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Chem 1310: Introduction to physical chemistry Part 2: Chemical Kinetics

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Chem 1310: Introduction to physical chemistry Part 2: Chemical Kinetics Peter H.M. Budzelaar Kinetics vs Thermodynamics Thermodynamics show why a reaction wants to ... – PowerPoint PPT presentation

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Title: Chem 1310: Introduction to physical chemistry Part 2: Chemical Kinetics


1
Chem 1310 Introduction to physical chemistry
Part 2 Chemical Kinetics
  • Peter H.M. Budzelaar

2
Kinetics vs Thermodynamics
  • Thermodynamics show why a reaction wants to
    proceed.
  • Kinetics can explain how it proceeds.
  • The two topics are complementaryyou will need
    both to understand chemistry.

3
Kinetics -the rates of chemical reactions
  • How fast does a reaction go?
  • Does the rate change over time?
  • Can it be influenced?
  • What does all this sayabout how the reaction
    proceeds?

4
Convoluted reaction paths
  • Most reactions follow a complicated course.
  • You might see (on paper)
  • CH4 Cl2 CH3Cl HCl
  • But what actually happens is
  • Cl2 hn 2 Cl start ("initiation")
  • Cl CH4 HCl CH3 production
  • CH3 Cl2 CH3Cl Cl ("propagation")
  • 2 Cl Cl2 stop
  • CH3 Cl CH3Cl ("termination")
  • 2 CH3 C2H6

5
Convoluted reaction paths
  • Kinetics, the study of the rates of reactions,
    can help us establish mechanisms and predict what
    will happen in "new" circumstances.

6
What is a rate?
  • Rate of a reactionthe number of moles L-1 s-1
    of reactants passing into products.
  • Why the unitstwice the volume, same
    concentration twice as many molecules go.same
    volume, concentrations, wait twice as long twice
    as many molecules go (approx...).

7
What is a rate?
  • We usually mean the rate at any given moment (the
    instantaneous rate, Dt very small) rather than
    over a whole second or other time spin (the
    average rate, a larger Dt).
  • Don't forget stoichiometry here! If not all
    components in a reaction have coefficient 1, we
    define the "rate of the reaction" as the rate of
    (dis)appearance of the components divided by
    their coefficients.

8
What is a rate? (2)
Example 4 NH3 3 O2 2 N2 6 H2O (using
smallest possible integer coefficients)
  • The rate belongs to the reaction as written.
  • Compare with DH calculations, where we also give
    the result for the equation as written.

9
Rates, rate lawsand elementary steps
  • A "rate law" expresses the dependence of the rate
    on the concentrations of reactants.
  • For an elementary step
  • A X ...
  • we would expect an expression like
  • rate k A
  • Similarly, for
  • A B X ...
  • one would expect
  • rate k AB

10
Rates, rate lawsand elementary steps
  • The k values ("rate constants") depend on the
    temperature, and are different for different
    reactions.
  • Most "real" reactions consist of many steps. If
    we would know them, we could construct an
    expression for the overall rate. This can depend
    on concentrations in a complicated fashion, as we
    will see.
  • Catalysts, compounds that accelerate reactions,
    work by enabling alternative paths (new
    elementary steps), not by affecting rate
    constants of existing elementary steps.

11
Measuring rate laws
  • The dependence of rate on concentrations may be
    simple, as in
  • rate k A
  • or it can be complicated, as in
  • In practice, you measure the dependence of the
    rate on concentrations, then fit to various
    reasonable "laws".

12
Measuring rate lawsfrom initial rates
  • Start reaction at a certain concentration.
  • Measure conversion in the first 0.1 or so
    seconds.
  • Þ initial rate at the original concentration
  • Do the same at 0.5, 0.25, 0.125 etc times the
    original concentration(s).
  • Þ initial rates at different concentrations
  • If there is more than one reactant, vary the
    concentration of each one independently.

13
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14
Measuring rate lawsfrom initial rates
  • Test for first-order kinetics
  • if rate µ A, then etc
  • or plot rate vs A Þ straight line
  • We write rate (in mol L-1 s-1) k A, k in s-1.
  • If there are more components, the reaction may be
    first-order in each rate k AB, k in L
    mol-1 s-1.

15
Measuring rate lawsfrom initial rates
  • If first-order doesn't work, test for
    second-order kinetics
  • if rate µ A2, then
    etc
  • or plot rate vs A2 (or Örate vs A) Þ straight
    line
  • We write rate (in mol L-1 s-1) k A2, k in L
    mol-1 s-1.
  • If the rate does not depend on a particular
    reactant, we say it is zero-order in that
    reactant.

16
Measuring rate lawsfrom reaction progress
17
Measuring rate lawsfrom reaction progress
  • Follow reaction in time, try to fit the curve
    with various models.
  • Do this for several initial concentrations, to
    verify the model!
  • For a first-order reaction
  • Limit for small Dt
  • This is a differential equation.
  • The solution is
  • (verify by differentiation)
  • To check for this model plot ln A vs t

18
Measuring rate lawsfrom reaction progress
  • For a second-order reaction
  • Solution
  • To check for this model plot 1/A vs t
  • For a zero-order reaction
  • Solution
  • (obviously not valid for large t!)

