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Rates of Chemical Reaction III

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					Rates of Chemical Reaction III. A Clock Reaction

        In the previous experiment we discussed the factors that influence the rate of a chemical
reaction and presented the terminology used in quantitative relations in studies of the kinetics of
chemical reactions. That material is also pertinent to this experiment and should be studied
before you proceed further.
       This experiment involves the study of the rate properties, or chemical kinetics, of the
following reaction between iodide ion and bromate ion under acidic conditions:
       6 I-(aq) + BrO3-(aq) + 6 H+(aq)  3 I2(aq) + Br-(aq) + 3 H2O                          (1)
        This reaction proceeds reasonably slowly at room temperature, its rate depending on the
concentrations of the I-, Br03-, and H+ ions according to a rate law. For this reaction the rate law
takes the form
       rate = k[I-]m[BrO3-]n[H+]p                                                            (2)
One of the main purposes of the experiment will be to evaluate the rate constant k and the
reaction orders m, n, and p for this reaction. We will also investigate the manner in which the
reaction rate depends on temperature and will evaluate the activation energy Ea. for the reaction.
        Our method for measuring the rate of the reaction involves what is frequently called a
clock reaction. In addition to Reaction 1, whose kinetics we will study, the following reaction
will also be made to occur simultaneously in the reaction flask:
       I2(aq) + 2 S2O32-(aq)  2 I-(aq) + S4062- (aq)                                        (3)
As compared with Equation 1 this reaction is essentially instantaneous. The I2 produced in (1)
reacts completely with the thiosulfate, S2O32-, ion present in the solution, so that until all the
thiosulfate ion has reacted, the concentration of I2 is effectively zero. As soon as the S2O32- is
gone from the system, the I2 produced by (1) remains in the solution and its concentration begins
to increase. The presence of I2 is made strikingly apparent by a starch indicator that is added to
the reaction mixture, since I2 even in small concentrations reacts with starch solution to produce
a blue color.
        By carrying out Reaction 1 in the presence Of S2O32- and a starch indicator, we introduce
a "clock" into the system. Our clock tells us when a given amount of BrO3- ion has reacted (1/6
mole BrO3- per mole S2O32-). which is just what we need to know, since the rate of reaction can
be expressed in terms of the time it takes for a particular amount of BrO3- to be used up. In all
our reactions, the amount of BrO3- that reacts in the time we measure will be constant and small
as compared to the amounts of any of the other reactants. This means that the concentrations of
all reactants will be essentially constant in Equation 2, and hence so will the rate during each
reaction.
        In our experiment we will carry out the reaction between BrO3-, I-, and H+ ions under
different concentration conditions. Measured amounts of each of these ions in water solution
will be mixed in the presence of a constant small amount of S2O32-. The time it takes for each
mixture to turn blue will be measured. The time obtained for each reaction will be inversely
proportional to its rate. By changing the concentration of one reactant and keeping the other


Chemistry 153                                                                                      4-1
concentrations constant, we can investigate how the rate of the reaction varies with the
concentration of a particular reactant. Once we know the order for each reactant we can
determine the rate constant for the reaction.
        In the last part of the experiment we will investigate how the rate of the reaction depends
on temperature. You will recall that in general the rate increases sharply with temperature. By
measuring how the rate varies with temperature we can determine the activation energy, Ea, for
the reaction by making use of the Arrhenius equation:
                           Ea
        log 10 k             constant                                                    (4)
                        2.30RT
In this equation, k is the rate constant at the Kelvin temperature T, Ea is the activation energy,
and R is the gas constant. By plotting log10k against 1/T we should obtain, by Equation 4, a
straight line whose slope equals -Ea/2.30R. From the slope of that line we can easily calculate
the activation energy.



