aq ) + H 2 O( l )C 4 H 9 OH( aq ) + HCl( aq )
    • We can calculate the average rate in terms of the disappearance of C 4 H 9 Cl.
    • The units for average rate are mol/Ls orM/s .
    • The average rate decreases with time.
    • We plot [C 4 H 9 Cl] versus time.
    • The rate at any instant in time (instantaneous rate) is the slope of the tangent to the curve.
    • Instantaneous rate is different from average rate.
    • We usually call the instantaneous rate the rate.

Properties of C 4 H 9 Cl Reaction 10. Instantaneous Reaction Rates for C 4 H 9 Cl 11.

  • For the reaction
  • C 4 H 9 Cl( aq ) + H 2 O( l )C 4 H 9 OH( aq ) + HCl( aq )
  • we know
  • In general for
  • a A+ b B c C+ d D

Reaction Rates and Stoichiometry 12.

  • In general rates increase as concentrations increase.
  • NH 4 + ( aq ) + NO 2 - ( aq )N 2 ( g ) + 2H 2 O( l )

Concentration and Rate Table 13.

  • For the reaction
  • NH 4 + ( aq ) + NO 2 - ( aq )N 2 ( g ) + 2H 2 O( l )
  • we note
    • as [NH 4 + ] doubles with [NO 2 - ] constant the rate doubles,
    • as [NO 2 - ] doubles with [NH 4 + ] constant, the rate doubles,
    • We conclude rate[NH 4 + ][NO 2 - ].
  • Rate law:
  • The constantkis the rate constant.

Concentration and Rate Equation 14.

  • For a general reaction with rate law
  • we say the reaction ism th order in reactant 1 andn th order in reactant 2.
  • The overall order of reaction ism + n +.
  • A reaction can be zeroth order ifm ,n , are zero.
  • Note the values of the exponents (orders) have to be determined experimentally.They are not simply related to stoichiometry.

Exponents in the Rate Law 15.

  • A reaction is zero order in a reactant if the change in concentration of that reactant produces no effect.
  • A reaction is first order if doubling the concentration causes the rate to double.
  • A reaction isn th order if doubling the concentration causes an 2 nincrease in rate.
  • Note that the rate constant does not depend on concentration.

Determining Order of Reactions 16.

  • First Order Reactions
  • Goal: convert rate law into a convenient equation to give concentrations as a function of time.
  • For a first order reaction, the rate doubles as the concentration of a reactant doubles.

Concentration Change with Time 17.

  • A plot of ln[A] tversus t is a straight line withslope - kandintercept ln[A] 0 .
    • Put into simple math language : y =m x +b
  • Plotting of this equation for given reaction should yield a straight line if the reaction is first order.
  • In the above we use the natural logarithm, ln, which is log to the basee .

Plotting First Order Reactions 18.

  • Left graph plotted with pressure vs. t, and right graph plotted with ln(pressure) vs. t.

Example Plots of a 1 stOrder Reaction 19.

  • Second Order Reactions
  • For a second order reaction with just one reactant
  • A plot of 1/[A] tversus t is a straight line withslopekandintercept 1/[A] 0
    • Put into simple math language:y =m x +b
  • For a second order reaction, a plot of ln[A] t vs.tis not linear.
    • However, a plot of 1/[A] tversus t is a straight line

Concentration Change with Time 20.

  • Left is Ln[NO 2 ] vs t, and right is 1/[NO 2 ] vs. t.

Example plots of a 2 ndOrder Reaction 21.

  • First Order Reactions
  • Half-life is the time taken for the concentration of a reactant to drop to half its original value.
  • For a first order process, half life,t is the time taken for [A] 0to reach [A] 0 .
  • Mathematically defined by:
  • The half-life for a 1 storder reactions depends only on k

Half-Life Reactions 22.

  • Second Order
  • Mathematically defined by:
  • A second order reactions half-life depends on the initial concentration of the reactants

Half-Life Reaction 23.

  • The Collision Model
  • Most reactions speed up as temperature increases.(E.g. food spoils when not refrigerated.)
  • When two light sticks are placed in water: one at room temperature and one in ice, the one at room temperature is brighter than the one in ice.
  • The chemical reaction responsible for chemiluminescence is dependent on temperature: the higher the temperature, the faster the reaction and the brighter the light.

Temperature and Rate 24. The Collision Model

  • As temperature increases, the rate increases.

25.

  • Goal: develop a model that explains why rates of reactions increase as concentration and temperature increases.
  • The collision model: in order for molecules to react they must collide.
  • The greater the number of collisions the faster the rate.
  • The more molecules present, the greater the probability of collision and the faster the rate.
  • Faster moving molecule collide with greater energy and more frequently, increasing reaction rates.

Collision Model: The Central Idea 26.

  • The Collision Model
  • The higher the temperature, the more energy available to the molecules and the faster the rate.
  • Complication: not all collisions lead to products.In fact, only a small fraction of collisions lead to product.
    • Why is this?
  • The Orientation Factor
  • In order for reaction to occur the reactant molecules must collide in the correct orientation and with enough energy to form products.

The Speed of a Reaction 27.