19
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20
First-order reactions and half-life
  • Half-life time to go from a certain initial
    reactant concentration to half this initial
    value.
  • This does not mean that after two half-lifes all
    reactant has been consumed! Rather, every
    half-life reduces the concentration by a factor
    of 2.

21
First-order reactions and half-life
This is called exponential decay.
22
Back to the microscopic model
  • Why are some reactions faster than others?
  • At any given moment, only asmall fraction of the
    moleculeshave enough energy to hopover the
    barrier. The fractionbecomes larger if the
    barrieris reduced (easier, faster reaction)or
    if the temperature is raised(more of the
    molecules haveenough energy).

23
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24
Intermezzo explosive reactions
  • Once a few molecules hop over, the reaction
    produces so much energy that many more can
    follow. Heat is produced faster than it can
    dissipate even via diffusion of the gaseous
    products.
  • A detonator is used to provide the initial bit of
    energy.
  • Ammonium nitrate can decompose explosively.
  • To what?
  • Why would adding kerosene improve the explosive
    power?

25
The Arrhenius expression
  • Nearly all reactions have a similar temperature
    dependence

Frequency factor what fraction of
collisions could in principle lead to a reaction.
Energy term what fraction of molecules will have
enough energy to pass the barrier.
26
The Arrhenius plot
  • Plot ln k vs 1/T
  • ln k ln A -Ea/RT

intercept ln A
slope -Ea/R
27
The rate-limiting step
  • The slowest step in a sequence of elementary
    steps
  • (the "bottleneck").
  • The overall rate is determined by this step.
  • Typically corresponds to the highest barrier
    (activation energy) on the path, since frequency
    factors are usually not too different.

28
The rate-limiting step
  • Overall rate rate of RLS

29
Rates and mechanisms
  • Single-step mechanism
  • rate rate of single step
  • e.g. rate kAB, no problem
  • Multi-step mechanism
  • if first step is RLS,treat as single-step
    (remaining steps don't matter)
  • if later step is RLS, things get complicated
  • rate k2B, but we don't know B. What now?

30
Steady-state approximation
  • When the reaction starts, we have B 0. It
    builds up from A, then starts to deplete as also
    A decreases. For a long time it will be nearly
    constant.

this is what we obtain as kin a first-order rate
law
Question what wouldan Arrhenius plotof "k"
produce now?
31
Catalysis
  • Not faster reactions but new reactions made
    possible by reaction with the catalyst.
  • The catalyst does react!It might be
    regeneratedat the end of the reaction,but it is
    not chemically inert.
  • Catalysis is not "homeopathic"!

32
The best catalyst?
  • What is the most common, most versatile
    catalystin chemistry?

33
The best catalyst?
  • What is the most common, most versatile
    catalystin chemistry?

H
34
Enzymes - the catalysts of nature
  • Enzymes are proteins containing one or more
    regions or groups specifically adapted to promote
    a chemical reaction. This may be a single H/OH-
    catalyzed reaction of something more complicated,
    including transition-metal catalysis.
  • The protein provides a semi-rigid environment
    where the reaction can take place. It is usually
    (shape-)selective, able to distinguish between
    subtly different substrates and to produce
    products with high selectivity.

35
Enzymes - the catalysts of nature
  • Enzymes require water to be stable.
  • They will irreversibly denature at high
    temperatures.
  • This is an example of catalyst deactivation.
  • Enzymes can get poisoned by molecules that bind
    to them but do not undergo the desired reaction.

36
Enzyme kinetics
  • Typical behaviour
  • Steady-state (assume ES is constant,
    EESE0)
  • For small S (like a
    regular two-step reaction)
  • For large S (all
    enzyme present as E-S
    complex, rate limited by
    E0)
  • "Saturation kinetics"

37
Man-made catalysts
  • Homogeneous in same phase as reactants(usually
    solution)
  • Heterogeneous different phases(usually solid
    catalyst, gaseous reactants)

38
Homogeneous catalysts
  • Typically small, well-defined transition-metal
    complexes. Designed to do a specific type of
    reaction on a specific type of functional group,
    for a class of substrates.
  • Not as specific/selective as enzymes.
  • Work under mild conditions (0-120C).
  • Used for synthesis of fine chemicals/pharmaceutica
    ls.

39
An example ofa homogeneous catalyst
40
Heterogeneous catalysts
  • Typically metals, metal oxides or silicate-like
    materials with a range of "active sites".
  • Tend to produces mixtures of products, with
    distributions determined by product stabilities.
  • Used in petrochemical industry and oil refining.

41
An example ofa heterogeneous catalyst
  • Via dissociation to atoms.
  • Works equally well for all NxOy compounds.
  • Products are N2 and O2 because N-O bond is weak.
  • Requires high temperature to work.
  • Can be poisoned lead in gasoline precipitates on
    surfaces, blocks active sites.
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