Energy and Kinetics

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					Ch. 6 Learning Goals: Kinetics
• Calculate average and instantaneous rates of
  reaction from data in tables and graphs.
• Sketch graphs of [R] vs. time and [P] vs. time.
• Use stoichiometric relationships to calculate rates
  of consumption and production.
• Propose methods of measuring the rate of a
  reaction.
• Explain, using collision theory and potential energy
  diagrams, how factors such as temperature, the surface
  area of the reactants, the nature of the reactants, the
  addition of catalysts, and the concentration of the
  solution control the rate of a chemical reaction.
• Draw and describe simple potential energy diagrams of
  chemical reactions (e.g., the relationships between the
  relative energies of reactants and products and the
  activation energy of the reaction).
• Explain the significance of the Maxwell-Boltzmann
  distribution for two different temperatures.
• Calculate the rate law, including the value and units for k.
• Use the rate law to calculate the rate of a reaction.
• Explain how the rate of a reaction is determined by
  the series of elementary steps that make up the
  overall reaction mechanism.
• Identify the rate determining step and relate the rate
  of the overall reaction to the rate law of the RDS.
• Use appropriate terminology related to rates of
  reaction, including, but not limited to: activation
  energy, endothermic, exothermic, potential energy
  diagram, orientation, reaction rate, elementary step,
  reaction mechanism, reaction intermediate, rate
  determining step
• Conduct an experiment to gather data for the
  purpose of determining the order of a reaction with
  respect to a particular reactant.
• Plan and conduct an inquiry to determine how
  various factors (e.g., change in temperature,
  addition of a catalyst, increase in surface area of a
  solid reactant) affect the rate of a chemical reaction.
Ch. 6: Chemical Kinetics
The fact that a reaction occurs tells us nothing about the rate at
which it occurs.
• rate of reaction – the speed at which a chemical
  reaction occurs
• usually expressed as change in concentration of
  reactants or products over time
• r = Δc
             units?
      Δt




           Sketch a graph of c vs. t for [product]
Graph for
[reactant]
vs t?
    Graphing and Related Analytical Skills
•   pencil, ruler, as large as possible
•   Title
•   concentration (y-axis)
•   time (x-axis)
•   plot pencil dot points
•   draw line of best fit for linear data, smooth
    curve for non-linear data
average rate = slope of secant
instantaneous rate = slope of
tangent
Consider:
Mg(s) + 2 HCl(aq) --> H2(g) + MgCl2(aq)

The rate of this reaction can be measured by:
a) measuring change in conductivity over time
   (MgCl2(aq)) or colour intensity
   (spectrophotometer)
b) measuring change in pH over time ([H+(aq)])
   (pH probe)
c) measuring volume and/or pressure of H2(g)
   over time
  Stoichiometry and Reaction Rate
2A+3B C + 4D
rate of consumption of A   rate of production of C
          -Δ[A]                     Δ[C]
          Δt                        Δt

            -Δ[A] :   Δ[C] = 2 : 1
             Δt       Δt

   -½ Δ[A] = -⅓ Δ[B] = Δ[C] = ¼ Δ[D]
       Δt         Δt      Δt       Δt
                   Try this!
• Lab 6.1.1 (p. 390)
• Complete a-f

HW:
• p. 360 #1,2
• p. 361 #1-5
     Factors Affecting Reaction Rate
1.   Nature of reactant (K is more reactive than Na.)
2.   Temperature (RαT i.e. ￿￿T R)
3.   Catalyst (MnO2 ￿ of decomposition of H2O2)
                    R
4.   Concentration (Skittle burns rapidly in 100% O2(g)
     compared to air at 20% O2(g))
5. Surface Area (A crushed antacid tablet reacts quicker
     than a tablet in one piece.)