  • Consider:
  • Cl + NOClNO + Cl 2
  • There are two possible ways that Cl atoms and NOCl molecules can collide; one is effective and one is not.

The Orientation Factor 28.

  • Arrhenius: molecules must posses a minimum amount of energy to react.Why?
    • In order to form products, bonds must be broken in the reactants.
    • Bond breakage requires energy.
  • Activation energy,E a , is the minimum energy required to initiate a chemical reaction.

Activation Energy 29. Energy Profile for Methly Isonitrile 30.

  • How does a methyl isonitrile molecule gain enough energy to overcome the activation energy barrier?
  • From kinetic molecular theory, we know that as temperature increases, the total kinetic energy increases.
  • We can show the fraction of molecules,f , with energy equal to or greater thanE ais
    • whereRis the gas constant (8.314 J/molK).

Fraction of Molecules Possessing E a 31. Activation Energy, E a , Plot 32.

  • Arrhenius discovered most reaction-rate data obeyed the Arrhenius equation:
    • kis the rate constant,E a is the activation energy,Ris the gas constant (8.314 J/K-mol) andTis the temperature in K.
    • Ais called the frequency factor.
      • Ais a measure of the probability of a favorable collision.
    • BothAandE aare specific to a given reaction.

The Arrhenius Equation 33.

  • If we have a lot of data, we can determineE aandAgraphically by rearranging the Arrhenius equation:
  • From the above equation, a plot of lnkversus 1/ Twill have slope of E a /Rand intercept of lnA .

Determing Activation Energy 34. Lnkversus 1/T

  • E acan be determined by finding the slope of the line.

35.

  • The balanced chemical equation provides information about the beginning and end of reaction.
  • The reaction mechanism gives the path of the reaction (i.e., process by which a reaction occurs).
    • Mechanisms provide a very detailed picture of which bonds are broken and formed during the course of a reaction.
  • Elementary Steps
  • Elementary step: any process that occurs in a single step.

Reaction Mechanisms 36.

  • Molecularity: the number of molecules present in an elementary step.
    • Unimolecular: one molecule in the elementary step,
    • Bimolecular: two molecules in the elementary step, and
    • Termolecular: three molecules in the elementary step.
  • It is not common to see termolecular processes (statistically improbable).

Properties of the Elementary Step Process 37.

  • Some reaction proceed through more than one step:
    • Consider the reaction of NO 2and CO
  • NO 2 ( g ) + NO 2 ( g )NO 3 ( g ) + NO( g )
  • NO 3 ( g )+ CO( g )NO 2 ( g ) + CO 2 ( g )
    • Notice that if we add the above steps, we get the overall reaction:
  • NO 2 ( g ) + CO( g )NO( g ) + CO 2 ( g )
  • The elementary steps must add to give the balanced chemical equation.
  • Intermediate : a species which appears in an elementary step which is not a reactant or product

Multistep Mechanisms 38.

  • Rate Laws for Elementary Steps
  • The rate law of an elementary step is determined by its molecularity:
    • Unimolecular processes are first order,
    • Bimolecular processes are second order, and
    • Termolecular processes are third order.
  • Rate Laws for Multistep Mechanisms
  • Rate-determining step: is the slowest of the elementary steps.

Reaction Mechanisms 39.

  • Rate Laws for Elementary Steps

Table 14.3, Page 551 40.

  • It is possible for an intermediate to be a reactant.
  • Consider
  • 2NO( g ) + Br2( g )2NOBr( g )
  • The experimentally determined rate law is
          • Rate = k[NO] 2 [Br2]
  • Consider the following mechanism

Initial Fast Step of a Mechanisms 41.

  • The rate law is (based onStep 2 ):
  • Rate =k 2 [NOBr 2 ][NO]
  • The rate law should not depend on the concentration of an intermediate (intermediates are usually unstable).
  • Assume NOBr 2is unstable, so we express the concentration of NOBr 2in terms of NOBr and Br 2assuming there is an equilibrium instep 1we have

Intermediate as a Reactant 42.

  • By definition of equilibrium:
  • Therefore, the overall rate law becomes
  • Note the final rate law is consistent with the experimentally observed rate law.

Intermediate as a Reactant Conts. 43.

  • A catalyst changes the rate of a chemical reaction.
    • Catalyst lower the overall E afor a chemical reaction.
  • There are two types of catalyst:
    • Homogeneous and heterogeneous
    • Example: Cl atoms are catalysts for the destruction of ozone.
    • Homogeneous Catalysis
  • The catalyst and reaction is in one phase.
  • Heterogeneous Catalysis
  • The catalyst and reaction exists in a different phase.

Catalysis 44. The Effects of a Catalyst 45.

  • Catalysts can operate by increasing the number of effective collisions (i.e., from the Arrhenius equation: catalysts increasekwhich results in increasingAor decreasingE a .
  • A catalyst may add intermediates to the reaction.
    • Example: In the presence of Br - , Br 2 ( aq ) is generated as an intermediate in the decomposition of H 2 O 2 .

Functions of the Catalysis 46. End of Chapter 14 Chemical Kinetics

Recommended

View more >