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Chemical Kinetics

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AP Chapter 14 Activation Energy Activated Complex Heat of Reaction Activation Energy Activated Complex Heat of Reaction Because the kinetic energy of molecules ... – PowerPoint PPT presentation

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Title: Chemical Kinetics


1
Chemical Kinetics
  • AP Chapter 14

2
Chemical Kinetics
  • Chemical kinetics is the area of chemistry that
    involves the rates or speeds of chemical
    reactions.
  • The more collisions there are between molecules,
    the faster the reaction occurs.

3
Factors that Affect Reaction Rates
  1. The physical state of the reactants gases and
    liquids tend to react faster than solids - think
    surface area.
  2. The concentrations of reactants.
  3. The temperature at which the reaction occurs.
  4. The presence of a catalyst.

4
Reaction Rates
  • The speed of a chemical reaction is its reaction
    rate.
  • Reaction rates are usually expressed as changes
    in concentration per unit of time. (M/s)

5
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6
Reaction Rates
  • The preceding slide shows that the rate of the
    reaction can be expressed as either the rate of
    disappearance of reactant A, or the appearance of
    product B.
  • The overall concentration does not change!

?A ?t
Avg rate of disappearance of A
7
Rate Data
  • C4H9Cl(aq) H2O(l) ? C4H9OH(aq) HCl(aq)
  • It is typical for rates to decrease as a reaction
    proceeds, because the concentration of the
    reactants decreases.

8
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9
Instantaneous Rate
  • An instantaneous rate is the rate at a particular
    moment in the reaction.
  • It is determined from the slope of the curve at
    the point of interest.

10
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11
Instantaneous Rate
?C4H9Cl ?t
Instantaneous rate
12
Reaction Rates and Stoichiometry
  • In the preceding reaction, 1 mol of C4H9OH is
    produced for every mol of C4H9Cl consumed. (this
    is a 11 ratio)
  • Not all equations have this ratio
  • 2HI(g) ? H2(g) I2(g)

13
Reaction Rates - Stoichiometry
  • For a general equation
  • a A b B ? c C d D

1 ?A 1 ?B 1 ?C 1 ?D a ?t
b ?t c ?t d ?t
Rate



14
Beers Law
  • The absorption of electromagnetic radiation by a
    substance at a particular wavelength is directly
    proportional to its concentration.

15
Concentration
  • The quantitative relationship between rate and
    concentration is expressed by rate law
  • Rate kreactant 1mreactant 2n . . .
  • The constant k is the rate constant
  • The exponents m, n, etc are the reaction orders
    for the reactants.
  • The sum of the reaction orders gives the overall
    reaction order.

16
Rate Laws
  • The exponents in a rate law determine how the
    rate is affected by the concentration of each
    reactant.
  • The values of these exponents must be determined
    experimentally
  • In most rate laws, reaction orders are 0, 1, and
    2.

17
Rate Laws
  • Rate laws can be used to determine the
    concentration of reactants or products at any
    time during a reaction.
  • A first-order reaction is a reaction whose rate
    depends on the concentration of a single
    reactant raised to the first power.
  • This form of a rate law is called a differential
    rate law.

18
First Order Reactions
  • This uses the following calculation
  • lnAt - lnA0 -kt OR
  • ln -kt

At A0
19
m is the slope and b is the y-intercept. A
reaction that is not first order will not yield a
straight line.
20
Kinetic data for the conversion of methyl
isonitrile. In (b) above, the plot of the
natural log as a function of time yields a
straight line, which confirms that it is a first
order reaction.
21
Second Order Reactions
  • A second order reaction is one whose rate depends
    on the reactant concentration raised to the
    second power or on the concentrations of 2
    different reactants, each raised to the first
    power.

22
Second-Order Reactions
?A ?t
  • Rate kA2
  • kt OR kt
  • One way to tell the difference between first-
    and second-order rate laws is to graph both
    lnAt and 1/At. If the lnAt plot is linear,
    the reaction is first order. If the 1/At plot
    is linear, the reaction is second order.

