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General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
General chemistry ii chapter 14
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General chemistry ii chapter 14

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  • 1. Chapter 14 Chemical Kinetics CHEMISTRY The Central Science 9th 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 the reaction rate .
    • For a reaction A  B
    • 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.
      • At t = 0 (time zero) there is 1.00 mol A (100 red spheres) and no B present.
      • At t = 10 min, there is 0.54 mol A and 0.26 mol B.
      • At t = 20 min, there is 0.30 mol A and 0.70 mol B.
      • Calculating,
    Calculating Reaction Rates Using Products
  • 6.
    • For the reaction A  B 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/L·s or M/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 constant k is the rate constant.
    Concentration and Rate Equation
  • 14.
    • For a general reaction with rate law
    • we say the reaction is m th order in reactant 1 and n th order in reactant 2.
    • The overall order of reaction is m + n + ….
    • A reaction can be zeroth order if m , 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 is n th order if doubling the concentration causes an 2 n increase 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] t versus t is a straight line with slope - k and intercept 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 base e .
    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 st Order Reaction
  • 19.
    • Second Order Reactions
    • For a second order reaction with just one reactant
    • A plot of 1/[A] t versus t is a straight line with slope k and intercept 1/[A] 0
      • Put into simple math language: y = m x + b
    • For a second order reaction, a plot of ln[A] t vs. t is not linear.
      • However, a plot of 1/[A] t versus 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 nd Order 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 ½ is the time taken for [A] 0 to reach ½[A] 0 .
    • Mathematically defined by:
    • The half-life for a 1 st order reactions depends only on k
    Half-Life Reactions
  • 22.
    • Second Order
    • Mathematically defined by:
    • A second order reaction’s 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 + NOCl  NO + 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 than E a is
      • where R is the gas constant (8.314 J/mol·K).
    Fraction of Molecules Possessing E a
  • 31. Activation Energy, E a , Plot
  • 32.
    • Arrhenius discovered most reaction-rate data obeyed the Arrhenius equation:
      • k is the rate constant, E a is the activation energy, R is the gas constant (8.314 J/K-mol) and T is the temperature in K.
      • A is called the frequency factor.
        • A is a measure of the probability of a favorable collision.
      • Both A and E a are specific to a given reaction.
    The Arrhenius Equation
  • 33.
    • If we have a lot of data, we can determine E a and A graphically by rearranging the Arrhenius equation:
    • From the above equation, a plot of ln k versus 1/ T will have slope of – E a /R and intercept of ln A .
    Determing Activation Energy
  • 34. Ln k versus 1/T
    • E a can 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 2 and 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 on Step 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 2 is unstable, so we express the concentration of NOBr 2 in terms of NOBr and Br 2 assuming there is an equilibrium in step 1 we 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 a for 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 increase k which results in increasing A or decreasing E 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

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