Experimental Procedure
A. Dependence of Reaction Rate on Concentration

Use the methods you have learned in the ealier kinetics experiments to design a procedure. You
need to determine the order of reaction with respect to each ion (I- and BrO3-).
      Plan to prepare about 50 mL of reaction mixture per run, you will need 0.0002 M
       Na2S2O3 and 0.02M HCl in the reaction vessel.
      To carry out the kinetics experiments, find a scheme to mix some of the reagents together
       in each of two flasks so that the reaction doesn’t start until you mix the contents of the
       two separate flasks. Be sure to add starch indicator to one of the flasks or you won’t see
       anything.
      A 10 mL graduated cylinder should provide the necessary precision for your
       measurements. Be sure to rinse the graduated cylinder when measuring different
       reagents!
      Be sure to measure the temperature of the solution at the end of each run.
      It would be wise to run replicate runs of each reaction condition. If two runs don’t agree
       for a given set of conditions, make a note and repeat the run to determine which value is
       correct.




Chemistry 153                                                                                      4-2
You will have access to the following reagents:
      0.010 M KI
      0.0010 M Na2S2O3
      0.040 M KBrO3
      0.10 M HCl
      Starch Indicator Solution


B. Dependence of Reaction Rate on Temperature

        Run the experiment with the reagents at about 20°C, 40°C, 10°C, and 0°C. The exact
temperature doesn’t matter as long it is constant. You will not need to vary the reagent
concentrations for this section – you already know what the reaction order is. Your room
temperature experiments are close enough to 20°C. To run the experiment at higher
temperatures, obtain a large beaker full of warm water (the faucet in the back instrument room
gets really hot!). Before starting the kinetics run, bathe the reagent mixtures in a large beaker of
hot water for several minutes. You can mix hot water, room temperature water, and ice water to
get the other required temperatures. Be sure to get replicate runs that agree at each temperature.


C. Dependence of the Reaction Rate on the Presence of Catalyst

        Some ions have a pronounced catalytic effect on the rates of many reactions in water
solution. Observe the effect on this reaction by once again making up Reaction Mixture 1.
Before mixing, add one drop of 0.5 M (NH4)2MoO4, ammonium molybdate and a few drops of
starch indicator to one of the flasks. Swirl the flask to mix the catalyst thoroughly before starting
the kinetics run.


D. Disposal of Reaction Products.

       The reaction products in this experiment are very dilute and may be poured into the sink
as you complete each part of the experiment.



Data and Calculations: Rates of Chemical Reactions III. A Clock Reaction
The following calculations should help guide you through the calculations necessary to get the
answers needed for the experiment. Be sure to report data for all of your 2nd week runs. If any
runs were identified as bad, include the data along with an explanation of why the run was not
used (don’t include that run in averages!)




Chemistry 153                                                                                    4-3
A. Dependence of Reaction Rate on Concentration

Reaction:
       6 I-(aq) + BrO3-(aq) + 6 H+(aq)  3 I2(aq) + Br-(aq) + 3 H2O                                        (1)
       rate = k[I-]m[BrO3-]n[H+]p =-[BrO3-]/t                                                             (2)
In all the reaction mixtures used in the experiment, the color change occurred when a constant
predetermined number of moles of BrO3- had been used up by the reaction. The color "clock"
allows you to measure the time required for this fixed number of moles of BrO3- to react. The
rate of each reaction is determined by the time t required for the color to change; since in
Equation 2 the change in concentration of BrO3- ion, [BrO3-], is the known for each mixture, the
rate of each reaction is [BrO3-]/t. Something similar to the following table may be helpful for
your calculations:


Run     Time t (sec)           Rate of Reaction (M/s)                  Reactant Concentrations in   Temperature
        for Color to               Rate=[BrO3-]/t                        Reaction Mixture (M)          (°C)
          Change
                                                                        [I-]   [BrO3-]     [H+]
  1
  2
3…


The reactant concentrations in the reaction mixture are not those of the stock solutions, since the
reagents were diluted by the other solutions. To calculate the concentration of a reagent in the
reaction mixture, do the following:
        C stock reagentVstock reagent  C reaction mixtureVreaction mixture
                               C stock reagentVstock reagent
        C reaction mixture 
                                    Vreaction mixture

If you used 8.0 mL of 0.010 M KI stock reagent and made a total volume of 50 mL (ignore the
volume due to indicator), this calculation would be:
                            Cstock reagentVstock reagent 0.01 M 8.0 mL
        Creaction mixture                                                 0.0016 M
                                 Vreaction mixture           50.0 mL
Calculate the rest of the concentrations in the table by the same approach.