1 At
1 A0
1 At
1 A0
23
Half-Life
  • Half-life of a reaction, t1/2 is the time
    required for the concentration of a reactant to
    reach ½ of its initial value, At1/2 ½A0.
  • In a first-order reaction, the concentration of
    the reactant decreases by ½ in each of a series
    of regularly spaced time intervals, t ½ .

24
Half Life of a First Order Reaction
25
Temperature and Rate
  • The rate of most chemical reactions increase as
    the temperature increases.

26
The Collision Model
  • Reactions occur as a result of collisions between
    molecules.
  • This is why the magnitude of rate constants
    increase with increasing temperature.
  • The greater the kinetic energy of the colliding
    particles, the greater the energy of the
    collision.

27
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28
Activation Energy
  • Activation energy is the minimum amount of energy
    needed for a reaction to occur.
  • Ea
  • A collision with energy Ea or greater can cause
    the atoms of the colliding molecules to reach
    the activated complex, or transition state.
  • The activated complex is a particular
    arrangement of atoms at the top of the barrier
    (or hill.)

29
Activation Energy
A.E.
Reactants
Products
Potential Energy
Reaction Coordinate
30
Activated Complex
Activated Complex
Reactant
Product
Potential Energy
Reaction Coordinate
31
Heat of Reaction
Reactant
?H
Product
Potential Energy
Reaction Coordinate
32
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33
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34
Arrhenius equation
  • Because the kinetic energy of molecules depends
    on the temperature, the rate constant of a
    reaction is very dependent on temperature.
  • This relationship is given by the Arrhenius
    equation
  • ln k - ln A

k is the rate constant, Ea R gas constant,
8.314 J/mol-K T is the absolute temperature
Ea RT
35
Arrhenius equation
  • The term A is called the frequency factor.
  • It relates to the number of collisions that are
    favorably oriented for a reaction.
  • It remains nearly constant as the temperature
    varies.
  • Reaction rates decrease as Ea increases.

36
Reaction Mechanisms
  • Reaction mechanisms detail the individual steps
    that occur in a chemical reaction.
  • Each step, called an elementary reaction, has a
    well-defined rate law that depends on the number
    of molecules (molecularity) in the step.
  • These are either unimolecular, (1 reactant
    molecule), bimolecular (2) or termolecular, 3.
  • The follow rate laws unimolecular follows first
    order overall, etc.

37
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38
Multistep Mechanisms
  • Multistep reactions involve two or more
    elementary reactions, or steps.
  • An intermediate that is produced in one
    elementary step is consumed in another later step
    and doesnt show up in the overall equation.

39
Rate Determining Step
  • The slowest elementary step is the one that
    limits the speed of the overall reaction and is
    called the rate-limiting step, or rate
    determining step.
  • The rate determining step governs the rate law
    for the overall reaction.

40
Catalysts
  • A catalyst is a substance that changes the speed
    of a chemical reaction without taking part in the
    chemical process itself.
  • It provides an alternate pathway for the reaction
    to occur.

41
Homogeneous catalysts
  • A homogeneous catalyst is a catalyst that is
    present in the same phase as the reacting
    molecules.

42
Heterogeneous catalyst
  • A heterogeneous catalyst is a catalyst that
    exists in a different phase from the reactant
    molecules, like a solid catalyst involved with
    gaseous reactant molecules.

43
Enzymes
  • Biological catalysts that are necessary for many
    large inter-related chemical reactions that occur
    in body systems are called enzymes.
  • These body system reactions must occur at
    suitable/specific rates in order to sustain life.

44
Enzymes
  • Active site the specific location on the enzyme
    where the reaction occurs.
  • Substrate - the substance that undergoes a
    reaction at the active site.
  • Lock and key model a model that shows how
    enzyme actions occur.
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