HW: p. 365 #1-5
 Collision Theory and Rate of Reaction
• R convert to P by successful collisions b/w R
• For a collision to be successful, reactant
  particles must collide with:
  – the minimum energy required to produce P
    particles (Activation Energy)
  – the proper spatial orientation (geometry)

  Rate = (frequency)(% of effective collisions)

  ￿frequency of effective collisions, ￿rate
Consider:   A + B-C        A---B---C          A-B + C
            reactants   activated complex         products
                        (aka transition state)
                    (bonds partially broken/formed)
Temperature of the Reaction System:
  Maxwell-Boltzmann Distribution
          T rate?
Why does ￿ ￿
Effect of Temperature

 as , particles collide more often (￿
ü KE￿                                frequency)
 at
ü a higher temperature, more particles have the required Ea
than at a lower temperature
Effect of Concentration and Surface Area

üincreases frequency of collisions
Effect of Chemical Nature of Reactant

üfewer bonds to be broken = faster rate
üweak bonds – lower threshold energy  lower Ea = faster rate
  stronger bonds – higher threshold energy  higher Ea = slower rate
ücollision geometry may be more difficult to achieve with more complex
molecules and ions
Effect of a Catalyst
üprovides an alternative energy pathway (lower Ea), thereby ￿   %
of effective collisions since more particles possess the required Ea
Maxwell-Boltzmann Distribution and
             Catalysts
Rate = (frequency) (% of effective collisions)

Factors affecting frequency?
                                  temperature
                                  surface area
                                  concentration

Factors affecting % of effective collisions?
                                   temperature
                                   nature of reactants
                                   catalyst
                Homework:
• p. 372 #1-5
Concept Check: Label i, ii, iii.
       Rate Law and Order of Reaction
Given: aA + bB  products
By expt, rate α [A]m[N]n
(Rate law) rate = k[A]m[N]n where k is the rate constant
                                   m, n indicate the sensitivity of the
                                   rate to changes in [A] or [B], not
                               related to coefficients

Order of reaction – the exponents in the rate law

Ex. Given 2 NO(g) + 2 H2(g)  N2(g) + 2 H2O(g)          R = k[NO(g)]2[H2(g)]

The reaction is 2nd order in NO(g) and 1st order in H2(g). The overall
reaction is 3rd order.
Ex 1: Given R= 1.1 x 104 M-2s-1 [BrO3-][HSO3-]2, determine
the rate when [BrO3-] = 0.0020 M and [HSO3-] = 0.0060 M.
Ex 2: Given 2A + B + 3C  products
 The reaction is 1st in [A], 2nd order in [B] and third order
   overall. What is the affect on the rate if:
a) [A] is doubled?
b) [B] is tripled?
c) [C] is doubled?
d) [B] is halved?
Relating Reaction Rate to Time (p. 378)
                 Homework
• p. 380 #1-5 or p. 382 #1-4

• Print Iodine Clock lab from website. Complete
  pre-lab.

  Pre-lab: Title, Question, Hypothesis, Variables,
  Complete concentration calculations in table,
  plan graphical analysis.
analogy   dirty dry dishes  clean wet dishes
          clean wet dishes  clean dry dishes
          dirty dry dishes  clean dry dishes
            Reaction Mechanisms
• Reaction mechanism – a series of elementary steps that
  predict how reactants are converted to products
• elementary steps – one-step process in which product
  particles are (in most cases) the result of collisions b/w 2
  reactant particles (bimolecular). Trimolecular collisions are
  rare.

Ex.    2 NO(g) + H2(g)  N2(g) + H2O2(g)
        H2O2(g) + H2(g)  2 H2O(g)
       2 NO(g) + 2H2(g)  N2(g) + 2 H2O(g)

Rate Determining Step – the slowest step in the mechanism
(determines the rate of the overall reaction)
  Ex.
     Cl2  2 Cl
  Cl + H2  HCl + H
  H + Cl2  HCl + Cl
   Cl + Cl  Cl2


Write the overall equation.
The rate law of the overall reaction is R=k[Cl2]
Which step is the RDS?
              Homework:
• p. 386 #1-3
• p. 387 #1-9
• Unit 3 Summary note

				
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posted:5/15/2014
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