Determination of the Orders of the Reaction
You should have designed the reagent concentrations in section A so that from one configuration
to the next, only one reagent’s concentration changes in a manner similar to rate determination
problems from the textbook. Be sure to describe the logic you use to determine in the text of


Chemistry 153                                                                                                     4-4
your report. Once the reaction order is determined, calculate k for each run using Equation 2.
These values should be the same within experimental error – report the average, standard
deviation, and relative standard deviation for this result.


B. Effect of Temperature on Reaction Rate: The Activation Energy
To find the activation energy for the reaction it will be helpful to complete the following table:
                                                                Approximate
                                                                Temperature (°C)
                                                                0     10    20    40
       a. Time, t, for color to appear (sec)
       b. Actual Temperature of the reaction mixture (°C)
       c. Rate constant, k ( units depend on rate law)
       d. Log10(k) (drop/ignore the units)
       e. Actual Temperature (K)
       f. 1/T (K)


The dependence of the rate constant, k' for a reaction is given by Equation 4:
                           Ea
        log 10 k             constant                                                      (4)
                        2.30RT
Find the slope of the line obtained by drawing the best straight line through the experimental
points. You may either plot the average value of k at each temperature or extend the table to
include all of your runs. The slope of the line equals –Ea/2.3R, where R = 8.314 joules/mole K
if Ea is to be in joules per mole. Find the activation energy, Ea, for the reaction.


C. Effect of a Catalyst on Reaction Rate
Report the rate constant for your catalyzed reaction runs. Would you expect the activation
energy, Ea, for the catalyzed reaction to be greater than, less than, or equal to the activation
energy for the uncatalyzed reaction? Why?




Chemistry 153                                                                                        4-5
Prelab Assignment: Rates of Chemical Reaction III. A Clock Reaction
Answer these questions in your laboratory notebook. The instructor will check and initial your
Prelab at the beginning of the experiment.
1. A student studied the clock reaction described in this experiment. She set up a Reaction
Mixture by mixing 20 mL 0.010 M KI, 10 mL 0.001 M NaS2O3, 10 mL 0.040 M KBrO3, and 10
mL 0.10 M HCl. It took about 45 seconds for the color to turn blue.
a. She found the concentrations of each reactant in the reacting mixture by realizing that the
   number of moles of each reactant did not change when that reactant was mixed with the
   others, but that its concentration did. For any reactant A,
      no. moles A = MA stock  Vstock = MA mixture  Vmixture
   The volume of the mixture was 50 mL. Revising the above equation, she obtained
      MA mixture = MA stock  Vstock (mL)/50 mL
   Find the concentrations of each reactant by using the above equation.




       [I-] = _________M; [BrO3-] = _________M; [H+] = _________M


b. What was the rate constant of the reaction (k)?


       k = __________
c. Knowing the rate of reaction for the reaction mixture and the concentrations of I-, BrO3- and
   H+ in that mixture, she was able to set up Equation 2 for the relative rate of the reaction. The
   only quantities that remained unknown were k, m, n, and p. Set up Equation 2 as she did,
   presuming she did it properly.


   Rate =


2. For another reaction mixture (made by mixing 10 mL 0.010 M KI, 10 mL 0.001 M NaS2O3,
10 mL 0.040 M KBrO3, and 10 mL H2O) the student found that 85 seconds were required. What
is the value of m (remember that reaction orders are often whole numbers)?




       m = __________




Chemistry 153                                                                                    4-6

				
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